38 Sr 87.62 56 Ba 137.3 88 Ra (226)
37 Rb 85.47 55 Cs 132.9 87 Fr (223)
104 Rf (261)
72 Hf 178.5
40 Zr 91.22
59 Pr 140.9 91 Pa 231.0
90 Th 232.0
106 Sg (266)
74 W 183.9
42 Mo 95.94
24 Cr 52.00
58 Ce 140.1
105 Db (262)
73 Ta 180.9
41 Nb 92.91
23 V 50.94
6 6B
92 U 238.0
60 Nd 144.2
107 Bh (264)
75 Re 186.2
43 Tc (98)
25 Mn 54.94
7 7B
93 Np (237)
61 Pm (145)
108 Hs (269)
76 Os 190.2
44 Ru 101.1
26 Fe 55.85
8
94 Pu (244)
62 Sm 150.4
109 Mt (268)
77 Ir 192.2
45 Rh 102.9
27 Co 58.93
9 8B
MD DALIM 1123973B 12/13/10 CYAN MAG YELO BLACK
30 Zn 65.39 48 Cd 112.4 80 Hg 200.6
29 Cu 63.55 47 Ag 107.9 79 Au 197.0 111 Rg (280)
64 Gd 157.3 96 Cm (247)
28 Ni 58.69 46 Pd 106.4 78 Pt 195.1
110 Ds (271)
63 Eu 152.0 95 Am (243)
97 Bk (247)
65 Tb 158.9
112 Cn (285)
12 2B
11 1B 10
98 Cf (251)
66 Dy 162.5
113
81 Tl 204.4
49 In 114.8
31 Ga 69.72
99 Es (252)
67 Ho 164.9
114
82 Pb 207.2
50 Sn 118.7
32 Ge 72.61
14 Si 28.09
13 Al 26.98
100 Fm (257)
68 Er 167.3
115
83 Bi 209.0
51 Sb 121.8
33 As 74.92
15 P 30.97
7 N 14.01
15 5A
101 Md (258)
69 Tm 168.9
116
84 Po (210)
52 Te 127.6
34 Se 78.96
16 S 32.07
8 O 16.00
16 6A
102 No (259)
70 Yb 173.0
117
85 At (210)
53 I 126.9
35 Br 79.90
17 Cl 35.45
9 F 19.00
17 7A
103 Lr (262)
71 Lu 175.0
118
86 Rn (222)
54 Xe 131.3
36 Kr 83.80
18 Ar 39.95
10 Ne 20.18
2 He 4.003
18 8A
The 1–18 group designation has been recommended by the International Union of Pure and Applied Chemistry (IUPAC) but is not yet in wide use. In this text we use the standard U.S. notation for group numbers (1A–8A and 1B–8B). No name has been assigned for element 112. Elements 113–118 have not yet been synthesized.
Nonmetals
Metalloids
89 Ac (227)
57 La 138.9
39 Y 88.91
22 Ti 47.88
5 5B
*Source: Adapted from the Twelve Principles of Green Chemistry, by Paul Anastas and John Warner.
21 Sc 44.96
4 4B
useful life.
Metals
20 Ca 40.08 19 K 39.10
3 3B
6 C 12.01
5 B 10.81
4 It is better to use less energy.
12 Mg 24.31
14 4A
6 It is better to use materials that degrade into innocuous products at the end of their 13 3A
5 It is better to use renewable materials.
Atomic mass
Atomic number
3 It is better to use and generate substances that are not toxic.
11 Na 22.99
24 Cr 52.00
2 It is better to minimize the amount of materials used in the production of a product.
4 Be 9.012
2 2A
1 It is better to prevent waste than to treat or clean up waste after it is formed.
3 Li 6.941
1 H 1.008
1 1A
Principles of Green Chemistry*
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A Project of the American Chemical Society
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Seventh Edition
Applying Chemistry to Society Catherine H. Middlecamp University of Wisconsin–Madison
Steven W. Keller University of Missouri
Karen L. Anderson Madison Area Technical College
Anne K. Bentley Lewis & Clark College
Michael C. Cann University of Scranton
Jamie P. Ellis The Scripps Research Institute A Project of the American Chemical Society
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CHEMISTRY IN CONTEXT: APPLYING CHEMISTRY TO SOCIETY, SEVENTH EDITION Published by McGraw-Hill, a business unit of The McGraw-Hill Companies, Inc., 1221 Avenue of the Americas, New York, NY 10020. Copyright © 2012 by the American Chemical Society. All rights reserved. Previous editions © 2009, 2006, and 2003. No part of this publication may be reproduced or distributed in any form or by any means, or stored in a database or retrieval system, without the prior written consent of The McGraw-Hill Companies, Inc., including, but not limited to, in any network or other electronic storage or transmission, or broadcast for distance learning. Some ancillaries, including electronic and print components, may not be available to customers outside the United States. This book is printed on acid-free paper. 1 2 3 4 5 6 7 8 9 0 DOW/DOW 1 0 9 8 7 6 5 4 3 2 1 ISBN 978–0–07–337566–3 MHID 0–07–337566–7 Vice President, Editor-in-Chief: Marty Lange Vice President, EDP: Kimberly Meriwether David Senior Director of Development: Kristine Tibbetts Publisher: Ryan Blankenship Sponsoring Editor: Todd L. Turner Developmental Editor: Jodi Rhomberg Executive Marketing Manager: Tamara L. Hodge Lead Project Manager: Sheila M. Frank Senior Buyer: Kara Kudronowicz Senior Media Project Manager: Sandra M. Schnee Designer: Tara McDermott Cover Designer: Christopher Reese Cover Image: © Getty Images/Grant Faint Lead Photo Research Coordinator: Carrie K. Burger Photo Research: Jerry Marshall/pictureresearching.com Compositor: Aptara, Inc. Typeface: 10/12 Times Roman Printer: R. R. Donnelley All credits appearing on page or at the end of the book are considered to be an extension of the copyright page. Library of Congress Cataloging-in-Publication Data Chemistry in context : applying chemistry to society / Catherine H. Middlecamp ... [et al.]. -- 7th ed. p. cm. Includes index. “A Product of the American Chemical Society.” ISBN 978–0–07–337566–3 — ISBN 0–07–337566–7 (hard copy : alk. paper) 1. Biochemistry. 2. Environmental chemistry. 3. Geochemistry. I. Middlecamp, Catherine. QD415.C482 2012 540--dc22 2010024045
www.mhhe.com
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Brief Contents 0 Chemistry for a Sustainable Future 2 1 The Air We Breathe 16 2 Protecting the Ozone Layer 64 3 The Chemistry of Global Climate Change 106 4 Energy from Combustion 154 5 Water for Life 198 6 Neutralizing the Threat of Acid Rain 242 7 The Fires of Nuclear Fission 282 8 Energy from Electron Transfer 330 9 The World of Polymers and Plastics 368 10 Manipulating Molecules and Designing Drugs 404 11 Nutrition: Food for Thought 446 12 Genetic Engineering and the Molecules of Life 490 Appendices 1 Measure for Measure: Conversion Factors and Constants A1 2 The Power of Exponents A2 3 Clearing the Logjam A3 4 Answers to Your Turn Questions Not Answered in Text A4 5 Answers to Selected End-of-Chapter Questions Indicated in Color in the Text A16
Glossary G1 Credits C1 Index I1 v
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The World of Polymers and Plastics
Contents Preface xi
Chapter 2
Chapter 0
Protecting the Ozone Layer 64
Chemistry for a Sustainable Future 2
2.1
Ozone: What and Where Is It? 66
2.2
Atomic Structure and Periodicity 68
0.1
The Choices We Make Today 4
2.3
Molecules and Models 71
0.2
The Sustainable Practices We Need for Tomorrow 5
2.4
Waves of Light 76
0.3
The Triple Bottom Line 7
0.4
Cradle-to-Where? 8
0.5
Your Ecological Footprint 10
0.6
Our Responsibilities as Citizens and Chemists 11
0.7
Back to the Blue Marble 13
Chapter 1
2.5 Radiation and Matter 78 2.6
The Oxygen–Ozone Screen 80
2.7
Biological Effects of Ultraviolet Radiation 82
2.8
Stratospheric Ozone Destruction: Global Observations and Causes 86
2.9
Chlorofluorocarbons: Properties, Uses, and Interactions with Ozone 89
2.10
The Antarctic Ozone Hole: A Closer Look 93
2.11
Responses to a Global Concern 94
2.12
Replacements for CFCs 97
The Air We Breathe 16 1.1 What’s in a Breath? 17 1.2 What Else Is in a Breath? 20 1.3
Air Pollutants and Risk Assessment 23
1.4
Air Quality and You 26
1.5
Where We Live: The Troposphere 29
1.6
Classifying Matter: Pure Substances, Elements, and Compounds 30
1.7
Atoms and Molecules 33
1.8
Names and Formulas: The Vocabulary of Chemistry 35
1.9
Chemical Change: The Role of Oxygen in Burning 37
1.10
Fire and Fuel: Air Quality and Burning Hydrocarbons 40
1.11
Air Pollutants: Direct Sources 42
1.12
Ozone: A Secondary Pollutant 47
1.13
The Inside Story of Air Quality 49
1.14
Back to the Breath— at the Molecular Level 54
Conclusion 57 Chapter Summary 57 Questions 58
Conclusion 101 Chapter Summary 101 Questions 102
Chapter 3 The Chemistry of Global Climate Change 106 3.1
In the Greenhouse: Earth’s Energy Balance 108
3.2
Gathering Evidence: The Testimony of Time 111
3.3 Molecules: How They Shape Up 116 3.4
Vibrating Molecules and the Greenhouse Effect 120
3.5
The Carbon Cycle: Contributions from Nature and Humans 123
3.6
Quantitative Concepts: Mass 125
3.7 Quantitative Concepts: Molecules and Moles 127 3.8
Methane and Other Greenhouse Gases 130
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Contents 3.9
How Warm Will the Planet Get? 133
5.12 Water Solutions for Global Challenges 233
3.10
The Consequences of Climate Change 138
3.11
What Can (or Should) We Do About Climate Change 142
Conclusion 237 Chapter Summary 238 Questions 238
Conclusion 148 Chapter Summary 148 Questions 149
Chapter 6 Neutralizing the Threat of Acid Rain 242 6.1
Chapter 4 Energy from Combustion 154
What Is an Acid? 244
6.2 What Is a Base? 245 6.3
Neutralization: Bases Are Antacids 247 Introducing pH 249
4.1
Fossil Fuels and Electricity 156
6.4
4.2
Efficiency of Energy Transformation 159
6.5
Ocean Acidification 250
6.6
The Challenges of Measuring the pH of Rain 252
6.7
Sulfur Dioxide and the Combustion of Coal 257
4.3
The Chemistry of Coal 161
4.4
Petroleum 166
4.5
Measuring Energy Changes 171
4.6
Energy Changes at the Molecular Level 175
6.8
Nitrogen Oxides and the Combustion of Gasoline 259 The Nitrogen Cycle 260
4.7
The Chemistry of Gasoline 178
6.9
4.8
New Uses for an Old Fuel 181
6.10
4.9
Biofuels I—Ethanol 183
4.10
Biofuels II—Biodiesel, Garbage, and Biogas 187
4.11 The Way Forward 190 Conclusion 193 Chapter Summary 193 Questions 194
Chapter 5 Water for Life 198
SO2 and NOx —How Do They Stack Up? 264
6.11 Acid Deposition and Its Effects on Materials 266 6.12 Acid Deposition, Haze, and Human Health 270 6.13 Damage to Lakes and Streams 274 Conclusion 276 Chapter Summary 276 Questions 277
Chapter 7
5.1
The Unique Properties of Water 200
5.2
The Role of Hydrogen Bonding 202
7.1
Nuclear Power Worldwide 283
7.2
How Fission Produces Energy 286
Water Use 204
7.3
How Nuclear Reactors Produce Electricity 292
7.4
What Is Radioactivity? 295
5.3
The Fires of Nuclear Fission 282
5.4 Water Issues 207 5.5
Aqueous Solutions 212
5.6 A Close Look at Solutes 215 5.7
Names and Formulas of Ionic Compounds 218
7.5 Looking Backward to Go Forward 298 7.6
Radioactivity and You 302
7.7
The Weapons Connection 308 Nuclear Time: The Half-Life 311
5.8
The Ocean—An Aqueous Solution with Many Ions 221
7.8
5.9
Covalent Compounds and Their Solutions 223
7.9 Nuclear Waste: Here Today, Here Tomorrow 315
5.10 Protecting Our Drinking Water: Federal Legislation 226
7.10 Risks and Benefits of Nuclear Power 319
5.11
7.11
Water Treatment 230
A Future for Nuclear Power 321
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Contents Conclusion 324 Chapter Summary 324 Questions 325
Chapter 8 Energy from Electron Transfer 330 8.1
Batteries, Galvanic Cells, and Electrons 332
8.2
Other Common Galvanic Cells 335
8.3
Battery Ingredients: Cradle-to-Cradle 338
8.4 Hybrid Vehicles 341 8.5 Fuel Cells: The Basics 344 8.6
Hydrogen for Fuel Cell Vehicles 348
8.7 Photovoltaic Cells: The Basics 352 8.8
Electricity from Renewable (Sustainable) Sources 360
Conclusion 363 Chapter Summary 363 Questions 364
Chapter 9 The World of Polymers and Plastics 368
10.6
Give These Molecules a Hand! 423
10.7
Steroids 427
10.8 Prescription, Generic, and Overthe-Counter Medicines 429 10.9
Herbal Medicine 432
10.10 Drugs of Abuse 435 Conclusion 439 Chapter Summary 440 Questions 440
Chapter 11 Nutrition: Food for Thought 446 11.1 Food and the Planet 448 11.2
You Are What You Eat 450
11.3
Fats and Oils 452
11.4
Fats, Oils, and Your Diet 457
11.5
Carbohydrates: Sweet and Starchy 461
11.6 How Sweet It Is: Sugars and Sugar Substitutes 464 11.7
Proteins: First Among Equals 467
11.8
Vitamins and Minerals: The Other Essentials 470
11.9 Energy from Food 474 11.10 Quality Versus Quantity: Dietary Advice 478
9.1
Polymers: Long, Long Chains 370
11.11 Eat Local? Eat Veggies? 480
9.2
Adding up the Monomers 371
11.12 Feeding a Hungry World 483
9.3 Polyethylene: A Closer Look 374 9.4
The “Big Six”: Theme and Variations 377
Conclusion 485 Chapter Summary 485 Questions 486
9.5
Condensing the Monomers 382
9.6
Polyamides: Natural and Nylon 385
Chapter 12
9.7
Recycling: The Big Picture 388
Genetic Engineering and the Molecules of Life 490
9.8 Recycling Plastics: The Details 391 Conclusion 398 Chapter Summary 398 Questions 398
12.1
Stronger and Better Corn Plants? 491
12.2
A Chemical That Codes Life 493
12.3
The Double Helix of DNA 497
Chapter 10
12.4 Cracking the Chemical Code 502
Manipulating Molecules and Designing Drugs 404
12.5 Proteins: Form to Function 503
10.1 A Classic Wonder Drug 405 10.2
The Study of Carbon-Containing Molecules 407
10.3
Functional Groups 411
10.4
How Aspirin Works: Function Follows Form 416
10.5
Modern Drug Design 418
12.6
The Process of Genetic Engineering 507
12.7
Making Chemical Synthesis Green from Genetic Modification 511
12.8
The New Frankenstein 513
Conclusion 515 Chapter Summary 516 Questions 516
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Contents
Appendix 1
Appendix 5
Measure for Measure: Conversion Factors and Constants A1
Answers to Selected End-of-Chapter Questions Indicated in Color in the Text A16
Appendix 2
Glossary G1 Credits C1 Index I1
The Power of Exponents A2
Appendix 3 Clearing the Logjam A3
Appendix 4 Answers to Your Turn Questions Not Answered in Text A4
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The World of Polymers and Plastics
Preface Dear Readers, When first published in 1993, Chemistry in Context was “the book that broke the mold.” Unlike the books of its time, it did not teach chemistry in isolation from people and the real-world issues they were facing. Similarly, it did not introduce a fact or concept for the sake of “covering it” as part of the curriculum. Rather, Chemistry in Context carefully matched each chemical principle to a real-world issue such as air quality, energy, or water use. Each was introduced on a need-to-know basis; that is, at the point in the book at which there was a demonstrated need for the principle. Most importantly, the book presented chemistry in the context of significant social, political, economic, and ethical issues. Context! The word derives from the Latin word meaning “to weave.” The spider web motif on the Chemistry in Context cover exemplifies the complex connections that can be woven between chemistry and society. In the absence of the real-world issues, there could be no Chemistry in Context. Similarly, without teachers and students who were willing (and brave enough) to engage in these issues, there could be no Chemistry in Context. Together we weave chemistry into the issues that we face in our lives. Context! Today we also know that teaching in context is a high-impact practice backed by the research on how people learn. Chemistry in Context uses real-world contexts that engage students on multiple levels: their individual health and well-being, the health of their local communities, and the health of wider ecosystems that sustain life on this planet. As we planned this edition, the writing team members questioned how a tradition of “breaking the mold” might best be continued today. The team raised the question not only for the sake of keeping with tradition (and for the fun of breaking molds), but also for a compelling reason: the needs of our readers. We wanted to continue to find ways of communicating chemistry that served our students, given the challenging issues they face today, the complex needs of the societies in which they live, and the changing landscapes on which they will work in the future.
Teaching (and Learning) in Context The organization of Chemistry in Context has remained the same in every edition. The first six chapters form a core through which basic chemical principles are introduced. These chapters provide a coherent strand of topics that focus on a single theme—the environment. They develop a foundation of chemical concepts that can be expanded in subsequent chapters. Chapters 7 and 8 consider alternative (non-fossil fuel) energy sources—nuclear power, batteries, fuel cells, and the hydrogen economy. The remaining chapters are carbon-based, focusing on polymers, drugs, food production, and genetic engineering. They provide students with the opportunity to explore interests, as time permits, beyond the core topics.
Sustainability—New Content Global sustainability is not just a challenge; rather, it is the defining challenge of our century. The seventh edition of Chemistry in Context is designed to help students better meet this challenge. With its new Chapter 0, “Chemistry for a Sustainable Future,” our intent was to establish sustainability as a core, normative part of the chemistry curriculum and part of the foundational learning. Sustainability adds a new degree of complexity to Chemistry in Context. In part, the complexity arises because sustainability can be conceptualized in two ways: as a topic worth studying and as a problem worth solving. As a topic, sustainability provides a new xi
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body of content for students to master. For example, the tragedy of the commons, the Triple Bottom Line, and the concept of cradle-to-cradle are part of this new body of content and are introduced in Chapter 0. As a problem worth solving, sustainability generates new questions for students to ask—ones that help them to imagine and achieve a sustainable future. For example, students will find questions about the risks and benefits of both acting and of not acting on behalf of future generations. To incorporate sustainability, then, requires more than a casual rethinking of the curriculum. How do you teach and learn about something as complicated as sustainability? In responding to this question, the author team realized that it was necessary both to update the material and to recast it in a new light. Here are some examples of the changes that the team made:
WAT E R F O R L I F E 2005 – 2015
Chapter 3, The Chemistry of Global Climate Change, was updated in the light of new developments in climate change science. It now clearly outlines the consequences of climate change, introducing the sustainability concept of external costs. Chapter 5, Water for Life, now connects to the “Water for Life” decade themes of the United Nations: the scarcity of fresh water, sustainable management of water resources, and water contamination. Chapter 8, Energy from Electron Transfer, was recast to better show the match between our energy needs and the available technologies. The sustainability concept of cradle-to-cradle, introduced earlier in Chapter 0, is connected to battery design. Chapter 11, Nutrition, Food for Thought, points out that what you eat affects both your health and the health of the planet. In addition, here is a listing of the sustainability concepts and the chapters in which they can be found: Tragedy of the commons: Triple Bottom Line: Cradle-to-cradle: External costs: Environmental footprints:
Chapters Chapters Chapters Chapters Chapters
0, 0, 0, 0, 0,
1, 2, 5, and 6 1, 4, 5, 6, and 7 7, and, 8 3, and, 8 1, 4, 5, 11, and 12
Green chemistry, a means to sustainability, continues to be an important theme in Chemistry in Context. As in previous editions, examples of green chemistry are highlighted in each chapter. In this new edition, look for even more examples. This expanded coverage offers the reader an even better sense of the need for and the importance of greening our chemical processes. For easier access, the principles of green chemistry are now listed on the inside front cover of the text.
Updates to Existing Content People sometimes ask us “Why do you release new editions so often?” We also hear the question “Would it work to keep using the older edition with our students?” Indeed, we are on a fast publishing cycle, turning out a new version every three years. We do this because Chemistry in Context, with its current real-world focus, risks being out-of-date the very day it is published. Given this, we strongly urge instructors to switch to the new edition immediately, given how sensitive the real-world content is to the passing of time. With each new addition, the author team reworks practically every chapter, updating its content and focus to reflect new scientific developments, changes in policies, energy trends, and current world events. These updates are nontrivial to implement. They involve writing new content as well as producing new graphs and data tables. The issues that we select to “hook” the reader at the start of the chapter also are recast from edition to edition. For example, in this new edition, the “Water Chapter,” Chapter 5, underwent a significant revision. This chapter has been on our list for a makeover for many years, and each successive author team has puzzled over how best to refocus it. Changes in earlier editions had switched the chapter hook to the issue of bottled water vs. tap water; but even with this change, we knew there was more work to be done. In the seventh edition, we finally reworked Chapter 5 from start to finish. The chapter now
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follows the United Nations theme of “Water for Life,” highlighting that fresh water is indeed limited on our planet and all need access to it. Safe drinking water, water footprints, and water for agriculture and industry are the keys to our health and prosperity. From this point of departure, all of the chemical principles fall nicely into place. The “Food Chapter,” Chapter 11, also underwent a major revision. In past editions, we have employed several different hooks to open the chapter. For example, we talked about different diets, including the high protein one. We also talked about the dizzying array of food recommendations. But for this new edition, we went with a new theme that couples food, personal health, and the health of the planet. This theme better connects to issues of energy consumption, water use, land use, and public health. The questions of sustainability also more naturally arise. In addition, this new theme well sets the stage for Chapter 12 on genetic engineering. Speaking of Chapter 12, we did a careful technical review of this chapter, reworking its content to better reflect what we know about the genetic code and recent developments in the field. Chapter 8 on batteries and alternative energy sources also received careful attention to its technical content. Chapter 3 was updated to reflect the latest climate change science, including that we changed its title from “The Chemistry of Global Warming” to “The Chemistry of Climate Change.” Chapter 6, the “Acid Rain” chapter, now opens with ocean acidification, connecting emissions from combustion to the coral reefs and changes in sea water. And Chapter 1 now has an expanded section on indoor air quality, including a green chemistry section about paints that emit fewer volatile organic compounds while drying.
CO2
CO2
H2O
H2CO3
H CO32
CaCO3 (coral)
H
HCO3
HCO3
Ca2 CO32
Figure 6.6 Chemistry of CO2 in the ocean.
Updates to the Laboratory Manual As we rethought the chapters of Chemistry in Context with an eye to sustainability, we recognized that we also needed to rethink the accompanying Laboratory Manual. Both the introductions to the experiments and the post-lab questions were revised in order to emphasize and reinforce the environmental issues and the sustainability concepts
xiii
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presented in the textbook. The experiments were redesigned to include “greener” procedures, such as a microwave synthesis of aspirin and an investigation of the properties of a biodegradable polymer. Additionally, pre-lab questions were added and expanded data sheets encourage students to form and test hypotheses during each experiment.
The New Edition—A Team Effort It is always a pleasure to bring a new textbook or new edition to fruition. But the work is not done by just one individual; rather, it is the work of many talented individuals. The seventh edition builds on the tradition of prior author teams led by A. Truman Schwartz, Macalester College, Conrad L. Stanitski, University of Central Arkansas, and Lucy Pryde Eubanks, Clemson University. This edition, we were fortunate to have the leadership and encouragement of Mary Kirchhoff, the Director of the ACS Education Division. We also recognize the able assistance of ACS staff members Marta Gmurczyk, Michael Mury, Jerry Bell, and Corrie Kuniyoshi. This new edition was prepared by a team of writers: Cathy Middlecamp, Steve Keller, Karen Anderson, Anne Bentley, Michael Cann, and Jamie Ellis. The laboratory manual to accompany it was revised and updated by Jennifer Tripp. Each of us brought different expertise to the project. In common, though we brought our goodwill, hopes, dreams, and seemingly boundless enthusiasm to bring real-world chemistry into the classroom. The McGraw-Hill team was superb in all aspects of this project. Marty Lange (Vice President, Editor-in-Chief), Ryan Blankenship (Publisher), Todd Turner (Sponsoring Editor), and Jodi Rhomberg (Developmental Editor) led this outstanding team. Tami Hodge served as the Executive Marketing Manger. The Lead Project Manager was Sheila Frank, who coordinated the production team of Carrie Burger (Lead Photo Research Coordinator), Jerry Marshall (Photo Research), Tara McDermott (Designer), Kara Kudronowicz (Senior Buyer), and Mary Jane Lampe (Project Coordinator). The Digital Product Manager was Daryl Bruflodt, and Sandra Schnee served as Senior Media Project Managers. The team also benefited from the careful editing of Linda Davoli. The author team truly benefitted from the expertise of a wider community. We extend our thanks to the following individuals for the technical expertise they provided to us in preparing the manuscript: Mark E. Anderson, University of Wisconsin–Madison David Argentar, Sun Edge, LLC Marion O’Leary, Carnegie Institution for Science Ross Salawitch, University of Maryland Kenneth A. Walz, Madison Area Technical College We also acknowledge those who served as reviewers for the new edition: Bill Blanken Michael Doescher Jeannine Eddleton Denise V. Greathouse Susan J. Glenn Narayan S. Hosmane Donna K. Howell Richard Kashmar John Kirk Amy Lumley Mya A. Norman Marion O’Leary Heather U. Price Victor Ryzhov William L. Schreiber Lee Alan Shaver Jeff Shen
Chapman University, Lemoore Campus Benedictine College Virginia Tech University of Arkansas, Fayetteville University of South Carolina, Aiken Northern Illinois University Park University Wesley College University of Iowa Coffeyville Community College University of Arkansas–Fayetteville Carnegie Institution for Science Highline Community College Northern Illinois University Monmouth University Washburn University Park University
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Preface
Carnetta Skipworth Laura K. Stultz Sara J. E. Sutcliffe Burton Tropp
Western Kentucky University Birmingham–Southern College University of Texas, Austin Queens College, CUNY
Finally, we would like to thank Andy Jorgensen, University of Toledo, for his expertise developing Connect problems and to Maggie Phillips, Great Lakes Bioenergy Research Center College of Agricultural, and Life Sciences, University of Wisconsin–Madison, for her work developing Figures Alive! Connect problems.
Wishing Our Readers Well We are very excited by the features of this new edition that continue to “break the mold” in bringing chemistry to you, our reader. We selected engaging and timely topics that we hope will serve you not only today, but also in the years to come. At the same time, we strove to be honest to the science behind these topics. We wish you well as you read, explore the issues, argue with each other (and with the authors) and, most importantly, as you use what you learn to bring your dreams to reality. Sincerely, and with all good wishes from the author team,
Cathy Middlecamp Senior Author and Editor-in-Chief October 2010
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From the Publisher McGraw-Hill Education, a division of The McGraw-Hill Companies, is proud to bring you this new edition of Chemistry in Context. We are a leading innovator in the development of teaching and learning solutions for the 21st century. Through a comprehensive range of traditional and digital education content and tools, McGraw-Hill Education empowers and prepares professionals and students of all ages to connect, learn, and succeed in the global economy. Chemistry in Context is supported by a complete package for instructors and students. Several print and media supplements have been prepared to accompany the text and make learning as meaningful and up-to-date as possible. McGraw-Hill Connect Chemistry
McGraw-Hill Connect™ Chemistry (www.mcgrawhillconnect.com), is a web-based assignment and assessment platform that gives students the means to better connect with their coursework, with their instructors, and with the important concepts that they will need to know for success now and in the future. With Connect Chemistry, your instructor can deliver assignments, quizzes, and tests online. Questions from the text are presented in an auto-gradable format and tied to the text’s learning objectives. Instructors can edit existing questions and author entirely new problems. They also can track individual student performance—by question, assignment, or in relation to the class overall—with detailed grade reports. Integrate grade reports easily with Learning Management Systems (LMS) such as WebCT and Blackboard. And much more. By choosing Connect Chemistry, instructors are providing their students with a powerful tool for improving academic performance and truly mastering course material. Connect Chemistry allows students to practice important skills at their own pace and on their own schedule. Importantly, students’ assessment results and instructors’ feedback are all saved online—so students can continually review their progress and plot their course to success. Like Connect Chemistry, McGraw-Hill ConnectPlus™ Chemistry provides students with online assignments and assessments, plus 24/7 online access to an eBook—an
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Preface
online edition of the text—to aid them in successfully completing their work, wherever and whenever they choose. McGraw-Hill Higher Education and Blackboard® McGraw-Hill Higher Education and Blackboard have teamed up! What does this mean for you?
1. Your life, simplified. Now you and your students can access McGraw-Hill’s Connect and Create™ right from within your Blackboard course – all with one single sign-on. Say goodbye to the days of logging in to multiple applications. 2. Deep integration of content and tools. Not only do you get single sign-on with Connect and Create, you also get deep integration of McGraw-Hill content and content engines right in Blackboard. Whether you’re choosing a book for your course or building Connect assignments, all the tools you need are right where you want them – inside of Blackboard. 3. Seamless Gradebooks. Are you tired of keeping multiple gradebooks and manually synchronizing grades into Blackboard? We thought so. When a student completes an integrated Connect assignment, the grade for that assignment automatically (and instantly) feeds your Blackboard grade center. 4. A solution for everyone. Whether your institution is already using Blackboard or you just want to try Blackboard on your own, we have a solution for you. McGraw-Hill and Blackboard can now offer you easy access to industry leading technology and content, whether your campus hosts it, or we do. Be sure to ask your local McGraw-Hill representative for details. Website to Accompany the Text Throughout Chemistry in Context, you will see references to the textbook website, www.mhhe.com/cic. For example, many of the Consider This activities embedded in each chapter reference technical websites with data, such as the stratospheric ozone data from NASA (U.S. National Aeronautics and Space Administration) and the air quality data from the EPA (U.S. Environmental Protection Agency). On the Chemistry in Context website, students and instructors can access these websites with a single click. Figures Alive! Figures Alive! is a set of interactive, web-based activities, also available on the textbook website. Each one is keyed to a figure in Chemistry in Context and leads the student through the discovery of various layers of knowledge inherent in the figure. Look for this icon in the textbook. The self-testing segments built into Figures Alive! are based on the same categories as the chapter-end problems—Emphasizing Essentials, Concentrating on Concepts, and Exploring Extensions.
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Digital Lecture Capture: Tegrity McGraw-Hill Tegrity Campus records and distributes your lecture with just a click of a button. Students can view anytime/anywhere via computer, iPod, or mobile device. Tegrity indexes as it records your slideshow presentations and anything shown on your computer, so students can use keywords to find exactly what they want to study.
Also for the Instructor Test Bank: The electronic test bank offers a bank of questions that can be used for homework assignments or the preparation of exams. The test bank can be utilized to quickly create customized exams. It allows instructors to sort questions by format or level of difficulty; edit existing questions or add new ones; and scramble questions and answer keys for multiple versions of the same test. McGraw-Hill Presentation Center: Build instructional material wherever, whenever, and however you want! McGraw-Hill Presentation Center is an online digital library containing assets such as photos, artwork, and other media types that can be used to create customized lectures, visually enhanced tests and quizzes, compelling course websites, or attractive printed support materials. The McGraw-Hill Presentation Center Library includes thousands of assets from many McGraw-Hill titles. This ever-growing resource gives instructors the power to utilize assets specific to an adopted textbook as well as content from all other books in the library. The Presentation Center can be accessed from the instructor side of your textbook’s Connect website, and the Presentation Center’s dynamic search engine allows you to explore by discipline, course, textbook chapter, asset type, or keyword. Simply browse, select, and download the files you need to build engaging course materials. All assets are copyrighted by McGrawHill Higher Education but can be used by instructors for classroom purposes.
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A Project of the American Chemical Society
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Chemistry for a Sustainable Future
The “blue marble,” our Earth, as seen from outer space. “The first day or so, we all pointed to our countries. The third or fourth day, we were pointing to our continents. By the fifth day, we were aware of only one Earth.” Prince Sultan bin Salman Al Sa’ud, Astronaut, Saudi Arabia, 1985.
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Only one Earth. From the vantage point of outer space, the planet we call home is truly magnificent—a “blue marble” of water, land, and clouds. In 1972, the crew of the Apollo 17 spacecraft photographed the Earth at a distance of about 28,000 miles (45,000 kilometers). In the words of Soviet cosmonaut Aleksei Leonov, “the Earth was small, light blue, and so touchingly alone.” Are we alone in the universe? Possibly. Clearly, though, we are not alone on our planet. We share it with other creatures large and small. Biologists estimate that upwards of 1.5 million species exist in addition to our own. Some help to feed and sustain us. Others contribute to our well-being. Still others (like mosquitoes) annoy and perhaps even sicken us. We also share the planet with just short of 7 billion other people. Over the past century, the human population on Earth has more than tripled, an unprecedented growth spurt in the history of our planet. By 2050, the population may grow by another 2 to 3 billion. Large and small, all of the species on our planet somehow connect. Exactly how this happens, however, may not be so obvious to us. For example, unseen microorganisms shuttle nitrogen from one chemical form to another, providing nutrients for green plants to grow. These plants harness the light energy of the Sun during the process of photosynthesis. Using this energy, they convert the compounds carbon dioxide and water to that of glucose, their food source. At the same time, they release the element oxygen into the air that we breathe. And we humans are host to countless microorganisms that have taken up residence in our skin and internal organs. Large and small, these connections are breaking at an alarming rate. Our grandparents used to be able to tell stories of “how it was” when they fished and hunted. Consider, for example, that salmon once ran the rocky streams of the North Atlantic coast. Sea captains once reported large fish in such abundance that they literally could be taken by hand. Abalone used to be present everywhere off the Pacific coasts of North America, rather than in the few good spots boats can find today. However, nobody now living remembers fish in this abundance. A change has occurred in our perspective. Today, we are accustomed to reading reports of declining fish populations and of endangered species. Have you run across the term shifting baselines? This refers to the idea that what people expect as “normal” on our planet has changed over time, especially with regard to ecosystems. The abundance of fish and wildlife that once was normal is no longer carried in the memories of those living today. Similarly, many of us have no memory of cities unclogged by vehicles. Fewer people remember clear summer days when it really did seem that you could see forever. Clearly we humans are industrious creatures. We grow crops, dam rivers, burn fuels, build structures, and jet across time zones. When we carry out such activities by the million, we change the quality of the air we breathe, the water we drink, and the land on which we live. Over time, our actions have changed the face of our planet. What was it once like? See what you can find out by doing the next activity.
Consider This 0.1
Shifting Baselines
Seek out one of the elders in your community. This person may be a friend, relative, or possibly even a community historian. a. Think about the current price of a loaf of bread, a gallon of gas, or a candy bar. Then inquire about what things used to cost. How has what people expect as “normal” shifted? b. Now a more difficult task. Think about the local rivers, air quality, vegetation, or wildlife. In talking with an elder, see if you can identify at least one case in which the perception of what is “normal” has shifted. It may be that nobody remembers.
Chapter 6 describes the nitrogen cycle. Chapter 4 describes photosynthesis.
Consider This activities appear in all chapters. These activities give you a chance to use what you are learning to make informed decisions. For example, they may require you to consider opposing viewpoints or to make and defend a personal decision. They may require additional research.
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The bottom line? The things we perceive today as “normal” were not normal in the past. Although we cannot turn back the clock, we still can make choices that promote our health and the health of our planet today and in the future. A knowledge of chemistry can help. The global problems that we face—and their solutions—are intimately linked with chemical expertise and good old human ingenuity.
0.1
Look for more about crude oil and how it is refined in Chapter 4.
Look for more about what comes out a tailpipe in Chapter 1.
Chapters 1, 3, 4, and 6 all explore the connections between fuels and the waste products they release when burned.
Nonrenewable resources, such as crude oil and metal ores, are finite in supply. Once we use them up, they are gone.
| The Choices We Make Today
Individually, it may seem that our actions have little effect on a system as large as our planet. After all, compared with a hurricane, a drought, or an earthquake, what we do every day can seem pretty inconsequential. What difference could it possibly make if we biked to work instead of driving, used a reusable cloth bag instead of discarding a plastic one, or ate foods grown locally instead of consuming those shipped from hundreds or even thousands of miles away? Most human activities—including biking, driving, using bags of whatever sort, and eating—have two things in common: They require the consumption of natural resources, and they result in the creation of waste. Driving a car requires gasoline (refined from crude oil), and burning gasoline sends waste products out the exhaust pipe. Although riding a bike is a more ecological choice, all bicycles, just like automobiles, still require the manufacture and disposal of metals, plastics, synthetic rubbers, fabrics, and paints. Shopping bags, whether paper or plastic, require the materials to produce them. Later down the line, these bags become a waste product. And growing food requires water and energy to harvest and transport to market. In addition, food production may require fertilizer and involve the use of insecticides and herbicides. You can see where this is going. Any time we manufacture and transport things, we consume resources and produce waste. Clearly, though, some activities consume fewer resources and produce less waste than others. Biking produces less waste than driving; reusing cloth bags produces less waste than continually throwing plastic ones away. Although what you do may be negligible in the grand scheme of things, what 7 billion people do clearly is not. Our collective actions not only cause local changes to our air, water, and soil but also hurt regional and global ecosystems. We need to think by the billion. A single cooking fire? No problem; well, unless it accidentally burns down a dwelling. But imagine a few billion people across the planet each tending an individual cooking fire. Add in fires from those who cook using stoves, brick ovens, and outdoor grills. Now you have a lot of fuel being burned! Each fuel releases waste products into the atmosphere as it is burned. Some of these waste products— better known as air pollutants—are highly unfriendly to our lungs, our eyes, and of course to our ecosystems. Today, the waste products we release are unprecedented in their scale and in their potential to lower our quality of life and even shorten it. For example, in a large city such as New York, Atlanta, Mexico City, or Beijing, you will find that hospital admissions and death rates correlate with air pollution levels. Although the health risks are smaller than those caused by obesity or smoking, the issues of public health loom large because people are exposed to air pollutants, both indoors and out, over a lifetime. Also worrisome is that our actions (by the billion) release waste products that destroy the habitats of other species on the planet. Extinction, of course, is a natural phenomenon. But today the rate is many times faster than would be expected from natural causes. Our destruction of local specialized habitats, particularly those of plants, has led to these extinctions. Underpinning much of our waste production is energy. The need to find energy sources that are both clean and sustainable is arguably the major challenge of our century. Currently we are consuming nonrenewable and renewable resources and adding waste to our air, land, and water at a rate that cannot be sustained. This should
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Chemistry for a Sustainable Future
5
come as no surprise. From what you have learned from your other studies, most likely you are well aware of this. With any problem comes the opportunity to find creative solutions. We hope you are asking yourself “What can I do?” and “How can I make a difference in my community?” As you ask questions such as these, remember to include chemistry in your deliberations. Indeed, chemistry is well named the “central science.” Today, chemists are at the center of the action when it comes to the sustainable use of resources. Chemists are challenged to use what they know, to do so responsibly, and to proceed with reasonable haste. The same, of course, is true for you. In this book, we will support you as you learn and encourage you to use what you learn to act responsibly and with reasonable haste. “. . . No matter what, the road to better health, better materials, and better energy sources goes through chemistry.” Bill Carroll, 2006 American Chemical Society Past President
0.2
|
The Sustainable Practices We Need for Tomorrow
What does it mean to use the resources of our planet in a sustainable manner? We hope that you can answer this question—at least in part—from what you already have learned in other classes. People who study many other disciplines, including economics, political science, engineering, history, nursing, and agriculture, have a stake in developing sustainable practices. And as you will learn in this text, those of us in chemistry also have a stake in developing sustainable practices. Because the term sustainability is used by so many groups of people, it has taken on different meanings. We have selected one that is frequently quoted: “Meeting the needs of the present without compromising the ability of future generations to meet their needs.” This definition is drawn from a statement written over two decades ago, a 1987 report, Our Common Future, written by the World Commission on Environment and Development of the United Nations. Even so, it has stood the test of time. In Table 0.1, we reprint excerpts from the foreword to Our Common Future so that you can read its challenging words in their original context. Brundtland’s words carry a message to those who teach and learn. She writes: “In particular, the Commission is addressing the young. The world’s teachers will have a crucial role to play in bringing this report to them.” We agree. To this end, we hope that your chemistry course will stimulate conversations both inside and outside of the classroom. One such conversation is about practices that are not sustainable. For example, you will study fossil fuels and learn why their use is not sustainable (Chapter 4). But don’t stop here. You also need to discuss what you can do to solve the problems we face today. Use what you learn about air quality to act to improve local air quality and to make informed decisions as a citizen to improve it more widely (Chapter 1). Similarly, use what you learn about aqueous solubility and waste water to evaluate public policies that relate to water quality (Chapter 5). Any discussion of sustainability needs to include conservation. Although people may equate conservation with undue self-sacrifice, the equation is far more complex than this. To conserve, we need the vigorous development of community-based technologies that improve
Our Common Future is also called the Brundtland report. It was named after Gro Harlem Brundtland, the woman who chaired the commission.
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Chapter 0
Table 0.1
Our Common Future (excerpts from the Foreword)
“A global agenda for change”—this was what the World Commission on Environment and Development was asked to formulate. It was an urgent call by the General Assembly of the United Nations. In the final analysis, I decided to accept the challenge. The challenge of facing the future, and of safeguarding the interests of coming generations. After a decade and a half of a standstill or even deterioration in global co-operation, I believe the time has come for higher expectations, for common goals pursued together, for an increased political will to address our common future. The present decade has been marked by a retreat from social concerns. Scientists bring to our attention urgent but complex problems bearing on our very survival: a warming globe, threats to the Earth’s ozone layer, deserts consuming agricultural land. The question of population—of population pressure, of population and human rights—and the links between these related issues and poverty, environment, and development proved to be one of the more difficult concerns with which we had to struggle. But first and foremost our message is directed towards people, whose well being is the ultimate goal of all environment and development policies. In particular, the Commission is addressing the young. The world’s teachers will have a crucial role to play in bringing this report to them. If we do not succeed in putting our message of urgency through to today’s parents and decision makers, we risk undermining our children’s fundamental right to a healthy, lifeenhancing environment. In the final analysis, this is what it amounts to: furthering the common understanding and common spirit of responsibility so clearly needed in a divided world. Gro Harlem Brundtland, Oslo, 1987.
Renewable resources are expected always to be available. They include sunlight, wind and wave energy, and timber.
efficiency, promote the use of renewable resources, and minimize or prevent waste. All of these require ingenuity and chemical know-how. Any discussion of sustainability usually is accompanied by a sense of urgency as well. We currently are not using the resources of our Earth in a sustainable manner. We need to change our practices and we cannot delay. This sense of urgency has put sustainability squarely on the radar screens of chemists, other scientists, and the professional societies of which they are members. For example, botanist Peter H. Raven, former president of the American Association for the Advancement of Science, spoke about sustainability in his 2002 address as president titled Science, Sustainability and the Human Prospect. He called for nothing short of new ways of thinking. “New ways of thinking—an integrated multidimensional approach to the problems of global sustainability—have long been needed, and it is up to us to decide whether the especially difficult challenges that we are facing today will jolt us into finding and accepting them.” Peter H. Raven, 2002
This report, The Chemistry Enterprise in 2015, also points out that most issues of sustainability ultimately come down to questions of energy, a topic of several chapters in this book.
Not only must we find these ways, but we will also be judged by our success— or lack thereof—in employing them. Bill Carroll, a recent president of the American Chemical Society, clearly points this out: “By 2015, the chemistry enterprise will be judged under a new paradigm of sustainability. Sustainable operations will become both economically and ethically essential.” These words came from a report he coauthored in 2005, The Chemistry Enterprise in 2015. Carroll also pointed out that “No matter what, the road to better health, better materials, and better energy sources goes through chemistry.” In this text, we will be following this road. The next section describes the Triple Bottom Line, an important way to set our bearings as we travel.
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Chemistry for a Sustainable Future
The Triple Bottom Line
Scientists aren’t the only ones. If you are a business or economics major, you may be well aware that people in the business sector have put sustainability on the corporate agenda. In fact, sustainable practices now offer a competitive advantage in the marketplace. In the world of business, the bottom line always has included turning a profit, preferably a large one. Today, however, the bottom line includes more than this. For example, corporations are judged to be successful when they are fair and beneficial to workers and to the larger society. Another measure of their success is how well they protect the health of the environment, including the quality of the air, water, and land. Taken together, this three-way measure of the success of a business based on its benefits to the economy, to society, and to the environment has become known as the Triple Bottom Line. One way to represent the Triple Bottom Line is with the overlapping circles shown in Figure 0.1. The economy must be healthy, that is, the annual reports need to show a profit. But no economy exists in isolation; rather, it connects to a community whose members also need to be healthy. In turn, communities connect to ecosystems that need to be healthy. Hence the figure includes not one, but three connecting circles. At the intersection of these circles lies the “Green Zone.” This represents the conditions under which the Triple Bottom Line is met. Harm that occurs in any of the circles of Figure 0.1 ultimately translates into harm for the business. Conversely, achieving success can provide a competitive advantage, both immediately and in the years to come. Businesses can turn a profit; at the same time, they can get good publicity (and minimize any harm) by using less energy, consuming fewer resources, and creating less waste. A triple win! Recent news articles document the changes that are occurring. For example, read this excerpt from a news article about Clorox, a company that produces the bleach that you are likely to find with the cleaning supplies in your local supermarket. The source is Chemical & Engineering News (C&EN), the weekly publication of the American Chemical Society. One sentence was underlined for emphasis. Greener Cleaners: Consumer demand for environmentally friendly cleaning products has changed the game for chemical suppliers “This month, Clorox, a company almost synonymous with the environmental lightning rod chlorine, is going national with what might seem like an unlikely product line: a family of natural cleaners sold under the Earth-friendly name Green Works. That a consumer products giant like Clorox would venture into the market for so-called green cleaning products says a lot about how much the home care industry has changed in the past two years. Once solely the province of fringe players, green or sustainable cleaners are attracting the interest of big corporations in America and elsewhere. In such products, companies see both a growing market and a way to burnish their environmental credentials.” (C&EN, Jan. 21, 2008)
Healthy ecosystems
Healthy economies
Green Zone
Healthy communities
Figure 0.1 A representation of the Triple Bottom Line. The “Green Zone” (where the Triple Bottom Line is achieved) lies at the intersection of these three circles.
The Triple Bottom Line sometimes is shortened to the 3Ps: Profits, People, and the Planet.
VIEW FROM THE HILL The pushes and pulls on legislators in 2008 P.35
JA N UA RY 2 1 , 20 0 8
0.3 |
7
PITTCON 2008 Program summary for New Orleans meeting P.50
SOAPS AND DETERGENTS The greening of home cleaning P.15
PUBLISHED BY THE AMERICAN CHEMICAL SOCIETY
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FE B RUA RY 18, 20 0 8
Look for more about sodium hypochlorite in Chapter 5 in connection with water purification.
FEDERAL R&D SUPPORT How 2009’s proposed $147 billion pie cuts up P.27 PROJECT SEED An ACS social action initiative turns 40 P.42
FUTURE SPENDING
Firms suit up for ’08 capital, R&D projects P.11 PUBLISHED BY THE AMERICAN CHEMICAL SOCIETY
At the time the article was published, Clorox posted a video on its corporate website about new green cleaning products. The sound track described the “natural plant-based cleaners . . . without harsh chemical fumes or residues.” The website also included strongly worded environmental statements indicating that the corporation is working toward “using as little packaging material as needed to do the job” and “using recycled material wherever it is environmentally and economically sound to do so.” Clearly Clorox is responding to the negative image alluded to in the news article. At issue here is the chlorine-containing bleach, sodium hypochlorite (NaClO). We will revisit this compound in the next section. Before we leave the green cleaning scene, consider one more news article, this time on laundry detergents. The article describes the placards in hotels that give you the option of not having the sheets and towels changed daily. What happens, though, when people who market soap products stay at these hotels? Do they choose the option that means selling less detergent? Again, one sentence was underlined for emphasis. Having It All: Chemical makers supplying the detergents industry seek both sales and sustainability “Although it’s lush to the point of excess, the Boca Raton Resort & Club in Florida does demonstrate a modicum of environmental responsibility with a card placed on nightstands informing guests that bedsheets won’t be removed and washed unless requested. This card no doubt presented a conundrum to chemical industry executives who were at the resort earlier this month for the Soap & Detergent Association’s annual conference: Should they or shouldn’t they participate in a program intended to cut consumption of the very products they are there to sell? But then, the challenges of good environmental behavior have been on the minds of all participants in the cleaning products industry lately. Everyone from the government to retailers to consumers seems to be demanding environmentally sustainable products. . . . Still, the chemical companies that supply ingredients to the cleaning products industry see robust sales and environmental stewardship as mutually obtainable. Rather than cut back on surfactants or other cleaning chemicals, they are advising their customers to formulate products with ingredients that have less of an environmental impact.” (C&EN, February 18, 2008)
Consider This 0.2
Green Conundrums
Put yourself in two different roles. First, you are a manufacturer attending the annual Soap & Detergent Association conference. You find a card on your pillow saying that your bedsheets will not be washed unless you request it. Argue for a course of action, either way. Second, you are the hotel manager. Do you remove the card from the pillows, knowing that Soap & Detergent Association people are coming to town? Again make your case.
Donella Meadows (1941–2001) was a scientist and writer. Her books include The Limits to Growth and The Global Citizen.
The point? The debates over how to be “green” are likely to continue over your lifetime. The issues are not new; they are not likely to be resolved anytime soon. Chemistry in Context will encourage you to explore these issues, arming you with the knowledge you need and enabling you to more creatively formulate your own response. “A sustainable society is one that is far-seeing enough, flexible enough, and wise enough not to undermine either its physical or its social systems.” Donella Meadows, 1992
0.4 |
Cradle-to-Where?
You may have heard the expression cradle-to-grave; that is, an approach to analyzing the life cycle of an item, starting with the raw materials from which it came and ending with its ultimate disposal someplace, presumably on Earth. This catchy phrase offers a frame of reference from which to ask questions about consumer items. Where did the item come from? And what will happen to it when you are finished with it? More than ever, individuals, communities, and corporations are recognizing the importance of asking these types of questions.
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Chemistry for a Sustainable Future
For example, the website for Clorox asserts that the bleach you purchase, sodium hypochlorite, “starts as salt and water and ends as salt and water.” We agree. However, we urge you to consider the entire life cycle of a product like Clorox. One chemical used in the manufacture of Clorox is an element named chlorine (Cl). Chlorine will receive mention in almost every chapter of Chemistry in Context! How was the chlorine produced and transported? What waste products were created in its manufacture? Cradleto-grave means thinking about every step in the process. Companies should take responsibility—as should you—for items from the moment the natural resources used to make them were taken out of the ground, air, or water to the point at which they were ultimately “disposed of.” Think of items such as batteries, plastic water bottles, T-shirts, cleaning supplies, running shoes, cell phones— anything that you buy and eventually discard. Cradle-to-grave thinking clearly has its limitations. As an illustration, let us follow one of the plastic bags that supermarkets provide for your groceries. The raw material for these bags is petroleum. Accordingly, the “cradle” of this plastic bag most likely was crude oil somewhere on our planet, for example, the oil fields of Canada. Let’s assume that the oil was pumped from a well in Alberta and then transported to a refinery in the United States. At the refinery, the crude oil was separated into fractions. One of the fractions was then cracked into ethylene, the starting material for a polymer. Ethylene next was polymerized and formed into polyethylene bags. These bags were packaged and then trucked (burning diesel fuel, another refinery product) to your grocery store. Ultimately, you purchased groceries and used one of the bags to carry them home. As stated, this is not a cradle-to-grave scenario. Rather, it was cradle-to-yourkitchen, definitely several steps short of any graveyard. So what happened to this plastic bag after you used it? Did it go into the trash? The term grave describes wherever an item eventually ends up. One trillion plastic bags, give or take, are used each year in supermarkets. Only about 5% are recycled. The rest end up in our cupboards, our landfills, or littered across the planet. As litter, these bags begin a 1000-year cycle, again give or take, of slow decomposition into carbon dioxide and water. Cradle-to-a-grave-somewhere-on-the-planet is a poorly planned scenario for a supermarket bag. If each of the trillion plastic bags instead were to serve as the starting material for a new product, we then would have a more sustainable situation. Cradleto-cradle, a term that emerged in the 1980s, refers to a regenerative approach to the use of things in which the end of the life cycle of one item dovetails with the beginning of the life cycle of another, so that everything is reused rather than disposed of as waste. In Chapter 9, we will examine different recycle-and-reuse scenarios for plastic bottles. But right now, you can do your own cradle-to-cradle thinking in the next activity.
Your Turn 0.3
The Can That Holds Your Beverage
People tend to think of an aluminum can as starting on a supermarket shelf and ending in a recycling bin. There is more to the story! a. Where on the planet is aluminum ore (bauxite) found? b. Once removed from the ground, the ore usually is refined to alumina (aluminum oxide) near the mining site. The alumina is then transported to a production facility. What happens next to produce aluminum metal? c. The metal is then shaped into a beverage can. See if you can find where the can was filled and how far it was transported to land on the shelves of your neighborhood store. d. What happens to the can after you recycle it? Answers a. Bauxite is mined in several places, including Australia, China, Brazil, and India. b. The ore must be refined electrolytically to produce aluminum metal. This process is energy-intensive and carried out in many locations worldwide. Note: the textbook’s website provides some helpful links.
9
Look for more about batteries in Chapter 8 and more about plastics in Chapter 9.
Chapter 4 explains how and why crude oil is separated into fractions.
Chapter 9 explains how polyethylene is made from ethylene (and why).
The term cradle-to-cradle caught on with the publication of a book of this same title published in 2002.
Your Turn activities appear in all chapters. They provide an opportunity for you to practice a new skill or calculation that was just introduced in the text. Answers are given either following the Your Turn activity or in Appendix 4.
Some activities are marked with a Web icon. This cues you to use the resources of the Internet as you work.
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As you can tell from these examples, it is not just the decisions of manufacturers that matter. Your decisions do as well. What you buy, what you discard, and how you discard it all warrant attention. The choices that we make—individually and collectively— matter. At the risk of repeating ourselves, we remind you that the current state of affairs in which we consume the nonrenewable resources of our planet and add waste to our air, land, and water is not sustainable. With a sense of urgency, we recall the words of botanist Peter Raven quoted earlier: “We must find new ways to provide for a human society that presently has outstripped the limits of global sustainability.” In the next section, we explain why, setting the stage for the topics in chemistry that you will explore in this text.
0.5 | Although you may not know how much air you breathe in a day, Chapter 1 will help you to estimate this.
Carbon footprint is a subset of ecological footprint. Look for more about carbon footprints in Section 3.9.
A hectare is 10,000 square meters, or 2.471 acres.
Your Ecological Footprint
You already may know how to estimate the gas mileage for a vehicle. Likewise, you can estimate how many calories you consume. How might you estimate how much of the Earth’s natural capital it takes to support the way in which you live? Clearly, this is a far more difficult task. Fortunately, other scientists already have grappled with how to do the math. They base the calculations on the way in which a person lives coupled with the available renewable resources needed to sustain this lifestyle. Consider the metaphor of a footprint. You can see the footprints that you leave in sand or snow. You also can see the muddy tracks that your boots leave on the kitchen floor. Similarly, one might argue that your life leaves a footprint on planet Earth. To understand this footprint, you need to think in units of hectares or acres. A hectare is a bit more than twice the area of an acre. The ecological footprint is a means of estimating the amount of biologically productive space (land, water, and sea surface) necessary to support a particular standard of living or lifestyle. For the average U.S. citizen, the ecological footprint was estimated in 2005 to be about 9.7 hectares (24 acres). In other words, if you live in the United States, on average it requires 9.7 hectares to provide the resources to feed you, clothe you, transport you, and give you a dwelling with the creature comforts to which you are accustomed. The people of the United States have relatively big feet, as you can see in Figure 0.2. The world average in 2003 was estimated to be 2.2 hectares per person; today the value is believed to be closer to 2.7. How much biologically productive land and water is available on our planet? We can estimate this by including regions such as croplands and fishing zones, and omitting regions such as deserts and ice caps. Currently, the value is estimated at about 11 billion hectares (roughly 27 billion acres) of land, water, and sea surface. This turns out to be about a quarter of the Earth’s surface. Is this enough to sustain everybody on the planet with the lifestyle that people in the United States have? The next activity allows you to see for yourself.
Your Turn 0.4
Your Personal Share of the Planet
As stated earlier, an estimated 11 billion hectares (~27 billion acres) of biologically productive land, water, and sea is available on our planet. a. Find the current estimate for the world population. Cite your source. b. Use this estimate together with the one for biologically productive land to calculate the amount of land theoretically available for each person in the world. Answers a. In 2010, the population of the Earth was between 6.8 and 6.9 billion. b. About 1.6 hectares or ~4 acres per person.
Bottom line? A nation whose people have an average footprint greater than about 1.6 hectares is exceeding the “carrying capacity” of the Earth. Using the United States as an example, let’s do one more calculation to see by how much.
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USA
China
Mexico
France
Global average
9.7
2.1
3.4
4.9
2.7
Figure 0.2 A comparison of ecological footprints, in global hectares per person. Source: Global Footprint Network, 2008 data tables.
Your Turn 0.5
How Many Earths?
In 2009, the United States had an ecological footprint of about 9.7 hectares (~24 acres) per person. a. Find an estimate of the current population of the U.S. Cite your source. b. Calculate the amount of biologically productive space (land, water, and sea surface) that the U.S. currently requires for this population. c. What percent is this amount of the biologically productive space (about 11 billion hectares or ~27 billion acres) available on our planet? Answer b. Estimating the U.S. population at 310 million and using the estimate of 9.7 hectares per person, the United States required about 3 billion hectares (7.3 billion acres).
Let’s now say that everyone on the planet lived like the average citizen in the United States. Here is the calculation based on the world population in 2009. 6.8 billion people 3 9.7 hectares/person 3 1 planet 5 6.0 planets 11 billion hectares Thus, to sustain this same standard of living for everybody on the planet we would need 5 more Earths in addition to the one we currently have! “Only One Earth.” On this Earth, the number of people has risen dramatically in the last few hundred years. So has economic development. As a result, the estimated global ecological footprint is rising, as shown in Figure 0.3. In 2003, we estimate that humanity used the equivalent of 1.25 Earths. By 2040, the projection is that we will be using 2 Earths. Clearly this rate of consumption cannot be sustained. Through your study of chemistry, we hope you will learn ways either to reduce your ecological footprint or to keep it low, if it already is. The next section describes how chemists can help make this process work in some ways you might not expect.
0.6
|
Our Responsibilities as Citizens and Chemists
We humans have a special responsibility to take care of our planet. Living out this responsibility, however, has proven to be no easy task. Each chapter in Chemistry in Context highlights an issue of interest such as air quality, water quality, or nutrition.
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Number of planet Earths
2.5 2.0
1960–2005 Ecological footprint 2005–2050, Scenarios Moderate business as usual Rapid reduction
1.5 1.0 0.5 0 1970
1980
1990
2000
2010 Years
2020
2030
2040
2050
2060
Figure 0.3 The ecological footprint of humans, current (solid line) and projected (dashed lines). Current estimates place the footprint above 1.0 Earth, that is, above what the Earth can sustain. Note: The upper projection assumes “moderate business as usual.” The lower projection assumes an increase in sustainable practices. Source: Adapted from data provided by the Global Footprint Network
Look for green chemistry examples throughout this book, designated with this icon.
These issues not only affect you personally, they also affect the health and well-being of the wider communities of which you are a part. For each issue, you will work on two related tasks: (1) learning about the issue and (2) finding ways to act constructively. Chemists are learning and working right along with you. Recall the words of Bill Carroll that we quoted earlier in the chapter: “By 2015, the chemistry enterprise will be judged under a new paradigm of sustainability. Sustainable operations will become both economically and ethically essential.” How do do cch hemi he mist istss me meet et tth he cch he hall hall llen enge g s off ssus ge usta taiin inab abil bil iliit ity? y? T Th he aans he nswe werr lies lies in in part p rt pa How chemists the challenges sustainability? The answer wiith “green “gr gree eenn chemistry,” ee chem ch hem emiis istr ist try,” try y,” a se sett off pprinciples rinc ri inc nciip iple ipl les originally les oriig or igin igi inal all lly ly articulated art rti tic icullatted d by by people peop pe oplle op le aatt th thee U.S U .S. S. with U.S. Env E nvir nv iron ir onme on ment me ntal nt al Protection Pro Pro rote tect te ctio ct ionn Agency io Agen Ag ency en cy (EP ((EPA) EPA) EP A) and and now now act aactively ctiv ct ivel iv elyy pu el purs rsue rs uedd by the ue the Ame A meri me rica ri cann ca Environmental pursued American Chemical Society (ACS). Green chemistry is the design of chemical products and processes to use less energy, produce fewer hazardous materials, and use renewable resources whenever possible. The desired outcome is to produce less waste, especially toxic waste, and to use fewer resources. Recognize that green chemistry is a tool in achieving sustainability, not an end in itself. As stated in the article “Color Me Green” from an issue of Chemical & Engineering News published in 2000: “Green chemistry represents the pillars that hold up our sustainable future. It is imperative to teach the value of green chemistry to tomorrow’s chemists.” Actually, we believe that it is imperative to teach the value of green chemistry to citizens as well. This is why so many applications of green chemistry are woven throughout Chemistry in Context. To get you started, we list six big ideas about green chemistry (Table 0.2). They also are printed on the inside cover of this book. Initiated under the EPA’s Design for the Environment Program, green chemistry leads to cleaner air, water, and land, and the consumption of fewer resources. Chemists are now designing new processes (or retooling older ones) to make them more environmentally friendly. We call this “benign by design.” Every green innovation does not necessarily have to be successful in achieving all six of these principles. But achieving several of them is an excellent step on the road to sustainability. For example, an obvious way to reduce waste is to design chemical processes that don’t produce it in the first place. One way is to have most or all of the atoms in the reactants end up as part of the desired product molecules. This “atom economy” approach, although not applicable to all reactions, has been used for the synthesis of many products, including pharmaceuticals, plastics, and pesticides. The approach saves
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Table 0.2
13
Principles of Green Chemistry
1. It is better to prevent waste than to treat or clean up waste after it is formed.
2. It is better to minimize the amount of materials used in the production of a product. 3. It is better to use and generate substances that are not toxic.
4. It is better to use less energy.
5. It is better to use renewable materials.
6. It is better to use materials that degrade into innocuous products at the end of their useful life. Source: Adapted from the Twelve Principles of Green Chemistry, by Paul Anastas and John Warner.
money, uses fewer starting materials, and minimizes waste. The connection between green chemistry and the Triple Bottom Line should be apparent! Many Ma nyy chemical che che hemi mica mi call manufacturing ca manu ma nufa nu fact fa ctur ct urin ur ing in g processes p oc pr oces esse es sess no se now w us usee in inno innovative nova no vati va tive ti ve ggreen reen re en che chemical che hemi mica mi call me ca meth meththth ods. ods od ds. For For example, exa exa xamp mpple le you le, y u will yo wiill see see app aapplications ppli pp lica li cati tion ion onss off ggreen reen re en ch cchemistry hemi he mist ist stry ry y th tthat hat ha hat ha have lled have ed d to to ch chea cheaper, hea eape p r, pe less le ss w was wasteful, aste as tefu te full, aand fu nd lles less esss to es toxi toxic xicc pr xi prod production oduc od ucti uc tion ti on ooff lo low w VO VOC C pa pain paint intt (C in (Cha (Chapter hapt ha pter pt er 11) 1). ). Y You ou aals also lsoo wi ls will ll ssee ee the principles of green chemistry applied to processing raw cotton and to dry-cleaning methods (Chapter 5). There are more economical and healthier ways to process vegetable oils (Chapter 11). “Green chemistry and engineering hold the key for our sustainable future.” Robert Peoples, 2009 Director Green Chemistry Institute
G Gree Green reenn chemistry chem ch hemiist istry try ef eff efforts ffo fort rts ts have have bbee been eenn re rewa rewarded! ward ded ed! d! A select selle se lectt group gro roup upp ooff re rese research sear arch h chemists cchhe hemiist sts ts and chemical and chem ch hemiic icall eng eengineers ng ginee ineers rs has has rec rreceived ecei eive ivedd the th he Presidential Presid Pres iden id entiiall Green Gre Green en Chemistry Ch Che hemi mist istry ry y Ch C Challenge hall hall llen enge g ge Award. Awar Aw ardd. Initiated ar Ini Ini niti tiat ti ated at ed in in 1996, 1996 19 96, this 96 this presidential-level pre pre resi side si dent de ntia nt iall-le ia leve le vell award ve awar aw ardd recognizes ar reco re cogn co gniz gn izes iz es chemists che che hemi mist mi stss and st and the the chemical industry for their innovations aimed at reducing pollution pollution. These awards recognize innovations in “cleaner, cheaper, and smarter chemistry.”
0.7
| Back to the Blue Marble
Before we send you off to Chapter 1, we revisit the 1987 United Nations document Our Common Future. Earlier, we drew from this document our definition of sustainability: “Meeting the needs of the present without compromising the ability of future generations to meet their own needs.” The foreword to this report was written by the chair, Gro Harlem Brundtland. She also wrote these words, ones that call us back to the image of Earth that opened this chapter: “In the middle of the 20th century, we saw our planet from space for the first time. Historians may eventually find that this vision had a greater impact on thought than did the Copernican revolution of the 16th century, which upset the human self-image by revealing that the Earth is not the centre of the universe. From space, we see a small and fragile ball dominated not by human activity and edifice but by a pattern of clouds, oceans, greenery, and soils. Humanity’s inability to fit its activities into that pattern is changing planetary systems, fundamentally. Many such changes are accompanied by life-threatening hazards. This new reality, from which there is no escape, must be recognized—and managed.”
VOC stands for volatile organic compounds. Look for more about VOCs in connection with air pollution in Chapter 1, “The Air You Breathe.”
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We agree. The new reality must be recognized. There is no escape. And all of us—students and teachers alike—have important roles to play. With Chemistry in Context, we will strive to provide you with the chemical information that can make a difference in your life and in the lives of others. We hope that you will use it to meet the challenges of today and tomorrow with understanding and good spirit.
Questions 1. This chapter opened with a famous quote from a Saudi astronaut, Prince Sultan. Here is what he said in a 2005 interview: “Being an astronaut has had an enormous impact on me. Looking at the planet from the perspective of the blackness of space, it makes you wonder, . . .” Prince Sultan went on to describe what he wondered about. We did not reprint his words. Rather, we hoped that you would write your own. a. Write a three-paragraph essay. In the first, introduce yourself briefly. In the second, describe what is important to you as you begin your study of chemistry. And in the third, describe what you most wonder about that relates to planet Earth. b. Share your self-introduction with others in your class, as indicated by your instructor. 2. Read the full text of the Foreword to the Brundtland report. It is only a few pages and contains some of the most compelling language ever written. Pick a small section and write a short piece that connects it to something you care about or is of concern to you. You may choose a stance of agreement or disagreement. Your textbook’s website contains a link to the document.
3. This chapter introduces the idea that the species on our planet are all connected, sometimes in ways that are not obvious. From your studies in other fields, give three examples of how organisms are linked or in some way depend on one another. 4. Figure 0.1 shows one possible representation of theTriple Bottom Line. Here is another. Comment onthe similarities and differences between these twofigures.
Economy Society Environment
5. Calculate your ecological footprint. Feel free to use one of the websites suggested at your textbook’s website. 6. A bus serving a university campus had this sign painted on its side: “Reduce your footprints. Take the bus.” Explain the double meaning.
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7. Select a corporation, visit its website, and look for the pages that describe its efforts at sustainability or corporate responsibility. Possibilities include CocaCola, WalMart, Exxon Mobil, and Burger King. Report on what you find. 8. Mel George, a retired mathematics professor and one of the architects of Shaping the Future, New Expectations for Education in Undergraduate Science, Technology, Engineering, and Mathematics (National Science Foundation), remarked: “Putting a man on the moon did not require me to do anything. In contrast, we all must do something to save the planet.” Describe five things that you could do, given who you are and your field of study. 9. The principles of green chemistry are not just for chemists. Perhaps you are an economics major. Or you are planning to become a nurse or a teacher. Perhaps you spend time gardening or you like to commute by bike. Pick two of the principles of green chemistry and show how they connect to your life or intended profession. 10. Newspaper and magazine advertisements often proclaim how “green” a business or corporation is. Find one and read it closely. What do you think—is it a case of “greenwashing,” in which a corporation is trying to use as a selling point one tiny green drop in an otherwise wasteful bucket? Or is it a case of a real improvement that significantly reduces the waste stream? Note: It may be difficult to tell. For example, removing 7 tons of waste may sound large unless you know that the actual waste stream is in the billions of tons. 11. The mission statement of the American Chemical Society, the world’s largest scientific society, is “To advance the broader chemistry enterprise and its practitioners for the benefit of Earth and its people.” Explain how the final six words of this statement connect to the definition of sustainability used in this chapter.
ACS Chemistry for Life
TM
15
12. Thomas Berry, theologian and cultural historian (1914– 2009), described in his book The Great Works (1999) how humanity is faced with the job of moving from the present geological age, the Cenozoic, to the coming age. He names this age as “Ecozoic,” reflecting the tremendous power of humans to shape the face of the world. “The universities must decide whether they will continue training persons for temporary survival in the declining Cenozoic Era or whether they will begin educating students for the emerging Ecozoic. . . . While this is not the time for continued denial by the universities or for attributing blame to the universities, it is the time for universities to rethink themselves and what they are doing.” a. What do you see as the coming age? Describe it. If you don’t like the term Ecozoic Age, suggest one of your own. b. How have your studies to date been preparing you for the future world in which you will live? c. In what ways do you think your study of chemistry might prepare you?
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The Air We Breathe
California blue skies, Lake Tahoe region. “Ancient Greeks saw air as one of nature’s basic elements, along with earth, fire, and water. Californians see it . . . Ah, perhaps those words offer clarification: Californians see too much of something that ought to be less visible. They also feel effects from breathing that air, which too often brings the routine act of respiration to their attention.” David Carle, Introduction to Air in California, 2006. page xiii.
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People have always noticed—and been curious about—the air they breathe. Together with earth, fire, and water, the ancient Greeks named air as a basic element of nature. Hundreds of years later, chemists experimented to learn more about the composition of air. Today, we can view the Earth’s atmosphere from outer space. And daily, just like the ancients, we can peer up through the night air to catch a glimpse of the twinkling stars. Our atmosphere is the thin veil between us and outer space. This chapter describes the atmospheric gases that support the life on Earth. The next chapter describes the ozone in the stratosphere that protects us from harmful ultraviolet radiation emitted by the Sun. The third chapter describes the greenhouse gases in our atmosphere that protect us from the bitter cold of outer space. Truly, our atmosphere is a resource beyond price. Our goal is to convey its beauty, its magnitude, and its frailty. As you will learn in this chapter, we humans have altered the composition of the atmosphere. This is not surprising, because about 6.9 billion humans currently live on the planet. In a few decades, the population may reach 9 or 10 billion. The next activity invites you to think about how our actions, both individually and collectively, can change the air we breathe.
Consider This 1.1
Footprints in the Air
Hiking boot treads, asphalt pavement, corn fields. Each of these is an example of a “ground print” left by humans because each one alters the lay of the land. Similarly, our activities leave “air prints” that alter the composition of our atmosphere. a. Name three things that leave an indoor air print. b. Does each (1) hurt the air quality, (2) improve the air quality, or (3) have some effect, but you don’t know what it is? Explain. c. Repeat parts a and b for an outdoor air print. Answers a. Indoor air prints are left by growing house plants, burning candles, and applying paint. b. Green plants remove carbon dioxide from the air and add oxygen, both a plus. Some plants emit pollen. For some people, this reduces the air quality.
As Homo sapiens, we have a special responsibility to guard the quality of the air on our planet. Living out this responsibility has proven to be no easy task. In fact, we have made some tragic errors that have killed people, animals, and vegetation. Ultimately, as we pointed out in Chapter 0, our responsibility is to live in ways today that will not compromise either our own health or the health of future generations. Keeping the air clean is part of this responsibility. In this chapter, you will learn more about the air you breathe and its importance to your well-being. We hope you will also learn how important it is to make choices about air quality—both as an individual and as a member of a larger society—that demonstrate wisdom now and in the years to come.
1.1
In 1948, smog killed 20 and sickened between 5000 and 7000 residents of Donora, Pennsylvania. In 1952, the Great London Smog killed thousands.
| What’s in a Breath?
Take a breath! Automatically and unconsciously, you do this thousands of times each day. Certainly you do not need us to tell you to breathe! Although a doctor or nurse may have encouraged your first breath, from then on nature took over. Even if you were to hold your breath in a moment of fear or suspense, you soon would involuntarily gasp a lungful of that invisible stuff we call air. Indeed, you could survive only minutes without a fresh supply. 17
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Consider This 1.2
Take a Breath
What total volume of air do you inhale (and exhale) in a typical day? Figure this out. First determine how much air you exhale in a single “normal” breath. Then determine how many breaths you take per minute. Finally, calculate how much air you exhale per day. Describe how you made your estimate, provide your data, and list any factors you believe may have affected the accuracy of your answer.
Some mixtures are composed of gases. But gasoline is a mixture of liquids (Section 1.10), and soil is a mixture of solids and liquids.
Were you surprised at how much you breathe? For an adult, the value typically is more than 11,000 liters (about 3000 gallons) of air per day. The value would be even higher had you spent the day on a bike trail or paddling a kayak. Although you cannot tell by looking, the air you are breathing is not a single pure substance. Rather, it is a mixture; that is, a physical combination of two or more pure substances present in variable amounts. Mixtures are one of the two forms of matter that we encounter on our planet (Figure 1.1). The other form is pure substances. In this section, we focus on the pure substances that are the major components of air: nitrogen, oxygen, argon, carbon dioxide, and water vapor. All are colorless gases, invisible to the eye. Matter
Pure substances
Mixtures
Figure 1.1 Matter can be classified either as a single pure substance or as a mixture.
The composition of the mixture that we call “air” depends on where you are. There’s less oxygen in a stuffy room and more pollutants in an urban area. Exhaled air is a slightly different mixture than inhaled air. Trace amounts of substances may vary in the air as well. For example, the scent of lilacs may permeate the air outside. Indoors, the aroma of freshly brewed coffee may beckon you to the kitchen. In fact, the human nose is an extremely sensitive odor detector. In some cases, only a minute trace of a substance is needed to trigger the olfactory receptors responsible for detecting odors. Thus, tiny amounts of substances can have a powerful effect on our noses, as well as on our emotions.
Consider This 1.3
Your Nose Knows
The air is different in a pine forest, a bakery, an Italian restaurant, and a dairy barn. Blindfolded, you could smell the difference. Our noses alert us to the fact that air contains trace quantities of many substances. a. Name three indoor and three outdoor smells that indicate small quantities of chemicals in the air. b. Our noses warn us to avoid certain things. Give three examples of when a smell indicates a hazard.
The composition of our atmosphere has not been constant over the millennia. For example, the concentration of oxygen has varied.
Look for more about atoms and molecules in Section 1.7.
Using a pie chart and a bar graph, Figure 1.2 represents the composition of air. The pie chart emphasizes the fractions of the whole, whereas the bar graph emphasizes the relative amounts of each substance. Regardless of how we present the data, the air you breathe is primarily nitrogen and oxygen. More specifically, the composition of air by volume is about 78% nitrogen, 21% oxygen, and 1% other gases. Percent (%) means “parts per hundred.” In this case, the parts are either molecules or atoms. The percents shown in Figure 1.2 are for dry air. Water vapor is not included, as its concentration varies by location. In dry desert air, the concentration of water vapor can be close to 0%. In contrast, it can reach 5% by volume in a warm tropical rain forest.
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80 70 60
Other gases (1%)
Percent
50 40 30
Oxygen (21%) 20
Nitrogen (78%)
10 0
Nitrogen
Oxygen
Other gases
Figure 1.2 The composition of dry air, by volume.
Figures Alive! Visit the textbook’s website to learn more about the molecules and atoms that make up air. Watch for the Figures Alive! icon throughout this text. Whether at high or low concentration, water vapor is a colorless gas that is invisible to the eye. Since you can see fog banks and clouds, they are not composed of water vapor. Rather, these consist of tiny droplets of liquid water or crystals of ice (Figure 1.3). Nitrogen is the most abundant substance in the air and constitutes about 78% of what we breathe. This gas is colorless, odorless, and relatively unreactive, passing in and out of our lungs unchanged (Table 1.1). Although nitrogen is essential for life and part of all living things, most plants and animals obtain it from sources other than the nitrogen in our atmosphere. Even though oxygen is less abundant than nitrogen in our atmosphere, it still plays a key role on our planet. Oxygen is absorbed into our blood via the lungs and reacts with the foods we eat to release the energy to power the chemical processes within our bodies. It is necessary for many other chemical reactions as well, including burning and
Figure 1.3 Clouds consist of minuscule droplets of water that remain suspended because of upward air currents. Clouds can weigh millions of pounds.
Water vapor is the gas that we think of as “humidity.”
Section 6.9 describes the cycle by which atmospheric nitrogen becomes part of living plants and animals.
Chapter 11 provides more information about the energy content of foods.
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Table 1.1 Substance
Typical Composition of Inhaled and Exhaled Air Inhaled Air (%)*
Exhaled Air (%)*
nitrogen
78.0
78.0
oxygen
21.0
16.0
0.9
0.9
argon carbon dioxide water vapor
0.04 variable
4.0 variable
*percents are by volume
We explain the term element in Section 1.6.
rusting. As the “O” in H2O (water), oxygen is the most abundant element in the human body. Present in many rocks and minerals, it is also the most abundant element in the Earth’s crust. Given this broad distribution, it is somewhat surprising that oxygen was not isolated as a pure substance until 1774. But once isolated, oxygen proved to be of great significance in establishing the young science of chemistry.
Consider This 1.4
More Oxygen . . . ?
We live in an atmosphere of 21% oxygen. A match burns in less than a minute, a fireplace consumes a small pine log in about 20 minutes, and we exhale about 15 times a minute. Life on Earth would be very different if the oxygen concentration were twice as high. List at least four differences.
Respiration also provides the energy to power chemical reactions in our bodies. Look for more about breathing and respiration in Chapter 4.
1 liter 5 1.06 quart. Appendix 1 contains this and many other conversion factors.
Look for more about atoms and molecules in Section 1.7.
Every time we exhale, we add carbon dioxide to the atmosphere. Table 1.1 indicates the difference between inhaled dry air and exhaled air. Clearly some changes have taken place that use up some of the oxygen and give off carbon dioxide and water. In the process of respiration, the foods we eat are metabolized to produce carbon dioxide and water. With each breath, some of the water in our bodies evaporates from the moist tissue in our lungs. Other gases are found in our atmosphere as well (see Figure 1.2). Argon, for example, is about 0.9% of the air. The name argon, meaning “lazy” in Greek, reflects the fact that argon is chemically inert. As you can see from Table 1.1, any argon that you inhale you simply exhale. The percentages we have been using to describe the composition of the atmosphere are based on volume; that is, the amount of space that each gas occupies. If we wanted to, we could closely approximate 100 liters (L) of dry air by mixing 78 L of nitrogen, 21 L oxygen, and 1 L argon (78% nitrogen, 21% oxygen, and 1% argon). The composition of air also can be represented in terms of the numbers of molecules and atoms present. This works because equal volumes of gases contain equal numbers of molecules, providing the gases are at the same temperature and pressure. Thus, if you took a sample of 100 of the molecules and atoms in air (an unrealistically small amount), 78 would be nitrogen molecules, 21 would be oxygen molecules, and 1 would be an argon atom. In other words, when we say that air is 21% oxygen, we mean that there are 21 molecules of oxygen per 100 of the molecules and atoms in the air. A bit later we’ll explain why nitrogen and oxygen are found as molecules. In contrast, argon is found as an atom. You now know that air contains nitrogen, oxygen, argon, carbon dioxide, and usually some water vapor. As you might imagine, there is more to the story.
1.2
| What Else Is in a Breath?
No matter where you live, each lungful of air you inhale contains tiny amounts of some substances. Many are present at concentrations of less than 1%, or one part per hundred. Such is the case with carbon dioxide, a gas you both inhale and exhale. In our atmosphere,
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the concentration of carbon dioxide currently is approximately 0.0390%. This value is slowly but steadily rising as we humans burn fossil fuels. Although we could express 0.0390% as 0.0390 molecules of carbon dioxide per 100 molecules and atoms in the air, it sounds strange to refer to 0.0390 of a molecule. For low concentrations, it is more convenient to use parts per million (ppm). One ppm is a unit of concentration 10,000 times smaller than 1% (one part per hundred). Here are the relationships. 0.0390% means means means means means
0.0390 parts per hundred 0.390 parts per thousand 3.90 parts per ten thousand 39.0 parts per hundred thousand 390 parts per million
21
Chapter 3 tells more about the rising CO2 concentrations in the atmosphere.
Changing between % and ppm involves moving the decimal point four places to the right. You can practice using the activities in Figures Alive!
Out of a sample of air containing 1,000,000 molecules (and some atoms), we now can say that 390 of them will be carbon dioxide molecules. The carbon dioxide concentration is 390 ppm, or 0.0390%.
Skeptical Chemist 1.5
Really One Part per Million?
It has been said that a part per million is the same as one second in nearly 12 days. Is this a correct analogy? How about one step in a 568-mile journey? Check the validity of these analogies, explaining your reasoning. Then come up with an analogy of your own.
Your Turn 1.6
Practice with Parts per Million
a. In some countries, the limit for the average concentration of carbon monoxide in an 8-hour period is set at 9 ppm. Express this as a percentage. b. Exhaled air typically contains about 78% nitrogen. Express this concentration in parts per million. Answers a. 0.0009%
b. 780,000 ppm
Even at tiny concentrations, some air pollutants are harmful and even lifethreatening. Although officials enact air pollution laws, 100% clean air is not attainable. Our human activities always add some waste to the air. Furthermore, our air gets contaminated by natural events such as wild fires and volcanic eruptions. There is no such thing as pure air! Nonetheless, dirty air definitely exists. Large metropolitan areas are one place you are likely to encounter it. For example, Figure 1.4 shows smog against the mountains near Santiago, Chile, a city of over 6 million inhabitants. Other large cities such as Los Angeles, Mexico City, Mumbai, and Beijing often have dirty air. Earlier, we noted how human activities leave “air prints,” both indoors and out. When large numbers of people cook meals and drive vehicles, they tend to dirty the air. In this chapter, we focus on four gases that contribute to air pollution at the surface of the Earth. One of these gases, carbon monoxide, is odorless; the other three—ozone, sulfur dioxide, and nitrogen dioxide—have characteristic odors. With sufficient exposure, all are hazardous to your health, even at concentrations well below 1 ppm. Together with particulate matter (PM), they represent the most serious air pollutants. Let’s now examine the health effects of each.
Today, over half of the people in the world live in cities. In contrast, in 1900 this value was 10–15%.
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Figure 1.4 Sunny November spring day in Santiago, Chile.
Figure 1.5 A propane camping stove. In the stratosphere, ozone absorbs some wavelengths of UV radiation. Look for more about this in Chapter 2.
Figure 1.6 White pine needles damaged by ozone. (Courtesy of Missouri Botanical Garden PlantFinder)
SO2 and NO2 dissolve in moist lung tissue to form sulfurous acid and nitric acid, respectively. Look for more about these acids in Chapter 6.
Chapter 6 provides more information about acid rain.
Carbon monoxide (CO) has earned the nickname “silent killer,” as it has no color, taste, or smell. When you inhale carbon monoxide, it passes “silently” into your bloodstream. Once there, it interferes with the ability of your hemoglobin to carry oxygen. If exposed to carbon monoxide, at first you may feel dizzy, nauseous, or your head may hurt, symptoms that easily could be mistaken for other illnesses. Continued exposure, however, can lead to severe illness, even death. Automobile exhaust is one source of carbon monoxide. Charcoal grills and propane camping stoves are others. These can be used both outdoors and indoors. However, if used indoors, they need proper venting (Figure 1.5). Ozone (O3) has a sharp odor, one that you may have detected around a photocopier, an electric motor, or welding equipment. Even at very low concentrations, ozone can reduce lung function in healthy people. Symptoms of exposure include chest pain, coughing, sneezing, and lung congestion. Ozone also mottles the leaves of crops and yellows pine needles (Figure 1.6). At the Earth’s surface, ozone is definitely a bad actor. At high altitudes, however, it plays an essential role in screening out ultraviolet radiation. Sulfur dioxide (SO2) has a sharp, unpleasant odor and dissolves in the moist tissue of your lungs to form an acid. The elderly, the young, and individuals with lung diseases such as emphysema or asthma are most susceptible to sulfur dioxide poisoning. At present, sulfur dioxide in the air comes primarily from the burning of coal. For example, the 1952 London smog that eventually killed over 10,000 people was in part caused by the emissions of coal-fired stoves. The causes of death included acute respiratory distress, heart failure (from preexisting conditions), and asphyxiation. Some who survived had permanent lung damage. Nitrogen dioxide (NO2) has a characteristic brown color and also damages lung tissue. Like sulfur dioxide, it can combine with fog, mist, snow, or rain to produce acid rain. Nitrogen dioxide is produced in our atmosphere from another pollutant (nitrogen monoxide, colorless) emitted both from the hot engines of vehicles and from the hot fires of coal-fired power plants. Oxides of nitrogen also can form naturally in grain silos, causing injury or death to the farmers who may inadvertently inhale the gases. Particulate matter (PM) is the least understood of these five air pollutants. Particulate matter is a complex mixture of tiny solid particles and microscopic liquid droplets. It is classified by size rather than composition, because the particle size determines the health consequences. PM10 includes particles with an average diameter of
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Figure 1.7 A 2004 wildfire near San Jose, California. This fire is releasing particulate matter, some of which is visible as soot.
10 μm or less, a length on the order of 0.0004 inches. PM2.5 is a subset of PM10 and includes particles with an average diameter of less than 2.5 μm. These tinier and more deadly particles are sometimes called fine particles. Particulate matter originates from many sources, including truck and car engines, coal-burning power plants, fires, and blowing dust. Sometimes particulate matter is visible as soot or smoke (Figure 1.7). However, of more concern are the particles too tiny to see: PM10 and PM2.5. These particles, when inhaled, go deep into your lungs and cause irritation. The smallest particles pass from your lungs into your bloodstream and can cause heart disease or aggravate an existing condition. We end this section with a fact that may surprise you. All of the air pollutants that we just listed can occur naturally! For example, a wildfire (see Figure 1.7) produces particulate matter and carbon monoxide, lightning produces ozone and nitrogen oxides, and volcanoes release sulfur dioxide. The pollutants have the same hazards, whether released from natural or human sources. What are the risks to your health? We now turn to this topic.
1.3
| Air Pollutants and Risk Assessment
Risk is part of living. Although we cannot avoid risk, we still try to minimize it. For example, certain practices are illegal because they carry risks that are judged to be unacceptable. Other activities carry high risks and we label them as such. For example, cigarette packages carry a warning about lung cancer. Wine bottles carry warnings about birth defects and about operating machinery under the influence of alcohol. The absence of a warning, however, does not guarantee safety. The risk may be too low to label, it may be obvious or unavoidable, or it may be far outweighed by other benefits. Warnings are just that. They do not mean that somebody will be affected. Rather, they report the likelihood of an adverse outcome. Let’s say that the odds of dying from an accident are one in a million for each 30,000 miles traveled. On average, this means that one person out of every million traveling 30,000 miles would die in an accident. Such prediction is not simply a guess, but the result of risk assessment, the process
A micrometer (μm) is a millionth (1026) of a meter (m). The term micron also is used for this unit.
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The U.S. EPA was formed in 1970 by President Richard Nixon. Senators from earlier years also played key roles in the legislation.
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of evaluating scientific data and making predictions in an organized manner about the probabilities of an outcome. When is it risky to breathe the air? Fortunately, existing air quality standards can offer you guidance. We say guidance, because standards are set through a complex interaction of scientists, medical experts, governmental agencies, and politicians. People may not necessarily agree on which standards are reasonable and safe. Standards also change over time, as new scientific knowledge is generated. In the United States, national air quality standards were established in 1970 as a result of the Clean Air Act. If pollutant levels fall below these standards, presumably the air is healthy to breathe. We say “presumably” because air quality standards change over time, usually becoming stricter. If you look worldwide, you will find that air quality regulations vary both in their strictness and in the degree to which they are enforced. The risks presented by an air pollutant are a function of both toxicity, the intrinsic health hazard of a substance, and exposure, the amount of the substance encountered. Toxicities are difficult to accurately assess for many reasons, including that it is unethical to run experiments on people. Even if data were available, we still would have to determine the levels of risk that are acceptable for different groups of people. In spite of the complexities, government agencies have succeeded in establishing limits of exposure for the major air pollutants. Table 1.2 shows the National Ambient Air Quality Standards established by the U.S. Environmental Protection Agency (EPA). Here, ambient air refers to the air surrounding us, usually meaning the outside air. As our knowledge grows, we modify these standards. For example, in 2006 these standards were lowered for PM2.5. Similarly, in 2008 they were lowered for ozone.
Table 1.2
U.S. National Ambient Air Quality Standards
Pollutant
Standard (ppm)
Approximate Equivalent Concentration (μg/m3)
Carbon monoxide 8-hr average
9
10,000
1-hr average
35
40,000
Nitrogen dioxide Annual average
0.053
100
8-hr average
0.075
147
1-hr average
0.12
235
Ozone
Particulates* PM10, annual average
—
50
PM10, 24-hr average
—
150
PM2.5, annual average
—
15
PM2.5, 24-hr average†
—
35
Annual average
0.03
80
24-hr average
0.14
365
3-hr average
0.50
1,300
Sulfur dioxide
*PM10 refers to all airborne particles 10 μm in diameter or less. PM2.5 refers to particles 2.5 μm in diameter or less. —The unit of ppm is not applicable to particulates. †
PM2.5 standards are likely to be revised after 2011.
Source: U.S. Environmental Protection Agency. Standards also exist for lead, but are not included here.
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Exposure is far more straightforward to assess than toxicity, because exposure depends on factors that we more easily can measure. These include: ■
■
■
Concentration in the air The more toxic the pollutant, the lower the concentration must be set. Concentrations are expressed either as parts per million or as micrograms per cubic meter (μg/m3), as shown in Table 1.2. Earlier, we used the prefix microwith micrometers (μm), meaning a millionth of a meter (1026 m). Similarly, one microgram (μg) is a millionth of a gram (g), or 1026 g. Length of time Higher concentrations of a pollutant can be tolerated only briefly. A pollutant may have several standards, each for a different length of time. Rate of breathing Physically active people, such as athletes or laborers, breathe at a higher rate. If the air quality is poor, reducing activity is one way to reduce exposure.
Suppose you collect an air sample on a city street. An analysis shows that it contains 5000 μg of carbon monoxide (CO) per cubic meter of air. Is this concentration of CO harmful to breathe? We can use Table 1.2 to answer this question. Two standards are reported for carbon monoxide, one for a 1-hour exposure and another for an 8-hour exposure. The 1-hour exposure is set at a higher level (4 3 104 μg CO/m3) because a higher concentration can be tolerated for a short time. Both concentrations are expressed in scientific notation, a system for writing numbers as the product of a number and 10 raised to the appropriate power. Scientific notation enables us to avoid writing strings of zeros either before or after the decimal point. For example, the value 1 3 104 is equivalent to 10,000. To understand this conversion, simply count the number of zeros to the right of the 1 in 10,000. There are four of them. The number 1 is then multiplied by 104 to obtain 1 3 104 μg CO/m3. Similarly, 4 3 104 μg CO/m3 is equivalent to 40,000 μg CO/m3. Scientific notation is even more useful for very large numbers, such as the 20,000,000,000,000,000,000,000 molecules in a typical breath. In scientific notation, this value is written as 2 3 1022 molecules. Using scientific notation, we now can express the value of 5000 μg CO/m3 as 5 3 103 μg CO/m3. Clearly, this value is less than either standard. In the case of an 8-hour exposure, 5 3 103 is less than 1 3 104. Similarly, for a 1-hour period, 5 3 103 is less than 4 3 104. For all values, the units are μg CO/m3. Table 1.2 also allows us to assess the relative toxicities of pollutants. For example, we can compare the 8-hour average exposure standards for carbon monoxide and ozone: 9 ppm vs. 0.075 ppm. Doing the math, ozone is about 130 times more hazardous to breathe than carbon monoxide! Nonetheless, carbon monoxide still can be exceedingly dangerous. As the “silent killer,” it may impair your judgment before you recognize the danger.
Your Turn 1.7
Estimating Toxicities
a. Which pollutant in Table 1.2 is likely to be the most toxic? Exclude particulate matter. b. Examine the particulate matter standards. Earlier, we stated that “fine particles,” PM2.5, are more deadly than the coarser ones, PM10. Do the values in Table 1.2 bear this out? Answer a. O3. This is a hard call, as no common exposure period exists on which to base the comparison. Clearly CO is not the most toxic, as all the standards are higher. It is not NO2, because SO2 has a lower annual standard. Between NO2 and O3, ozone has the tighter standards.
Although the standards for air pollutants are expressed in parts per million, the concentrations of sulfur dioxide and nitrogen dioxide could conveniently be reported
1 μg is approximately the mass of a period printed on a page.
If you need help with exponents, consult Appendix 2.
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in parts per billion (ppb), meaning one part out of one billion, or 1000 times less concentrated than one part per million. sulfur dioxide 0.030 ppm 5 30 ppb nitrogen dioxide 0.053 ppm 5 53 ppb As these values reveal, converting from parts per million to parts per billion involves moving the decimal point three places to the right.
Your Turn 1.8
Living Downwind
Copper metal can be recovered from copper ore by smelting, a process that releases sulfur dioxide (SO 2). Let’s assume that a woman living downwind of a smelter inhaled 1050 μg of SO 2 in a day. a. If she inhaled 15,000 liters (15 m3) of air per day, would she exceed the 24-hour average for the U.S. National Ambient Air Quality Standards for SO2? Support your answer with a calculation. b. If she were exposed at this rate every day, would she exceed the annual average?
To end this section, we note that our perception of a risk also plays an important role. For example, the risks of traveling by car far exceed those of flying. Each day in the United States, more than 100 people die in automobile accidents. Yet some people avoid taking a flight because of their fear of a plane crash. Similarly, some people fear living near a nuclear power plant. Yet as some notorious hurricanes have demonstrated, living in a coastal area can be a far riskier proposition. Whether perceived as a risk or not, air pollution presents real hazards, both to present and future generations. In the next section, we offer you the tools to assess these hazards.
1.4
This act was the impetus for green chemistry, a topic introduced in Chapter 0.
| Air Quality and You
Depending on where you live, you will breathe air of different quality. Some locations always have good air; others have air of moderate quality, and still others have unhealthy air much of the time. As we will see, the differences arise because of the number of people living in a region, their activities, the geographical features of the region, the prevailing weather patterns, and the activities of people in neighboring regions. To improve air quality, many nations have enacted legislation. For example, we already have cited the U.S. Clean Air Act (1970) that led to the establishment of air quality standards. Like many environmental laws, this one focused on limiting our exposure to hazardous substances. It has been named as a “command and control law” or an “end of the pipe solution” because it tries to limit the spread of hazardous substances or clean them up after the fact. The Pollution Prevention Act (1990) was a significant piece of legislation that followed the Clean Air Act. It focused on preventing the formation of hazardous substances, stating that “pollution should be prevented or reduced at the source whenever feasible.” The language shift is significant. Rather than attempting to regulate existing pollutants, people should not produce them in the first place! With the Pollution Prevention Act, it became national policy to employ practices that reduce pollutants at their source.
Your Turn 1.9
The Logic of Prevention
Take off your muddy shoes at the door rather than cleaning up the carpet later! List three “common sense” examples that prevent air pollution rather than cleaning it up after the fact. Hint: Revisit the first activity in this chapter on “air prints.”
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40% Ozone
20% 0%
National standard
⫺20%
PM2.5
CO
⫺40%
NO2 PM10
⫺60% SO2
⫺80% ⫺100% 80
82
84
86
88
90
92
94
96
98
00
02
04
06
Figure 1.8 U.S. average levels of air pollutants (at selected sites) compared with national ambient air quality standards, 1980–2006. Note: Data for PM10 and PM2.5 start in 1990 and 1999, respectively. Source: U.S. Environmental Protection Agency, Latest Findings on National Air Quality, www.epa.gov/air/airtrends/2007
The de The Th ddecrease decr ecr crea ease ea se in i the h concentration of air pollutants in the United States has been dramatic (Figure 1.8). Some improvements occurred through a combination of laws and regulations, such as the ones we just mentioned. Others stemmed from local decisions. For example, a community may have built a new public transportation system or an industry may have installed more modern equipment. Still others occurred because of the ingenuity of chemists, most notably via a set of practices called “green chemistry.” Look for green chemistry examples throughout this book that are designated with the Green Chemistry icon. However, aggregate data such as those in Figure 1.8 hide the fact that people in some places still breathe dirty air. Although air quality may have improved on average, people in some metropolitan areas breathe air that contains unhealthy levels of pollutants. Check the data for the United States presented in Table 1.3. The label “unhealthy” means just that. As we described earlier, air pollutants are the perpetrators of biological mischief. To help you more quickly assess the hazards,
Table 1.3
The green chemistry examples in this text include some of the winners of the prestigious Presidential Green Chemistry Challenge Awards.
Air Quality Data for Selected U.S. Metropolitan Areas # o f U n h e a l t h y D a y s / Ye a r *
Metropolitan Area
O3
PM2.5
Boston
10
5
Chicago
11
10
Cleveland
17
9
Houston
35
2
Los Angeles
62
38
Pittsburgh
14
42
Phoenix
10
2
Sacramento
28
10
Seattle Washington, DC
1
4
19
8
*Days for which the AQI for the pollutant exceeds 100, five year average (2001–2005). Source: EPA, Latest Findings on National Air Quality.
Seattle has few bad air days because of its rainy climate. Section 1.12 describes the connections between climate and smog formation.
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Table 1.4
Levels for the Air Quality Index
Air Quality Index (AQI) Values
Levels of Health Concern
Colors
When the AQI is in this range:
…air quality conditions are:
…as symbolized by this color.
Good
Green
Moderate
Yellow
101–150
Unhealthy for sensitive groups
Orange
151–200
Unhealthy
Red
201–300
Very unhealthy
Purple
301–500
Hazardous
Maroon
0–50 51–100
EPA air quality logo
The AQI for a particular day is set by the highest pollutant. To quote the EPA, “If a certain area had AQI values of 90 for ozone and 88 for sulfur dioxide, the AQI value would be 90 on that day.”
the U.S. EPA developed the color-coded Air Quality Index (AQI) shown in Table 1.4. This Index is scaled from 1–500, with the value of 100 pegged to the national standard for the pollutant. Green or yellow (#100) indicates air of good or moderate quality. Orange indicates that the air has become unhealthy for some groups. Red, purple, or maroon (.150) indicates that the air is unhealthy for everybody to breathe. Some newspapers provide only a general air quality report. For example, the air quality may be listed as “moderate” for a city. This means that at least one pollutant was moderate, but possibly others were as well. Sometimes the report is given in general terms, but numerical. For example, if two pollutants were present, one with a value of 85 and the other with a value of 91, the daily value would be reported as 91. More and more, though, metropolitan areas are beginning to report each pollutant separately. This is helpful, because how you act depends on which pollutants are present. For example, Figure 1.9 shows the air quality forecast for carbon monoxide, ozone, and particulates on a sunny spring day in Phoenix, Arizona. Ozone clearly was the pollutant of concern. TODAY THU 05/14/2009
TOMORROW FRI 05/15/2009
EXTENDED SAT 05/16/2009
O3
61 CAVE CREEK & FOUNTAIN HILLS
80 MODERATE
97 MODERATE
104 UNHEALTHY FOR SENSITIVE GROUPS
CO
07 GREENWOOD
09 GOOD
11 GOOD
09 GOOD
PM10
53 WEST FORTY THIRD
55 MODERATE
61 MODERATE
48 GOOD
PM2.5
45 PHOENIX SUPERSITE
46 GOOD
49 GOOD
45 GOOD
FORECAST DATE
YESTERDAY WED 05/13/2009
AIR POLLUTANT
Highest AQI Reading/SITE
Figure 1.9 Air quality forecast for Phoenix on May 14, 2009. Values .100 are unhealthy. Source: Arizona Department of Environmental Quality.
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1.5
29
| Where We Live: The Troposphere
As we pointed out earlier, our atmosphere is a mixture that varies slightly in composition at different locations. About 75% of the air is found within about 6 miles (,10 kilometers) of the surface of our planet. This is the troposphere, the lower region of the atmosphere in which we live that lies directly above the surface of the Earth. Tropos is Greek for “turning” or “changing.” The troposphere contains the air currents and turbulent storms that turn and mix our air. The warmest air in the troposphere usually lies at ground level because the Sun heats the ground, which in turn warms the air above it. Cooler air is found higher up, a phenomenon you may have observed as you hike or drive to higher elevations. However, air inversions occur when cooler air gets trapped beneath warmer air. Air pollutants can accumulate in an inversion layer, especially if the layer remains stationary for an extended period. This often occurs in cities ringed by mountains, such as Salt Lake City (Figure 1.10). Air pollutants might better be termed “people fumes.” One hundred years ago, our Earth was home to fewer than 2 billion people. Soon we will reach the 7-billion mark, with the majority of people living in urban regions. This growth in population has been accompanied by a massive growth in both the consumption of resources and the production of waste. The waste that we stash in our atmosphere is called air pollution. Air pollution provides the first context in which we can discuss sustainability, the topic that set the stage for this book. As we pointed out in Chapter 0, we need to make decisions with an eye not only for today’s outcomes but also for the needs of generations to come. It makes sense to avoid actions that produce pollutants that can compromise our health and well-being. Does this sound familiar? Again, this is the logic behind the Pollution Prevention Act of 1990, legislation that calls for preventing pollution rather than controlling our exposure to it. The chemistry, a set T he Pollution Th Poll Poll Po llut utiion ut io Prevention P revention Act provided the impetus for green chemistry, of principles to guide all in the chemical community, including teachers and students. Green chemistry is “benign by design.” It calls for designing chemical products and processes that reduce or eliminate the use or generation of hazardous substances. Begun under the EPA Design for the Environment Program, green chemistry reduces pollution through the design or redesign of chemical processes. The goal is to use less energy, create less waste, use fewer resources, and use renewable resources. Green chemistry is a tool for achieving sustainability, rather than an end in itself. Innovative “green” chemical methods already have decreased or eliminated toxic substances used or created in chemical manufacturing processes. For example,
Sun
Warm air
Trapped pollution Cold air
(a)
(b)
Figure 1.10 (a) An air inversion can trap pollution. (b) A air inversion, trapping a smoggy layer of air over Salt Lake City, Utah.
The depth of the troposphere varies from the mid-latitudes (,12 miles) to the poles (,5 miles).
Hurricanes in the troposphere clearly illustrate the meaning of the Greek word tropos (5 turning or changing).
Principles of green chemistry are listed on the inside cover of this book.
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The green chemistry awards are given in five areas: greener synthetic pathways, greener reaction conditions, designing greener chemicals, small business, and academia.
we now have cheaper and less wasteful ways to produce ibuprofen, pesticides, disposable diapers, and contact lenses. We have new dry cleaning methods and recyclable silicon wafers for integrated circuits. The research chemists and chemical engineers who developed these methods have received Presidential Green Chemistry Challenge Awards. Begun in 1995, these presidential-level awards recognize chemists for their innovations on behalf of a less polluted world. Each year since 1996, five awards have been given with the theme “Chemistry is not the problem, it’s the solution.”
Consider This 1.10
Green Chemistry
Recall these two green chemistry principles: It is better to prevent waste than to treat or clean up waste after it is formed. It is better to use less energy. a. Why is it better to use less energy? Give two examples that demonstrate the connection between using energy and putting waste in the air. b. Now choose an air pollutant. Give two examples that demonstrate why it makes more sense to prevent its formation rather than to try to clean it up once in the air.
Bottom line: Nobody wants dirty air. It makes you sick, reduces the quality of your life, and may hasten your death. However, the problem is that many people have become so accustomed to breathing dirty air, that they don’t notice it. Recall the concept of shifting baselines mentioned in Chapter 0. Haze is now so common that we have forgotten the clear days when it seemed we really could see forever. Burning eyes and breathing disorders have become so common, that we have forgotten that they once were not. We have become accustomed to living in megacities, urban areas with 10 million people or more. Tokyo, New York City, Mexico City, and Mumbai are examples. Pollutants such as wood smoke, car exhaust, and industrial emissions all concentrate in the troposphere around megacities. Clearly, we have some problems! As promised, your knowledge of chemistry can lead you to making better choices to deal with these problems, both as an individual and in your local community. The next two sections give you a better grasp of the language of chemistry. We then use this language to approach the issues of air pollution in more detail.
1.6
The periodic table lists the known elements, as we will see momentarily.
|
Classifying Matter: Pure Substances, Elements, and Compounds
In describing air and its quality, we already have employed several chemical names. For example, in Section 1.2 we listed nitrogen, oxygen, argon, and usually water vapor as four of the pure substances that make up the majority of our atmosphere. We named some pollutants found at very low concentrations: ozone, sulfur dioxide, carbon monoxide, and nitrogen dioxide. We included their chemical formulas as well: O3, SO2, CO, and NO2. We also mentioned the terms atom and molecule. For example, we noted that air is a mixture that contains different molecules (and a few atoms). In this section, we explain more about elements and compounds. In the next, we will focus on atoms and molecules that these elements and compounds contain. Matter consists of elements and compounds, as shown in Figure 1.11. An element is one of the 100 or so pure substances in our world from which compounds are formed. As we will see shortly, they contain only one type of atom. Nitrogen (N2), oxygen (O2), and argon (Ar) are examples of elements. So is ozone (O3), another form of oxygen. Over 100 elements are known. In contrast, a compound is a pure substance made up of two or more different elements in a fixed, characteristic chemical combination. Compounds
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Matter
Pure substances
Elements
Mixtures
Compounds
Figure 1.11 One way to classify matter.
contain two or more different types of atoms. For example, water (H2O) is a compound of the elements oxygen and hydrogen. Similarly, carbon dioxide (CO2) is a compound of the elements oxygen and carbon. In CO2, the two elements are chemically combined and are no longer in their elemental forms. Sulfur dioxide (SO2) and methane (CH4) are other examples of compounds. About 90 elements occur naturally on Earth and, as far as we know, elsewhere in the universe. The others have been created from existing elements through nuclear reactions. Plutonium is probably the best known of the human-made elements, although it does occur in trace amounts in nature. The vast majority of elements are solids. Nitrogen, oxygen, argon, and eight other elements are gases; and only bromine and mercury are liquids at room temperature. An alphabetical list of the elements and their chemical symbols, one- or two-letter abbreviations for the elements, appears facing the inside back cover of the text. These symbols, established by international agreement, are used throughout the world. Some of them make immediate sense to those who speak English or related languages. For example, oxygen is O, nitrogen is N, sulfur is S, and nickel is Ni. Other symbols have their origin in other languages. For example, Fe is iron, Pb is lead, Au is gold, and Hg is mercury. These metals were known to the ancients and given Latin names long ago. For example, ferrum is iron, plumbum is lead, aurum is gold, and hydrargyrum is mercury. Elements have been named for properties, planets, places, and people. Hydrogen (H) means “water former,” because hydrogen gas (H2) burns in oxygen (O2) to form the compound water (H2O). Neptunium (Np) and plutonium (Pu) were named after two planets in our solar system. Berkelium (Bk) and californium (Cf) honor the Berkeley lab in which a team of researchers first produced them. Darmstadtium (Ds) and roentgenium (Rg) were the most recently produced elements with names at the time this book went to press. The former was named after Darmstadt, the city in Germany in which it was discovered; the latter was named after Wilhelm Roentgen, Only a few atoms of each have been produced. A new element, #117, was reported in 2010 and currently is the heaviest element known.
Your Turn 1.11
Plutonium can fuel both nuclear reactors and nuclear bombs. See Chapter 7 for the details.
Chemical symbols sometimes also are referred to as atomic symbols.
In 2006, Pluto lost its status as a planet.
William Roentgen discovered X-rays and received the first Nobel Prize in physics.
Pure Substances in Air
a. Hydrogen (0.54 ppm), helium (5 ppm), and methane (17 ppm) are found in our atmosphere. Two of these are elements. Which ones? b. List five other substances found in the air and classify them as elements or compounds. c. Express each concentration in part a as a percent.
It is fitting that the 19th-century Russian chemist Dmitri Mendeleev has his own element (Md), because the most common way of arranging the elements—the
Lothar Meyer, a German chemist, also developed a periodic table at the same time as Mendeleev.
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1A
8A
1 H
2 He
2A
3A
4A
5A
6A
7A
3 Li
4 Be
5 B
6 C
7 N
8 O
9 F
10 Ne
11 Na
12 Mg
13 Al
14 Si
15 P
16 S
17 Cl
18 Ar
19 K
3B
4B
5B
6B
7B
8B
1B
2B
20 Ca
21 Sc
22 Ti
23 V
24 Cr
25 Mn
26 Fe
27 Co
28 Ni
29 Cu
30 Zn
31 Ga
32 Ge
33 As
34 Se
35 Br
36 Kr
37 Rb
38 Sr
39 Y
40 Zr
41 Nb
42 Mo
43 Tc
44 Ru
45 Rh
46 Pd
47 Ag
48 Cd
49 In
50 Sn
51 Sb
52 Te
53 I
54 Xe
55 Cs
56 Ba
57 La
72 Hf
73 Ta
74 W
75 Re
76 Os
77 Ir
78 Pt
79 Au
80 Hg
81 Tl
82 Pb
83 Bi
84 Po
85 At
86 Rn
87 Fr
88 Ra
89 Ac
104 Rf
105 Db
106 Sg
107 Bh
108 Hs
109 Mt
110 Ds
111 Rg
112
113
114
115
116
117
118
58 Ce
59 Pr
60 Nd
61 Pm
62 Sm
63 Eu
64 Gd
65 Tb
66 Dy
67 Ho
68 Er
69 Tm
70 Yb
71 Lu
90 Th
91 Pa
92 U
93 Np
94 Pu
95 Am
96 Cm
97 Bk
98 Cf
99 Es
100 Fm
101 Md
102 No
103 Lr
Metals Metalloids Nonmetals
Figure 1.12 A simplified period table that shows the locations of metals, metalloids, and nonmetals.
Metals, nonmetals, and the ions they form are discussed in Section 5.6.
Semiconductors are explained in Section 8.7.
periodic table—reflects the system he developed. This is an orderly arrangement of all the elements based on similarities in their properties. The inside back cover has a copy for handy reference, and we will explain more about it in Chapter 2. Figure 1.12 shows a simplified version of the periodic table that lists the elements by number but does not include their masses. The light green shading indicates the metals, elements that are shiny and conduct electricity and heat well. These include familiar substances such as iron, gold, and copper. Far fewer are the nonmetals, elements that do not conduct heat or electricity well and have no one characteristic appearance. These elements are indicated by the light blue shading and include sulfur, chlorine, and oxygen. A mere eight elements fall into a category known as metalloids, elements that lie between metals and nonmetals on the periodic table and do not fall cleanly into either category. Metalloids also are called semimetals and are indicated in light red shading. The semiconductors silicon and germanium are examples of metalloids. The elements fall into vertical columns called groups. These organize elements according to important properties that they have in common and are numbered from left to right. Some groups are given names as well. For example, a halogen is one of the reactive nonmetals in Group 7A, such as fluorine (F), chlorine (Cl), bromine (Br), or iodine (I). Similarly, a noble gas is one of the inert elements in Group 8A that undergoes few, if any, chemical reactions. We already mentioned argon, a noble gas that is a constituent of our atmosphere. You may recognize helium as the noble
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gas used to make balloons float, as it is less dense than air. Radon is the one noble gas that is radioactive, a characteristic that distinguishes it from the other elements in Group 8A. Although only 100 or so elements exist, over 20 million compounds have been isolated, identified, and characterized. Some are very familiar naturally occurring substances such as water, table salt, and sucrose (table sugar). Many known compounds were chemically synthesized by men and women across our planet. You might be wondering how 20 million compounds could possibly be formed from so few elements. In short, elements have the ability to combine in many different ways. For example, carbon dioxide is a chemical combination of the elements carbon and oxygen. All pure samples of carbon dioxide contain 27% carbon and 73% oxygen by mass. Thus, a 100-g sample of carbon dioxide always consists of 27 g of carbon and 73 g of oxygen, chemically combined to form this particular compound. These values never vary, no matter the source of the carbon dioxide. This illustrates the fact that every compound exhibits a constant characteristic chemical composition. Carbon monoxide (CO) is a different compound of carbon and oxygen. Pure samples of carbon monoxide contain 43% carbon and 57% oxygen by mass. Thus, 100 g of carbon monoxide contain 43 g of carbon and 57 g of oxygen, a composition different from that of carbon dioxide. This is not surprising, because carbon monoxide and carbon dioxide are two different compounds. As we will see, each compound has its own set of properties. For example, water (H2O) is a compound that is 11% hydrogen and 89% oxygen by weight. At room temperature, water is a colorless, tasteless liquid. At sea level, it boils at 100 8C and freezes at 0 8C. All samples of pure water have these same properties. Water is composed of molecules, a term that we now have used several times. The next section will help you to use the terms atom and molecule with more confidence.
1.7
33
Look for more about density in Chapter 5.
Radon affects the quality of indoor air, as we will see in Section 1.13. Chapter 7 tells more about other radioactive substances.
A small paper clip weighs about a gram.
Water is not tasteless if it contains dissolved gases or minerals. Look for more about water as a solvent in Chapter 5. As we will see, water is rarely “pure.”
| Atoms and Molecules
The definitions we just gave for elements and compounds made no assumptions about the nature of matter. We now know that elements are made up of atoms, the smallest unit of an element that can exist as a stable, independent entity. The word atom comes from the Greek for “uncuttable.” Although today it is possible to “cut” atoms using specialized processes, atoms remain indivisible by ordinary chemical or mechanical means. Atoms are extremely small. Because they are so tiny, we need huge numbers of atoms in order to see, touch, or weigh them. For example, the molecules in a single drop of water contain about 5 3 1021 atoms. This is roughly a trillion times greater than the 6.9 billion people on Earth, almost enough to give each person a trillion atoms. As Figure 1.13 reveals, atoms now can be photographed. Using a scanning tunneling microscope, scientists at the IBM Almaden Research Center lined up iron atoms on a copper surface to create the kanji (Japanese character) for “atom.” Nanotechnology refers to the creation of materials at the atomic and molecular (nanometer) scale: 1 nanometer (nm) 5 1 3 1029 m. This kanji is a few nanometers high and wide. At this size, about 250 million nanoletters could fit on a cross section of a human hair, equivalent to about 90,000 pages of text! Using the concept of atoms, we can better explain the terms element and compound. Elements are made up of only one kind of atom. For example, the element carbon is made up only of carbon atoms. By contrast, compounds are made up of two or more different kinds of atoms. For example, the compound carbon dioxide contains carbon and oxygen atoms. Similarly, water is made up of hydrogen and oxygen atoms. But we need to be careful with our language. The carbon and oxygen atoms in carbon dioxide are not present as such. Rather, the carbon and oxygen atoms are
Figure 1.13 Iron atoms arranged on a copper surface, as imaged with a scanning tunneling microscope. This is the Japanese kanji for the word genshi (atom).
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Atoms are color-coded:
chemically combined to form a carbon dioxide molecule, two or more atoms held together by chemical bonds in a certain spatial arrangement. More specifically, two oxygen atoms (red ) are combined with one carbon atom (black) to form a carbon dioxide molecule. Similarly, the water molecule contains two hydrogen atoms (white) combined with one oxygen atom (red ).
carbon hydrogen oxygen nitrogen sulfur
water molecule
Section 3.3 explains why these two molecules have different shapes.
carbon dioxide molecule
A chemical formula is a symbolic way to represent the elementary composition of a substance. It reveals both the elements present (by chemical symbols) and the atom ratio of those elements (by the subscripts). For example, in the compound CO2, the elements C and O are present in a ratio of one carbon atom for every two oxygen atoms. Similarly, H2O indicates two hydrogen atoms for each oxygen atom. Note that when an atom occurs only once, such as the O in H2O or the C in CO2, the subscript of “1” is omitted. Some elements exist as single atoms, such as helium or radon. We represent these as He and Rn, respectively. Other elements exist as molecules. For example, nitrogen and oxygen are found in our atmosphere as N2 and O2 molecules. Each is a diatomic molecule, meaning that it is a molecule consisting of two atoms. These representations clearly show the difference.
oxygen molecule
nitrogen molecule
helium atom
Table 1.5 summarizes our discussion of elements, compounds, and mixtures, listing both what we can observe experimentally and what exists at the atomic level that we cannot see. We now apply these concepts to the mixture we call air. Some of its components, such as nitrogen, oxygen, and argon, are elements. Others, most notably water vapor and carbon dioxide, are compounds. All of the compounds discussed so far contain molecules (e.g., CO2 and H 2O). But with the elements, it is not so simple. In the troposphere, the elements nitrogen and oxygen exist primarily as diatomic molecules (N2 and O2). In contrast, elements such as argon and helium exist as uncombined atoms.
Table 1.5
Types of Matter
Substance
Example
Definition
Contains . . .
element
O2, oxygen Ar, argon
cannot be broken down into simpler substances, although different forms of an element are possible
one type of atom only
compound
H2O, water
can be broken down into elements and has a fixed composition
two or more different types of atoms
mixture
air
can be separated into two or more components and has a variable composition of elements, compounds, or both
many types of atoms, molecules, or both
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Dry air is composed mainly of nitrogen and oxygen, that is, molecules of N2 and O2. If the air is humid, add in some water vapor in the form of H2O molecules. Dry air contains just under 1% of Ar (argon) atoms, as well as tiny amounts of He (helium) atoms, Xe (xenon) atoms, and extremely small amounts of Rn (radon) atoms. Remember also to include the 390 ppm of carbon dioxide, meaning 390 CO2 molecules per 1 3 106 of the molecules and atoms in the air.
Skeptical Chemist 1.12
The Chemistry of Lawn Care
News reports and advertisements should be viewed with a critical eye for their accuracy. For example, a lawn care service ad reports its fertilizers as “a balanced blend of nitrogen, phosphorus, and potassium. They have an organic nature, made up of carbon molecules. These fertilizers are biodegradable and turn into water.” Edit the text of this ad to fix the chemical glitches.
1.8
|
Names and Formulas: The Vocabulary of Chemistry
If chemical symbols are the alphabet of chemistry, then chemical formulas are the words. The language of chemistry, like any other language, has rules of spelling and syntax. In this section we help you to “speak chemistry” using chemical formulas and names. As you’ll see, each name corresponds uniquely to one chemical formula. However, chemical formulas are not unique and may correspond to more than one name. In addition, some compounds are known by several names. In this section, we follow a need-to-know philosophy. We help you learn what you need to know to understand the topic at hand and omit other naming rules until the need arises. Right now, you need to know the chemical names and formulas of compounds that relate to the air you breathe. So, for now, we work on these. We already named some of the pure substances found in air, including carbon monoxide, carbon dioxide, sulfur dioxide, ozone, water vapor, and nitrogen dioxide. Although it may not be apparent, this list includes two types of names: systematic and common. Systematic names for compounds follow a reasonably straightforward set of rules. Here are the rules for compounds composed of two nonmetals, such as carbon dioxide (CO2) and carbon monoxide (CO): ■
■
■
Name each element in the chemical formula, modifying the name of the second element to end in -ide. For example, oxygen becomes oxide, and sulfur becomes sulfide. Use prefixes to indicate the numbers of atoms in the chemical formula (Table 1.6). For example, di- means 2, and thus the name carbon dioxide means two oxygen atoms for each one carbon atom. Omit the prefi x mono- if there is only one atom for the fi rst element in the chemical formula. For example, CO is carbon monoxide, not monocarbon monoxide.
If instead you are writing a chemical formula from a name, keep in mind that the subscript of 1 is not used in chemical formulas. Thus the chemical formula for carbon dioxide is CO2 and not C1O2. Similarly, carbon monoxide is CO and not C1O1. The next activity gives you a chance to practice.
Section 5.7 explains another set of naming rules, those for ionic compounds.
Remember that ozone, O3, is an element, not a compound.
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Table 1.6
Prefixes Used in Naming Compounds
Prefix
Meaning
Prefix
mono-
1
hexa-
6
di- or bi-
2
hepta-
7
tri-
3
octa-
8
tetra-
4
nona-
9
penta-
5
deca-
10
Your Turn 1.13
Meaning
Oxides of Sulfur and Nitrogen
a. Write chemical formulas for nitrogen monoxide, nitrogen dioxide, dinitrogen monoxide, and dinitrogen tetraoxide. b. Give chemical names for SO2 and SO3. Answer a. NO, NO2, N2O, and N2O4. Note: NO and N2O also are called nitric oxide and nitrous oxide.
Prefixes in hydrocarbon names methethpropbut-
1 2 3 4
C atom C atoms C atoms C atoms
Look for more about hydrocarbons in Section 4.4.
Some names (“common names”) do not follow a set of rules. Water is one example. You might have expected H2O to be called dihydrogen monoxide. This makes sense! But water was given its name long before anybody knew anything about hydrogen and oxygen. Chemists, being reasonable folks, did not rename water. Rather, they call the stuff they swim in and drink by the common name water, just like everybody else. Ozone (O3) is another common name, as is ammonia (NH3). Common names cannot be figured out; you have to know them or look them up. In the next two sections, we explore the connection between air quality and the fuels we burn. Following our need-to-know philosophy, we need to introduce the names of several hydrocarbons; that is, compounds made up only of the elements hydrogen and carbon. Hydrocarbons follow a very different set of naming rules from the ones just presented. Methane (CH4) is the smallest hydrocarbon. Other small hydrocarbons include ethane, propane, and butane. Although methane may not appear to be a systematic name, it indeed is one if you are willing to accept that meth- means 1 carbon atom. Similarly, ethmeans 2 carbon atoms, and C2H6 is ethane. Prop- means 3 carbon atoms, and but- means 4. So propane is C3H8, and butane is C4H10. Just as mono-, di-, tri-, and tetra- are used to count, so are meth-, eth-, prop-, and but-. In Chapter 4, we will explain the suffix -ane as well as explain the different numbers of C atoms and H atoms in the chemical formulas. These new prefixes are very versatile. They can be used not only at the beginning of chemical names, but also within them, as the next activity shows.
Your Turn 1.14
Mother Eats Peanut Butter
Many generations of students have used the memory aid “mother eats peanut butter” for meth-, eth-, prop-, but-. Use this or another memory aid of your choice to tell how many carbon atoms are in each of these compounds. a. ethanol (a gasoline additive) b. methylene chloride (a component of paint strippers and sometimes an indoor air pollutant) c. propane (the major component in LPG, liquid petroleum gas) Answer b. The meth- in methylene indicates 1 C atom in the chemical formula.
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As we’ll see, hydrocarbon molecules can contain over 50 carbon atoms! For the smaller molecules, use the prefixes shown in Table 1.6. For example, the octane molecule contains 8 carbon atoms. That’s it for names and chemical formulas, at least for now. In the next section, we put this chemical vocabulary to work.
1.9
|
Chemical Change: The Role of Oxygen in Burning
Life on Earth bears the stamp of oxygen. Compounds containing oxygen occur in the atmosphere, in your body, and in the rocks and soils of the planet. Why? The answer is that many different elements combine chemically with oxygen. One such element is carbon. You were already introduced to the compound carbon monoxide, CO, a pollutant listed in Table 1.2. Fortunately, CO is relatively rare in our atmosphere. In contrast, carbon dioxide, CO2 , is far more abundant, but still only a mere 390 ppm. Even so, at this concentration CO2 plays an important role as a greenhouse gas. In this section, we explain how both CO2 and CO are emitted into our atmosphere. As you know, humans exhale CO2 with each breath. Breathing is one natural source of CO2 in our atmosphere. Carbon dioxide also is produced when humans burn fuels. Combustion is the chemical process of burning; that is, the rapid combination of fuel with oxygen to release energy in the form of heat and light. When carboncontaining compounds burn, the carbon combines with oxygen to produce carbon dioxide (CO2). When the oxygen supply is limited, carbon monoxide (CO) is likely to form as well. Combustion is a major type of chemical reaction, a process whereby substances described as reactants are transformed into different substances called products. A chemical equation is a representation of a chemical reaction using chemical formulas. To students, a chemical equation is probably better known as “the thing with an arrow in it.” Chemical equations are the sentences in the language of chemistry. They are made up of chemical symbols (corresponding to letters) that often are combined in the formulas of compounds (the words of chemistry). Like a sentence, a chemical equation conveys information, in this case about the chemical change taking place. A chemical equation also must obey some of the same constraints that apply to a mathematical equation. At the most fundamental level, a chemical equation is a qualitative description of this process: reactant(s)
The concentration of CO2 in the atmosphere is increasing, with serious consequences for planet Earth, as we will describe in Chapter 3.
See Section 4.1 for more about combustion.
product(s)
By convention, the reactants are always written on the left and the products on the right. The arrow represents a chemical transformation and can be read as “is converted to.” The combustion of carbon (charcoal) to produce carbon dioxide as shown in Figure 1.14 can be represented in several ways. One is with chemical names. carbon ⫹ oxygen
carbon dioxide Figure 1.14
Another more common way is to use chemical formulas. C ⫹ O2
CO2
Charcoal burns in air.
[1.1]
This compact symbolic statement conveys a good deal of information. It might sound something like this: “One atom of the element carbon reacts with one molecule of the element oxygen to yield one molecule of the compound carbon dioxide.” Using black for carbon and red for oxygen, we also can represent the molecules and atoms involved.
The colors displayed here for atoms reflect the standard used in molecular modeling software and many model kits.
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These equations are similar to a mathematical expression in that the number and kind of atom on the left side of the arrow must equal the number and kind of each atom on the right: Left side: 1 C and 2 O
Here is an analogy. The building materials used to construct a warehouse (reactants) can be disassembled and rearranged to build three houses and a garage (products).
Right side: 1 C and 2 O
Atoms are neither created nor destroyed in a chemical reaction. The elements present do not change their identities when converted from reactants to products, although they may be bonded in different ways. This relationship is known as the law of conservation of matter and mass: In a chemical reaction, matter and mass are conserved. The mass of the reactants consumed equals the mass of the products formed. Let’s look at another example. Using yellow for sulfur, we can represent how sulfur burns in oxygen to produce the air pollutant sulfur dioxide. S ⫹ O2
SO2
[1.2]
This equation is balanced: the same number and types of atoms are present on each side of the arrow. These atoms, however, were rearranged. This is what a chemical reaction is all about! It is possible to pack even more information into a chemical equation by specifying the physical states of the reactants and products. A solid is designated by (s), a liquid by (l), and a gas by (g). Because carbon and sulfur are solids, and oxygen, carbon dioxide, and sulfur dioxide are gases at ordinary temperatures and pressures, equations 1.1 and 1.2 become
C(s) ⫹ O2(g)
CO2(g)
S(s) ⫹ O2(g)
SO2(g)
We designate the physical states when this information is particularly important, but otherwise for simplicity we will omit them. In a correctly balanced chemical equation, some things must be equal, others need not be. Table 1.7 summarizes our discussion so far. Equation 1.1 describes the combustion of pure carbon in an ample supply of oxygen. But this is not always the case. If the oxygen supply is limited, CO may be one of the products. Let’s take the extreme case in which carbon monoxide is the sole product.
C ⫹ O2
CO (unbalanced equation)
This equation is not balanced because there are 2 oxygen atoms on the left but only 1 on the right. You might be tempted to balance the equation by simply adding an additional oxygen atom to the right side. But once we write the correct chemical formulas for the reactants and products, we cannot change them. We only can use whole-number coefficients
Table 1.7
Characteristics of Chemical Equations
Always Conserved identity of atoms in reactants 5 identity of atoms in products number of atoms of each element in reactants 5 number of atoms of each element in products mass of all reactants 5 mass of all products May Change number of molecules in reactants may differ from those in products physical states (s, l, or g) of reactants may differ from those in products
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(or occasionally fractional ones) in front of the given chemical formulas. In simple cases like this, the coefficients can be found by trial and error. If we place a 2 in front of CO, it signifies two molecules of carbon monoxide. This balances the oxygen atoms. C ⫹ O2
2 CO (still not balanced)
But now the carbon atoms do not balance. Fortunately, this is easily corrected by placing a 2 in front of the C on the left side of the equation. 2 C ⫹ O2
39
A subscript follows a chemical symbol, as in O2 or CO2. A coefficient precedes a symbol or a formula, as in 2 C or 2 CO.
[1.3]
2 CO (balanced equation)
By comparing equations 1.1 and 1.3, you can see that more O2 is required to produce CO2 from carbon than is needed to produce CO. This matches the conditions we stated for the formation of carbon monoxide; namely, that the supply of oxygen was limited. You may be surprised to learn the origin of the air pollutant nitrogen monoxide (also called nitric oxide). It comes from the nitrogen and oxygen found in the air! These two gases chemically combine in the presence of something very hot, such as an automobile engine or a forest fire. N2 ⫹ O2
high temperature
NO (unbalanced equation)
The equation is not balanced, as 2 oxygen atoms are on the left side, but only one is on the right. The same is true for nitrogen atoms. Placing a 2 in front of NO supplies 2 N and 2 O atoms on the right. The equation is now balanced. N2 ⫹ O2
high temperature
[1.4]
2 NO
high temperature
Your Turn 1.15
Chemical Equations
Balance these chemical equations and draw representations of all reactants and products, analogous to equation 1.4. In H2O and NO2, O and N are the central atoms, respectively. a. H2 ⫹ O2
H2O
b. N2 ⫹ O2
NO2
Note: Both H2O and NO2 are bent molecules. We will explain why in the next chapter. Answer a. 2 H2 ⫹ O2
2 H2O
Consider This 1.16
Advice from Grandma
A grandmother offered this advice to rid the garden of pesky caterpillars. “Hammer some iron nails about a foot up from the base of your trees, spacing them every 3 to 5 inches.” According to this grandmother, the iron nails convert the sugary tree sap (a compound (continued)
Nitrogen and oxygen both are diatomic molecules.
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Consider This 1.16
Advice from Grandma
(continued)
containing the elements carbon, hydrogen, and oxygen) into ammonia (NH3), a substance the caterpillars cannot stand. Comment on the accuracy of grandma’s chemistry (allowing for the possibility that the nails may still work, regardless of her explanation).
1.10 Look for other examples of burning fuels throughout Chapter 4.
|
Fire and Fuel: Air Quality and Burning Hydrocarbons
As we mentioned earlier, hydrocarbons are compounds of hydrogen and carbon. The hydrocarbons that we use today are primarily obtained from crude oil. Methane (CH4), the simplest hydrocarbon, is the primary component of natural gas. Both gasoline and kerosene are mixtures of many hydrocarbons. Given an ample supply of oxygen, hydrocarbon fuels burn completely. You may hear this called “complete combustion.” In essence, all of the carbon atoms in the hydrocarbon molecule combine with O2 molecules from the air to form CO2. Similarly, all the hydrogen atoms combine with O2 to form H2O. For example, here is the chemical equation for the complete combustion of methane. This equation is your first peek at why burning carbon-based fuels releases carbon dioxide into the atmosphere. CH4 ⫹ O2
When balancing chemical equations, it is fine to use fractional coefficients.
CO2 ⫹ H2O (unbalanced equation)
Note that O appears in both products: CO2 and H2O. To balance it, start with an element that appears in only one substance on each side of the arrow. In this case, both H and C qualify. No coefficients need to be changed for carbon, because both sides contain one C atom. Balance the H atoms by placing a 2 in front of the H2O. CH4 ⫹ O2
CO2 ⫹ 2 H2O (still not balanced)
Balance the oxygen atoms last. Four O atoms are on the right side and two O atoms are on the left, so we need 2 O2 to balance the equation. CH4 ⫹ 2 O2
CO2 ⫹ 2 H2O (balanced equation)
[1.5]
A nice feature of chemical equations is that counting the atoms on both sides of the arrow tells you if it is balanced. Here, the equation is balanced because each side has 1 C atom, 4 H atoms, and 4 O atoms. Most automobiles run on the complex mixture of hydrocarbons that we call gasoline. Octane, C8H18, is one of the pure substances in this mixture. With sufficient oxygen, octane burns to form carbon dioxide and water. 2 C8H18 ⫹ 25 O2
16 CO2 ⫹ 18 H2O
[1.6]
Both products travel from the engine out the exhaust pipe and into the air. Are these combustion products visible? Usually not. Water, in the form of water vapor, and carbon dioxide are both colorless gases. But if you happen to be outside on a winter day, the water vapor condenses to form clouds of steam or tiny ice crystals that you can see. Occasionally, the frozen vapor gets trapped in an inversion layer and forms an ice fog (Figure 1.15). With less oxygen, the hydrocarbon mixture we call gasoline burns incompletely (“incomplete combustion”). Water is still produced together with both CO2 and CO. The extreme case occurs when only carbon monoxide is formed, as is shown here for the incomplete combustion of octane. 2 C8H18 ⫹ 17 O2
16 CO ⫹ 18 H2O
[1.7]
Compare the coefficient of 17 for O2 in this equation with that of 25 for O2 in equation 1.6. Less oxygen is needed for incomplete combustion, as CO contains less oxygen than CO2.
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Figure 1.15 A winter ice fog in Fairbanks, Alaska.
Your Turn 1.17
Balancing Equations
Demonstrate that equations 1.6 and 1.7 are balanced by counting the number of atoms of each element on both sides of the arrow. Answer Equation 1.6 contains 16 C, 36 H, and 50 O on each side.
What actual mixture of products is formed when gasoline is burned in your car? This is not a simple question, as the products vary with the fuel, the engine, and its operating conditions. It is safe to say that gasoline burns primarily to form H2O and CO2. However, some CO and soot also are produced. The amounts of soot, CO, and CO2 that go out the tailpipe indicate how efficiently the car burns the fuel, which in turn indicates how well the engine is tuned. Some regions of the United States monitor auto emissions with a probe that detects CO (Figure 1.16). The CO concentrations in the exhaust are compared with established standards, for example, 1.20% in the state of Minnesota. If the vehicle fails the emissions test, it must be serviced.
Consider This 1.18
Auto Emissions Report
a. Figure 1.16 reports NOx emissions in grams per mile. NOx is a way to collectively represent the oxides of nitrogen. If x 5 1 and x 5 2, write the corresponding chemical formulas. Also give the chemical names. b. NO is the primary oxide of nitrogen emitted. What is the source of this compound? Hint: Revisit equation 1.4. c. The green line is missing on the CO2 graph, but present on the others. Explain. Answers a. NO, nitrogen monoxide and NO2, nitrogen dioxide c. In the year that this graph was produced, CO2 was not classified as an air pollutant in the United States. Accordingly, it has no green line indicating an acceptable range.
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Second-By-Second Emissions Report Hydrocarbons (grams per mile) MPH 6.0 60 5.0 4.0 40 3.0 2.0 20 1.0 0 0 0 20 40 60 80 100 120 140 160 180 200 220 240 SEC. CO (grams per mile) 80 60 40 20 0 0 20 40 60
MPH 60 40 20 80
NOX (grams per mile) 12.0 10.0 8.0 6.0 4.0 2.0 0 0
20
40
MPH 60 40 20 60
CO2 (grams per mile) 2000 1500 1000 500 0 0 20 40 60
CO2 emissions are measured but not yet regulated. Look for more about CO2 emissions in Chapter 3.
0 100 120 140 160 180 200 220 240 SEC.
80
0 100 120 140 160 180 200 220 240 SEC. MPH 60 40 20
80
0 100 120 140 160 180 200 220 240 SEC.
Figure 1.16 A U.S. auto emissions report. The blue line shows the change in engine speed; the red line shows the change in emissions. Any emissions below the green line are in the acceptable range.
1.11
| Air Pollutants: Direct Sources
In this section, we examine two major sources of air pollutants: motor vehicles and coal-fired plants that generate electricity. These sources directly emit SO2, CO, NO, and PM, and we will revisit each of these pollutants in turn. We also digress to discuss VOCs (volatile organic compounds), pollutants that are not regulated but are still intimately connected with the ones that are. We tackle ozone in the section that follows.
Your Turn 1.19
Tailpipe Gases
What comes out of the tailpipe of an automobile? Start your list now and build it as you work through this section. Hint: Some of the air that enters the engine also comes out the tailpipe.
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Sulfur dioxide emissions are linked to the coal that is burned to generate electric power. Although coal consists mostly of carbon, it may contain 1–3% sulfur together with small amounts of minerals. The sulfur burns to form SO2, and the minerals end up as fine ash particles. If not contained, the SO2 and ash go right up the smokestack. The hundreds of millions of tons of coal burned by the United States translates into millions of tons of waste in the air. As we will see in Chapter 6, the SO2 produced by burning coal can dissolve in the water droplets of clouds and fall to the ground as acid rain. The story does not end with SO2. Once in the air, sulfur dioxide can react with oxygen to form sulfur trioxide, SO3. 2 SO2 ⫹ O2
2 SO3
H2SO4
SO2 from the Mining Industry
Burning coal is not the only source of sulfur dioxide. As you saw in Your Turn 1.8, smelting is another. For example, silver and copper metal can be produced from their sulfide ores. Write the balanced chemical equations. a. Silver sulfide (Ag2S) is heated in air to produce silver and sulfur dioxide. b. Copper sulfide (CuS) is heated in air to produce copper and sulfur dioxide. Answer a. Ag2S ⫹ O2
Section 6.12 describes how sulfuric acid aerosols contribute to haze.
[1.9]
If inhaled, the droplets of the sulfuric acid aerosol are small enough to become trapped in the lung tissue and cause severe damage. The good news? Sulfur dioxide emissions in the United States are declining (see Figure 1.8). For example, in 1985, approximately 20 million tons of SO2 was emitted from the burning of coal. Today the value is closer to 9 million tons. This impressive decrease can be credited to the Clean Air Act of 1970 that mandated many reductions, including those from coal-fired electric power plants. More stringent regulations were established in the Clean Air Act Amendments and the Pollution Prevention Act of 1990. For example, gasoline and diesel fuel both once contained small amounts of sulfur, but the allowable amounts were drastically lowered in 1993 and in 2006, respectively. But continued progress has a price tag. Cleaning up the smaller older and dirty power plants will not come cheaply. Then again, allowing emissions to continue is costly in terms of human and environmental health.
Your Turn 1.20
See Section 4.6 for more about coal and its chemical composition.
[1.8]
Although normally quite slow, this reaction is faster in the presence of small ash particles. The particles also aid another process. If the humidity is high enough, they help condense water vapor into an aerosol of tiny water droplets. Aerosols are liquid and solid particles that remain suspended in the air rather than settling out. Smoke, such as from a campfire or a cigarette, is a familiar aerosol made up of tiny particles of solids and liquids. The aerosol of concern here is made up of tiny droplets of sulfuric acid, H2SO4. It forms because sulfur trioxide dissolves readily in water droplets to produce sulfuric acid. H2O ⫹ SO3
43
2 Ag ⫹ SO2
With more than 250 million vehicles (and over 300 million people), the United States has more vehicles per capita than any other nation. Do these vehicles emit
Look for more about the economic and societal costs of atmospheric SO2 in Chapter 6.
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sulfur dioxide? Fortunately, the answer is no, because cars have internal combustion engines primarily fueled by gasoline. We already mentioned that the combustion of hydrocarbons in gasoline produces—at best—carbon dioxide and water vapor (see equation 1.6). Because gasoline contains little or no sulfur, burning it produces little or no sulfur dioxide. Nonetheless, each tailpipe puffs out its share of air pollutants. The ubiquitous motor car adds to the atmospheric concentrations of carbon monoxide, volatile organic compounds, nitrogen oxides, and particulate matter. We discuss each of these in turn. Carbon monoxide pollution comes primarily from automobiles. But think in terms of all the tailpipes out there, not just those attached to cars. Some are attached to heavy trucks, SUVs, and the three m’s: motorcycles, minibikes, and mopeds. Others are on equipment such as farm tractors, bulldozers, and motor boats. The tailpipes attached to all gasoline and diesel engines emit carbon monoxide.
Your Turn 1.21
Other Tailpipes
Visit “Nonroad Engines, Equipment, and Vehicles” courtesy of the EPA. A link is provided at the textbook’s website. a. The text mentioned tractors, bulldozers, and boats. Name five other enginepowered machines or vehicles that do not run on roads. b. Select a machine or vehicle of interest to you. How are emissions from its engine being reduced? What is the time scale for the reduction?
Wildfires contribute additional CO, increasing emissions as much as 10% each year.
Even though the number of cars has risen, a dramatic reduction in CO emissions has occurred. Based on measurements by the EPA at over 250 sites in the United States, since 1980 the average CO concentration has decreased almost 60% (see Figure 1.8). If wildfires are excluded, today’s levels are the lowest reported in three decades. The decrease is due to several factors, including improved engine design, computerized sensors that better adjust the fuel–oxygen mixture, and most importantly, the requirement that all cars manufactured since the mid-1970s have catalytic converters (Figure 1.17). Catalytic converters reduce carbon monoxide in the exhaust stream by catalyzing the combustion of CO to CO2. They also lower NOx emissions by catalyzing the conversion of nitrogen oxides back to N2 and O2, the two atmospheric gases that formed them. In general, a catalyst is a chemical substance that participates in a chemical reaction and influences its rate without itself undergoing permanent change. Catalytic converters typically use metals such as platinum and rhodium as catalysts.
Catalytic converter (a)
(b)
Figure 1.17 (a) Location of catalytic converter in car. (b) Cutaway view of a catalytic converter. Metals such as platinum and rhodium serve as catalysts and form a coating on the surface of ceramic beads.
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Cars not only emit carbon in the form of carbon monoxide but also in the form of unburned and partially burned hydrocarbons. This leads us to the topic of VOCs, volatile organic compounds. A volatile substance readily passes into the vapor phase; that is, it evaporates easily. Gasoline and nail polish remover are both volatile. If you were to spill either of these, the puddle would soon evaporate. When you apply varnish to a surface, you can smell the volatile compounds that evaporate with each brush stroke. An organic compound always contains carbon, almost always contains hydrogen, and may contain elements such as oxygen and nitrogen. Organic compounds include methane and octane, hydrocarbons mentioned earlier. They also include alcohol and sugar, compounds that contain oxygen in addition to carbon and hydrogen. Accordingly, volatile organic compounds (VOCs) are carbon-containing compounds that pass easily into the vapor phase. They originate from a variety of sources. For example, you can smell naturally occurring VOCs in a spruce or pine forest. VOCs from tailpipes are not so pleasant, as they are vapors of incompletely burned gasoline molecules or fragments of these molecules. The exhaust gas still contains oxygen, as not all of it is consumed in the engine. Catalytic converters utilize this oxygen to burn VOCs to form carbon dioxide and water. Section 1.12 describes the connection between VOCs and ozone formation. But right now, we want to connect VOCs with the formation of NO2. Nitrogen monoxide and nitrogen dioxide are collectively known as NOx, as mentioned in Consider This 1.18. NO2 is brown in color, giving smog its characteristic brownish tinge. Recall that N2 and O2 combine to produce NO which is a colorless gas (see equation 1.4). But what is the origin of NO2? Here is a balanced equation that appears to be a likely candidate. 2 NO ⫹ O2
2 NO2
45
Look for more about organic compounds in Chapter 4.
[1.10]
However, this is not what actually occurs. Instead, NO2 is formed by other pathways that are more complex. Here is the one that predominates in urban settings where you are likely to find NO. In some cities, this actually lowers the ozone concentrations along highways congested with vehicles emitting NO. NO ⫹ O3
NO2 ⫹ O2
[1.11]
To further complicate things, on a sunny day, some of the NO2 converts back into NO, as we will see in the next section. Again, this is why people refer to NOx, rather than to either NO or NO2. The conversion of NO to NO2 connects to the breakdown of VOCs in the air. A new player is involved, the reactive hydroxyl radical (?OH). This reactive species is present in tiny amounts in air, polluted or otherwise. VOC ⫹ OH A ⫹ O2 A⬘ ⫹ NO
A A⬘ A⬙ ⫹ NO2
[1.12]
Here, A, A9, and A0 represent reactive molecules that can form in the air from ?OH and VOCs. The bottom line? Atmospheric chemistry is complex and involves many players. You have met some of them, including NO, NO2, O2, O3, VOCs, and ?OH. The United States has had limited success in curbing NOx emissions. In turn, as we’ll see in the next section, this means limited success in curbing ozone. Nonetheless, given the increasing number of vehicles, any decrease in NOx is impressive. Despite early claims from the auto industry that it would be impossible (or too costly) to meet new emissions standards, the industry is curbing emissions by improving catalytic converters, engine designs, and gasoline formulations.
The dot in (?OH) indicates an unpaired electron. In Chapter 2, you will encounter other reactive species with unpaired electrons.
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Chapter 8 discusses fuel cells and other alternatives to gasolinepowered vehicles.
Consider This 1.22
Forget Road Rage
Burning less gasoline equates to fewer tailpipe emissions. Which driving practices conserve fuel? Which practices expend it more than necessary? Think about the behavior of motorists on highways, city streets, and in parking lots. For each of these venues, list at least three ways that drivers could burn less gasoline. Hint: Consider how you accelerate, coast, idle, brake, and park. Answer Some possibilities relate to parking. For example, if you are able-bodied, take a spot further away and walk, rather than cruising around to find a closer spot. Find a parking space that you can pull through. This way, when you exit you don’t need to back and turn, thus conserving fuel.
Ammonia (NH3) is a colorless gas with a pungent odor. It is condensed to the liquid phase and applied to soil as a fertilizer. Look for more about ammonia in Chapter 6.
Particulate matter comes in a range of sizes, but only the tiny particles (PM10 and PM2.5) are regulated as pollutants. This size can penetrate deeply into your lungs, pass into your bloodstream, and inflame your cardiovascular system. In terms of regulation, particles are the new pollutant on the block. Data collection in the United States for PM10 and PM2.5 started in 1990 and 1999, respectively (see Figure 1.8). In 2006, the daily air quality standard for PM2.5 was lowered from 65 to 35 μg/m3 because these particles proved to be more hazardous than originally thought. Particulate matter has many different sources. In the summer, wildfires may raise the concentration of particulate matter to a hazardous level. In the winter, wood stoves may produce exactly the same effect. At any time of the year in almost any urban environment, older diesel engines on trucks and buses emit clouds of black smoke. Diesel engines on tractors similarly can pollute. Construction sites, mining operations, and the unpaved roads that serve them also loft tiny particles of dust and dirt into the atmosphere. Particulate matter can even form right in the atmosphere. For example, the compound ammonia, used in agriculture, is a major player in forming ammonium sulfate and ammonium nitrate in the air, both PM2.5. Given all these sources, particulate matter has proven a tough pollutant to control. Even so, in 2007 the EPA reported a decrease of 14% in the annual PM2.5 concentrations from 2000 to 2006. However, 10% of the sites monitored still showed an increase in particle pollution. Again, what you breathe very much depends on where you live.
Consider This 1.23
Particles Where You Live
Here is a map of the continental U.S. that shows PM2.5 data for August 30, 2008. a. In terms of air quality, what do the green, yellow, and orange colors indicate? b. Which groups of people are most sensitive to particulate matter? c. Visit State of the Air, a website posted by the American Lung Association. How many days a year does your state have “orange days” and “red days” for particle pollution? What is the difference?
August 30, 2008 12:00 am EDT
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1.12
47
| Ozone: A Secondary Pollutant
Ozone definitely is a bad actor in the troposphere. Even at very low concentrations, it reduces lung function in healthy people who are exercising outdoors. Ozone also damages crops and the leaves of trees. But ozone does not come out of a tailpipe and is not produced when coal is burned. How is it produced? Before we fill you in on the details, do this activity.
Consider This 1.24
Today, background tropospheric levels for ozone are about 40 parts per billion (ppb). In pre-industrial times, the level was about 10 ppb.
Ozone Around the Clock
Ozone concentrations vary during the day, as shown in Figure 1.18. a. Near which cities is the air hazardous to one or more groups? Hint: Refer back to the color-coded AQI (see Table 1.4). b. At about what time does the ozone level peak? c. Can moderate levels (shown in yellow) of ozone exist in the absence of sunlight? Assume sunrise occurs around 6 AM and sunset at about 8 PM. Answer c. In the absence of sunlight, ozone does not persist for long. After sundown, the ozone levels drop.
The previous activity raises several related questions. Why is ozone more prevalent in some areas than others? What role does sunlight play in ozone production? We now address these.
Eureka
Eureka
Redding Reno
Ukiah
San Francisco
Eureka
Redding Reno
Ukiah
Sacramento Stockton San Jose
San Francisco Las Vegas
Fresno
Bakersfield
San Diego
6
Bakersfield
Barstow Santa Barbara Los Angeles
Palm Springs
Barstow Santa Barbara Los Angeles
Palm Springs
San Diego
San Diego
10
AM
Eureka
Redding
Ukiah
Ukiah
San Francisco
noon
AM
Eureka
Redding
Reno Sacramento Stockton
San Francisco
Reno Sacramento Stockton San Jose
San Jose Las Vegas
Fresno
Bakersfield Barstow
Barstow Palm Springs
Santa Barbara Los Angeles San Diego
San Diego
4
PM
Figure 1.18 Ozone level maps for a summer day in July 2006 in California.
Las Vegas
Fresno
Bakersfield
Santa Barbara Los Angeles
Las Vegas
Fresno
Bakersfield
Barstow Santa Barbara Los Angeles
Sacramento Stockton San Jose
San Jose Las Vegas
Fresno
Reno
Ukiah
Sacramento Stockton
San Francisco
Redding
10
PM
Palm Springs
Palm Springs
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Recall from the previous section that ·OH is the hydroxyl radical.
Unlike the pollutants described in the previous section, ozone is a secondary pollutant. It is produced from chemical reactions involving one or more other pollutants. For ozone, these other pollutants are VOCs and NO2. Recall from the previous section that NO, rather than NO2, comes directly out of a tailpipe (or a smokestack). But over time and in the presence of VOCs and ?OH, NO in the atmosphere is converted to NO2. Nitrogen dioxide meets several fates in the atmosphere. The one of most interest to us occurs when the Sun is high in the sky. The energy provided by sunlight splits one of the bonds in the NO2 molecule: NO2
sunlight
NO ⫹ O
[1.13]
The oxygen atoms produced then can react with oxygen molecules to produce ozone. O ⫹ O2
“Good” ozone is in the stratosphere. “Bad” ozone is in the troposphere.
O3
[1.14]
This explains why ozone formation requires sunlight. Sunlight splits NO2 to release O atoms. These in turn react with O2 to form O3. Thus once the Sun goes down, the ozone concentrations drop off sharply, as you can see in Figure 1.18. What happened to the ozone? In just a matter of hours, the ozone molecules react with many things, including animal and plant tissue. Note that equation 1.14 contains three different forms of elemental oxygen: O, O2, and O3. All three are found in nature, but O2 is the least reactive and by far the most abundant, constituting about one fifth of the air we breathe. Our atmosphere naturally contains tiny amounts of protective ozone up in the stratosphere. Oxygen atoms also exist in our upper atmosphere and are even more reactive than ozone.
Consider This 1.25
O3 Summary
Summarize what you have learned about ozone formation by developing your own way to arrange these chemicals sequentially and in relation to one another: O, O2, O3, VOCs, NO, NO2. Chemicals may appear as many times as you like. You also may wish to include sunlight.
Because sunlight is involved in ozone formation, the concentration of ground-level ozone varies with weather, season, and latitude. High levels of O3 are much more likely to occur on long sunny summer days, especially in congested urban areas. Stagnant air also favors the buildup of air pollution. For example, revisit the air quality data for cities shown in Table 1.3. Ozone was usually the culprit responsible for pollution in cities with sunny summer days. In contrast, windy and rainy cities have lower levels of ozone.
Consider This 1.26
Ozone and You
The AIRNOW website, courtesy of the EPA, provides a wealth of information about ground-level ozone levels in the United States. a. Let’s say that the ozone level is “orange,” actually a common occurrence in many U.S. cities during the summer months. Does air of this quality affect you if you have no health concerns, but are actively exercising out of doors? Hint: Check the link to AIRNOW provided at the textbook’s website. b. Again assume you are active out of doors. How does the air quality in your state compare with others?
Canada also publishes daily ozone maps (Figure 1.19). Some of Canada’s polluted air originates in the United States, blown northeastward from population centers in Ohio, Pennsylvania, and New York. Pollution knows no boundaries! This is an example of the tragedy of the commons. The tragedy arises when a resource is common to all and used by many, but has no one in particular responsible
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for it. As a result, the resource may be destroyed by overuse to the detriment of all who use it. For example, we cannot lay individual claim to the air; it belongs to all of us. If the air we breathe has waste dumped into it, this leads to an unhealthy situation for everyone. Individuals whose activities have little or no effect on the air still suffer the same consequences as those who pollute. The costs are shared by all. In later chapters, we will see other examples of the tragedy of the commons that relate to water, energy, and food. Air pollution, once primarily a local concern, is now a serious international issue. Many cities worldwide have high ozone levels. Couple vehicles with a sunny location anywhere on the planet, and you are likely to find unacceptable levels of ozone. Some places, however, are worse than others. London with its cool foggy days has low ozone levels. In contrast, ozone is a serious problem in Mexico City. Ozone attacks rubber and so it damages the tires of the vehicles that led to its production in the first place. Should you park your car indoors in the garage to minimize possible rubber damage? In fact, should you park yourself indoors if the levels of ozone outside are unhealthy? The next section speaks to the quality of indoor air.
Garrett Hardin is credited with coining the term “tragedy of the commons.” In an article published in 1968, he pointed out how individuals using a common resource may destroy it such that ultimately nobody can use it.
Figure 1.19 Tropospheric ozone map for September 7, 2007. Source: Environment Canada.
1.13
| The Inside Story of Air Quality
In the Wizard of Oz, Dorothy hugged her dog Toto and exclaimed, “There’s no place like home!” Of course she was right, but when it comes to air quality, your home may not always be the best place to be. Indoors, the levels of air pollution may far exceed those outside. Given that most of us sleep, work, study, and play indoors, we should learn about the air in the place we call home. Indoor air may contain up to a thousand substances at low levels. If you are in a room where somebody is smoking, add another thousand or so. Indoor air contains some familiar culprits: VOCs, NO, NO2, SO2, CO, ozone, and PM. These pollutants are present either because they came in with the outside air or because they were generated right inside your dwelling. Let’s begin our discussion with the question posed in the previous section. Should you move indoors to escape the ozone present outside? In general, if a pollutant is
Some copy machines and air cleaners generate ozone, which may increase indoor ozone levels.
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Ponderosa pine bark emits compounds with a scent reminiscent of butterscotch.
Sick building syndrome has many causes. One way or another, most relate to air quality. The source of the bad air may come both from inside the building, outside the building, or both.
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highly reactive, it does not persist long enough to be transported indoors. Thus, for highly reactive molecules such as O3, NO2, and SO2, you expect lower levels indoors. Indeed, this is the case, as indoor air is typically 10–30% lower in ozone concentration than outdoor air. Similarly, sulfur dioxide and nitrogen dioxide levels are lower indoors, although the decrease is not as dramatic as that for ozone. Carbon monoxide is a different story. As a relatively unreactive pollutant, CO has a long enough atmospheric lifetime to move freely in and out of buildings through doors, windows, or a ventilation system. The same is true for some VOCs, but not for the more reactive ones such as those that give pine forests their scent. If you want to inhale the delicious volatile compounds emitted from the bark of Ponderosa pines, it is best to remain outside near the trees. Some pollutants are trapped by the filters in the heating or cooling system of the building. For example, many air-handling systems contain filters that remove largersized particulate matter and pollen. As a result, people who suffer from seasonal allergies can find relief indoors. Similarly, those near a wildfire can go inside to escape some of the irritating smoke particles. However, gas molecules such as O3, CO, NO2, and SO2 are not trapped by the filters used in most ventilation systems. These days, many buildings are constructed with an eye to increasing their energy efficiency. This is a win–win, lowering your heating bills and lowering the amounts of pollutants generated in producing the heat. But there can be a down side. A building that is air-tight with a limited intake of fresh air may have unhealthy levels of indoor air pollutants. Therefore what appeared initially to be a benefit (better energy efficiency) can turn into a higher risk (increased pollutant levels). In some cases, poor ventilation can cause indoor pollutants to reach hazardous levels, creating a condition known as “sick building syndrome.” Clearly this is an undesirable outcome. Today, architects and builders are finding ways to make buildings more energy-efficient while maintaining a good air exchange. Even with good ventilation, indoor activities can compromise air quality. For example, tobacco smoke is a serious indoor air pollutant, containing over a thousand chemical substances. Nicotine is one that you may recognize; others include benzene and formaldehyde. Taken as a whole, tobacco smoke is carcinogenic, meaning capable of causing cancer. Combustion of carbon-containing fuels also can generate carbon monoxide and nitrogen oxides. For example, carbon monoxide from cigarette smoking in bars can reach 50 ppm, a value well within the unhealthy range. In cigarette smoke, NO2 levels can exceed 50 ppb. Fortunately, smokers take puffs rather than constantly breathing cigarette smoke.
Your Turn 1.27
Indoor Cigar Party
In 2007, student researchers attended a cigar party and trade show in Times Square, New York. They carried concealed detectors and found levels of 1193 μg of particulate matter per cubic meter of air inside the ballroom. a. The news article did not report whether the particulate matter was PM2.5 or PM10. For either, does the value they measured exceed the air quality standards in the United States? Hint: See Table 1.2. b. Assume that the students were measuring PM2.5. What are the health implications?
People also burn candles, perhaps to soften the lighting or set a mood. However, candles deplete the oxygen in a room. They also can produce soot, carbon monoxide, and VOCs. Similarly, people may burn incense in their homes for one reason or another. Atmospheric scientist Stephen Weber, a researcher who studied incense burning in churches in Europe, found that “the pollutants in smoke from incense and candles may be more toxic than fine-particle pollution from sources such as vehicle engines.” Burning candles or incense can generate fumes more rapidly than these can be removed by your ventilation system or by the breezes that pass through open windows. The next activity gives you the opportunity to further investigate sources of indoor pollutants.
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Your Turn 1.28
Indoor Activities
Name 10 activities that add pollutants or VOCs to indoor air. To get you started, two are pictured in Figure 1.20. Remember that some pollutants have no detectable odors.
As Figure 1.20 suggests, paints and varnishes are a source of VOCs. Your nose alerts you to these while you paint; so does your head if it starts aching. Although the amount of VOCs released per volume of paint applied varies widely, all oil-based paints and varnishes rank higher than water-based ones. Check out the VOC emission values printed right on the paint can. These range from zero for low-VOC paints to more than 600 grams VOCs per liter for some outdoor oil paints. You would be fortunate to find an oil paint with less than 350 grams VOCs per liter. Consumers C Cons Co ons nsum umer um erss tod er ttoday oday can purchase high-quality paints that are nontoxic and odorfree. For example, check out the label on the paint can shown in Figure 1.21. This “zero-VOC” water-based paint emits less than 5 grams of VOCs per liter applied. This paint has earned a “Green Seal” certification, meaning that not only does it emit less than 50 grams of VOCs per liter, but also that it contains no toxic metals such as lead, mercury, or cadmium. Low- or zero-VOC paints are equally important to use outside. The paint applied to buildings, bridges, and railings in the United States once added up to over 7 million pounds of VOCs per year. In 2005, the value was reported at less than 4 million pounds, thanks to changes in the formulation of paints. Before we can explain how volatile compounds have been removed from paint, we first need to address why paints contained them in the first place. Some VOCs are additives that evaporate as the paint dries. For example, you may recognize the two antifreeze additives listed in Table 1.8. Antifreeze (“glycols”) allows the folks who live up north to store paint in their basements without worrying that it will be damaged by repeated freezing and thawing. Antifreeze also allows paint to be applied in colder weather and provides longer “open time,” so that the paint does not dry out too quickly. Coalescents are another additive used in latex paints. Coalescents are chemicals added to soften the latex particles in paint so that these particles spread to form a continuous film of uniform thickness. After all, you want your paint to brush on evenly! As the paint dries and hardens, the coalescents evaporate into the air. Between 2 and
(a)
Figure 1.20 Activities that can pollute indoor air.
(b)
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Figure 1.21 All ingredients in YOLO paints are selected as “zero VOC.”
Look for more about the oils found in plants in Chapter 11.
3 gallons of volatile coalescents are used for every 100 gallons of paint. But do the math. In the United States this additive translates to over 100 million pounds of coalescents emitted each year, and approximately 3 times this amount worldwide. Oil-based paints contain just that, oils. Most are oils derived from plants, such as linseed oil. These oils slowly react with the oxygen in the air and over time release a host of volatile compounds to the air. Oil paints also may contain solvents (thinners) that evaporate as the paint dries.
Skeptical Chemist 1.29
Varnish Fumes
A can of clear satin floor varnish claims a maximum of 450 grams/liter of VOCs. A consumer group computes this to just under 4 pounds of VOCs emitted per gallon as the varnish is applied. Did this group do the math correctly?
Table 1.8
VOCs Emitted by Some Paints
Glycols (antifreezes) ethylene glycol propylene glycol Coalescents (used in latex paints) 2,2,4-trimethyl-1,3-pentanediol monoisobutyrate (trade name: Texanol) Hazardous air pollutants (solvents and preservatives) benzene formaldehyde ethylbenzene methylene chloride vinyl chloride
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Stricter government regulations on VOC emissions have prompted paint manufacturers to devise new formulations for latex paints. In 2005, the Archer Daniels Midland Company won a Presidential Green Chemistry Challenge award for developing nonvolatile coalescents. The coalescents developed by the company react with the oxygen in the air, enabling them to chemically bond to the latex. Thus, the coalescents become part of the paint film rather than evaporating into the atmosphere. Another advantage of these new coalescents is that they are produced from vegetable oils (a renewable resource) as opposed to crude oil (nonrenewable). Their production also creates less waste and requires less energy, a plus for both environmental and cost savings. There is no loss in quality, as the paints formulated with vegetable oil-based coalescents meet or exceed the performance of traditional paints. They have lower odors, better scrub resistance, and better opacity. Do these environmental, economic, and societal benefits sound reminiscent of the Triple Bottom Line? Again, this is the heart of sustainability. If you have ever been to an auto body shop, you most likely have smelled the odors from the VOCs in the paints. More stringent laws and regulations in many countries have forced manufacturers to reformulate automotive primers (undercoats) and finishes. Most traditional primers are made by mixing two components that have a limited shelf life. The primer is then applied and cured in an oven that requires large amounts of energy to heat. BASF Corporation won a 2005 Presidential Green Chemistry Challenge award for the development of a one-component primer that reduces VOCs by 50% and that is rapidly cured in sunlight or with a UV-A lamp. This greatly reduces the time required to make and cure the primer. Here again, the reduced costs associated with less waste, reduced energy consumption, and greater throughput equate to an improved Triple Bottom Line.
Your Turn 1.30
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UV-A lamps are longer-wave ultraviolet light, similar to tanning lamps. Look for more about UV light in Chapter 2.
Bring Home the Green
Revisit the principles of Green Chemistry listed on the inside front cover of this book. Which of the principles are met by the new coalescents developed by the Archer Daniels Midland Company? By the new undercoat developed by BASF? Prepare a list for each company.
We end our discussion of indoor air quality by turning to radon, a noble gas (Group 8A) that we mentioned earlier. Radon is a special case of indoor air pollution. It occurs naturally in tiny amounts and usually is no problem. But it may reach hazardous levels in basements, mines, and caves. Like all noble gases, radon is colorless, odorless, tasteless, and chemically unreactive. But unlike the others, it is radioactive. Radon is generated in the decay series of uranium, another naturally occurring radioactive element. Because uranium occurs at a concentration of about 4 ppm in the rocks of our planet, radon is ubiquitous. Depending on how your apartment or dorm is constructed, the radon produced from uranium-containing rocks may fi nd entry into the basement. Radon causes lung cancer and is the second leading cause behind tobacco smoke. As is the case for other pollutants, the threshold for danger can be estimated but is not precisely known. Radon test kits, such as the one shown in Figure 1.22 , are used to measure the radon concentration in living spaces. Indoors or out, we need to breathe healthy air. And with each breath, we inhale a truly prodigious number of molecules and Figure 1.22 atoms. We end this chapter by revisiting these molecules and A home radon test kit. atoms.
Look for more about uranium and its natural decay series in Chapter 7.
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1.14 Check Figures Alive! for more about your breathing and lung capacity.
Figure 1.23 A spirometer is used for measuring an individual’s breathing capacity.
|
Back to the Breath—at the Molecular Level
The maximum concentrations of pollutants allowed by the air quality standards seem very small (see Table 1.2). Indeed, an exposure to 9 ppm CO is a tiny amount! But even this low concentration of CO contains a staggering number of carbon monoxide molecules. This seeming contradiction is a consequence of the minuscule mass of molecules. Recall Consider This 1.2: Take a Breath. If you are an averagesized adult, the capacity of your lungs is between 5 and 6 L. You do not empty your lungs each time you take a breath. Rather, as you are at rest reading this, you are inhaling about 500 milliliters of air or approximately half a quart with each breath. Accurately measuring the volume of air you inhale and exhale can be done with the help of a spirometer (Figure 1.23). Determining the number of molecules and atoms in this volume of air is a harder task, but it can be done. From experiments, we know that a typical breath of 500 mL contains about 2 3 1022 molecules and atoms. Remember that air is primarily N2 and O2 molecules together with a small amount of Ar atoms and a varying amount of H2O molecules (humidity). Using this number of molecules and atoms in the air (2 3 1022), we now can calculate the number of CO molecules in the breath you just inhaled. We assume the breath contained 2 3 1022 molecules and atoms, and that the CO concentration in the air was at the air quality standard of 9 ppm. Thus, out of every million (1 3 106) molecules and atoms in the air, nine will be CO molecules. To compute the number of CO molecules in a breath, multiply the total number of molecules and atoms in the air by the fraction that are CO molecules. # of CO molecules 1 breath of air
⫽ ⫽ ⫽
2 ⫻ 1022 molecules and atoms in air 1 breath of air
2 ⫻ 9 ⫻ 10
CO molecules
1 ⫻ 10
breath of air
22
18 1
⫻
6
⫻
9 CO molecules 1 ⫻ 10 molecules and atoms in air 6
1022 CO molecules 106 breath of air
In writing this out, we carefully retain the units on the numbers. Not only does this remind us of the physical entities involved, but also it guides us in setting up the problem correctly. The units “molecules and atoms in the air” cancel, and we are left with the unit we want: CO molecules per breath of air. However, we need to divide 1022 by 106 to determine a final answer. To divide powers of 10, simply subtract the exponents. In this case, 1022 106
⫽ 10(22⫺6) ⫽ 1016
Thus, a breath contains 18 3 1016 CO molecules. The preceding answer is mathematically correct, but in scientific notation it is customary to have only one digit to the left of the decimal point. Here we have two: 1 and 8. Therefore, our last step is to rewrite 18 3 1016 as 1.8 3 1017. We can make this conversion because 18 5 1.8 3 10, which is the same as 1.8 3 101. We add exponents to multiply powers of 10. Thus, 18 3 1016 CO molecules equals (1.8 3 101) 3 1016 CO molecules, which equals 1.8 3 1017 CO molecules in that last breath you inhaled. If this use of exponents is coming at you a little too fast, consult Appendix 2. It may sound surprising, but it is more accurate to round off the answer and report it as 2 3 1017 CO molecules. Certainly 1.8 3 1017 looks more accurate, but the data that went into our calculation were not very exact. The breath contains about 2 3 1022
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molecules, but it might be 1.6 3 1022, 2.3 3 1022, or some other number. We say that 2 3 1022 expresses a physically based property to “one significant figure.” A significant figure is a digit that is included (or excluded) to correctly represent the accuracy with which an experimental quantity is known. Only one digit, the 2 from the value 2.3 is used, and so 2 3 1022 has only one significant figure. Accordingly, the number of molecules in the breath is closer to 2 3 1022 than to 1 3 1022 or to 3 3 1022, but anything beyond this we cannot say with certainty. Similarly, the concentration of carbon monoxide is known to only one significant figure, 9 ppm. That 2 3 9 equals 18 is certainly correct mathematically, but our question about CO is based on physical data. The answer, 1.8 3 1017 CO molecules, includes two significant figures: the 1 and the 8. Two significant figures imply a level of knowledge that is not justified. The accuracy of a calculation is limited by the least accurate piece of data that goes into it. In this case, both the concentration of CO and the number of molecules and atoms in the breath were known only to one significant figure (9 and 2, respectively). Thus two significant figures in the answer are unjustified. The common-sense rule is that you cannot improve the accuracy of experimental measurements by manipulations like multiplying and dividing. Therefore, the answer must also contain only one significant figure and is 2 3 1017.
Your Turn 1.31
Ozone Molecules
The local news reports that today’s ground-level ozone readings are at the unacceptable level, 0.12 ppm. How many ozone molecules do you inhale in each breath? Assume that one breath contains 2 3 1022 molecules and atoms. Answer Start with the number of molecules and atoms in a breath. If the ozone concentration is 0.12 ppm, this gives the ratio 0.12 O3 molecules per 106 molecules and atoms in air. 0.12 O3 molecules 2 ⫻ 1022 molecules and atoms in air ⫻ 1 breath of air 1 ⫻ 106 molecules and atoms in air ⫽ 2.4 ⫻ 1015 O3 molecules / breath ⫽ 2 ⫻ 1015 O3 molecules / breath (to one significant figure)
You may well question the significance of all of this talk about significant figures. It has been observed that “figures don’t lie, but liars can figure.” Numbers often lend an air of authenticity to newspaper or television stories, so popular press accounts are full of numbers. Some are meaningful; others are not. Informed citizens can discriminate one from the other. For example, the assertion that the concentration of carbon dioxide in the atmosphere is 390.2537 ppm should be taken with a rather large grain of sodium chloride (salt). Values such as 390 ppm or 390.6 ppm (three or four significant figures) better represent what we actually can measure; any assertion with seven significant figures simply is not valid.
Your Turn 1.32
CO Monitors
Carbon monoxide monitors are available for homes and businesses. Figure 1.24 shows a convenient handheld CO detector that reads 35 ppm. a. Would it be more helpful to have a meter that read 35.0388217 ppm? Explain. b. Would 35.0388217 ppm be more valid? Explain. Answer a. No, it wouldn’t be more helpful. The issue is whether the concentration of CO exceeds a certain value, such as 9 ppm over an 8-hour period or 35 ppm over a 1-hour period. The extra decimal places are of no use.
Figure 1.24 Carbon monoxide meter showing 35 ppm.
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Absence of evidence is not the same as evidence of absence. The substance may be present but in undetectable amounts.
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Recall that we started with the concentration of CO in an air sample of 9 ppm. Even so, the number of CO molecules in a breath is enormous, about 2 3 1017. From these numbers, you can see that it is impossible to completely remove all the CO molecules from the air. “Zero pollutants” is an unattainable goal. At present, our most sensitive methods of chemical analysis are capable of detecting one target molecule out of a trillion. One part per trillion is analogous to moving 6 inches in the 93 million-mile trip to the Sun, a single second in 320 centuries, or a pinch of salt in 10,000 tons of potato chips. Again, air always has trace levels of contaminants, ones that we cannot detect. A breath of air contains molecules of hundreds, perhaps thousands of different compounds, most in minuscule concentrations. Their origin could be either natural or related to human activity. As with all chemicals, “natural” is not necessarily good, and “human-made” is not necessarily bad. As you learned earlier, exposure and toxicity are what matter.
Your Turn 1.33
CO Molecules in Perspective
To help you comprehend the magnitude of the 2 3 1017 CO molecules in one breath, assume that they were equally distributed among the 6.9 billion (6.9 3 109) inhabitants of the Earth. Calculate each person’s share of the 2 3 1017 CO molecules you just inhaled. Answer You are trying to distribute the huge number of molecules in a breath among all the human inhabitants of the Earth. This can be found by dividing the total number of CO molecules by the total number of humans: 17 Each person’s share is 2 ⫻ 10 CO molecules 6.9 ⫻ 109 people
Thus, to one significant figure, each person’s share is 3 3 107 (or 30,000,000) molecules of CO.
In addition to being extremely small, the molecules and atoms you breathe possess other remarkable characteristics. They are in constant motion. At room temperature and pressure, a nitrogen molecule travels at about 1000 feet per second and experiences approximately 400 billion collisions with other molecules in that time interval. Nevertheless, relatively speaking, the molecules are quite far apart. The actual volume of the extremely tiny molecules making up the air is only about 1/1000 of the total volume of the gas. If the particles in your half-liter breath were squeezed together, their volume would be about 0.5 mL, less than a quarter teaspoon. Sometimes people mistakenly think that air is empty space. It is 99.9% empty space, but the matter it contains is literally a matter of life and death! Moreover, it is matter that we continuously exchange with other living things. The carbon dioxide we exhale is used by plants to make the food we eat. The oxygen that plants release is essential for our existence. Our lives are linked by the elusive medium of air. With every breath, we exchange millions of molecules with one another. As you read this, your lungs contain 4 3 1019 molecules that have been previously breathed by other human beings, and 6 3 108 molecules that have been breathed by some particular person, say Julius Caesar, Mahatma Gandhi, or Joan of Arc. In fact, the odds are very good that right now your lungs contain one molecule that was in Caesar’s last breath. The consequences are breathtaking!
Skeptical Chemist 1.34
Caesar’s Last Breath
We just claimed that your lungs currently contain one molecule that was in Caesar’s last breath. That assertion is based on some assumptions and a calculation. Are these assumptions reasonable? We are not asking you to reproduce the calculation, but rather to identify some of the assumptions and arguments we might have used. Hint: The calculation assumes that all of the molecules in Caesar’s last breath have been uniformly distributed throughout the atmosphere.
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Consider This 1.35
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Air Quality Today
The addition of waste to our atmosphere has not occurred overnight. Rather, air pollution has been a growing concern since at least the time of the Industrial Revolution. Why have nations and the larger world community become more concerned about air quality? Identify at least four factors that have brought air quality to the attention of citizens and voters.
CONCLUSION The air we breathe affects both our health and the health of the planet. Our atmosphere contains the essentials for life, including two elements (oxygen and nitrogen) and two compounds (water and carbon dioxide). Our very existence on this planet depends on having a large supply of relatively clean, unpolluted air. But the air you breathe may be polluted with carbon monoxide, ozone, sulfur dioxide, and the oxides of nitrogen. Polluted air is more common in large cities, the very places where most people now live. Emergency room visits correlate with bad air quality. So do shortness of breath, scratchy throats, and stinging eyes. The pollutants that cause us harm are, for the most part, relatively simple chemical substances. They largely are produced as consequences of our dependence on coal for electricity production in power plants, gasoline in internal combustion engines, and the fuels we burn to heat and cook. Over O Ov ver ver er tthe he ppas he past astt 30 as 30 yyea years, ears ea rs,, go rs gove governmental v rn ve rnme ment ntal all rreg regulations, egul ulat atio ions ns,, in indu industry duust stry ry iinitiatives, niti ni tiat ativ i es, and modern technology tech te chno ch nolo logyy hhave avee reduced av redu re duce c d pollutant ce poll po llut utan antt levels. leve le vels ls.. Both Both catalytic ccat atal alyt ytic ic converters con o ve v rters on cars and emissions controls on smokestacks have been important players. But it makes more sense not to generate “people fumes” in the first place. Here is where green chemistry plays an important role. By designing new processes that do not produce air pollutants, we do not later have to clean them up. Indoors or out, the oxygen-laden air we breathe is very close to the surface of the Earth. However, the Earth’s atmosphere extends upward for considerable distance and contains other gases that also are essential for life on this planet. Chapters 2 and 3 will describe two of these: stratospheric ozone and carbon dioxide. We will see that our human footprints and “air prints” on planet Earth connect in surprising ways to both of these gases.
Recall the first green chemistry principle listed on the inside cover: “It is better to prevent waste than to treat or clean up waste after it is formed.”
Chapter Summary The numbers in parentheses indicate the sections in which the topics are introduced. Having studied this chapter, you should be able to: ■ Explain the connection between your health and what you breathe (entire chapter) ■ Describe air in terms of its major components, their relative amounts, and the local and regional variations in the composition of air (1.1, 1.5) ■ List the major air pollutants and describe the health effects of each (entire chapter) ■ Compare and contrast indoor and outdoor air in terms of which pollutants are likely to be present and their sources (1.3, 1.13)
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Interpret local air quality data, including why air quality standards are set separately for each pollutant (1.3) Evaluate the risks and benefits of a particular activity (1.3) Discuss the green chemistry initiative and why it makes sense to prevent pollution rather than to clean it up afterward (1.5) Relate these terms: matter, pure substances, mixtures, elements, compounds, metals, nonmetals (1.6) Discuss the features of the periodic table, including the groups it contains (1.6)
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Explain the difference between atoms and molecules, giving examples of each (1.7) Name chemical elements and compounds that relate to air quality (1.7) Write and interpret chemical formulas that relate to air quality (1.8) Balance and interpret chemical equations that relate to air quality (1.9–1.10) Understand oxygen’s role in combustion, including how hydrocarbons burn to form carbon dioxide, carbon monoxide, and soot (1.9–1.10)
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Describe how ozone forms, including how sunlight, NO, NO2, and VOCs are involved (1.12) Identify the sources and nature of indoor air pollution (1.13) Explain the unreasonableness of “pollution-free” air (1.14) Use scientific notation and significant figures in performing basic calculations (1.4, 1.14) Apply what you know about air pollution to ways of living that result in cleaner air (entire chapter)
Questions The end-of-chapter questions are grouped in three ways: ■ Emphasizing Essentials questions give you the opportunity to practice fundamental skills. They are similar to the Your Turn exercises in the chapter. ■ Concentrating on Concepts questions are more difficult and may relate to societal issues. They are similar to the Consider This activities in the chapter. • ■ Exploring Extensions questions challenge you to go beyond the information presented in the text. Appendix 5 contains the answers to questions with numbers in blue.
4.
5.
6.
Questions marked with this icon require the resources of the Internet. Questions marked with this icon relate to green chemistry. Emphasizing Essentials 1. a. Calculate the volume of air in liters that you might inhale (and exhale) in an 8-hour working day. Assume that each breath has a volume of about 0.5 L, and that you are breathing 15 times a minute. b. From this calculation, you can see that breathing exposes you to large volumes of air. Name five things that you can do to improve the quality of the air that you and others breathe. 2. Our atmosphere can be characterized both as a thin veil that supports life and as a few vertical miles of chemicals. Explain what makes each description accurate. Also state which feature(s) of our atmosphere each description emphasizes and which ones it omits. 3. These gases are found in the troposphere: Rn, CO2, CO, O2, Ar, N2. a. Rank them in order of their abundance in the troposphere. b. For which of these gases is it convenient to express its concentration in parts per million? c. Which of these gases is/are currently regulated as an air pollutant where you live?
7.
d. Which of these gases are found in Group 8A of the periodic table, the noble gases? Give three examples of particulate matter found in air. Explain the difference between PM2.5 and PM10 in terms of size and health effects. Radon is one of the noble gases, found in Group 8A on the periodic table. Which properties does it share with the other inert gases? In which way is it distinctly different? a. The concentration of argon in air is approximately 0.9%. Express this value in ppm. b. The air exhaled from the lungs of a smoker has a concentration of 20–50 ppm CO. In contrast, air exhaled by nonsmokers is 0–2 ppm CO. Express each concentration as a percent. c. In a tropical rain forest, the water vapor concentration may reach 50,000 ppm. Express this as a percent. d. In the dry polar regions, water vapor may be a mere 10 ppm. Express this as a percent. In these diagrams, two different types of atoms are represented by color and size. Characterize each sample as an element, a compound, or a mixture. Explain your reasoning.
(a)
(b)
(c)
(d)
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8. Consider this representation of the reaction between nitrogen and hydrogen to form ammonia (NH3).
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a. What combustion products would you expect from the burning of wood? b. This fire is emitting at least three pollutants. Which are visible and which are not? 13. Consider this portion of the periodic table and the groups shaded on it.
a. Are the masses of reactants and products the same? Explain. b. Are the numbers of molecules of reactants and of products the same? Explain. c. Are the total number of atoms in the reactants and the total number of atoms in the products the same? Explain. 9. Express each of these numbers in scientific notation. a. 1500 m, the distance of a foot race b. 0.0000000000958 m, the distance between O and H atoms in a water molecule
a. What is the group number for each shaded region? b. Name the elements that make up each group. c. Give a general characteristic of the elements in each of these groups. 14. Consider the following blank periodic table.
c. 0.0000075 m, the diameter of a red blood cell d. 150,000 mg of CO, the approximate amount breathed daily 10. Write each of these values as a “regular” number. a. 8.5 3 104 g, the mass of air in an average room b. 1.0 3 107 gallons, the volume of crude oil spilled by the Exxon Valdez c. 5.0 3 1023%, the concentration of CO in the air on acity street d. 1 3 1025 g, the recommended daily allowance of vitamin D 11. The threshold for detecting NO2 by smell is 0.00022 g/m3 of air. a. Express this value in scientific notation. b. Would you expect a similar value for the threshold of CO? c. Name another pollutant that has a sharp, easily detected odor. 12. Wildfires occur all across our planet. The one shown here was photographed on a commercial flight north of Phoenix, AZ.
a. Shade the region of the periodic table in which metals are found. b. Common metals include iron, magnesium, aluminum, sodium, potassium, and silver. Give the chemical symbol for each. c. Give the name and chemical symbol for five nonmetals (elements that are not in your shaded region). 15. Classify each of these substances as an element, a compound, or a mixture. a. a sample of “laughing gas” (dinitrogen monoxide, also called nitrous oxide) b. steam coming from a pan of boiling water c. a bar of deodorant soap d. a sample of copper e. a cup of mayonnaise f. the helium filling a balloon 16. These gases are found in the atmosphere in small amounts: CH4, SO2 and O3. a. What information does each chemical formula convey about the number and types of atoms present? b. Give the names of these gases. 17. Hydrocarbons are important fuels that we burn for many different reasons. a. What is a hydrocarbon? b. Rank these hydrocarbons by the number of carbons they contain: propane, methane, butane, octane, ethane.
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c. We suggested “mother eats peanut butter” as a memory aid for the names of the fi rst four hydrocarbons. Propose a new one that includes pent-, the prefi x that indicates five carbon atoms. 18. Write balanced chemical equations to represent these reactions. Hint: Nitrogen and oxygen are both diatomic molecules. a. Nitrogen reacts with oxygen to form nitrogen monoxide. b. Ozone decomposes into oxygen and atomic oxygen (O). c. Sulfur reacts with oxygen to form sulfur trioxide.
Concentrating on Concepts 26. “Air prints” were mentioned in the opening activity of this chapter. Examine these two photographs. The first is a beautiful view from a lodging on the Hilo coast of the island Hawaii. The second shows the tarmac on a hazy summer day at the Narita International Airport in Tokyo. List three ways in which each photo shows the air print of humans. Hint: Some may not be visible but rather implied by the photograph.
19. Analogous to equation 1.8, draw models to represent the chemical equations from the previous question. 20. These questions relate to combustion of hydrocarbons. a. LPG (liquid petroleum gas) is mostly propane, C3H8. Balance this equation. C3H8(g) ⫹ O2(g)
CO2(g) ⫹ H2O(g)
b. Cigarette lighters burn butane, C4H10. Write a balanced equation, assuming complete combustion, that is, plenty of oxygen. c. With a limited supply of oxygen, both propane and butane can burn incompletely to form carbon monoxide. Write balanced equations for both reactions. 21. Balance these equations in which ethane (C2H4) burns in oxygen. a. C2H4(g) ⫹ O2(g) b. C2H4(g) ⫹ O2(g) c. C2H4(g) ⫹ O2(g)
C(s) ⫹ H2O(g) CO(g) ⫹ H2O(g) CO2(g) ⫹ H2O(g)
22. Examine the coefficients for oxygen in the balanced equations from the previous question. Explain why they vary, depending on whether C, CO, or CO2 is formed. 23. Count the atoms on both sides of the arrow to demonstrate that these equations are balanced. a. 2 C3H8(g) ⫹ 7 O2(g) 6 CO(g) ⫹ 8 H2O(l) b. 2 C8H18(g) ⫹ 25 O2(g) 16 CO2(g) ⫹ 18 H2O(l ) 24. Platinum, palladium, and rhodium are used in the catalytic converters of cars. a. Give the chemical symbol for each metal. b. Locate each metal on the periodic table. c. What can you infer about the properties of these metals, given that they are useful in this application? 25. Nail polish remover containing acetone was spilled in a room 6 m 3 5 m 3 3 m. Measurements indicated that 3600 mg of acetone evaporated. Calculate the acetone concentration in micrograms per cubic meter.
27. The AIRNOW website (EPA) states that “Quality of air means quality of life.” Demonstrate the wisdom of this statement for two air pollutants of your choice. 28. In Consider This 1.2, you calculated the volume of air exhaled in a day. How does this volume compare with the volume of air in your chemistry classroom? Show your calculations. Hint: Think ahead about the most convenient unit to use for measuring or estimating the dimensions of your classroom. 29. According to Table 1.1, the percentage of carbon dioxide in inhaled air is lower than it is in exhaled air, but the percentage of oxygen in inhaled air is higher than in exhaled air. How can you account for these relationships? 30. Cars don’t inhale and exhale like humans do. Nonetheless, the air that goes into a car is different from what comes out. In Your Turn 1.19 you listed what comes out of a tailpipe. Now comment on the differences between the air that goes into the car engine and what comes out the tailpipe. For which chemicals have the concentrations noticeably increased or decreased? 31. A headline from the Anchorage Daily News in Alaska (January 17, 2008): “Family in car overcome by carbon monoxide. Fire department saves five after slide into snow bank.”
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34.
35.
36.
37. 38.
a. In general, which groups of people are the most sensitive to ozone? 50 40 30 20 10 0 8
10
12 14 16 18 Date in June, 2005
20
b. Air rated above 30 is hazardous for some or all groups by the Canadian Environmental Assessment Agency. For the data shown, how many days was the air hazardous? c. Ozone levels drop off sharply at night. Explain why. d. During the daytime, the ozone dropped off sharply after June 10. Propose two different reasons that could account for this observation. e. Data for the month of December is not shown. Would you expect the ozone levels to be higher or lower than those in June? Explain. 39. Here are air quality data for April 1–10, 2005 in Beijing. The primary pollutant was PM10. a. In general, which groups of people are the most sensitive to particulate matter? 400 350 300 Chinese AQI
33.
Canadian AQI
32.
a. If your car is in a snow bank and the engine is running, CO may accumulate inside the car. Normally, however, CO does not accumulate in the car. Explain. b. Why didn’t the occupants detect the CO? A headline from the Pioneer Press in St. Paul, Minnesota (January 8, 2008): “Man dies after exposure to gas; carbon monoxide sickens five others.” a. Name two possible sources of CO inside a home. b. The level measured was 4700 ppm. Express this value as a percent. c. How does this level compare with the U.S. ambient air quality standards set by the EPA? d. Name three symptoms that the survivors most likely experienced. e. Where in a home should you install CO detectors? Note: Adjacent to a furnace is not usually recommended. In Consider This 1.4, you considered how life on Earth would change if the concentration of oxygen were twice as high. Now consider how life would change if the concentration of O2 were cut in half. Give two examples of things that what would be affected. Explain why CO is named the “silent killer.” Select two other pollutants for which this name would not apply and explain why not. Undiluted cigarette smoke may contain 2–3% carbon monoxide. a. How many parts per million is this? b. How does this value compare with the National Ambient Air Quality Standards for CO in both a 1-hour and an 8-hour period? c. Propose a reason why smokers do not die from carbon monoxide poisoning. For many states, the ozone season runs from May 1 to October 1. Why are ozone levels typically not reported in the winter months? The EPA characterizes ozone as “good up high, bad nearby.” Explain. Here are ozone air quality data for London, Ontario, from June 8–20, 2005.
61
250 200 150 100 50 0 1
3 5 7 9 Date in April, 2005
b. Air rated above 100 is hazardous for some or all groups by the National Environmental Monitoring Centre in China. For the data shown, how many days was the air hazardous? c. The levels of PM do not necessarily drop off at night the way they do for ozone. Explain. d. The levels of particulate matter increased sharply on April 5. Propose two different reasons that could account for this observation. 40. Prior to 1990, diesel fuel could contain as much as 2% sulfur. New regulations have changed this, and today most diesel fuel is ultra-low sulfur diesel (ULSD) containing a maximum of 15 ppm sulfur.
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a. Express 15 ppm as a percent. Likewise, express 2% in terms of ppm. How many times lower is the ULSD than the older formulation of diesel fuel? b. Write a chemical equation that shows how burning diesel fuel containing sulfur contributes to air pollution. c. Diesel fuel contains the hydrocarbon C12H26. Write a chemical equation that shows how burning diesel adds carbon dioxide to the atmosphere. d. Comment on burning diesel fuel as a sustainable practice, both in terms of how things have improved and in terms of where they still need to go. 41. A certain city has an ozone reading of 0.13 ppm for 1 hour, and the permissible limit is 0.12 for that time. You have the choice of reporting that the city has exceeded the ozone limit by 0.01 ppm or saying that it has exceeded the limit by 8%. Compare these two methods of reporting. 42. Here is the air quality outlook for the United States for early August, 2005. Source: www.airnow.gov
45.
46.
47.
48.
a. This forecast is typical in that most of the ozone pollution is expected in California, Denver, Texas, the Midwest, and the East Coast. Why these regions of the country? b. Phoenix typically has high ozone levels in the summer, but not on this particular day. Offer a possible explanation. c. The air quality is forecasted to improve in Illinois and Wisconsin. Offer a possible explanation. d. Why are inland areas in California, such as the Sacramento Valley which is shown as unhealthy for sensitive groups, likely to have worse air quality than the California coast? 43. Look up the ozone air quality data for two states, one in the Sun Belt and one not. Account for any difference you find. Hint: Use State of the Air, a website posted by the American Lung Association. 44. At certain times of the year, inhabitants of the beautiful city of Santiago, Chile, breathe some of the worst air on the planet.
a. Driving private cars has been severely restricted in Santiago. How specifically does this improve air quality? b. Although the population of Santiago is comparable to that in other cities, its air quality is much worse. Suggest geographical features that might be responsible. a. Explain why jogging outdoors (as opposed to sitting outdoors) increases your exposure to pollutants. b. Jogging indoors at home can decrease your exposure to some pollutants, but may increase your exposure to others. Explain. Consumers now can purchase paints that emit only low amounts of VOCs. However, these consumers may not know why it matters to purchase this paint. a. What would you print on the label of a paint can to make the point that a low-VOC paint is a good idea? b. We apply paint to many outdoor surfaces, such as buildings, bridges, and fence posts. Comment on the environmental effects of the VOCs that these paints emit. c. Explain how producing low-VOC paint meets the Triple Bottom Line. One can purchase a carbon monoxide monitor that immediately sounds an alarm if the concentration of CO reaches a threshold. In contrast, most radon detection systems sample the air over a period of time before an alarm sounds. Why the difference? Select a profession of your choice, possibly the one you intend to pursue. Name at least one way that a person in this profession could have a positive effect on air quality.
Exploring Extensions 49. “Air pollution is a diffuse problem, the shared fault of many emitters. It is a classic example of the tragedy of the commons.” Source: Introduction to Air in California, by David Carle, 2006. Explain the phrase “tragedy of the commons” and how air pollution is a classic example. 50. Mercury, another serious air pollutant, is not described in this chapter. If you were a textbook author, what would you include about mercury emissions? How would you connect mercury emissions to the sustainable use of resources? Write several paragraphs in a style that would match that of this textbook. 51. The EPA oversees the Presidential Green Chemistry Challenge Awards. Use the EPA website to find when the program started and to find the list of the most recent winners of the award. Pick one winner and
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o ur 3h
s
urs 2 ho
60 % blood saturation
summarize in your own words the green chemistry advance that merited the award. 52. Recreational scuba divers usually use compressed air that has the same composition as normal air. A mixture being used is called Nitrox. What is its composition, and why is it being used? 53. Here are two scanning electron micrograph images of particulate matter, courtesy of the National Science Foundation and researchers at Arizona State University. The first is of a soil particle and the second of a rubber particle, and each is about 10 μm in diameter.
63
Dangerous to life Coma, collapse
50
r 1 hou
40
Severe headache, nausea, dizziness
30 Slight headache 20 10
a. What is a likely source of the rubber particle? Name two other substances that might contribute PM to the air. b. The soil particle is composed mainly of silicon and oxygen. What other elements are commonly present in the rocks and minerals in Earth’s crust? c. What about these photographs suggests that these particles would inflame your blood vessels? 54. Ultrafi ne particles have diameters less than 0.1 μm. In terms of their sources and health effects, how do these particles compare with PM 2.5 and PM10? Use the Internet to locate the most up-to-date information. 55. Most lawnmowers do not have catalytic converters (at least as this book went to press). What comes out of the tailpipe of a gasoline-powered lawn mower? Why has adding a catalytic converter been so controversial? What are the immediate benefits to curbing these emissions, as well as the longer term ones? 56. Consider this graph that shows the effects of carbon monoxide inhalation on humans. a. Both the amount of exposure and the duration of exposure have an effect on CO toxicity in humans. Use the graph to explain why. b. Use the information in this graph to prepare a statement to include with a home carbon monoxide detection kit about the health hazards of carbon monoxide gas.
400
800 1200 1600 Parts per million
2000
57. Consider This 1.4 asks you to consider how our world would be different if the oxygen content of the atmosphere were doubled. Develop your answer into an essay. Title your essay “An Hour in the Life of . . .” and describe how things would be different for a person of your choice. If an hour is too short to make your point, substitute “A Morning . . .” or “A Day . . .”. 58. You may have admired the beauty of hardwood floors. Polyurethane is the finish of choice for floors because it is more durable than varnishes and shellacs. Until recently, polyurethane was always an oil-based paint. But recently, the Bayer Corporation developed a water-based polyurethane that reduces the amount of VOCs by 50–90%. In 2000, Bayer was awarded a Presidential Green Chemistry Challenge award for this development. Prepare a summary of this work. Also check stores to see if any water-based polyurethanes are available in your area. 59. Composite wood is made by gluing smaller pieces of wood (often waste scraps of wood) together. Examples include plywood, particle board, and fiber board. a. Many glues release formaldehyde, a volatile compound. What are its hazards? b. Professor Kaichang Li of Oregon State University and Columbia Forest Products developed a new soy-based adhesive glue, winning a Presidential Green Chemistry Challenge Award in 2007. Prepare a summary of his accomplishments.
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The stratospheric ozone “hole” in 2009 (the purple and dark blue areas) over Antarctica reached a maximum area of 24.0 million km2 on September 17, 2009. The record 2000 hole area peaked at 28.4 million km2. Source: NASA.
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Good up high; bad nearby. To understand ozone, think location. Down here in the troposphere, where we live, ozone is a pollutant that forms in the presence of sunlight from other pollutants in our atmosphere. Once the Sun goes down, the generation of ozone ceases. Any ozone that is present quickly reacts with something else, and the concentrations drop in the twilight hours. Were the Sun not to rise again, we wouldn’t have to worry about breathing ozone (but we’d face a few other problems). But up in the stratosphere, the story of ozone is altogether different. Unlike ground-level ozone, the ozone up high is formed naturally. Also unlike ground-level ozone, stratospheric ozone plays a vital role in protecting us from damaging solar radiation. One might say it acts like the Earth’s sunglasses. In the 1970s, chemists discovered that certain chemicals could make their way into the upper atmosphere and partially destroy the protective ozone found there. Ever since, scientists, policy makers, and indeed concerned citizens worldwide have participated in efforts to control and reverse ozone destruction. Somewhat surprisingly, the most severe depletion has been over Antarctica, and the yearly images of the ozone hole have become some of the most widely recognized scientific graphics. Later in this chapter, you will have the opportunity to examine past trends and update the Antarctic ozone-hole story. You may be wondering what this story has to do with you, because last time we checked, not many college students were living in Antarctica. Even though ozone depletion was first documented in that faraway region, it also has been observed to a lesser extent in many other locations on Earth, including over North America. Where you live and the season of the year both influence the amount of stratospheric ozone overhead and how well it provides its protective effects. Take a look at some of the important data for yourself.
Consider This 2.1
Ozone Levels Above You
As you read this, an instrument onboard a satellite is measuring ozone levels in the stratosphere. The textbook’s website provides a link to the data. a. What is the current total column ozone (in Dobson units) at your location? Request data for three different years and compute an average. b. Now retrieve data for Antarctica for the same dates. How does the ozone measurement where you live compare? Note: If you are making the comparison in September or October, you are likely to see the biggest difference.
One DU (Dobson unit) corresponds to about one ozone molecule for every billion molecules and atoms present in air.
What caused this stratospheric ozone depletion? Why is this depletion so serious? Look for the answers to these questions and more in the sections that follow. As we explore the topic of ozone depletion, we also highlight the precautionary principle. This principle stresses the wisdom of acting, even in the absence of complete scientific data, before the adverse affects on human health or the environment become significant or irrevocable. As you will learn, the world community did act. The wisdom of the collective actions is evident, as the measures taken to protect the ozone layer appear to be working. Even so, another warning bell will sound in the final section of this chapter. Listen for it in regards to global climate change—the topic of Chapter 3.
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2.1
If you have ever been near a sparking electric motor or in a severe lightning storm, most likely you have smelled ozone. Its odor is unmistakable but difficult to describe. Some compare it to that of chlorine gas; others think the odor reminds them of newly mown grass. It is possible for humans to detect concentrations as low as 10 parts per billion (ppb), that is, 10 molecules out of 1 billion. Appropriately enough, the name ozone comes from a Greek word meaning “to smell.” The ozone and oxygen molecules differ by only one atom.
Altitude miles km 34.1
| Ozone: What and Where Is It?
55 Mesosp phere
31.0
50
oxygen molecule 27.9
45
24.8
40
21.7
35 Stra St rato tosp sphe here re
18.6
30 Weather balloons
15.5
25
12.4
20
9.3
15
6.2
10
3.1
5
Figure 2.1 The regions of the atmosphere.
0.075 parts per million is equivalent to 75 ozone molecules for every billion molecules and atoms found in air.
2 O3
[2.1]
This chemical equation helps to explain why ozone is formed from oxygen in the presence of an electrical discharge, whether from an electric spark or lightning. Ozone is reasonably rare in the troposphere, the region of the atmosphere closest to the Earth’s surface (Figure 2.1). Only somewhere between 20 and 100 ozone molecules typically occur for each billion molecules and atoms that make up the air. Unhealthy concenJet airplanes trations are sometimes found near Earth’s surface, the result of chemical reactions that produce it as a component of photochemical smog. Mt. Everest, But these concentrations are very low. As was noted in the previous Nepal chapter, the air quality standard in the United States for ground-level ozone was set at 0.075 ppm for an 8-hr average as of 2009. But what is detrimental in one region of the atmosphere, even at very low concentrations, can be essential in another. The stratoLa Paz, Bolivia sphere, at an altitude of 15 to 30 km, is where ozone does most of Denver, CO, USA its filtering of some types of ultraviolet light from the Sun. The Sea level concentration of ozone in this region is several orders of magnitude greater than in the troposphere, but still very low. As an upper limit, about 12,000 ozone molecules are present for every billion molecules and atoms of gases that make up the atmosphere at this level. Most of the ozone on our planet, about 90% of the total, is found in the stratosphere. The term ozone layer refers to a designated region in the stratosphere of maximum ozone concentration. Figure 2.2 shows the ozone concentrations in the troposphere and stratosphere. Tops of intense thunderstorms
Troposphere
As we will see, this difference in molecular structure translates to significant differences in chemical properties. One difference is that ozone is far more chemically reactive than O2. As you will learn in Chapter 5, ozone can be used to kill microorganisms in water. Ozone is also used to bleach paper pulp and fabrics. At one time, ozone was even advocated as a deodorizer for air. This use only makes sense, however, if nobody breathes the air during the deodorizing process. In contrast, you safely can (and must) breathe oxygen day in and day out. Although oxygen is still quite chemically reactive, it is not reactive enough to bleach paper or purify water. Ozone forms both naturally and as a result of human activities. Given the high reactivity of ozone, it does not usually persist very long. If it were not for the fact that ozone is formed anew on our planet, you would not find it except as a curiosity in the chemistry lab. Ozone can be formed from oxygen, but the process requires energy. A simple chemical equation summarizes the process: energy ⫹ 3 O2
Ozone layer
ozone molecule
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35
21.7
30
18.6
Altitude (km)
15.5
Ozone layer 20
12.4
15
9.3
10 5
Ozone increases from pollution
Altitude (miles)
Stratosphere
25
67
6.2 Troposphere
Ozone amount
3.1 0
Figure 2.2 The ozone layer is a region of maximum ozone concentration in the stratosphere. Source: Global Ozone Research and Monitoring Project Report No. 44, 1998. Reprinted with permission of World Meteorological Organization.
Your Turn 2.2
The Ozone Layer
Use Figure 2.2 and values given in the text to answer these questions. a. What is the approximate altitude of maximum ozone concentration? b. What is the maximum number of ozone molecules per billion molecules and atoms of all types found in the stratosphere? c. What is the maximum number of ozone molecules per billion molecules and atoms of all types found in ambient air just meeting the EPA limit for an 8-hr average? Answer a. About 23 km (14 miles).
Because the range of altitudes is so broad, the concept of an “ozone layer” can be a little misleading. No thick, fluffy blanket of ozone exists in the stratosphere. At the altitudes of the maximum ozone concentration, the atmosphere is very thin, so the total amount of ozone is surprisingly small. If all the O3 in the atmosphere could be isolated and brought to the average pressure and temperature at Earth’s surface (1.0 atm and 15 °C), the resulting layer of gas would have a thickness of less than 0.5 cm, or about 0.25 inch. On a global scale, this is a minute amount of matter. Yet, this ozone shield protects the surface of the Earth and its inhabitants from the harmful effects of ultraviolet radiation. The total amount of ozone in a vertical column of air of known volume can be determined fairly easily. The determination can be done from Earth’s surface by measuring the amount of UV radiation reaching a detector; the lower the intensity of the radiation, the greater the amount of ozone in the column. G. M. B. Dobson, a scientist at Oxford University, pioneered this measurement method. In 1920, he invented the first instrument to quantitatively measure the concentration of ozone in a column of the Earth’s atmosphere. Therefore, it is fitting that the unit of such measurements is named for him.
Consider This 2.3
Interpreting Ozone Values
A classmate used the National Aeronautics and Space Administration (NASA) website to check the atmospheric ozone above her hometown in Ohio. She found the readings to be 417 DU (Dobson units) on April 10 and 386 DU on May 10. The student, reassured by these findings, concluded there had been an improvement in protection from damaging UV radiation. Do you agree? Explain.
One Dobson unit (DU) is equivalent to about 3 3 1016 O3 molecules in a vertical column of air with a cross section of 1 cm2.
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NASA’s EOS Aura mission also is collecting data about tropospheric air quality (Chapter 1) and key climate parameters (Chapter 3).
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Chapter 2
Scientists continue to measure and evaluate ozone levels using ground observations, weather balloons, and high-flying aircraft. However, since the 1970s, measurements of total column ozone have also been made from the top of the atmosphere. Satellite-mounted detectors record the intensity of the UV radiation scattered by the upper atmosphere. The results are then related to the amount of O3 present. The Space Shuttle, Columbia, tested a new approach for monitoring ozone. Rather than looking directly downward toward Earth from a satellite, the equipment aboard the Shuttle looked sideways through the thin blue haze (see photo) that rises above the denser regions of the troposphere and follows the curve of the Earth. This region is known as the Earth’s “limb” and is responsible for the name of this new technique, “limb viewing.” Reliable information can be gathered at each level of the atmosphere, particularly allowing scientists to better understand chemistry taking place in the lower regions of the stratosphere. In January 2004, NASA launched a new mission called Earth Observing System (EOS) Aura that also uses a variety of viewing geometries, including limb viewing, to gather additional data about changes in Earth’s stratospheric ozone layer. The process by which ozone protects us from damaging solar radiation involves the interaction of matter and energy from the Sun. To help you to understand this, we turn first to a submicroscopic view of matter.
2.2 |
Atomic Structure and Periodicity
Both the O2 and O3 molecules are composed of oxygen atoms. What do we know about these atoms? During the 20th century, scientists probed the inner workings of the atom. The physicists were almost too successful in their endeavors, finding more than 200 subatomic particles. Fortunately, most chemical behavior can be explained with only three. Every atom has a nucleus, a minuscule and highly dense center of an atom composed of protons and neutrons. Protons are positively charged particles, and neutrons are electrically neutral particles. Both have almost exactly the same mass. Indeed, the protons and neutrons in the nucleus account for almost all the mass of an atom. Outside the nucleus are the electrons that define the boundary of the atom. An electron has a mass much smaller than that of a proton or neutron and a negative electric charge equal in magnitude to that of a proton, but opposite in sign. Therefore, in any electrically neutral atom, the number of electrons equals the number of protons. The properties of these particles are summarized in Table 2.1. The number of protons in the nucleus determines the identity of the atom. The term atomic number refers to the number of protons in the nucleus of an atom. For example, all hydrogen (H) nuclei contain 1 proton; hydrogen has an atomic number of 1. Similarly, all helium (He) nuclei contain 2 protons and have an atomic number of 2. With each successive element in the periodic table, the atomic number increases. For example, the nucleus of element #92 (U, uranium) contains 92 protons.
Table 2.1
Properties of Subatomic Particles
Particle
Relative Charge
Relative Mass
Actual Mass, kg
proton
11
1
1.67 3 10 −27
neutron
1
1.67 3 10 −27
electron
−1
0*
9.11 3 10 −31
*This value is zero when rounded to the nearest whole number. The electron does indeed have mass, though very small!
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Your Turn 2.4
69
Atomic Bookkeeping
Using the periodic table as a guide, specify the number of protons and electrons in a neutral atom of each of these elements. a. carbon (C)
b. calcium (Ca)
c. chlorine (Cl)
d. chromium (Cr)
Answers a. 6 protons, 6 electrons b. 20 protons, 20 electrons
We wish we could show you a picture of a typical atom. However, atoms defy easy representation, and depictions in textbooks are at best oversimplifications. Electrons are sometimes pictured as moving in “orbits” about the nucleus, but the modern view of electrons is a good deal more complicated and abstract. For one thing, the relative size of the nucleus and the atom creates serious problems for the illustrator. If the nucleus of a hydrogen atom were the size of a period on this page, the atom’s single electron would most likely be found at a distance of about 10 feet from that period. Moreover, electrons do not follow specific circular orbits. In spite of what you may have learned early in your education, an atom is really not very much like a miniature solar system. Rather, the distribution of electrons in an atom is described best using concepts of probability and statistics. If this sounds rather vague to you, you are not alone. Common sense and our experience of ordinary things are not particularly helpful in our efforts to visualize the interior of an atom. Instead, we are forced to resort to mathematics and metaphors. The mathematics required (a field called quantum mechanics) can be formidable. Chemistry majors do not normally encounter this field until rather late in their undergraduate study. Although we cannot fully share with you the strange beauties of the peculiar quantum world of the atom, we can provide some useful generalizations. The periodic table lists elements in order of increasing atomic number. The table also has elements arranged so that those with similar chemical properties fall in the same column (group). For example, lithium (Li, atomic number 3), sodium (Na, 11), potassium (K, 19), rubidium (Rb, 37), and cesium (Cs, 55) all fall in the same column and all are highly reactive metals. What fundamental feature accounts for this? Today we know that the chemical properties of elements are the consequence of the distribution of electrons in the atoms of these elements. When chemical properties repeat themselves, this signals a repeat in electronic arrangement. As we will see, the electrons farthest from the nucleus are the main determinant of chemical properties. Both experiment and calculation demonstrate that the electrons are arranged in certain energy levels about the nucleus. The electrons in the innermost level are the most strongly attracted by the positively charged protons in the nucleus. The greater the distance between an electron and the nucleus, the weaker the attraction between them. We say that the more distant electron is in a higher energy level, which means that the electron itself possesses more potential energy. Each energy level has a maximum number of electrons that can be accommodated and is particularly stable when fully occupied. The innermost level, corresponding to the lowest energy, can hold only two electrons. The second level has a maximum capacity of eight, and the higher levels are also particularly stable when they contain eight electrons. Table 2.2 shows some important information about electrons in neutral atoms of the first 18 elements. The total number of electrons in each atom is printed in blue and the number of outer electrons is printed in maroon. Outer (valence) electrons are found in the highest energy level and help to account for many of the observed trends in chemical properties. Observe that the group designation (1A, 2A, etc.) corresponds to the number of outer electrons for the A group elements, one of the most useful organizing features of the periodic table.
What we call “levels” used to be referred to as “shells,” using the earlier solar system model of atomic structure.
Look for more about potential energy in Chapter 4.
The periodic table also contains B group elements. Table 2.2 does not show these, as these elements start with the fourth row.
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Chapter 2
Table 2.2
The smallest noble gas, helium, has 2 valence electrons rather than 8.
Group 1A
Lithium (stored in oil)
Atoms of the First 18 Elements (Total and Outer Electrons)
2A
3A
4A
5A
6A
7A
8A
1
2
H
He
1
2
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
1
2
3
4
5
6
7
8
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
1
2
3
4
5
6
7
8
• Above the atomic symbol is the atomic number, the number of protons in the nucleus. For a neutral atom, this also is the number of electrons. • Below the atomic symbol is the number of outer electrons in a neutral atom.
Sodium (removed from oil, being cut)
Potassium (in a sealed glass tube)
Take another look at the first column in Table 2.2. Lithium and sodium atoms both have one outer electron per atom, despite having different total numbers of electrons. This fact explains much of the chemistry that these two alkali metals have in common. It places them in Group 1A of the periodic table (the 1 indicates one outer electron). Moreover, we would be correct in assuming that potassium, rubidium, and the other elements in column 1A of the periodic table also have a single outer electron in each of their atoms. They are all metals that react readily with oxygen, water, and a wide range of other chemicals. Figure 2.3 shows photographs of some Group 1A elements. The periodic table is a useful guide to electron arrangement. In the families (another name for groups) of elements marked “A,” the number that heads the column indicates the number of outer electrons in each atom. We introduced the terms alkali metal, alkaline earth metal, halogen, and noble gas in Chapter 1. We now connect these terms with their group number. ■ ■ ■ ■
Alkali metals (Group 1A)—highly reactive metals with one outer electron Alkaline earth metals (Group 2A)—reactive metals with two outer electrons Halogens (Group 7A)—reactive nonmetals with 7 outer electrons Noble gases (Group 8A)—unreactive nonmetals with 8 outer electrons
Your Turn 2.5
Outer Electrons
Rubidium (in a sealed glass tube)
Figure 2.3 Selected Group 1A elements.
Using the periodic table as a guide, specify the group number and number of outer electrons in a neutral atom of each element. a. sulfur (S)
The group number does not necessarily indicate the number of outer electrons for elements in B groups.
b. silicon (Si)
c. nitrogen (N)
d. krypton (Kr)
Answers a. Group 6A; 6 outer electrons b. Group 4A; 4 outer electrons
Your Turn 2.6
Family Features
a. What electronic feature do fluorine (F), chlorine (Cl), bromine (Br), and iodine (I) have in common? b. The element beryllium (Be), like the other elements in Group 2A, has two outer electrons. Give the names and symbols for the other Group 2A elements. Answer a. All have seven outer electrons. They belong to Group 7A, the halogens.
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Table 2.3
Isotopes of Hydrogen # of Protons (atomic number)
# of Neutrons
#n1#p (mass number)
Name
Isotope
hydrogen
H-1 or 11H
1
1
deuterium
H-2 or 21H
1
1
2
tritium
H-3 or 31H
1
2
3
71
Elements have both stable and radioactive isotopes. For example, H-3 (tritium) is radioactive, but H-1 and H-2 are not. Look for more about radioisotopes in Chapter 7.
In addition to electrons and protons, atoms also contain neutrons. The one (and only) exception is an atom of the most common form of hydrogen, which consists of only one proton in its nucleus. But the nucleus of 1 out of every 6700 hydrogen atoms also contains a neutron. This naturally occurring form of hydrogen is called deuterium. Tritium, a radioactive form of hydrogen that is quite rare in nature, has two neutrons in its nucleus. Hydrogen, deuterium, and tritium are examples of isotopes, two or more forms of the same element (same number of protons) whose atoms differ in number of neutrons, and hence in mass. An isotope is identified by its mass number, the sum of the number of protons and neutrons in the nucleus of an atom. The mass number can vary for the same element. In contrast, the atomic number cannot vary for the same element. For example, the full atomic symbol 11H represents the most common isotope of hydrogen. Because the atomic number of 1 for hydrogen is invariant, the subscript is sometimes omitted. Thus you also may see 1H, hydrogen-1, or H-1. Table 2.3 summarizes information about the isotopes of hydrogen.
Your Turn 2.7
Protons and Neutrons
Specify the number of protons and electrons in the nucleus of each of these. a. carbon-14 (146 C)
b. uranium-235 ( 235 92U)
c. iodine-131 (131 53 I)
Answers a. 6 protons, 8 neutrons b. 92 protons, 143 neutrons
All elements have more than one isotope, but the number of stable ones varies considerably. Each element’s atomic mass, the number you see on every periodic table, takes the relative natural abundance of isotopes, as well as their masses, into account. Following our general rule of introducing information on a need-to-know basis, we will return to a discussion of atomic masses in Chapter 3.
2.3 |
Molecules and Models
Having taken a short excursion into the atomic realm, we now move to the topic of bonding in molecules so that in turn, we can understand the ozone hole. Let’s begin with the simplest molecule, H2. Each hydrogen atom has one electron. If two hydrogen atoms bond, the two electrons become common property. If we represent each electron by a dot, the two hydrogen atoms might look something like this: H
and
H
Bringing the two atoms together yields a molecule that can be represented this way. H H
Mass number is the total number of protons and neutrons in a specific isotope. Atomic mass refers to a weighted average of all naturally occurring isotopes of that element.
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You are unlikely to run into HF in your chemistry laboratory. It is a highly reactive compound, and in aqueous solution it is used to etch glass.
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Chapter 2
Each atom effectively has a share in both electrons. The resulting H2 molecule has a lower energy than the sum of the energy in the two individual H atoms, and consequently the molecule with its bonded atoms is more stable than the separate atoms. The two electrons that are shared constitute a covalent bond. Appropriately, the name covalent implies “shared strength.” A Lewis structure is a representation of an atom or molecule that shows its outer electrons. The name honors Gilbert Newton Lewis (1875–1946), an American chemist who pioneered its use. Lewis structures, also called dot structures, can be predicted for many simple molecules by following a set of straightforward steps. We first illustrate the procedure with hydrogen fluoride, HF, another simple molecule. 1. Note the number of outer electrons contributed by each of the atoms. Hint: The periodic table is a useful guide for Group A elements. 1 H atom (H ) 3 1 outer electron per atom 5 1 outer electron 1 F atom ) F ) 3 7 outer electrons per atom 5 7 outer electrons 2. Add the outer electrons contributed by the individual atoms to obtain the total number of outer electrons available. 1 1 7 5 8 outer electrons 3. Arrange the outer electrons in pairs. Then distribute them so as to maximize stability by giving each atom a share in enough electrons to fully fill its outer level: 2 electrons in the case of hydrogen, 8 electrons for most other atoms. H F We surrounded the F atom with 8 dots, organized into 4 pairs. The pair of dots between the H and the F represents the electron pair that forms the bond uniting the hydrogen and fluorine atoms. The other 3 pairs of electrons are not shared with other atoms. As such, they are called nonbonding electrons, or “lone pairs.” A single covalent bond is formed when two electrons (one pair) are shared between two atoms. A line may be used to represent the two electrons in the bond. H
F
Sometimes the nonbonding electrons are removed from a Lewis structure, simplifying it still more. The result is called a structural formula. H
F
Remember that the single line represents one pair of shared electrons. These 2 electrons plus the 6 electrons in the 3 nonbonding pairs mean that the fluorine atom is associated with 8 outer electrons, whether or not the electrons are specifically shown. Remember that the hydrogen atom has no additional electrons other than the single pair shared with fluorine. It is at maximum capacity with two electrons. The fact that electrons in many molecules are arranged so that every atom (except hydrogen) shares in eight electrons is called the octet rule. This generalization is useful for predicting Lewis structures and the formulas of compounds. Consider the Cl2 molecule, the diatomic form of elemental chlorine. From the periodic table, we can see that chlorine, like fluorine, is in Group 7A, which means that its atoms each have 7 outer electrons. Using the scheme given for HF earlier, we first count and add up the outer electrons for Cl2. 2 Cl atoms () Cl )) 3 7 outer electrons per atom 5 14 outer electrons For the Cl2 molecule to exist, a bond must connect the two atoms. The remaining 12 electrons constitute 6 nonbonding pairs, distributed in such a way as to give each chlorine atom a share in 8 electrons (2 bonding and 6 nonbonding). Here is the Lewis structure. Cl
Cl
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Your Turn 2.8
73
Lewis Structures for Diatomic Molecules
Draw the Lewis structure for each molecule. a. HBr b. Br2 Answer a. 1 H atom (H ( 3 1 outer electron per atom 5 1 outer electron 1 Br atom ( Br ) 3 7 outer electrons per atom 5 7 outer electrons Total 5 8 outer electrons Here is the Lewis structure: H Br
or H
Br
So far we have dealt only with molecules having just two atoms. But the octet rule applies to larger molecules as well. Let’s use a water molecule, H2O, as an example. Just as with two-atom molecules, first tally the outer electrons. 2 H atoms (H ( 3 1 outer electron per atom 5 2 outer electrons 1 O atom ( O ) 3 6 outer electrons per atom 5 6 outer electrons Total 5 8 outer electrons In molecules like water that have a single atom bonded to two or more atoms of a different element (or elements), the single atom is the central one. You’ll encounter exceptions, but this is a useful rule. Since oxygen is the “single atom” in H2O, we place it in the center of the Lewis structure. Each of the H atoms bonds to the O atom, using 4 electrons. The remaining 4 electrons go on the O atom as 2 nonbonding pairs. H O H
Each hydrogen atom forms only one bond (two shared electrons). Oxygen can form two bonds and is the central atom in H2O.
A quick count confirms that the O atom is surrounded by 8 electrons, as predicted by the octet rule. Alternatively, we could use lines for the single bonds. H
O
H
Chemical formulas show the types and ratio of atoms present. In contrast, Lewis structures also indicate how the atoms are connected and show the nonbonding pairs of electrons, if present. Note that Lewis structures do not directly reveal the shape of a molecule. For example, from the Lewis structure we drew it might appear that the water molecule is linear. In fact, the molecule is bent.
H
O
H or
Another molecule to consider is methane, CH4. Again, we begin by tallying the valence electrons. 4 H atoms (H ( 3 1 outer electron per atom 5 4 outer electrons 1 C atom ) C ) 3 4 outer electrons per atom 5 4 outer electrons Total 5 8 outer electrons The central carbon atom is surrounded by the 8 electrons, giving carbon an octet of electrons. In the Lewis structure, each H atom uses 2 of the electrons to bond with the C atom, for a total of 4 single covalent bonds. H H C H or H H
H C
H
H
Remember that H can only accommodate a pair of electrons. The next activity gives you the opportunity to practice with other molecules.
The space-filling model of water was shown in Section 1.7. We will explain why the water molecule is bent in Chapter 3.
The combustion of methane was discussed in Section 1.10. Look in Chapter 3 for an explanation of the shape of the methane molecule.
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Chapter 2
Your Turn 2.9
More Lewis Structures
Draw the Lewis structure for each of these molecules. Both obey the octet rule. a. hydrogen sulfide (H2S) b. dichlorodifluoromethane (CCl2F2) Answer a. 2 H atoms ( H ) 3 1 outer electron per atom 5 2 outer electrons 1 S atom ( S ) 3 6 outer electrons per atom 5 6 outer electrons Total 5 8 outer electrons The Lewis structure is H S H or H S H. The Lewis structures for H2S and H2O differ only in the central atom.
In some structures, single covalent bonds do not allow the atoms to follow the octet rule. Consider, for example, the O2 molecule. Here we have 12 outer electrons to distribute, 6 from each of the oxygen atoms. There are not enough electrons to give each of the atoms a share in eight electrons if only one pair is held in common. However, the octet rule can be satisfied if the two atoms share four electrons (two pairs). A covalent bond consisting of two pairs of shared electrons is called a double bond. This bond is represented by four dots or by two lines. O O or O
O
Double bonds are shorter, stronger, and require more energy to break than single bonds involving the same atoms. The experimentally measured length and strength of the bond in the O2 molecule correspond to a double bond. However, oxygen has a property that is not fully consistent with the Lewis structure just drawn. When liquid oxygen is poured between the poles of a strong magnet, it sticks there like iron filings. Such magnetic behavior implies the presence of unpaired electrons rather than the paired arrangement shown in the preceding Lewis structures. But this is hardly a reason to discard the useful generalizations of the octet rule. After all, simple scientific models seldom if ever explain all phenomena, but they can be helpful approximations. There are other common examples in which the straightforward application of the octet rule leads to discrepancies in interpreting experimental evidence. Coming across data that do not seem to fit has led to the development of more sophisticated models. A triple bond is a covalent linkage made up of three pairs of shared electrons. For the same atoms, triple bonds are even shorter, stronger, and harder to break than double bonds. For example, the nitrogen molecule, N2, contains a triple bond. Each Group 5A nitrogen atom contributes 5 outer electrons for a total of 10. These 10 electrons can be distributed in accordance with the octet rule if 6 of them (three pairs) are shared between the two atoms, leaving 4 of them to form two nonbonding pairs, one on each nitrogen atom. N The stability of the triple bond linking N atoms in N2 gas helps explain nitrogen’s relative inertness in the troposphere.
N
or
N
N
The ozone molecule introduces another structural feature. We again start with the octet rule. Each of the three oxygen atoms contributes 6 outer electrons for a total of 18. These 18 electrons can be arranged in two ways; each way gives a share in 8 outer electrons to each atom. O O O a
O O O b
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Structures a and b predict that the molecule should contain one single bond and one double bond. In structure a, the double bond is shown to the left of the central atom; in b it is shown to the right. But experiments reveal that the two bonds in the O3 molecule are identical, being intermediate between the length and strength of a single and double bond. Structures a and b are called resonance forms, Lewis structures that represent hypothetical extremes of electron arrangements in a molecule. For example, no single resonance form represents the electron arrangement in the ozone molecule. Rather, the actual structure is something like a hybrid of the two resonance forms. A double-headed arrow linking the different forms is used to represent the resonance phenomenon. O
O
O
O
O
O
Resonance is just another modeling concept invented by chemists to represent the complex microworld of molecules. It is not intended to be “the truth” but rather just a way to describe the structures of molecules that do not exactly fit the octet rule model. Figure 2.4 compares the Lewis structures of several different oxygen-containing species relevant to the chemistry in this and other chapters.
O
O
oxygen atom
O
oxygen molecule
O
O
O
ozone molecule
O H hydroxyl free radical
Figure 2.4 Lewis structures for several oxygen-containing species. Only one resonance form of ozone is shown.
A closer experimental inspection of that microworld reveals that the O3 molecule is not linear as the simple Lewis structures just drawn seems to indicate. Remember that Lewis structures tell us only what is connected to what and do not necessarily show the shape of the molecule. The O3 molecule is actually bent, as in this representation. O
O
O
O
O
O
An explanation of why the O3 molecule is bent must wait until Chapter 3. At this point, we only need to know how the bonding in the O2 and O3 molecules relates to their interaction with sunlight.
Your Turn 2.10
Lewis Structures with Multiple Bonds
Draw the Lewis structure for each compound. Both follow the octet rule. a. carbon monoxide (CO)
b. sulfur dioxide (SO2)
Answer a.
1 C atom ( C ) 3 4 outer electrons per atom 5 4 outer electrons 1 O atom ( O ) 3 6 outer electrons per atom 5 6 outer electrons Total 5 10 outer electrons
The Lewis structure is C O or C O and has 10 outer electrons. The N2 molecule also has 10 outer electrons and similarly forms a triple bond.
The hydroxyl radical was mentioned in Chapter 1 in connection with smog formation.
Observe that both H2O and O3 are bent molecules, with O as the central atom.
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Chapter 2
2.4 |
Waves of Light
Every second, light is emitted by the Sun and, after some time, reaches our planet. Some of this light we can see; some we cannot. Prisms and raindrops break the light we see into a spectrum of colors. Sometimes we name these colors simply as violet, indigo, blue, green, yellow, orange, and red. Other times, we distinguish between the hues with more descriptive names such as cherry red or forest green. Another way to describe a color is with a numerical value that corresponds to its wavelength. The word wavelength correctly suggests that light behaves something like a wave in a body of water. Wavelength is the distance between successive peaks. It is expressed in units of length and symbolized by the Greek letter lambda (l). Waves are also characterized by a certain frequency, the number of waves passing a fixed point in 1 second. Frequency is symbolized by the Greek letter nu (n). Figure 2.5 shows two waves of different wavelength and frequency. The relationship between frequency and wavelength can be summarized in a simple equation in which n is the frequency and c is the constant speed at which visible light and other forms of electromagnetic radiation travel, 3.00 3 108 m·s−1. frequency () ⫽ As wavelength ↑, frequency ↓.
speed of light (c) wavelength ()
[2.2]
Equation 2.2 indicates that wavelength and frequency are inversely related. As the l decreases, the n increases, and vice versa. It is both interesting and humbling to realize that out of the vast array of radiant energies, our eyes are sensitive only to the tiny portion between roughly 700 3 10 −9 meters (red light) and 400 3 10 −9 meters (violet light). These wavelengths are very short, so we typically express them in nanometers. One nanometer (nm) is defi ned as one billionth of a meter (m). 1 nm ⫽
1 1 m⫽ m ⫽ 1 ⫻ 10–9 m 1,000,000,000 1 ⫻ 109
We can use this equivalence to convert meters to nanometers. For example, this calculation shows how many nanometers are in 700 3 10 −9 m. wavelength () ⫽ 700 ⫻ 10–9 m ⫻
1 nm 1 ⫻ 10–9 m
⫽ 700 nm
The units of meters cancel, leaving nanometers. , wavelength
Longer wavelength Lower frequency
Shorter wavelength Higher frequency
Figure 2.5 Comparison of two different waves.
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Consider This 2.11
77
Analyzing a Rainbow
Water droplets in a rainbow act as prisms to separate visible light into its colors. a. In Figure 2.6, which color has the longest wavelength? The highest frequency? b. Green light has a wavelength of 500 nm. Express this value in meters. Answer b. 500 3 10 −9 m. In scientific notation, this is 5 3 10 −7 m.
The electromagnetic spectrum is a continuum of waves that ranges from short, high-energy X-rays and gamma rays to long, low-energy radio waves. Visible light is only a narrow band in this spectrum. The term radiant energy refers to the entire collection of different wavelengths, each with its own energy. Figure 2.7 shows the electromagnetic spectrum, the relative wavelengths (not drawn to scale), and some examples to help you develop perspective on the range of wavelengths represented. In this chapter, we consider the ultraviolet (UV) region that lies adjacent to the violet end of the visible region of the electromagnetic spectrum, but at shorter wavelengths. At still shorter wavelengths are the X-rays used in medical diagnosis and the determination of crystal structures, and gamma rays that are given off in processes of nuclear decay. At wavelengths longer than those of red visible light lies the infrared (IR) region. We cannot see these wavelengths, but can feel their heating effect. The microwaves used in radar and to cook food quickly have wavelengths on the order of centimeters. At still longer wavelengths are the regions of the spectrum used to transmit your favorite AM and FM radio and television programs.
Figure 2.6 A rainbow of color.
We will consider the IR region of the spectrum in Chapter 3.
Your Turn 2.12
Relative Wavelengths
Consider these four types of radiant energy from the electromagnetic spectrum: infrared, microwave, ultraviolet, visible. a. Arrange them in order of increasing wavelength. b. Approximately how many times longer is a wavelength associated with a radio wave than one associated with an X-ray? Hint: See Figure 2.7. Answer a. ultraviolet , visible , infrared , microwave
wavelength () in meters 10–14
10–12
Gamma rays diameter of atomic nucleus
10–10 X-rays
10–8
UV Visible
diameter diameter of atom of virus
400
10–6
10–4 IR
10–2 Microwave
diameter of diameter of animal cell period (.)
450 500 550 600 650 wavelength () in nanometers
1
102 Radio
diameter height of height of of CD human skyscraper
700
Figure 2.7 The electromagnetic spectrum. The wavelength variation from gamma rays to radio waves is not drawn to scale.
Figures Alive! Visit the textbook’s website to learn more about relationships in the electromagnetic spectrum.
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Chapter 2 Relative intensity of radiation
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20 15 39%
10 5
53% 8%
0 0
500 UV Visible
1000
1500
2000 2500 IR Wavelength (nm)
3000
3500
4000
Figure 2.8 Wavelength distribution of solar radiation above Earth’s atmosphere. Source: From An Introduction to Solar Radiation by Muhammad Iqbal, Academic Press, 1983. Copyright Elsevier 1983.
Our local star, the Sun, emits many types of radiant energy but not with equal intensity. This is evident from Figure 2.8, a plot of the relative intensity of solar radiation as a function of wavelength. The curve represents the spectrum as measured above the atmosphere, before there has been opportunity for interaction of radiation with the molecules found in air. The peak indicating the greatest intensity is in the visible region. However, 53% of the total energy emitted by the Sun is radiated to Earth as infrared radiation. This is the major source of heat for the planet. Approximately 39% of the energy comes to us as visible light and only about 8% as ultraviolet. (The areas under the curve give an indication of these percentages.) But in spite of its small percentage, the Sun’s UV radiation can be the most damaging to living things. To understand why, we need to look at electromagnetic radiation in terms of its energy.
2.5 |
Planck and Einstein were both amateur violinists who played duets together.
Radiation and Matter
The idea that radiation can be described in terms of wave-like character is well established and very useful. However, around the beginning of the 20th century, scientists found several phenomena that seemed to contradict this model. In 1909, a German physicist named Max Planck (1858–1947) argued that the shape of the energy distribution curve pictured in Figure 2.8 could only be explained if the energy of the radiating body were the sum of many energy levels of minute but discrete size. In other words, the energy distribution is not really continuous, but consists of many individual steps. Such an energy distribution is called quantized. An often-used analogy is that the quantized energy of a radiating body is like steps on a staircase, which are also quantized (no partial steps allowed), not like a ramp that allows any size stride. Albert Einstein (1879–1955), in the work that won him the 1921 Nobel Prize in physics, suggested that radiation itself should be viewed as constituted of individual bundles of energy called photons. One can regard these photons as “particles of light,” but they are definitely not particles in the usual sense. For example, they have no mass. These ideas form the basis of modern quantum theory. The wave model remains useful, even with the development of the quantum theory. Both are valid descriptions of radiation. The dual nature of radiant energy seems to defy common sense. How can light be described in two different ways at the same time, both waves and particles? There is no obvious answer to that very reasonable question—that’s just the way nature is. The two views are linked in a simple relationship that is one of the most important equations in modern science. It is also an equation relevant to the role of ozone in the atmosphere. energy (E) ⫽
hc
[2.3]
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Here E represents the energy of a single photon. Both symbols h and c represent constants. The symbol h is called Planck’s constant and c is the speed of light. This equation therefore shows that energy, E, is inversely proportional to the wavelength, l. Consequently, as the wavelength of radiation gets shorter, its energy increases. This qualitative relationship is important in the story of ozone depletion.
Your Turn 2.13
79
As wavelength ↓, energy ↑.
Color and Energy Relationships
Arrange these colors of the visible spectrum in order of increasing energy per photon: green, red, yellow, violet. Answer red , yellow , green , violet
Using equation 2.3, one can calculate that the energy associated with a photon of UV radiation is approximately 10 million times larger than the energy of a photon emitted by your favorite radio station. A consequence of this large difference in energy is that you can damage your skin with exposure to UV radiation, but not by listening to the radio—unless you happen to be listening to it outside in the sunlight. Whether or not your radio is turned on, you are continuously bombarded by radio waves. Your body cannot detect them, but your radio can. The energy associated with each of the radio photons is very low and not sufficient to produce a local increase in the concentration of the skin pigment, melanin, as happens with exposure to UV. Producing melanin involves a quantum jump, an electronic transition between energy levels that requires far more energy than radio wave photons can supply. The Sun bombards Earth with countless photons—indivisible packages of energy. The atmosphere, the planet’s surface, and Earth’s living things all absorb these photons. Radiation in the infrared region of the spectrum warms Earth and its oceans. The cells of our retinas are tuned to the wavelengths of visible light. Photons associated with different wavelengths are absorbed, and the energy is used to “excite” electrons in biological molecules. Some electrons jump to higher energy levels, triggering a series of complex chemical reactions that ultimately lead to sight. Compared with animals, green plants capture photons in an even narrower region of the visible spectrum (corresponding to red light) and use the energy to convert carbon dioxide and water into food, fuel, and oxygen in the process of photosynthesis. Remember that as the wavelength of light decreases, the energy carried by each photon increases. Photons in the UV region of the spectrum are sufficiently energetic to displace electrons from neutral molecules, converting them into positively charged species. Even shorter UV wavelength photons break bonds, causing molecules to come apart. In living things, such changes disrupt cells and create the potential for genetic defects and cancer. The interaction of UV radiation with chemical bonds is shown schematically in Figure 2.9. It is part of the fascinating symmetry of nature that this interaction of radiation with matter explains both the damage ultraviolet radiation can cause and the atmospheric mechanism that protects us from it. We turn next to understanding the ultraviolet shield provided by oxygen and ozone in our stratosphere.
Ultraviolet Bond breaks
Figure 2.9 Ultraviolet radiation can break some chemical bonds. Bonds are represented as springs that hold the atoms together but allow the atoms to move relative to each other.
From Figure 2.8 you can see that the photons from the Sun with energies in the visible light range are the most intense.
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2.6 |
The Oxygen–Ozone Screen
We know the colors of visible light by their names, red, blue, yellow and so on. Similarly, we call ultraviolet light by different names. Admittedly, however, these names are not as colorful: UV-A, UV-B, and UV-C. UV-A lies closest to the violet region of visible light and is the lowest in energy; you may know it as “black light.” In contrast, UV-C has the highest energy and lies next to the X-ray region of the electromagnetic spectrum. Table 2.4 shows the characteristics of the different types of UV light.
Your Turn 2.14
The ABCs of Solar UV
a. Arrange UV-A, UV-B, and UV-C in order of increasing wavelength. b. Is the order for increasing energy the same as for wavelength? Explain. c. Should you use a sunscreen that claims to protect against UV-C? Explain. Answer c. No, you should not. No protection is needed for UV-C, because this set of UV wavelengths is absorbed up in the stratosphere.
As you saw from the previous activity, UV-C radiation from the Sun is absorbed in the upper atmosphere before it ever reaches the ground. Both oxygen and ozone absorb light of these wavelengths. As we noted in Chapter 1, about 21% of the atmosphere consists of oxygen, O2. Photons with energy corresponding to 242 nm or less have sufficient energy to break the bonds in an O2 molecule. These wavelengths are found in the UV-C region. O2
UV photon 艋 242 nm
[2.4]
2O
If O2 were the only molecule absorbing UV light from the Sun, Earth’s surface and the creatures that live on it would still be subjected to damaging radiation in the range of 242–320 nm. It is here that O3 plays its important protective role. The O3 molecule is more easily broken apart than O2. Recall that the atoms in the O2 molecule are connected with a double bond, but each of the bonds in O3 is somewhere between a single and double bond in length and in strength. Accordingly, the bonds in O3 are weaker than the double bonds in O2. Therefore, photons of a lower energy (longer wavelength) are sufficient to separate the atoms in O3. Indeed, photons of wavelength 320 nm or less break the O-to-O bond in ozone. O3
Table 2.4
UV photon 艋 320 nm
O2 ⫹ O
[2.5]
Types of UV Radiation
Type
Wavelength
Relative Energy
Comments
UV-A
320–400 nm
Lowest energy
Least damaging and reaches the Earth’s surface in greatest amount
UV-B
280–320 nm
Higher energy than UV-A but less energetic than UV-C
More damaging than UV-A but less damaging than UV-C. Most UV-B is absorbed by O3 in the stratosphere
UV-C
200–280 nm
Highest energy
Most damaging but not a problem because it is totally absorbed by O2 and O3 in the stratosphere
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Consider This 2.15
81
Energy and Wavelength
We just stated that it takes photons in the UV-C range (< 242 nm) to break the double bond in O2. The bonds in O3 are somewhat weaker than those in O2, so lower energy photons (< 320 nm) can break those bonds. Just how much greater is the energy of a 242-nm photon than that of a 320-nm photon? Hint: One approach could be to calculate the ratio of the energies for a 242-nm photon and that of a 320-nm photon and then to compare that with the ratio of their wavelengths. Look for values of Planck’s constant and the speed of light in Appendix 1.
Equations 2.4 and 2.5, together with earlier equation 2.1 (that showed the formation of O3 from O2), are part of a set of chemical reactions in the stratosphere. Every day, 300,000,000 (3 3 108) tons of stratospheric O3 forms, and an equal mass decomposes. New matter is neither created nor destroyed but merely changes its chemical form. So the overall concentration of ozone remains constant in this natural cycle. The process is an example of a steady state, a condition in which a dynamic system is in balance so that there is no net change in concentration of the major species involved. A steady state arises when a number of chemical reactions, typically competing reactions, balance each other. The Chapman cycle, as shown Figure 2.10, represents the first set of natural steady-state reactions proposed for stratospheric ozone. This natural cycle includes chemical reactions for both ozone formation and decomposition.
Your Turn 2.16
The Ozone Layer
a. Ozone is formed by the reaction of oxygen atoms with oxygen molecules. Write the chemical reaction. b. Up in the stratosphere, what is the source of the oxygen atoms? c. Up in the stratosphere, the lifetime of a given ozone molecule ranges from days to years. For example, in the ozone layer, an O3 molecule can persist for several months. What does stratospheric ozone break down into? d. In contrast, at ground level, ozone molecules react in a matter of minutes rather than months. Why the difference? Answer d. Down in the troposphere, the air is much denser (“thicker”). Ozone molecules quickly bump into other molecules and react with them. For example, if you happen to breathe air containing ozone pollution, the O3 quickly reacts with your lung tissue, potentially damaging it.
In a later section, we consider what happens when something disturbs the steady state of the Chapman cycle, leading to destruction of the protective stratospheric ozone.
O2
UV photons ( ⭐ 242 nm)
new O fed into cycle
2O
collisions, fast
O ⫹ O2
O3
subcycle
UV photons ( ⭐ 320 nm)
O3 ⫹ O
collisions slow
2 O2
(O3 removed from cycle)
Figure 2.10 The Chapman cycle.
This set of reactions is named after Sydney Chapman, a physicist who first proposed it in 1929.
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Because of the presence of O2 and O3 in the stratosphere, only certain UV wavelengths reach the surface of the Earth. However, these wavelengths still can cause harm, which is the topic of the next section.
2.7
For animals and plants, the effects of UV radiation depend primarily on the energy of the radiation and the sensitivity of the organism. These relationships are evident from Figure 2.11, where biological sensitivity is plotted versus wavelength. As defined here, biological sensitivity is based on experiments in which the damage to deoxyribonucleic acid (DNA), the chemical basis of heredity, is measured at various wavelengths. In the figure, the biological sensitivity is expressed in relative units using a logarithmic scale. On this scale, each mark on the y-axis represents a biological sensitivity that is 10 times the value corresponding to the mark immediately below it on the vertical axis. Biological sensitivity at 320 nm is about 1 3 10−5, or 0.00001 units. But at 280 nm, the sensitivity is 1 3 100, or 1 unit. This means that radiation at 280 nm is 100,000 times more damaging than radiation at 320 nm. However, shorter wavelength solar radiation (,320 nm) is absorbed by O3 in the stratosphere. This is most fortunate, because radiation in this region of the spectrum is particularly damaging to living things.
Consider This 2.17
Relative Biological Sensitivity
DNA sensitivity decreases with increasing UV wavelength, as shown by Figure 2.11. a. Propose an explanation for this. b. Why is there no need to include the UV-C region in this figure?
As we will see in the next section, around 1980 the average concentrations of stratospheric ozone unexpectedly began to decrease. Although this has happened to varying extents in different regions of the globe, generally speaking living things are
100 Biological sensitivity (relative values)
Look for more about DNA in Chapter 12.
|
Biological Effects of Ultraviolet Radiation
10–2 UV-B
UV-A
10–4
10–6 280
300
320
340
Absorption by O3 in this region Wavelength (nm)
Figure 2.11 Variation of biological sensitivity of DNA with UV wavelength. The UV-C region is below 280 nm. Source: John E. Fredrick, University of Chicago Reprinted by permission.
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83
30.0 27.5 25.0
White / Male White / Female Black / Male Black / Female
22.5
Rate per 100,000
20.0 17.5 15.0 12.5 10.0 7.5 5.0 2.5 0
1974 1976 1978 1980 1982 1984 1986 1988 1990 1992 1994 1996 1998 2000 2002
Year of diagnosis
Figure 2.12 Increase in incidence of melanoma skin cancer in the United States, 1973–2003. Source: Surveillance, Epidemiology, and End Results (SEER) Program of the National Cancer Institute.
now exposed to greater intensities of damaging radiation than in the past. According to calculations, a given percent decrease in stratospheric ozone is expected to increase the biological damage done by UV radiation by twice that percentage. For example, from a 6% decrease in stratospheric ozone we predict a 12% rise in skin cancer, especially the more easily treated forms such as basal cell and squamous cell carcinomas. These conditions are considerably more common among people with light-colored skin than among people whose skin is more heavily pigmented (Figure 2.12).
Skeptical Chemist 2.18
Skin Cancer in Men and Women
Did Figure 2.12 mislabel the cancer rates for white men and women? That is, should the rate be higher for women than for men? Argue the case either way and then check out the data, either from the source given or from one of your choice. Who or what did you consult and what did you learn?
At least to date, geography seems to affect skin cancer more than ozone depletion. For example, evidence links the incidence of melanomas, the most deadly form of skin cancer, with the latitude at which you live. In general, the disease becomes more prevalent as one moves south in the Northern Hemisphere. Those who endure the long nights and short days of winter are compensated by a level of skin cancer that is only about half that of those who live near the equator. Skin cancer rates are high in the Southern Hemisphere as well. For example, two of every three Australians will develop skin cancer sometime during their lifetime; as of 2005, 1600 Australians died from this form of cancer per year. The Australian government has acted in several ways to reverse this trend, including banning tanned models from all advertising media. National Skin Cancer Action Week is held each year at the start of their summer season in November, with the goals of raising awareness and urging the use of clothing and lotions for sun protection. However, geographic location is not the only factor influencing development of skin cancer. Skin cancer rates generally continue to rise in all countries, despite
According to the Australian Department of Health and Aging (2007), Australia has the highest skin cancer rates in the world.
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increased awareness of the dangers of exposure to UV radiation. Changes in the natural protection afforded by the stratospheric ozone layer are only partially responsible for higher rates of skin cancer. About a million new cases of skin cancers occur each year in the United States, almost as many as the total number of cases of all other cancers. Although those with lighter skin tones are about 40 times more likely to develop melanoma, those with darker skin have a significantly lower long-term survival rate, according to a 2005 report issued by the American Academy of Dermatology. Skin cancers can develop many years after repeated, excessive exposure has stopped. Skin cancers even have been linked to a single episode of extreme sunburn in adolescence with the effects showing up many years later. Public service campaigns center on early detection and promote regular inspection of suspicious moles for people of all ethnic backgrounds. Other possible causes need to be considered. One of these is tanning, either naturally or in a tanning bed. This practice presents a risk–benefit activity for everyone, because skin cancer can strike people of all skin tones.
Consider This 2.19 Figure 2.13 Blue Lizard Suncream.
Nanotechnology was defined in Section 1.7. Look for more about this topic in Chapter 8.
Tanning Spas
The indoor tanning industry runs a public relations campaign that highlights positive findings about indoor tanning, promoting it as part of a healthy lifestyle. Countering these claims are the studies published in scientific journals that support the view of dermatologists that there is no such thing as a “safe tan” for any skin type. Investigate at least two websites that present different points of view. List the arguments on both sides. What do you conclude?
Wearing protective sunscreen is one way to reduce the risk of skin cancer. Such products contain compounds that absorb UV-B to some extent together with others for absorbing UV-A. The American Academy of Dermatology recommends a sunscreen with a skin protection factor (SPF) of 15 to 30. But wearing a sunscreen does not mean that you are without risk from UV rays. Because sunscreens allow you to be exposed for a longer time without burning, they may ultimately cause greater skin damage. The Australian product Blue Lizard Suncream (Figure 2.13) uses “smart bottle” technology for containing and marketing their product. The bottle itself changes color from white to blue in UV light, sending an extra reminder that the dangers of UV light are still present, even if a sunscreen is being used. Wearing protective sunblock is another. These products physically block the light from reaching your skin, much as tightly woven clothing would. Sunblocks reflect the incident light; some absorb UV as well. A familiar example may be the white opaque cream used by lifeguards (“lifeguard nose”) at a pool or beach. This cream contains small white particles of ZnO (zinc oxide) or TiO2 (titanium dioxide) (or both) and has a long-term track record of safety. However, the “see-through” formulations of ZnO and TiO2 that contain these compounds in nanoparticle form are more controversial. Because the particles of ZnO and TiO2 are so microscopically tiny, they do not scatter light. As a result, the cream is clear rather than opaque, definitely a plus to those who wear it. These nanoparticle products spread more evenly, are cost-effective, and are extremely effective at both absorbing and reflecting UV radiation. However, nanoparticles may present a risk if they penetrate the skin; the verdict is not yet in. As of 2009, both consumers and government agencies were calling for further studies to better quantify the risks. Controversial or not, sunscreens play an important role in protecting us from UV radiation. For example, the U.S. National Weather Service issues an Ultraviolet Index forecast that appears on the web. UV Index values range from 0 to 15 and are based on how long it takes for skin damage to occur (Table 2.5). The UV Index is color-coded for ease of interpretation. Specific steps are suggested to protect eyes and skin from
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Table 2.5 Exposure Category LOW
The UV Index Index ,2
Tips to Avoid Harmful Exposure to UV Wear sunglasses on bright days. Be aware that snow and water can reflect the Sun’s rays. If you burn easily, cover up and use sunscreen.
MODERATE
3–5
Take precautions, such as covering up, if you will be outside. Stay in shade near midday when the Sun is strongest.
HIGH
6–7
Protect against sunburn. Reduce time in the Sun between 10 AM and 4 PM . Cover up, wear a hat and sunglasses, and use sunscreen.
VERY HIGH
8–10
Take extra precautions against sunburn, as unprotected skin will be damaged and can burn quickly. Minimize Sun exposure between 10 AM and 4 PM . Otherwise, seek shade, cover up, wear a hat and sunglasses, and use sunscreen.
EXTREME
111
Take all precautions against sunburn, as unprotected skin can burn in minutes. Be aware that white sand and other bright surfaces reflect UV and will increase UV exposure. Avoid the Sun between 10 AM and 4 PM . Seek shade, cover up, wear a hat and sunglasses, and use sunscreen.
Source: U.S. EPA, 2009.
Sun damage. As you might guess, one is to not expose yourself in the first place. Stay in the shade and wear protective clothing!
Consider This 2.20
UV Index Forecasts
The UV Index indicates the amount of UV radiation reaching Earth’s surface at solar noon, the time when the Sun is highest in the sky. a. The UV Index depends on the latitude, the day of the year, time of day, amount of ozone above the city, elevation, and the predicted cloud cover. How is the UV Index affected by each of these? b. The U.S. EPA provides a UV Index forecast. The textbook’s website provides a direct link. Suggest reasons for values that you see on today’s map.
Although the UV Index focuses on skin damage, this is not the only biological effect of UV radiation. Your eyes can be damaged as well. For example, all people, no matter what the pigmentation level of their skin, are susceptible to retinal damage caused by UV exposure. Another effect is cataracts, a clouding of the lens of the eye caused by excessive exposure to UV-B radiation. It has been estimated that a 10% decrease in the ozone layer could create up to 2 million new cataract cases globally. However, just as proper clothing and sunscreen can cut down on skin damage, wearing optical-quality sunglasses capable of blocking at least 99% of UV-A and UV-B can protect your eyes. Learn more about sunglasses in this next activity.
Consider This 2.21
Protecting Your Eyes
Sunglasses make far more than a fashion statement. They offer protection from harmful UV rays. Check out several manufacturers to learn the virtues of their products. What activities especially require good UV eye protection? Report on your findings.
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Other species on the globe are affected by UV radiation as well. For example, increases in UV radiation can harm young marine life, such as floating fish eggs, fish larvae, juvenile fish, and shrimp larvae. Experimental evidence also exists for damage to the DNA of the eggs of Antarctic ice fish. Plant growth is suppressed by increased UV radiation, and experiments have measured the negative effect that increased UV-B radiation has on phytoplankton. These photosynthetic microorganisms live in the oceans, where they occupy a fundamental niche in the food chain. Any significant decrease in their number could have a major effect globally. In 1999, an international panel of scientists confirmed that exposure to elevated levels of UV-B affected both the movement of phytoplankton up and down in water as well as their ability to move through the water. If these microorganisms are not able to achieve proper position in the water, they cannot carry out photosynthesis effectively. Moreover, these tiny plantlike organisms play an important role in the carbon dioxide balance of the planet by absorbing atmospheric carbon dioxide. Given the harmful effects of too much UV radiation, you now can see why the decreasing stratospheric ozone concentrations observed in the 1980s set off planetary alarm bells. The next two sections tell the tale of the ozone hole and how it unexpectedly appeared on our planet.
2.8
|
Stratospheric Ozone Destruction: Global Observations and Causes
Switzerland holds the record for the longest continuous set of ozone level measurements. Since 1926, stratospheric O3 concentrations have been measured at the Swiss Meteorological Institute. More recently, starting in 1979, satellite-mounted detectors have been beaming down data on ozone levels at many points. These measurements show both that the natural concentration of stratospheric O3 is not uniform across the globe and that the levels have changed over time. On average, the total O3 concentration is higher the closer one gets to either pole, with the exception of the seasonal ozone “hole” over the Antarctic. The formation of ozone in the Chapman cycle is triggered when an O2 molecule absorbs a photon of UV-C light, splitting into two O atoms. These in turn react with O2 to form O3. O2
UV-C ) 艋 242 nm)
O ⫹ O2
In conjunction with the Earth’s energy balance and global warming, look for information about solar irradiance in Section 3.9.
Recall that a Dobson unit corresponds to about one ozone molecule for every billion molecules and atoms of air.
2O
[2.6a]
O3
[2.6b]
Therefore, ozone production increases with the intensity of the radiation striking the stratosphere, which in turn depends primarily on the angle of the Earth with respect to the Sun and the distance between the Sun and the Earth. At the equator, the period of highest intensity occurs at the equinox (March and October) when the Sun is directly overhead. Outside the tropics, the Sun is never directly overhead, so the maximum intensity occurs at the summer solstice (June in the Northern Hemisphere, December in the Southern). The angle of Earth with respect to Sun dominates both ozone production and the seasons. There is a slight (~7%) increase, however, in solar power reaching Earth in early January, when the Earth is nearest the Sun, compared with July, when the Earth is farthest away. In addition, the amount of radiation emitted by the Sun changes over an 11- to 12-year cycle related to sunspot activity. This variation also influences O3 concentrations, but only by a percent or two. The wind patterns in the stratosphere cause other variations in ozone concentrations, some on a seasonal basis and others over a longer cycle. To further complicate matters, seemingly random fluctuations often occur. Extraordinary images of the Earth, such as the one that opens this chapter, are color-coded to show stratospheric ozone concentrations. The dark blue and purple regions indicate where the lowest concentrations of O3 are observed. Total ozone levels above Earth’s surface are expressed in Dobson units (DU). A value of 250–270 DU is
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220 Antarctic ozone minimum (60⬚–90⬚ S) 200
194
195
1979–1992 Nimbus 7 TOMS 1993–1994 Meteor 3 TOMS 1995 (no TOMS in orbit) 1996–2004 Earth Probe TOMS 2005–2007 OMI
192
Minimum ozone (DU)
180
177 162
160 154 140
144
140
131 124
120
111 105 109
100
107
103
108
102 102 97
99 94 86
80 1980
1985
1990
1995 Year
89 91
2000
91
2005
94 83
2010
Figure 2.14 Minima in spring stratospheric ozone in Antarctica, 1979–2009. The number is the minimum reading in Dobson units. TOMS (Total Ozone-Measuring Spectrometer) and OMI (Ozone Monitoring Instrument) are analytical instruments. Source: NASA.
typical at the equator. As one moves away from the equator, values range between 300 and 350 DU, with seasonal variations. At the highest northern latitudes, values can be as high as 400 DU. Of special interest is the thinning of ozone (the “ozone hole”) that occurs seasonally over the South Pole. Indeed, these changes were so pronounced that, when the British monitoring team at Halley Bay in Antarctica first observed it in 1985, they thought their instruments were malfunctioning! The area over which ozone levels are reported to be less than 220 DU is usually considered to be the “hole.” From the mid1990s on, the annual size of the ozone hole has nearly equaled the total area of the North American continent, and in some cases exceeded it. Check out the dramatic decline in stratospheric ozone levels observed near the South Pole shown in Figure 2.14. In recent years, the minimum has been around 100 DU. Keep in mind that seasonal variation has always occurred in ozone concentration over the South Pole, with a minimum in late September or early October—the Antarctic spring. Unprecedented, however, is the striking decrease in this minimum that has been observed in recent decades.
Consider This 2.22
This Year’s Ozone Hole
Starting in September, citizens and scientists alike examine the data for the ozone hole over Antarctica. What is happening this year? NASA posts the data, and the textbook’s website provides a link. For the most recent year: a. What is the area of the hole? How does this compare with recent years? b. What is the lowest reading observed for ozone? Again, how does this compare?
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Free radicals appear in several other contexts. Chapter 1: ?OH, formation of NO2 and then tropospheric ozone Chapter 6: ?OH, formation of SO3 in acid rain Chapter 7: ?OH and H2O?1, damage to cells by nuclear radiation Chapter 9: R?, polymerization of ethylene
The air pollutant NO (nitrogen monoxide or nitric oxide) is highly reactive, which is one of the reasons it causes harm to your lungs. In contrast, N2O (nitrous oxide or “laughing gas”) is a very stable molecule that persists for decades.
The major natural cause of ozone destruction, wherever it takes place around the globe, is a series of reactions involving water vapor and its breakdown products. The great majority of the H2O molecules that evaporate from the oceans and lakes fall back to Earth’s surface as rain or snow. But a few molecules reach the stratosphere, where the H2O concentration is about 5 ppm. At this altitude, photons of UV radiation trigger the dissociation of water molecules into hydrogen (H ( and hydroxyl ( OH) free radicals. A free radical is a highly reactive chemical species with one or more unpaired electrons. An unpaired electron is often indicated with a dot: H2O
photon
H ⫹ OH
[2.7]
Because of the unpaired electron, free radicals are highly reactive. Thus, the H? and OH radicals participate in many reactions, including some that ultimately convert O3 to O2. This is the most efficient mechanism for destroying ozone at altitudes above 50 km. Water molecules and their breakdown products are not the only agents responsible for natural ozone destruction. Another is the free radical ?NO, also called nitrogen monoxide. Most of the ?NO in the stratosphere is of natural origin. It is formed from nitrous oxide, N2O, a naturally occurring compound of nitrogen that is produced in the soil and oceans by microorganisms and gradually drifts up to the stratosphere. Although N2O is quite stable in the troposphere, up in the stratosphere it can react with O atoms to produce ?NO. Little can or should be done to control this process. It is part of a natural cycle involving compounds of nitrogen, as we will see in Chapter 6.
Your Turn 2.23
Free Radicals
a. Draw the Lewis structure for the ? OH free radical. The unpaired electron goes on the O atom. b. Draw the Lewis structure for the ?NO free radical. The unpaired electron goes on the N atom. Note: In Chapter 1, in the context of air pollution, we wrote simply NO. c. In contrast, N2O has no unpaired electrons. Draw its Lewis structure. Hint: Place one of the N atoms in the middle.
Human activities also can alter steady-state concentrations of NO. In the 1970s, people became concerned about the increase in NO concentration that would result from building a fleet of Concorde SSTs (supersonic transports). These planes were designed to fly at an altitude of 15–20 km, the region of the ozone layer. As you learned in Chapter 1, hot engines produce NO as part of the exhaust gas stream. The NO is produced by the hot engines of a jet plane on takeoff, landing, and during flight. N2 ⫹ O2
high temperature
2 NO
[2.8]
Scientists carried out experiments and calculations to predict the effects of a fleet of SSTs. They concluded that the risks would outweigh the benefits. So people decided, partly on scientific grounds, not to build an American fleet of SSTs. Until 2003, the Anglo–French Concorde was the only commercial plane that operated at this altitude. The Concorde took its last flight on October 24, 2003. Both safety concerns and economic factors played roles in ending the flights of these remarkable jets. Even when the effects of water, nitrogen oxides, and other naturally occurring compounds are included in stratospheric models, the measured ozone concentration is still lower than predicted. Measurements worldwide indicate that the ozone concentration has been decreasing over the past 20 years. There is a good deal of fluctuation in the data, but the trend is clear. The stratospheric ozone concentration at midlatitudes (60° south to 60° north) has decreased by more than 8% in some cases. These changes cannot be correlated with changes in the intensity of solar radiation, so we must look elsewhere for an explanation. Thus it is time to turn our attention to chlorofluorocarbons.
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2.9
|
89
Chlorofluorocarbons: Properties, Uses, and Interactions with Ozone
A major cause of stratospheric ozone depletion was uncovered through the masterful scientific sleuthing of F. Sherwood Rowland, Mario Molina, and Paul Crutzen. For their work, the trio jointly won the 1995 Nobel Prize in chemistry. They analyzed vast quantities of atmospheric data and studied hundreds of chemical reactions. As with most scientific investigations, some uncertainties remained. Nonetheless, their evidence all pointed to an unlikely group of compounds: the chlorofluorocarbons. As the name implies, chlorofluorocarbons (CFCs) are compounds composed of the elements chlorine, fluorine, and carbon (but do not contain the element hydrogen). Fluorine and chlorine are members of the same chemical group, the halogens (Figure 2.15). In their elemental state, all of the halogens are diatomic molecules, but only fluorine and chlorine are gases. Fluorine is not shown in the figure because it is so reactive that it would react with the glass vessel. In contrast, CFCs are highly unreactive. To get started with CFCs, let’s examine two examples. F Cl
C
Bromine
F Cl
Cl
and
C
F
Cl
Cl
CCl3F trichlorofluoromethane Freon-11
CCl2F2 dichlorodifluoromethane Freon-12
Note how the names show the connection of CFCs to methane, CH4. The prefixes di- and tri- specify the number of chlorine and fluorine atoms that substitute for hydrogen atoms of methane. These two CFCs also are known by their trade names, Freon-11 and Freon-12. You may also hear them called CFC-11 and CFC-12, respectively, following a naming scheme developed in the 1930s by chemists at DuPont. CFCs do not occur in nature; we humans synthesized them for a variety of uses. This is an important verification point in the debate over the role of CFCs in stratospheric ozone depletion. As we saw in the previous section, other contributors to the destruction of ozone, such as the ?OH and ?NO free radicals, are formed in the atmosphere both from natural sources and human activities. Rightly, the introduction of CCl2F2 as a refrigerant gas in the 1930s was hailed as a great triumph of chemistry and an important advance in consumer safety. It replaced ammonia or sulfur dioxide, two toxic and corrosive refrigerant gases. In many respects, CCl2F2 was (and still is) an ideal substitute. It is nontoxic, odorless, colorless, and does not burn. In fact, the CCl2F2 molecule is so stable that it does not react with much of anything! Given the desirable nontoxic properties of CFCs, they soon were put to other uses. For example, CCl3F was often blown into polymer mixtures to make foams for cushions and foamed insulation. Other CFCs served as propellants in aerosol spray cans and as nontoxic solvents for oil and grease. Halons are close cousins of CFCs. Like them, halons are inert, nontoxic compounds that contain chlorine or fluorine (or both, but no hydrogen). But in addition, they contain bromine. For example, here is the Lewis structure for bromotrifluoromethane, CBrF3, also known as Halon-1301. Br F
C F
F
Chlorine
Iodine
Figure 2.15 Selected elements from Group 7A, the halogen family.
Methane, the smallest hydrocarbon, was described in Chapter 1.
For more about Freons and how they are named, see end-of-chapter question #53.
Polymers and plastics are the topic of Chapter 9. Gases that “puff up” the plastic, making it into a foam, are called blowing agents.
You will find different definitions for the chemical composition of halons, depending on where you look. Sometimes halons are defined by their use (in suppressing fires) rather than by their chemical composition.
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In the United States, CFCs helped spur the growth of cities with hot, humid weather, including Atlanta, Houston, Tampa, and Memphis.
Halons are used as fire suppressants. They are especially helpful when a fire hose or sprinkler system would be inappropriate, for instance in libraries (especially rare book rooms), grease fires (where water might spread the fire), chemical stockrooms (where some chemicals react with water), and aircraft (where hosing down the cockpit would definitely be a bad idea). For better or worse, the synthesis of CFCs has had a major effect on our lives. Because CFCs are nontoxic, nonflammable, cheap, and widely available, they revolutionized air conditioning, making it readily accessible for homes, office buildings, shops, schools, and automobiles. Beginning in the 1960s and 1970s, CFCs helped to spur the growth of cities in hot and humid parts of the world. In effect, a major demographic shift occurred because of CFC-based technology that transformed the economy and business potential of entire regions of the globe. Ironically, the very property that makes CFCs ideal for so many applications— their chemical inertness—ended up doing harm to our atmosphere. The C–Cl and C–F bonds in the CFCs are so strong that the molecules are virtually indestructible. For example, it has been estimated that an average CCl2F2 molecule can persist in the atmosphere for 120 years before it meets some fate that decomposes it. In contrast, it only takes about five years for atmospheric wind currents to bring molecules up to the stratosphere, which is exactly where some of the CFC molecules ended up. In 1973, Rowland and Molina, motivated largely by intellectual curiosity, set out to study the fate of stratospheric CFC molecules. They understood that with increasing altitude, the concentrations of oxygen and ozone decrease, but the intensity of UV radiation increases. They reasoned photons of high-energy UV-C light (,220 nm) would break C–Cl bonds. Here is the chemical reaction that releases chlorine atoms from dichlorodifluoromethane. Cl F
C
Cl Cl
UV photon 艋 220 nm
F
F
The Br? free radical undergoes a comparable reaction, starting another cycle of ozone destruction. Br? is up to 10 times more effective than Cl? in destroying O3.
C ⫹ Cl
[2.9]
F
A chlorine atom has seven outer electrons, one of them unpaired. We depict it as Cl? or ?Cl to emphasize this unpaired electron. The chlorine atom exhibits a strong tendency to achieve a stable octet by combining and sharing electrons with another atom. Rowland and Molina and subsequent researchers hypothesized that this reactivity would result in a series of reactions. Although CFCs are known to destroy stratospheric ozone via several pathways, we illustrate with a typical one known to take place in polar regions. First, the Cl? free radical pulls an oxygen atom away from an O3 molecule to form ClO ?, chlorine monoxide and an O2 molecule. The coefficient 2 is not canceled because we anticipate using it the next step. 2 Cl ⫹ 2 O3
2 ClO ⫹ 2 O2
[2.10]
The ClO ? species is another free radical; it has 13 outer electrons (7 1 6). Recent experimental evidence indicates that 75–80% of stratospheric ozone depletion involves joining two ClO? radicals to form ClOOCl. 2 ClO
ClOOCl
[2.11]
In turn, ClOOCl decomposes in a two-step sequence. ClOOCl
UV photon
ClOO
ClOO ⫹ Cl Cl ⫹ O2
[2.12a] [2.12b]
We can treat this set of chemical equations as if they were mathematical equations. If we add them together, here is the result. 2 Cl ⫹ 2 O3 ⫹ 2 ClO ⫹ ClOOCl ⫹ ClOO 2 ClO ⫹ 2 O2 ⫹ ClOOCl ⫹ ClOO ⫹ Cl ⫹ Cl ⫹ O2
[2.13]
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Just as is done with mathematical equations, we can eliminate the duplicate Cl?, ClO ?, and ClOOCl species from both sides of the chemical equation. The terms for O2 on the right side of the equation, 2 O2 and O2, can be combined into 3 O2. What remains is the net equation showing the conversion of ozone into oxygen gas. 2 O3
3 O2
[2.14]
Thus, the complex interaction of ozone with atomic chlorine provides a pathway for the destruction of ozone. Notice that Cl? appears both as a reactant in equation 2.13 and as a product in equations 2.12a and 2.12b. This means that Cl? is both consumed and regenerated in the cycle, with no net change in its concentration. Such behavior is characteristic of a catalyst, a chemical substance that participates in a chemical reaction and influences its speed without undergoing permanent change. Atomic chlorine acts catalytically by being regenerated and recycled to remove more ozone molecules. On average, a single atom can catalyze the destruction of as many as 1 3 105 ozone molecules before it is carried back to the lower atmosphere by winds. Interestingly, the mechanism just described for ozone destruction by CFCs in the stratosphere was not the one first proposed by Rowland and Molina. Their initial hypothesis was that Cl? reacted with O3 to form ClO? and O2. The second step proposed was that ClO? reacted with oxygen atoms to form O2 and regenerate radicals. Cl ⫹ O3
ClO ⫹ O2
[2.15]
Cl ⫹ O2
[2.16]
ClO ⫹ O
Although this mechanism did not prove to be the major one in the formation of the ozone hole, it did provide a reasonable explanation for why recycling a limited number of chlorine atoms could be responsible for the destruction of a large number of ozone molecules. But it is the cycle that primarily accounts for the destruction of ozone in the tropical and midlatitudes, regions in which the incident sunlight is more intense. As is often true in science, hypotheses need to be recast in light of experimental evidence. Thankfully, almost all of the chlorine in the stratosphere is not in the active form of Cl? or ClO?. Rather, chlorine is incorporated into stable compounds that do not destroy ozone. Hydrogen chloride (HCl) and chlorine nitrate (ClONO2) are two such compounds. These form quite readily at altitudes below 30 km. Thus, chlorine atoms are fairly effectively removed from the region of highest ozone concentration (about 20–25 km). These gases, as we will see in Chapter 5, are water-soluble. So in the troposphere, they are removed from the air when they wash out in the rain.
Your Turn 2.24
Bromine, too!
Although we have been casting the discussion in terms of chlorine atoms, bromine atoms also play a role. a. Write chemical reactions involving bromine analogous to equations 2.10 and 2.15. b. Bromine concentrations are much lower than those of chlorine. Propose a reason why. Answer b. Fewer of the substances that deplete ozone contain the element bromine. These compounds, such as CBrF3 (Halon-1301) have been manufactured in smaller amounts.
Rowland, a professor at the University of California at Irvine, and Molina, then a postdoctoral fellow in Rowland’s laboratory, published their first paper on CFCs and ozone depletion in 1974 in the scientific journal Nature. At about the same time, other scientists were obtaining the first experimental evidence of stratospheric ozone depletion and CFCs in the stratosphere. The conclusions were troublesome; the implications were that the use of CFCs should be discontinued. These initial reports were met with
The term catalyst also was mentioned in Section 1.11 in connection with catalytic converters.
Although HCl and ClONO2 do not destroy ozone, they still are potential sources of Cl?. For example, HCl can react with the hydroxyl radical (?OH) to produce Cl?.
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1.5
Stratospheric ozone (parts per billion)
2500
Stratospheric ozone
2000
1.0
1500 1000
0.5
Stratospheric chlorine (parts per billion)
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500 Stratospheric chlorine Dr. Susan Solomon, a chemist, headed the team that first gathered stratospheric ClO? and ozone data over Antarctica. The data solidified the connection between CFCs and the ozone hole. She was just 30 years old at the time.
0 63
64
65
66
67 68 69 Degrees south latitude
70
71
72
Figure 2.16 Antarctic stratospheric concentrations of ozone and reactive chlorine (from a flight into the Antarctic ozone hole, 1987). Source: United Nations Environment Programme.
skepticism, as might be expected when much was at stake economically. But the precautionary principle ultimately prevailed. Action was taken to mitigate the loss of ozone before even worse ozone destruction took place. Over the years, the correctness of the Rowland–Molina hypothesis has been well established. Perhaps the most compelling evidence for the involvement of chlorine and chlorine monoxide is presented in Figure 2.16. It shows two sets of data from the Antarctic, one for O3 concentration and the other for ClO?. Both are plotted versus the latitude at which samples were measured. As stratospheric O3 concentration decreases, the concentration increases; the two curves mirror each other almost perfectly. The major effect is a decrease in ozone and an increase in chlorine monoxide as the South Pole is approached. Because equation 2.10 links ClO?, Cl?, and O3, the conclusion is compelling. Figure 2.16 is sometimes described as the “smoking gun” for stratospheric ozone depletion. Not all of the chlorine implicated in stratospheric ozone destruction comes from CFCs. Other chlorinated carbon compounds come from natural sources, such as sea water and volcanoes. However, most chlorine from natural sources is in water-soluble forms. Therefore, any natural chlorine-containing substances are washed out of the atmosphere by rainfall long before they can reach the stratosphere. Of particular significance are the data gathered by NASA and by international researchers that establish that high concentrations of HCl (hydrogen chloride) and HF (hydrogen fluoride) always occur together. Although some of the HCl might conceivably arise from a variety of natural sources, the only reasonable origin of stratospheric concentrations of HF is CFCs.
Consider This 2.25
Radio Talk Show Opinions
“And if prehistoric man merely got a sunburn, how is it that we are going to destroy the ozone layer with our air conditioners and underarm deodorants and cause everybody to get cancer? Obviously we’re not . . . and we can’t . . . and it’s a hoax. Evidence is mounting all the time that ozone depletion, if occurring at all, is not doing so at an alarming rate.”* Consider the first thing you would ask this talk-show host about these statements. Remember that you need to formulate a short and focused question to get any airtime! *Source: Limbaugh, R. 1993. See, I Told You So. New York: Pocket Books.
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2.10
|
The Antarctic Ozone Hole: A Closer Look
Ozone-depleting gases are present throughout the stratosphere. Furthermore, as a result of global wind patterns, CFCs are present in comparable abundance in lower parts of the atmosphere over both hemispheres. Why, then, have the greatest losses of stratospheric ozone occurred over Antarctica? And given that more ozone-depleting gases are emitted in the Northern Hemisphere, why are their effects felt most strongly in the Southern Hemisphere? A special set of conditions exist in Antarctica, ones that relate to the fact that the lower stratosphere over the South Pole is the coldest spot on Earth. From June to September (Antarctic winter), the winds that circulate around the South Pole form a vortex that prevents warmer air from entering the region. As a result, the temperature may drop as low as 290 °C. Under these conditions, polar stratospheric clouds (PSCs) can form. These thin clouds are composed of tiny ice crystals formed from the small amount of water vapor present in the stratosphere. The chemical reactions that occur on the surface of these ice crystals convert molecules that do not deplete ozone, such as ClONO2 and HCl that we mentioned previously, to the more reactive species that do: HOCl and Cl2. Neither HOCl nor Cl2 causes any harm in the dark of winter. But when sunlight returns to the South Pole in late September, the light splits HOCl and Cl2 to release chlorine atoms. Given this increase in Cl?, a species that destroys vast quantities of ozone, the hole starts to form. Notice the conditions required: extreme cold, a circular wind pattern (vortex), enough time for ice crystals to form and provide a surface for the reactions, and darkness followed by rapidly increasing levels of sunlight. Figure 2.17 shows the seasonal variation and compares the minimum temperatures above the Arctic and the Antarctic. As you can see, the necessary conditions more often are found in Antarctica.
Arctic Winter Nov
Dec
Jan
Feb
March
April –80
–65
40 to 90 Latitude
Temperature (degrees Celsius)
–100
Arctic
–75
PSC formation temperature
–80
–110
–120
–85 Antarctic
–90 –95 –100 –105
May
–130
Range of Values
–140
Average winter value Arctic 1978–79 to 2001–2002 Antarctic 1979 to 2001
–150
June
July
August
Sep
Temperature (degrees Fahrenheit)
–90 –70
Oct
Antarctic Winter
Figure 2.17 Minimum air temperatures in the polar lower stratosphere. Polar stratospheric clouds (PSCs) are thin clouds of ice crystals that form at very low temperatures. Source: Scientific Assessment of Ozone Depletion: 2002, World Meteorological Organization, United Nations Environment Programme.
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Changes in ozone above Antarctica closely follow the seasonal temperatures. Typically a rapid ozone decline takes place during spring at the South Pole, that is, from September to early November. As the sunlight warms the stratosphere, the polar stratospheric clouds dissipate, halting the chemistry that occurs on the surfaces of their ice crystals. Then air from lower latitudes flows into the polar region, replenishing the depleted ozone levels. By the end of November, the hole is largely refilled. Although the deepest decrease in the ozone layer over Antarctica occurs during the spring, recent discoveries by British Antarctic Survey researchers indicate that the ozone depletion may begin earlier, as early as midwinter at the edges of the Antarctic, including over populated southern areas of South America. This situation presents us with another example of the tragedy of the commons—a resource that is common to all and used by many, but has no one in particular responsible for it. As a result, the resource may be harmed to the detriment of all. Here, the resource is our protective ozone layer. Decreased stratospheric ozone over the South Pole leads to increased UV-B levels reaching the Earth. In turn, skin cancer rates increase in Australia and southern Chile. Australian scientists believe that wheat, sorghum, and pea crop yields have decreased a result of increased UV radiation. Similar effects are also being felt in southern Chile in the area around Punta Arenas, and on the island of Tierra del Fuego at the southernmost tip of South America. Chile’s health minister has Figure 2.18 warned the 120,000 residents of Punta Arenas not to Arctic polar stratospheric clouds in the northern part of Sweden. be out in the Sun during the noon hours in the spring, Photo credit: Ross J. Salawitch, University of Maryland. when ozone depletion is greatest. It turns out that the depletion in the Northern Hemisphere is not nearly as severe In the context of air quality, the as it is in the Southern. The difference stems mainly from the fact that the air above tragedy of the commons was first the North Pole is not as cold. Even so, polar stratospheric clouds have been repeatedly mentioned in Chapter 1. observed in the Arctic. For example, the “mother-of-pearl” polar stratospheric clouds shown in Figure 2.18 were photographed above Porjus, a village in Swedish Lapland. The colors in PSCs are caused by However, these clouds do not lead to the formation of an ozone hole, as the air trapped diffraction of light by the ice particles over the Arctic generally begins to diffuse out of the region before the Sun gets bright in the clouds. enough to trigger as much ozone destruction as has been observed in Antarctica.
2.11
| Responses to a Global Concern
Once people understood the role of CFCs in ozone destruction, the response was surprisingly rapid. Some of the first steps toward reversing ozone depletion were taken by individual countries. For example, the use of CFCs in spray cans was banned in the United States and Canada in 1978; their use as foaming agents for plastics was discontinued in 1990. The problem of CFC production and subsequent release, however, was a global one. As such, it required international cooperation. In 1977, in response to growing experimental evidence, the United Nations Environment Programme (UNEP) convened a meeting. Those in attendance adopted a World Plan of Action on the Ozone Layer and established a coordinating committee to guide future international actions. In 1985, world leaders participated in the Vienna Convention on the Protection of the Ozone Layer. Through action taken at the convention, the nations represented committed themselves to protecting the ozone layer and to conducting scientific research to better understand atmospheric processes. A major
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breakthrough came in 1987 with the signing of a treaty: the Montreal Protocol on Substances That Deplete the Ozone Layer. Each nation that signed the treaty then needed to ratify it. Many did so immediately, but it was not until 2009 until the last four nations joined the other 182 nations who had ratified the treaty.
Consider This 2.26
Graffiti with a Message
a. This cartoon dates from the mid-1970s. Explain the basis of its humor. b. Is this cartoon still relevant to the problem of ozone depletion today? Explain. c. Create a cartoon of your own that deals with ozone depletion. Be sure that the chemistry is correct!
The key initial strategy for reducing chlorine in the stratosphere was to stop production of CFCs. The United States and 140 other countries agreed to a complete halt in CFC manufacture after December 31, 1995. Figure 2.19 indicates that the decline in global CFC production has been dramatic. By 1996, production of CFCs had fallen to 1960 levels. Without the international action required by the Montreal Protocol, stratospheric abundances of chlorine found in the stratosphere could have tripled by the middle of the 21st century. The Protocol included a provision to hold future meetings to revise goals as new scientific knowledge evolved. These meetings turned out to be of key importance, because all soon agreed—atmospheric scientists, environmentalists, chemical manufacturers, and government officials—that the Montreal Protocol was not stringent enough. Subsequent meetings have been held in locations worldwide. Currently, the goal is to phase out not just CFCs, but also 96 other substances that deplete ozone!
1400 Montreal Protocol signed, 1987
Thousand tons of CFCs
1200 First ozone depletion reported, 1974
1000 800 600 400 200 0 1950
1955
1960
1965
1970
1975 Year
Figure 2.19 Global production of CFCs, 1950–2004. Source: United Nations Environment Programme (UNEP).
1980
1985
1990
1995
2000
2005
95
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Effective stratospheric chlorine (parts per trillion)
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2000 1500 1000 500 0 1950
2000
2050
2100
Year
Figure 2.20 Concentrations of effective chlorine, 1950–2100. The height of the yellow band for any year is an estimate of the uncertainty in the prediction. Source: Scientific Assessment of Ozone Depletion: 2002, World Meteorological Organization, United Nations Environment Programme.
Inhalers no longer use CFCs as propellants.
In 2007, the 20th Anniversary of the Montreal Protocol found cause for much celebration. Starting in the 1990s, a dramatic decrease in global production of CFCs occurred, as shown in Figure 2.19. Over the course of 20 years, CFC production was halted, their use gradually phased out by developed nations, and a complete phase out occurred in 2010. In 2008, one of the last uses to be eliminated was as a propellant in inhalers, such as those used by people with asthma. Similarly, halons were phased out stepwise. Stopping production of CFCs and restricting their uses did not immediately translate into a drop in the stratospheric concentration of chlorine. Most Earth systems, including those in the atmosphere, are complex and slow to respond to changes. In fact, the atmospheric concentrations of ozone-depleting gases continued to rise steadily through the 1990s, despite the restrictions of the Montreal Protocol and its subsequent amendments. One reason for the slow rate of change is that many of the CFCs have long atmospheric lifetimes, some estimated to be over 100 years. Even so, the signs are encouraging that the Montreal Protocol has had a beneficial effect. Decreases are now being observed in the amount of effective stratospheric chlorine, a term reflecting both chlorine- and bromine-containing gases in the stratosphere. The values take into account the greater effectiveness but lower concentration of bromine relative to chlorine in depleting stratospheric ozone. Figure 2.20 shows one prediction of the future abundance of effective chlorine. Analysis of stratospheric chlorine levels indicates that it peaked in the late 1990s and then started slowly decreasing. But we are not yet in the clear. Scientists estimate that even with the most stringent international controls on the use of ozone-depleting chemicals, the stratospheric chlorine concentration would not drop to 2 ppb (2000 ppt, parts per trillion) for some years to come. That concentration is significant because the Antarctic ozone hole first appeared when effective stratospheric chlorine reached that level.
Consider This 2.27
Past and Future Effective Chlorine Levels
Use Figures 2.19 and 2.20 to answer these questions. a. In approximately what year did effective chlorine concentration peak? What was the reading in that year? b. Is the peak year for effective chlorine concentration the same as the peak year for CFC production? Why or why not? c. According to predictions, in approximately what year will the effective chlorine level return to 1980 levels? What will the reading be in that year?
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Although the Montreal Protocol and its subsequent adjustments set dates for the halt of all CFC production, the sale of existing stockpiles and recycled materials will remain legal for some time into the future. This is necessary because appliances that were designed to operate using CFCs still remain in use. For example, air conditioning units produced in the United States for cars and homes before 1996 were designed to run on CFCs. But as you might surmise, both the paperwork and the price of legally obtained CFCs have risen sharply. As of 2010, the production of CFCs is banned for all signatories of the Montreal Protocol, with the possible exception of a small number of essential uses, such as those in medicine.
2.12
| Replacements for CFCs
No one seriously advocates returning to ammonia and sulfur dioxide in home refrigeration units or giving up air conditioning as solutions to necessary restrictions on CFCs. Instead, chemists are synthesizing new compounds, concentrating on those similar to the CFCs but without their problematic long-term effects on stratospheric ozone. Substitute molecules might include one or two carbon atoms, at least one hydrogen atom, and fluorine or chlorine atoms. The rules of molecular structure limit the options. In synthesizing substitutes for CFCs, chemists initially weighed three undesirable properties—toxicity, flammability, and extreme stability—and attempted to achieve the most suitable compromise. For example, introducing hydrogen atoms in place of one or more of the halogen atoms reduces molecular stability and promotes destruction of the compounds at low altitudes, long before they enter the ozone-rich regions of the atmosphere. However, too many hydrogen atoms increase flammability. Moreover, if a hydrogen atom replaces a halogen atom, the total mass of the molecule is decreased. This results in a decrease in boiling point, making the compounds less than ideal for use as refrigerants. A boiling point in the 210 to 230 °C range is an important property for a refrigerant gas. Too many chlorine atoms seem to increase toxicity. Chloroform, CHCl3, would therefore not be a good CFC substitute both because of its toxicity and its higher boiling point (61 °C). The relationships among composition, molecular structure, boiling point, and proposed use must all be considered along with toxicity, flammability, and stability for any substitute. Fortunately, chemists were able to use their knowledge to synthesize some promising replacements for CFCs. Here are two examples of hydrochlorofluorocarbons, (HCFCs), compounds of hydrogen, chlorine, fluorine, and carbon. F H
C
F
Cl CHClF2 chlorodifluoromethane HCFC-22
and
H
H
Cl
C
C
H
Cl
Starting with the 1996 model year, automobiles in the United States use HCFCs in their air conditioning units rather than CFCs. Even so, older cars are still on the roads.
F
C2H3Cl2F dichlorofluoroethane HCFC-141b
The first of these, CHClF2 (also called R-22) was once the most widely used HCFC. Its ozone-depleting potential is about 5% that of CFC-12 and its estimated atmospheric lifetime is only 20 years, compared with 111 years for CFC-12. It is suitable both for air conditioners and as a blowing agent to make fast-food containers. The second, C2H3Cl2F, is also used as a blowing agent to make foam insulation. To their credit, HCFCs decompose in the troposphere more readily than CFCs and hence do not accumulate to the same extent in the stratosphere. However, HCFCs still contain chlorine. As a result, they still have adverse effects on the ozone layer and thus can be regarded only as an interim solution. The Montreal Protocol and its amendments are tackling the problems associated with HCFCs. This situation is complex in that developing nations want and need to use these compounds as refrigerant gases. In 2007, an adjustment was made concerning the production and consumption of HCFCs. Little change was made for the developed nations that
Look for more about foam containers for fast food and blowing agents in Section 9.4.
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are Parties to the Montreal Protocol, as they already were subject to caps on their consumption and production, ultimately leading to a phase out by 2030. For example, by 2015, the production or importation of all HCFCs will end in the United States. Any continued demand to service refrigerating equipment manufactured prior to those deadlines must be met with recovered HCFCs. But with the 2007 adjustment, a timetable was launched for the developing nations, one that also would lead to a 100% reduction by 2030. During the phaseout period, the price of available HCFCs will undoubtedly rise dramatically. One factor will be the limited supply of HCFCs for existing equipment but another is change in air-conditioning energy-efficiency regulations mandated by the Department of Energy’s new Seasonal Energy Efficiency Rating (SEER) standards for residential air conditioner manufacturers. Starting in January 2006, SEER ratings of new air-conditioning units were required to show a 30% increase in energy efficiency, a performance that HCFCs helped manufacturers achieve. If HCFCs are being phased out, what options do we have? At first glance, hydrofluorocarbons, (HFCs), compounds of hydrogen, fluorine, and carbon, would seem to be likely candidates. Here are two examples.
H
F
F
C
C
F
F
F F
and
H
C
F
H
C2HF5 pentafluoroethane HFC-125
CH2F2 difluoromethane HFC-32
From these structures, you can see that HFCs have no chlorine atoms to deplete ozone. In addition, their hydrogen atoms facilitate decomposition in the lower atmosphere, so they do not have excessively long atmospheric lifetimes. Furthermore, compounds containing only C and F (fluorocarbons) are nontoxic and nonflammable under normal conditions. The two HFCs for which we just drew Lewis structures, C2HF5 and CH2F2, often are blended to produce the refrigerant known as R-410A. Newer designs for air conditioners will be engineered to use this blend as a replacement for HCFC-22.
Consider This 2.28
Blended HFCs
R-407c is a blend of HFC-125, HFC-32, and HFC-134a. The formula, name, and Lewis structure for the first two components of R-407c were given earlier. The formula for HFC-134a is C2H2F4 and its name is tetrafluoroethane. a. In what ways are HFC-125, HFC-32, and HFC-134a different from CFCs? b. Draw the Lewis structure for HFC-134a. Hint: Place two F atoms on each C atom.
Global climate change is the topic of Chapter 3.
At the risk of getting ahead of the story of global warming, which we will tell in Chapter 3, we now mention that HFCs are greenhouse gases. Like CO2, HFCs absorb infrared radiation, trap heat in the atmosphere, and contribute to global warming. Actually, the same is true for CFCs, as you can see from Figure 2.21. Thus, over the long haul, HFCs cannot serve as replacements either. Also to be phased out under the Montreal Protocol are the bromine-containing halons, compounds discussed earlier in Section 2.9. Although very effective in fighting fires, molecule for molecule, halons are even more effective than CFCs in harming the ozone layer. In addition and just like HFCs, halons contribute to global warming as you can see in Figure 2.21. In 1998, 1199 9988, Pyrocool P Pyr y oc yr ocoo ooll Technologies Tech Te chno nolo logi g es of gi of Monroe, Monr Mo nroe oe,, Virginia, Virg Vi rg gin inia ia,, won won a Presidential Pres Pr esid iden enti tial al Green Gre Green en Chemistry Chem Ch hemiist istry t r y Ch Challenge hal all lleng leng ge Awar A Award wardd for for it its deve its ddevelopment evellopm lopm p en entt off ffoa foam oam m th that hatt is is en environmentally viro vi i ronm nmen enttall ta ll llyy benign beni be nign ni gn and a nd yet yet mor m more oree ef or effe effective fect fe ctiv ct ivee th iv than an the the hal hhalons alon al onss it rep on rreplaces. epla ep lace la cess. The ce T he product, pro pro rodu duct du ct, Py ct Pyro Pyrocool roco ro cool co ol
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CFC-11 CFC-12 Carbon tetrachloride (CCl4) CFC-113 HCFC-22 Halon-1211 Halon-1301 Methyl bromide (CH3Br) (5) HFC-134a HFC-23 HFC-125 0
4000 8000 12000 Global warming potential (100–yr)
16000
Figure 2.21 Relative importance of equal amounts (by mass) of selected CFCs, HCFCs, HFCs and halons in terms of their global warming potential. The values are for a 100-year time interval after emission. Source: Taken from D.W. Fahey, 2006, Twenty Questions and Answers about the Ozone Layer—2006 Update, a supplement to Scientific Assessment of Ozone Depletion: 2006, the World Meteorological Organization Global Ozone Research and Monitoring Project—Report No. 50, released 2007, and reproduced here with the kind permission of the United Nations Environment Programme.
Fire-Extinguishing Foam (FEF), can replace halons in fighting even large-scale fires such as those on oil tankers or jet aircraft. A 0.4% solution of Pyrocool FEF dissolved in water to produce a foam was used to control the spread of fires in the sublevels beneath the collapsed towers of the World Trade Center towers following the terrorist attack of September 11, 2001 (Figure 2.22). Many of the hot spots buried in the debris were dangerous to any rescue operations. The foam also helped protect the huge tanks that stored Freon for the air-conditioning systems. The Pyrocool FEF foam also has a cooling effect that helps firefighters, a useful feature when fighting brush fires. Other companies across the globe are working on the challenge of replacing halons.
Figure 2.22 An aqueous foam of Pyrocool FEF is being applied to subterranean fires of the north tower of the World Trade Center, September 30, 2001.
The x-axis of Figure 2.21 is global warming potential. Look for an explanation of the term in Section 3.8.
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Your Turn 2.29
Another Replacement for Halon-1301
Fluoroform (CHF3), also called HFC-23, can be used as a replacement for Halon-1301 (CBrF3). a. Examine the chemical formulas of Halon-1301 and HFC-23. From this information, what can you conclude about the ability of these two fire suppressants to deplete stratospheric ozone? b. What is the current status of Halon-1301? c. Why is fluoroform (HFC-23) unlikely to be used long term? Hint: See Figure 2.21. Answer b. All halons, including Halon-1301, were phased out in 2010.
The phaseout of CFCs and the continuing development of alternative materials are not without major economic concerns. At its peak, the annual worldwide market for CFCs reached $2 billion, but that was only the tip of a very large financial iceberg. In the United States alone, CFCs were used in or used to produce goods valued at about $28 billion per year. Although the conversion to CFC replacements has had some additional costs associated with it, the overall effect on the U.S. economy actually was minimal. Companies that produce refrigerators, air conditioners, insulating plastics, and other goods have adapted to using new compounds. Some substitutes for CFC refrigerants are less energy-efficient, hence increasing energy consumption somewhat. Butt the Bu the conversions ccon onve on vers ve rsio rs ions io ns provide pro rovi vide vi de a market m mar arke ar kett opportunity ke opppo p rt rtun unit un ityy for it for innovative inno in nova no vati va tive ti ve syntheses ssyn y th yn thes eses es es based bbas ased as ed on on thee principles th prin pr inci in cipl ci ples pl es of of green gree gr eenn chemistry ee chem ch emis em istr is tryy to produce tr ppro roddu ro duce environmentally duce eennv nvir nvir iron onme on ment me ntal nt allly al ly benign bben enig en ignn substances, ig subs subs su bsta tanc ta nces nc es,, a es win for both present and future generations. Developing countries face a different set of economic problems and priorities. CFCs have played an important role in improving the quality of life in the industrialized nations. Few of their citizens would be willing to give up the convenience and health benefits of refrigeration or the comfort of air conditioning. It is understandable that millions of people over the globe aspire to the lifestyle of the industrialized nations. But, if the developing nations are banned from using the relatively inexpensive CFC-based technology, they may not be able to afford alternatives. “Our development strategies cannot be sacrificed for the destruction of the environment caused by the West,” asserts Ashish Kothari, a member of an Indian environmental group. Both India and China originally refused to sign the Montreal Protocol because they felt that it discriminated against developing countries. To gain the participation of these highly populated nations, the industrially developed nations created a special fund that is administered through the World Bank. The goal of the fund is to help countries phase out their use of ozone-depleting materials without hurting their economic development.
Consider This 2.30
History Still Being Written
Each year, nations gather to continue conversations about the Montreal Protocol and its amendments. The textbook’s website links you to the recent conferences. Where was this year’s gathering held? Summarize the accomplishments.
Clearly, an understanding of chemistry is necessary to protect the ozone layer, but it is not sufficient. Chemists can help unravel the causes of ozone depletion and develop alternative materials to replace CFCs. Ultimately, though, the debate among governments and their citizens about how best to protect the stratospheric ozone layer continues in the global political arena.
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Conclusion Chemistry is intimately entwined with the story of ozone depletion. Chemists created the chlorofluorocarbons, the near-perfect properties of which only later revealed their dark side as predators of stratospheric ozone. Chemists worked internationally to discover the mechanism by which CFCs destroy stratospheric ozone and warned of the dangers of increased ultraviolet radiation reaching the Earth. And chemists will continue to synthesize the substitutes necessary to replace CFCs and other related compounds. Although chemistry is a necessary part of the solution, it was only part of the solution. At the 2005 meeting in Dakar, Senegal, of parties to the Montreal Protocol on Substances That Deplete the Ozone Layer, Executive Secretary Marco González reminded delegates that the final 20% of any global cooperative effort can be the hardest. Fundamental differences in domestic regulatory approaches have the potential to deplete the stockpiles of goodwill, placing long-term goals in jeopardy. In their complexity, the economic, social, and political issues rival that of the scientific and technological ones. Thus the problem of ozone depletion brought together an array of different participants in pursuit of a common end. Chemists provided the information on the causes and effects of ozone depletion. People in industry, responding to the stimulus provided by the control measures, developed alternatives far more rapidly and more cheaply than anyone initially thought possible, participating fully in the debates over further reductions. Nongovernmental organizations (NGOs) and the media served as essential channels of communication with the peoples of the world in whose name the measures had been taken. Governments worked well together in patiently negotiating agreements acceptable to a range of countries with widely varying circumstances, aims, and resources—and showed courage and foresight in applying the precautionary principle before the scientific evidence was entirely clear. Clearly our problems are global. Ultimately, our solutions must be global as well. As we mentioned earlier in this chapter, in 2007 a symposium was held on the 20th Anniversary of the Montreal Protocol. Georgios Souflias from Greece gave the opening address, pointing out that we cannot remain indifferent toward the environment, as the environment is our home. “The environment does not only provide people a better quality of life, but life itself.” Souflias also pointed out the connection between CFCs and their replacements and global climate change. His call to action was unequivocal: “We all, breathing on this planet today and having the potential, must guarantee its future, rapidly and decisively. We have no right to delay; we have no luxury of losing time.” We urge you to carry these words with you as we turn to our next topic, the chemistry of climate change. As we mentioned at the outset of this chapter, CFCs, HCFCs, and HFCs all are greenhouse gases.
Chapter Summary Having studied this chapter, you should be able to: ■ Differentiate between harmful ground-level ozone and beneficial stratospheric ozone layer (2.1) ■ Describe the chemistry of ozone, including how it is formed in our atmosphere. (2.1, 2.6, 2.8–2.10) ■ Describe the ozone layer, characterizing it in several different ways (2.1, 2.6, 2.8–2.10) ■ Apply the basics of atomic structure to atoms of certain elements (2.2) ■ Understand what it means when elements fall into the same group of the periodic table (2.2)
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Differentiate atomic number from mass number and apply the latter to isotopes (2.2) Write Lewis structures for small molecules with single, double, and triple covalent bonds (2.3) Describe the electromagnetic spectrum in terms of frequency, wavelength, and energy (2.4, 2.5) Interpret graphs related to wavelength and energy, radiation and biological damage, and ozone depletion (2.4–2.8) Understand the natural Chapman cycle of stratospheric ozone depletion (2.6)
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Understand how the stratospheric ozone layer protects against harmful ultraviolet radiation (2.6, 2.7) Compare and contrast UV-A, UV-B, and UV-C radiation along several different lines (2.6, 2.7) Discuss the interaction of radiation with matter and changes caused by such interactions, including biological sensitivity (2.6, 2.7) Relate the meaning and the use of the UV Index (2.7) Write Lewis structures for chlorine and bromine atoms, as well as for some other free radicals. Be able to explain why these free radicals are so reactive (2.8)
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Recognize the complexities of collecting accurate data for stratospheric ozone depletion and interpreting them correctly (2.8, 2.9) Understand the chemical nature and role of CFCs in stratospheric ozone depletion (2.9, 2.10) Explain the unique circumstances responsible for seasonal ozone depletion in the Antarctic (2.10) Summarize the outcomes of the Montreal Protocol and its amendments (2.11, 2.12) Evaluate articles on green chemistry alternatives to stratospheric ozone-depleting compounds (2.12) Discuss factors that will help lead to the recovery of the ozone layer (2.11, 2.12)
Questions Emphasizing Essentials 1. How does ozone differ from oxygen in its chemical formula? In its properties? 2. Explain why it is possible to detect the pungent odor of ozone after a lightning storm or around electrical transformers. 3. The text states that the odor of ozone can be detected in concentrations as low as 10 ppb. Would you be able to smell ozone in either of these air samples? a. 0.118 ppm of ozone, a concentration reached in an urban area b. 25 ppm of ozone, a concentration measured in the stratosphere 4. A journalist wrote “Hovering 10 miles above the South Pole is a sprawling patch of stratosphere with disturbingly low levels of radiation-absorbing ozone.” a. How big is this sprawling patch? b. Is the figure of 10 miles correct? Express this value in kilometers. c. What type of radiation does ozone absorb? 5. It has been suggested that the term ozone screen would be a better descriptor than ozone layer to describe ozone in the stratosphere. What are the advantages and disadvantages to each term? 6. Assume there are 2 3 1020 CO molecules per cubic meter in a sample of tropospheric air. Furthermore, assume there are 1 3 1019 O3 molecules per cubic meter at the point of maximum concentration of the ozone layer in the stratosphere. a. Which cubic meter of air contains the larger number of molecules? b. What is the ratio of CO to O3 molecules in a cubic meter? 7. a. What is a Dobson unit? b. Does a reading of 320 DU or 275 DU indicate more total column ozone overhead?
8. Using the periodic table as a guide, specify the number of protons and electrons in a neutral atom of each of these elements. a. oxygen (O) b. nitrogen (N) c. magnesium (Mg) d. sulfur (S) 9. Consider this representation of a periodic table.
10.
11.
12.
13.
a. What is the group number of the shaded column? b. Which elements make up this group? c. What is the number of electrons for a neutral atom of each element in this group? d. What is the number of outer electrons for a neutral atom of each element of this group? Give the name and symbol for the element with this number of protons. a. 2 b. 19 c. 29 Give the number of protons, neutrons, and electrons in each of these neutral atoms. a. oxygen-18 (188O) b. sulfur-35 (35 16 S) 238 c. uranium-238 ( 92U) d. bromine-82 (82 35 Br) 19 e. neon-19 (10 Ne) f. radium-226 (226 88 Ra) Give the symbol showing the atomic number and the mass number for the isotope that has: a. 9 protons and 10 neutrons (used in nuclear medicine). b. 26 protons and 30 neutrons (the most stable isotope of this element). c. 86 protons and 136 neutrons (the radioactive gas found in some homes). Draw the Lewis structure for each of these atoms. a. calcium b. nitrogen c. chlorine d. helium
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Wave 1
17.
18.
19.
20.
21.
22. 23.
Wave 2
a. wavelength b. frequency c. speed of travel Use Figure 2.7 to specify the region of the electromagnetic spectrum where radiation of each wavelength is found. Hint: Change each wavelength to meters before making the comparison. a. 2.0 cm b. 400 nm c. 50 μm d. 150 mm Arrange the wavelengths in question 17 in order of increasing energy. Which wavelength possesses the most energetic photons? Arrange these types of radiation in order of increasing energy per photon: gamma rays, infrared radiation, radio waves, visible light. The microwaves in home microwave ovens have a frequency of 2.45 3 109 s−1. Is this radiation more or less energetic than radio waves? Than X-rays? Ultraviolet radiation is categorized as UV-A, UV-B, or UV-C. Arrange these types in order of increasing: a. wavelength b. energy c. potential for biological damage Draw Lewis structures for any three different CFCs. CFCs were used in hair sprays, refrigerators, air conditioners, and plastic foams. Which properties of CFCs made them desirable for these uses?
150 100 50 0
Percent increase in UV-B
16. Consider these two waves representing different parts of the electromagnetic spectrum. How do they compare in terms of:
24. a. Can a molecule that contains hydrogen be classified as a CFC? b. What is the difference between an HCFC and an HFC? 25. a. Most CFCs are based either on methane, CH4, or ethane, C2H6. Use structural formulas to represent these two compounds. b. Substituting chlorine atoms, fluorine atoms, or both for all of the hydrogen atoms on a methane molecule, you obtain CFCs. How many possibilities exist? c. Which of the substituted CFC compounds in part b has been the most successful? d. Why weren’t all of these compounds equally successful? 26. These free radicals all play a role in catalyzing ozone depletion reactions: Cl?, ?NO2, ClO?, and ?OH. a. Count the number of outer electrons available and then draw a Lewis structure for each free radical. b. What characteristic is shared by these species that makes them so reactive? 27. a. How were the original measurements of increases in chlorine monoxide and the stratospheric ozone depletion over the Antarctic obtained? b. How are these measurements made today? 28. Which graph shows how measured increases in UV-B radiation correlate with percent reduction in the concentration of ozone in the stratosphere over the South Pole? Percent increase in UV-B
14. Assuming that the octet rule applies, draw the Lewis structure for each of these molecules. a. CCl4 (carbon tetrachloride, a substance formerly used as a cleaning agent) b. H2O2 (hydrogen peroxide, a mild disinfectant; the atoms are bonded in this order: H–O–O–H) c. H2S (hydrogen sulfide, a gas with the unpleasant odor of rotten eggs) d. N2 (nitrogen gas, the major component of the atmosphere) e. HCN (hydrogen cyanide, a molecule found in space and a poisonous gas) f. N2O (nitrous oxide, “laughing gas”; the atoms are bonded N–N–O) g. CS2 (carbon disulfide, used to kill rodents; the atoms are bonded S–C–S) 15. Several oxygen species play important chemical roles in the stratosphere, including oxygen atoms, oxygen molecules, ozone molecules, and hydroxyl radicals. Draw Lewis structures for each.
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150 100
0 20 40 60 Percent reduction in ozone
(a)
50 0
0 20 40 60 Percent reduction in ozone
(b)
Concentrating on Concepts 29. The EPA has used the slogan “Ozone: Good Up High, Bad Nearby” in some of its publications for the general public. Explain the message. 30. Nobel laureate W. Sherwood Rowland referred to the ozone layer as the Achilles heel of our atmosphere. Explain the metaphor. 31. In the abstract of a talk he gave in 2007, Nobel laureate W. Sherwood Rowland wrote “Solar UV radiation creates an ozone layer in the atmosphere which in turn completely absorbs the most energetic fraction of this radiation.” a. What is the most energetic fraction? Hint: See Figure 2.8. b. How does solar UV radiation “create an ozone layer”?
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32. In the conclusion of this chapter, we reported the words that Georgios Souflias spoke at the Symposium for the 20th Anniversary of the Montreal Protocol: “We all, breathing on this planet today and having the potential, must guarantee its future, rapidly and decisively. We have no right to delay; we have no luxury of losing time.” a. What danger in delaying was he referring to? b. Look back at the definitions of sustainability in the prologue. How do his words connect to these definitions? 33. Consider the Chapman cycle in Figure 2.10. a. Explain the source of the oxygen atoms. b. Can this cycle take place in the troposphere as well? Explain. 34. What are some of the reasons that the solution to ozone depletion proposed in this Sydney Harris cartoon will not work?
Source: ScienceCartoonsPlus.com. Reprinted with permission.
35. “We risk solving one global environmental problem while possibly exacerbating another unless other alternatives can be found.” The date of this quote by a U.S. official is 2009, and the context is phasing out the use of HCFCs. a. What compounds were HCFCs being replaced with in 2009? b. What is the risk of this replacement? 36. It is possible to write three resonance structures for ozone, not just the two shown in the text. Verify that all three structures satisfy the octet rule and offer an explanation of why the triangular structure is not reasonable. O
O
O
O
O
O
O
O
O
37. The average length of an O–O single bond is 132 pm. The average length of an O–O double bond is 121 pm. What do you predict the O–O bond lengths will be in ozone? Will they all be the same? Explain your predictions.
38. Consider the Lewis structures for SO2. How do they compare with the Lewis structures for ozone? 39. Even if you have skin with little pigment, you cannot get a tan from standing in front of a radio. Why? 40. The morning newspaper reports a UV Index Forecast of 6.5. Given the amount of pigment in your skin, how might this affect how you plan your daily activities? 41. All the reports of the damage caused by UV radiation focus on UV-A and UV-B radiation. Why is there no attention on the damaging effects of UV-C radiation? 42. If all 3 3 108 tons of stratospheric ozone that are formed every day are also destroyed every day, how is it possible for stratospheric ozone to offer any protection from UV radiation? 43. How does the chemical inertness of CCl2F2 (Freon-12) relate both the usefulness and the problems associated with this compound? 44. Explain how the small changes in concentrations (measured in parts per billion) can cause the much larger changes in O3 concentrations (measured in parts per million). 45. Development of the stratospheric ozone hole has been most dramatic over Antarctica. What set of conditions exist over Antarctica that help to explain why this area is well-suited to studying changes in stratospheric ozone concentration? Are these same conditions not operating in the Arctic? Explain. 46. The free radical CF3O? is produced during the decomposition of HFC-134a. a. Propose a Lewis structure for this free radical. b. Offer a possible reason why CF3O? does not cause ozone depletion. 47. One mechanism that helps break down ozone in the Antarctic region involves the BrO? free radical. Once formed, it reacts with ClO? to form BrCl and O2. BrCl, in turn, reacts with sunlight to break into Cl? and Br?, both of which react with O3 and form O2. a. Represent this information with a set of equations similar to those shown for the Chapman cycle. b. What is the net equation for this cycle? 48. Polar stratospheric clouds (PSCs) play an important role in stratospheric ozone depletion. a. Why do PSCs form more often over Antarctica than in the Arctic? b. Reactions occur more quickly on the surface of PSCs than in the atmosphere. One such reaction is the reaction of hydrogen chloride and chlorine nitrate (ClONO2), two species that do not deplete ozone, to produce a chlorine molecule and nitric acid (HNO3). Write the chemical equation. c. The chlorine molecule produced does not deplete ozone either. However, when the Sun returns to the Antarctic in the springtime, it is converted to a species that does. Show how with a chemical equation.
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49. Consider this graph that shows the atmospheric abundance of bromine-containing gases from 1950 to 2100.
4
Halon-1211
Atmospheric abundance (parts per trillion)
3
Halon-1301
2 1 12
10 Methyl bromide (CH3Br) 8 1950 Estimates of historical abundance
2000
Year
Observations
2050
6 2100 Future projections
Source: Taken from D.W. Fahey, 2006, Twenty Questions and Answers about the Ozone Layer—2006 Update, a supplement to Scientific Assessment of Ozone Depletion: 2006, the World Meteorological Organization Global Ozone Research and Monitoring Project—Report No. 50, released 2007, and reproduced here with the kind permission of the United Nations Environment Programme.
a. Halon-1301 is CBrF3 and Halon-1211 is CClBrF2. Why were these compounds once manufactured? b. Compare the patterns for Halon-1211 and Halon-1301. Why doesn’t Halon-1301 drop off as quickly? c. In 2005, methyl bromide was phased out in the United States except for critical uses. Why is its future use predicted as a straight line, rather than tailing off? Exploring Extensions 50. Chapter 1 discussed the role of nitrogen monoxide (NO) in forming photochemical smog. What role, if any, does NO play in stratospheric ozone depletion? Are NO sources the same in the stratosphere as in the troposphere? 51. Resonance structures can be used to explain the bonding in charged groups of atoms as well as in neutral molecules, such as ozone. The nitrate ion, NO3−, has one additional electron plus the outer electrons contributed by nitrogen and oxygen atoms. That extra electron gives the ion its charge. Draw the resonance structures, verifying that each obeys the octet rule.
105
52. Although oxygen exists as O2 and O3, nitrogen exists only as N2. Propose an explanation for these facts. Hint: Try drawing a Lewis structure for N3. 53. The chemical formulas for a CFC, such as CFC-11 (CCl3F), can be figured out from its code number by adding 90 to it to get a three-digit number. For example, with CFC-11 you get 90 1 11 5 101. The first digit is the # of C atoms, the second is the # of H atoms, and the third is the # of F atoms. Accordingly, CCl3F has 1 C atom, no H atoms, and 1 F atom. All remaining bonds are assumed to be chlorine. a. What is the chemical formula for CFC-12? b. What is the code number for CCl4? c. Does this “90” method work for HCFCs? Use HCFC-22 (CHClF2) in explaining your answer. d. Does this method work for halons? Use Halon-1301 (CF3Br) in explaining your answer. 54. Many different types of ozone generators (“ozonators”) are on the market for sanitizing air, water, and even food. They are often sold with a slogan such as this one from a pool store. “Ozone, the world’s most powerful sanitizer!” a. What claims are made for ozonators intended to purify air? b. What risks are associated with these devices? 55. The effect a chemical substance has on the ozone layer is measured by a value called its ozone-depleting potential, ODP. This is a numerical scale that estimates the lifetime potential stratospheric ozone that could be destroyed by a given mass of the substance. All values are relative to CFC-11, which has an ODP defined as equal to 1.0. Use those facts to answer these questions. a. Name two factors that affect the ODP value of a compound and explain the reason for each one. b. Most CFCs have ODP values ranging from 0.6 to 1.0. What range do you expect for HCFCs? Explain your reasoning. c. What ODP values do you expect for HFCs? Explain your reasoning. 56. Recent experimental evidence indicates that ClO? initially reacts to form Cl2O2. a. Predict a reasonable Lewis structure for this molecule. Assume the order of atom linkage is Cl–O–O–Cl. b. What effect does this evidence have on understanding the mechanism for the catalytic destruction of ozone by ClO??
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The Chemistry of Global Climate Change
“Global warming is a misnomer, because it implies something that is gradual, something that is uniform, something that is quite possibly benign. What we are experiencing with climate change is none of those things.” John Holdren, Meeting the Climate Change Challenge, National Council for Science and the Environment, 2008.
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Is global warming really a misnomer? Could those two little words misrepresent three large aspects of the issue we are facing? Let’s examine each of the assertions that Holdren has made. First, he asserts that global warming isn’t gradual. By this he means that in comparison with the past, the climate changes we are seeing today are occurring much more rapidly. Natural climate changes are part of our planet’s history. Glaciers, for example, have advanced and retreated numerous times, and global temperatures have been both much higher and much lower than the temperatures we currently experience. But the geologic evidence indicates these past changes occurred over millennia, not decades as they are today. So Holdren is correct. Global warming is not gradual, at least not in comparison with the geologic time frames of the past. Second, he asserts that global warming does not occur uniformly across the globe. Holdren is right again. To date, the most dramatic effects have been observed at the poles. These include quickly receding glaciers, shrinking sea ice, and melting permafrost. So far, the more densely populated lower latitudes have experienced far smaller effects from climate change. His third assertion, that global warming might not be benign, is the most difficult to assess. The issue is complicated in part because we cannot predict with certainty which aspects of our planet global warming will affect and to what degree. It is further complicated because we cannot easily understand why only a couple of degrees of warming might be catastrophic. It is important not to take one person’s word regarding a topic as complex as climate change. Therefore, in this chapter, we delve into Holdren’s assertions in detail but do so by taking things one step at a time. The first step is to examine how our Earth maintains its energy balance. We consider incoming and outgoing solar radiation and explain how our atmosphere functions much like a greenhouse. The second step in our discussion of climate change examines scientific data on both the current and past state of the planet. In particular, we examine the concentrations of atmospheric greenhouse gases, both past and present. Carbon dioxide resides at the center of this discussion, yet the reason is far from obvious. After all, CO2 is an essential component of the atmosphere, a gas that all animals exhale and green plants absorb. Key to understanding climate change is studying the molecular mechanism by which CO2 and other compounds absorb the infrared radiation emitted by the planet, helping to keep it warm. Some knowledge of molecular structure and shape is necessary to understand this mechanism. Climate change also has a significant quantitative component; we need numbers to help grasp the magnitude of the problem. Therefore, our third step is to introduce the way that chemists count the unimaginably numerous (and unimaginably small) particles we call atoms and molecules. Lastly, we examine the limits to which climate scientists can predict what the future may hold and describe several of the more dangerous consequences of a warmer planet Earth. As with many of the issues addressed in this book, we must consider not only what is happening today, but also the effects our actions (and our inactions) will have on future generations. Developing an understanding of these issues will lead us on a journey into the realm of chemical knowledge and its connections with public policy around the globe.
Consider This 3.1
The terms climate change and global warming are both used, both in the popular press and by scientists. Although they are not the same, they are closely related. We will use both in this chapter. See Consider This 3.1.
What’s in a Name?
Sometimes people, including scientists, talk about global warming. Other times, folks refer to global climate change. a. Interview two friends or family members and ask them to list what comes to mind when they hear the term “global warming.” Do the same for the term “global climate change.” Comment on the two lists. b. Do you prefer one term over the other? Explain.
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3.1
Figure 3.1 Venus, as photographed by the Galileo spacecraft.
These and other types of electromagnetic radiation were introduced in Section 2.4
|
In the Greenhouse: Earth’s Energy Balance
The brightest and most beautiful body in the night sky, after our own Moon, is considered by many to be Venus (Figure 3.1). It is ironic that the planet named for the goddess of love is a most unlovely place by earthly standards. The Venetian atmosphere has a pressure 90 times greater than that of Earth, and it is 96% carbon dioxide, with clouds of sulfuric acid. It makes the worst smog-bound day anywhere on Earth seem like a breath of fresh country air. Spacecraft have revealed a desolate, eroded surface with an average temperature of about 450 °C (840 °F). In contrast, the beautiful blue-green ball we inhabit has an average annual temperature of 15 °C (59 °F). The point of this little astronomical digression is that both Venus and Earth are warmer than one would expect based solely on their distances from the Sun and the amount of solar radiation they receive. If distance were the only determining factor, the temperature of Venus would average approximately 100 °C, the boiling point of water. Earth, on the other hand, would have an average temperature of 218 °C (0 °F), and the oceans would be frozen year-round. Processes that keep the energy of our Earth in balance are shown in Figure 3.2. The Earth receives nearly all of its energy from the Sun (orange arrows), primarily in the form of ultraviolet, visible, and infrared radiation. Accounting for what happens to the incoming radiation is relatively straightforward. Some of it is reflected back to space (blue arrows), either by the molecules, dust, and aerosol particles that are suspended in our atmosphere (25%) or by the surface of the Earth itself, especially those regions white with snow (6%). But most of the incoming radiation warms the Earth, either by being absorbed by the atmosphere (23%) or by being absorbed by the landmasses and oceans (46%). The numbers tally quite nicely: 25% reflected 1 6% reflected 1 23% absorbed 1 46% absorbed 5 100% of the incoming radiation.
Your Turn 3.2
Light from the Sun
Consider these three types of radiant energy, all emitted by the Sun: infrared (IR), ultraviolet (UV), and visible. a. Arrange them in order of increasing wavelength. b. Arrange them in order of increasing energy. Answer a. ultraviolet, visible, infrared
Accounting for what happens to the outgoing radiation is more complicated. Observe that our Earth, like the Sun, gives off radiation. If this were not the case, our planet would quickly become unbearably hot! The energy absorbed by the Earth must be reemitted to maintain the energy balance. But unlike the Sun, the Earth emits primarily in the infrared region (red arrows). A small amount of this IR radiation (9%) passes directly from the surface of the Earth out into space. However, most of the heat radiated by the Earth (37%) is absorbed by the atmosphere and returned to Earth rather than being lost to space. Heat is transferred by collisions between neighboring molecules, and these molecules are found in greater abundance in the denser regions of the lower atmosphere. Check the math: 46% absorbed 5 9% 1 37% emitted. The percent of solar radiation striking the Earth that remains in the atmosphere is about 80% (37% ÷ 46%). The greenhouse effect is the natural process by which atmospheric gases trap a major portion (about 80%) of the infrared radiation radiated by the Earth. Again, Earth’s average annual temperature of 15 °C (59 °F) is a result of the heat trapping gases in our atmosphere.
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Reflected from surface Emitted from surface 6 9
Emitted from atmosphere
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Atmosphere
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Absorbed in atmosphere
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Figure 3.2 The Earth’s energy balance. Orange represents a mixture of wavelengths. Shorter wavelengths of radiation are shown in blue, longer ones in red. The values are given in percentages of the total incoming radiation.
Your Turn 3.3
Earth’s Energy Balance
Refer to Figure 3.2 to answer these questions. a. Incoming solar radiation (100%) is either absorbed or reflected. Outgoing radiation from the Earth into space also can be accounted for (100%), as required for energy balance. Show how. b. What percent of the outgoing energy is absorbed in the Earth’s atmosphere? Calculate this by adding the percentage of incoming solar energy absorbed in the atmosphere to that absorbed in the atmosphere after being radiated from Earth’s surface. How does this value compare with the percentage emitted from the atmosphere? c. Suggest reasons why the different colors were used for incoming and outgoing radiation.
If you have ever parked a car with its windows closed on a sunny day, you probably have experienced firsthand how a greenhouse can trap heat. The car, with its glass windows, operates much the same way as does a greenhouse for growing plants. The glass windows transmit visible and a small amount of UV light from the Sun. This
Visit the textbook’s website to learn more about the electromagnetic spectrum, Earth’s energy balance, and the greenhouse effect. Look for the Figures Alive! icon elsewhere in this chapter.
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Water vapor is the most abundant greenhouse gas in our atmosphere. However, contributions of H2O from human activity are negligible compared with those from natural sources.
Another steady-state process, the Chapman cycle, was discussed in Section 2.6.
Section 1.9 described the chemistry of combustion. Look for more about coal (a fossil fuel) in Chapter 4.
energy is absorbed by the car’s interior, particularly by any dark surfaces. Some of this radiant energy is then reemitted as longer wavelength IR radiation (heat). Unlike visible light, infrared light is not easily transmitted through the glass windows and so becomes “trapped” inside the car. When you reenter the vehicle, a blast of hot air greets you. The temperature inside of a car can exceed 49 °C (120 °F) in the summer in certain climates! Although the physical barrier of the windows is not an exact analogy to the Earth’s atmosphere, the effect of warming the car’s interior is similar to the warming of the Earth. Although the ability of the atmosphere to trap heat was hypothesized by the French mathematician Jean-Baptiste Joseph Fourier (1768–1830) around 1800, it took another 60 years for scientists to identify the molecules that were responsible. Irish physicist John Tyndall (1820–1893) first demonstrated that both carbon dioxide and water vapor absorb infrared radiation. Greenhouse gases are those gases capable of absorbing and trapping infrared radiation, thereby warming the atmosphere. Examples include water vapor, carbon dioxide, methane, nitrous oxide, ozone, and chlorofluorocarbons. The presence of those gases is essential in keeping our planet at habitable temperatures. Because of the constant, dynamic energy exchange between Earth, its atmosphere, and outer space, a steady state is established that results in a more or less constant average terrestrial temperature. However, the buildup of greenhouse gases that is taking place today is changing the energy balance and causing increased warming of the planet. The term enhanced greenhouse effect refers to the process in which atmospheric gases trap and return more than 80% of the heat energy radiated by the Earth. An increase in the concentration of greenhouse gases will very likely mean that more than 80% of the radiated energy will be returned to Earth’s surface, with an accompanying increase in average global temperature. The popular term global warming often is used to describe the increase in average global temperatures that results from an enhanced greenhouse effect. We need look no further than ourselves to find the cause of the buildup of certain greenhouse gases in our atmosphere. Anthropogenic influences on the environment stem from human activities, such as industry, transportation, mining, and agriculture. In the late 19th century, Swedish scientist Svante Arrhenius (1859–1927) considered the problems that increased industrialization might cause by building up CO2 in the atmosphere. He calculated that doubling the concentration of CO2 would result in an increase of 5–6 °C in the average temperature of the planet’s surface. Writing in the London, Edinburgh, and Dublin Philosophical Magazine, Arrhenius dramatically described the phenomenon: “We are evaporating our coal mines into the air.” At the end of the 19th century, the Industrial Revolution was already well under way in Europe and America, and it was “picking up steam” as well as generating it.
Consider This 3.4
Evaporating Coal Mines
Although the Arrhenius statement about “evaporating our coal mines into the air” certainly was effective in grabbing attention in 1898, what process do you think he really was referring to in discussing the amount of CO2 being added to the air? Explain your reasoning.
To further investigate global climate change, we need answers to several important questions. For example, how have the atmospheric concentrations of greenhouse gases changed over time? Similarly, how has the average global temperature changed and how did we measure the changes? Can we determine if the changes in greenhouse gases and temperature are correlated? Can we distinguish natural climate variability from human influences? In the following section, we provide some data to help answer these questions.
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3.2
|
Gathering Evidence: The Testimony of Time
Over the past 4.5 billion years, both Earth’s climate and its atmosphere have varied widely as a result of astronomical, chemical, biological, and geological processes. Earth’s climate has been directly affected by periodic astronomical changes in the shape of Earth’s orbit and the tilt of Earth’s axis. Such changes are thought to be responsible for the ice ages that have occurred regularly during the past million years. Even the Sun itself has changed. Its energy output half a billion years ago was 25–30% less than it is today. In addition, changes in atmospheric greenhouse gas concentrations affect the Earth’s energy balance, and hence its climate. Carbon dioxide was once 20 times more prevalent in the atmosphere than it is today. Chemical processes lowered that level by dissolving much of the CO2 in the oceans, or incorporating it in rocks such as limestone. The biological process of photosynthesis also radically altered the composition of our atmosphere by removing CO2 and producing oxygen. Certain geological events like volcanic eruptions add millions of tons of CO2 and other gases to the atmosphere. Although these natural phenomena will continue to influence Earth’s atmosphere and its climate in the coming years, we must also assess the role that human activities are playing. With the development of modern industry and transportation, humans have moved huge quantities of carbon from terrestrial sources like coal, oil, and natural gas into the atmosphere in the form of CO2. To evaluate the influence humans are having on the atmosphere, and hence on any enhanced greenhouse effect, it is important to investigate the fate of this large unnatural influx of carbon dioxide. Indeed, CO2 concentrations in the atmosphere have increased significantly in the past half century. The best direct measurements are taken from the Mauna Loa Observatory in Hawaii (Figure 3.3). The red zigzag line shows the average monthly concentrations, with a small increase each April followed by a small decrease in October. The black line is a 12-month moving average. Notice the steady increase in average annual values from 315 ppm in 1960 to about 388 ppm in 2009. Later in this chapter, we will examine the evidence linking much of the added carbon dioxide to the burning of fossil fuels, combustible substances derived from the remnants of prehistoric organisms, the most common of which are coal, petroleum, and natural gas.
380 CO2 concentration (ppm)
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320 Jan Apr Jul Oct Jan 1960
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1980 Year
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Figure 3.3 Carbon dioxide concentrations from 1958 to 2009, as measured at Mauna Loa, Hawaii. Inset: One year of the monthly variations. Source: Scripps Institution of Oceanography, NOAA Earth System Research Laboratory, 2009.
The ionic compounds calcium carbonate (CaCO3) and magnesium carbonate (MgCO3) are both insoluble. Look for more about solubility in Chapter 5.
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Your Turn 3.5
The Cycles of Mauna Loa
a. Calculate the percent increase in CO2 concentration during the last 50 years. b. Estimate the variation in parts per million (ppm) CO2 within any given year. c. On average, the CO2 concentrations are higher each April than each October. Explain. Answer c. Photosynthesis removes CO2 from the atmosphere. Spring begins in the northern latitudes in April; October is the start of spring in the southern latitudes. But the landmasses (and number of green plants) are greater in the northern hemisphere sothe seasons in the northern hemisphere control the fluctuations.
In the history of human civilization, 50 years of direct measurements are not much to rely on. How can we obtain data about the composition of our atmosphere farther back in time? Much relevant information comes from the analysis of ice core samples. Regions on the planet that have permanent snow cover contain preserved histories of the atmosphere, buried in layers of ice. Figure 3.4a shows a dramatic example of annual ice layers from the Peruvian Andes. The oldest ice on the planet is located in Antarctica, and scientists have been drilling and collecting ice core samples there for over 50 years (Figure 3.4b). Air bubbles trapped in the ice (Figure 3.4c) provide a vertical timeline of the history of the atmosphere; the deeper you drill, the farther back in time you go. Relatively shallow ice core data show that for the first 800 years of the last millennium the CO2 concentration was relatively constant at about 280 ppm. Figure 3.5 combines the Mauna Loa data (red dots) with data from a 200-meter ice core from the Siple station in Antarctica (green triangles), and a deeper core from the Law Dome, also in Antarctica (blue squares). Beginning about 1800, CO2 began accumulating in the atmosphere at an ever-increasing rate, corresponding to the beginning of the Industrial Revolution and the accompanying combustion of fossil fuels that powered that transformation.
Skeptical Chemist 3.6
Checking the Facts on CO2 Increases
a. A recent government report states that the atmospheric level of CO2 has increased 30% since 1860. Use the data in Figure 3.5 to evaluate this statement. b. A global warming skeptic states that the percent increase in the atmospheric level of CO2 since 1957 has been only about half as great as the percent increase from 1860 to the present. Comment on the accuracy of that statement and how it could affect potential greenhouse gas emissions policy.
(a)
(b)
(c)
Figure 3.4 (a) Quelccaya ice cap (Peruvian Andes) showing the annual layers. (b) Ice core that can be used to determine changes in concentrations of greenhouse gases over time. (c) Microscopic air bubbles in ice.
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150 1000
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1400
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Figure 3.5 Carbon dioxide concentrations over the last millennium as measured from Antarctic ice cores (blue squares and green triangles) and the Mauna Loa observatory (red dots). Source: “Climatic Feedbacks on the Global Carbon Cycle,” in The Science of Global Change: The Impact of Human Activities on the Environment, American Chemical Society Symposium Series, 1992.
What about further back in time? Drilling by a team of Russian, French, and U.S. scientists at the Vostok Station in Antarctica yielded over a mile of ice cores taken from the snows of 400 millennia. The atmospheric carbon dioxide concentrations going back over 400,000 years are shown in Figure 3.6, with the data from Figure 3.5 in the inset. Most obvious from the graph are the periodic cycles of high and low carbon dioxide concentrations, which occur roughly in 100,000-year intervals. Although not shown on the graph, analysis of other ice cores indicate these regular cycles go back at least 1 million years. Two important conclusions can be drawn from these data. First, the current atmospheric CO2 concentration is 100 ppm higher than any time in the last 400 400
CO2 concentration (ppm)
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Ice age cycles
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Figure 3.6 Carbon dioxide concentrations for the last 400,000 years. Inset: Data from Figure 3.5 for comparison.
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El Niño and La Niña are names given to natural cyclical changes in the ocean–atmosphere system in the tropical Pacific. El Niño events lead to warmer ocean temperatures in the middle latitudes, and La Niña cycles produce cooler ocean temperatures.
Isotopes of hydrogen (and other elements) were discussed in Section 2.2.
million years. Also during that time, never has the CO2 concentration risen as rapidly as it is rising today. What about the global temperature? Measurements indicate that during the past 120 years or so, the average temperature of the planet has increased somewhere between 0.4 and 0.8 °C. Figure 3.7 shows the changes in surface air temperature from 1880 to 2006. Some scientists correctly point out that a century or two is an instant in the 4.5-billion-year history of our planet. They caution restraint in reading too much into short-term temperature fluctuations. Short-term changes in atmospheric circulation patterns like El Niño and La Niña events are certainly implicated in some of observed temperature anomalies. Figure 3.7 also shows the temperature ranges within each year (black error bars) as well as the longer term trend (blue line). Although the general trend in temperatures over the last 50 years generally follows the increases in carbon dioxide concentrations, the temperature data from year to year are much less consistent. Furthermore, determining whether the temperature increase is a consequence of the increased CO2 concentration cannot be concluded with absolute certainty. Nevertheless, as we shall see, experimental evidence implicates carbon dioxide from human-related sources as a cause of recent global warming. It is important to realize that an increase in global average temperature does not mean that across the globe every day is now 0.6 °C warmer than it was in 1970. A map of the temperatures for 2006 compared to the average temperature between 1951 and 1980 is displayed in Figure 3.8. Many regions have experienced just a little warming, and some others have even cooled (blue areas). Yet there are other regions (dark red areas), particularly in the higher latitudes, that have experienced much more than the average warming. The increases are most drastic in the Arctic, where not surprisingly, much of the tangible effects of climate change have already been observed. Ice cores also can provide data for estimating temperatures further back in time because of the hydrogen isotopes found in the frozen water. Water molecules containing the most abundant form of hydrogen atoms, 1H, are lighter than those that contain deuterium, 2H. The lighter H2O molecules evaporate just a bit more readily than the heavier ones. As a result, there is more 1H than 2H in the water vapor of the atmosphere than in the oceans. Likewise, the heavier H2O molecules in the atmosphere condense just a bit more readily than the lighter ones. Therefore, snow that condenses from atmospheric water vapor is enriched in 2H. The degree of enrichment depends on temperature. The ratio of 2H to 1H in the ice core can be measured and used to estimate the temperature at the time the snow fell. When we look back into the past, we see that the global temperature has undergone fairly regular cycles, matching the highs and lows in CO2 concentration quite
1.0 0.8 Temperature (⬚C) relative to 1901–2000 average
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Figure 3.7 Global surface temperatures (1880–2006). The red bars indicate the average temperature for each year, and the ranges for each year are shown as the black error bars. The blue line shows the 5-year moving average.
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Figure 3.8 Global temperatures for 2006 (in °C) relative to the 1951–1980 average. Source: NASA
remarkably (Figure 3.9). Other data show that periods of high temperature also have been characterized by high atmospheric concentrations of methane, another significant greenhouse gas. The precision of these data do not allow an assignment of cause and effect. It is difficult to conclude whether increasing greenhouse gases caused the temperature increases, or vice versa. What is clear, however, is that the current CO2 and methane levels are much higher than any time in the last million years. Notice that the variation from hottest to coldest is only about 10 °C, yet that is the difference between the moderate climate we have today, and ice covering much of northern North America and Eurasia, as was the case during the last glacial maximum 20,000 years ago. Over the past million years, Earth has experienced 10 major periods of glacier activity and 40 minor ones. Without question, mechanisms other than greenhouse gas concentrations are involved in the periodic fluctuations of global climate. Some of this temperature variation is caused by minor changes in Earth’s orbit that affect the distance from Earth to the Sun and the angle at which sunlight strikes the planet. However,
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Figure 3.9 Carbon dioxide concentration (blue) and global temperatures (red ) over the last 400,000 years from ice core data.
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this hypothesis cannot fully explain the observed temperature fluctuations. Orbital effects most likely are coupled with terrestrial events such as changes in reflectivity, cloud cover, and airborne dust, as well as CO2 and CH4 concentrations. The feedback mechanisms that couple these effects together are complicated and not completely understood, but it is likely that the effects from each are additive. In other words, the existence of natural climate cycles doesn’t preclude the effect that increased concentrations of greenhouse gases would have on global climate. We are a long way from the out-of-control hothouse of Venus, but we face difficult decisions. These decisions will be better informed with an understanding of the mechanisms by which greenhouse gases interact with electromagnetic radiation to create the greenhouse effect. For that we must again take a submicroscopic view of matter.
3.3 | Remember, the atmosphere is composed of 78% nitrogen and 21% oxygen.
Molecules: How They Shape Up
Carbon dioxide, water, and methane are greenhouse gases; in contrast, nitrogen and oxygen are not. Why the difference? The answer relates in part to molecular shape. In this section, we’ll help you to put your knowledge of Lewis structures to work to predict shapes of molecules. In the next, we connect these shapes to molecular vibrations, which can help us to explain the difference between greenhouse gases and nongreenhouse gases. In Chapter 2 you used Lewis structures to predict how electrons are arranged in atoms and molecules. Shape was not the primary consideration. Even so, in a few cases the Lewis structure did dictate the shape of the molecule. One example is for diatomic molecules such as O2 and N2. Here, the shape is unambiguous because the molecule must be linear. N
N
or
N
N
or N
N
O O
or
O
O
or O
O
Even though different geometries are possible with larger molecules, Lewis structures still can help us with the process of predicting the shape. Therefore, the first step in predicting the shape of a molecule is to draw its Lewis structure. If the octet rule is obeyed throughout the molecule, each atom (except hydrogen) will be associated with four pairs of electrons. Some molecules include nonbonding lone-pair electrons, but all molecules must contain some bonding electrons or they would not be molecules! A basic rule of physics is that opposite charges attract and like charges repel. Negatively charged electrons are attracted to a positively charged nucleus in every case. However, the electrons all have the same charge and therefore are found as far from each other in space as possible while still maintaining their attraction to the positively charged nucleus. Groups of negatively charged electrons repel one another. The most stable arrangement is the one in which the mutually repelling electron groups are as far apart as possible. In turn, this determines the atomic arrangement and the shape of the molecule. We illustrate the procedure for predicting the shape of a molecule with methane, a greenhouse gas. 1. Determine the number of outer electrons associated with each atom in the molecule. The carbon atom (Group 4A) has four outer electrons; each of the four hydrogen atoms contributes one electron. This gives 4 1 (4 3 1), or 8 outer electrons. 2. Arrange the outer electrons in pairs to satisfy the octet rule. This may require single, double, or triple bonds. For the methane molecule, use the eight outer electrons to form four single bonds (four electron pairs) around the central carbon atom. This is the Lewis structure. H H H C H or H C H H H
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Although this structure seems to imply that the CH4 molecule is flat, it is not. In fact, the methane molecule is tetrahedral, as we will see in the next step. 3. Assume that the most stable molecular shape has the bonding electron pairs as far apart as possible. (Note: In other molecules we need to consider nonbonding electrons as well, but CH4 has none.) The four bonding electron pairs around the carbon atom in CH4 repel one another, and in their most stable arrangement they are as far from one another as possible. As a result, the four hydrogen atoms also are as far from one another as possible. This shape is tetrahedral, because the hydrogen atoms correspond to the corners of a tetrahedron, a four-cornered geometric shape with four equal triangular sides, sometimes called a triangular pyramid. One way to describe the shape of a CH4 molecule is by analogy to Figure 3.10 the base of a folding music stand. The four C–H bonds correspond to the The legs and the shaft of a music stand approximate the geometry of the bonds in a tetrahedral molecule such as three evenly spaced legs and the vertical shaft of the stand (Figure 3.10). methane. The angle between each pair of bonds is 109.5°. The tetrahedral shape of a CH4 molecule has been experimentally confirmed. Indeed, it is one of the most common atomic arrangements in nature, particularly in carbon-containing molecules.
Consider This 3.7
Methane: Flat or Tetrahedral?
a. If the methane molecule really were flat, as the two-dimensional Lewis structure seems to indicate, what would the H–C–H bond angle be? b. Offer a reason why the tetrahedral shape, not the two-dimensional flat shape, is more advantageous for this molecule. c. Consider the music stand shown in Figure 3.10. In the analogy of shape using a music stand, where would the carbon atom be located? Where would each of the hydrogen atoms lie? Answer a. 90° (at right angles). The two H atoms across from each other would be at 180°.
Chemists represent molecules in several different ways. The simplest, of course, is the chemical formula itself. In the case of methane, this is simply CH4. Another is the Lewis structure, but again this is only a two-dimensional representation that gives information about the outer electrons. Figure 3.11 shows these two representations as well as two others that are three-dimensional in appearance. One has a wedge-shaped line that represents a bond coming out of the paper in a direction generally toward the reader. The dashed wedge in the same structural formula represents a bond pointing away from the reader. The two solid lines lie in the plane of the paper. The other, a space-filling model, was drawn with the help of a molecular modeling program. Space-filling models enclose the volume occupied by electrons in an atom or molecule. Seeing and manipulating physical models, either in the classroom or laboratory, can also help you visualize the structure of molecules.
H H C H H
H H
C
H H
C
H
H H
109.5⬚ H
(a)
Figure 3.11 Representations of CH4. (a) Lewis structures and structural formula; (b) Space-filling model.
(b)
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Section 2.9 discussed replacement of NH3 as a refrigerant gas by CFCs. The role of ammonia in the nitrogen cycle is the subject of Section 6.9 and Section 11.8 describes the importance of NH3 in agriculture.
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H
N
H
H
N H
107.3⬚ H
H (a)
(b)
Figure 3.12 Representations of NH3. (a) Lewis structures and structural formula; (b) Space-filling model.
Not all outer electrons reside in bonding pairs. In some molecules, the central atom has nonbonding electron pairs, also called lone pairs. For example, Figure 3.12 shows the ammonia molecule in which nitrogen completes its octet with three bonding pairs and one nonbonding pair. A nonbonding electron pair effectively occupies greater space than a bonding pair of electrons. Consequently, the nonbonding pair repels the bonding pairs somewhat more strongly than the bonding pairs repel one another. This stronger repulsion forces the bonding pairs closer to one another, creating an H–N–H angle slightly less than the predicted 109.5° associated with a regular tetrahedron. The experimental value of 107.3° is close to the tetrahedral angle, again indicating that our model is reasonably reliable. The shape of a molecule is described in terms of its arrangement of atoms, not electrons. The hydrogen atoms of NH3 form a triangle with the nitrogen atom above them at the top of the pyramid. Thus, ammonia is said to have a trigonal pyramidal shape. Going back to the analogy of the folding music stand (see Figure 3.10), you could expect to find hydrogen atoms at the tip of each leg of the music stand. This places the nitrogen atom at the intersection of the legs with the shaft, with the nonbonded electron pair forming around the shaft of the stand. The water molecule is bent, illustrating yet another shape. There are eight outer electrons on the central oxygen atom: one from each of the two hydrogen atoms plus six from the oxygen atom (Group 6A). These eight electrons are distributed in two bonding and two lone pairs of electrons (Figure 3.13a). If these four pairs of electrons were arranged as far apart as possible, we might predict the H–O–H bond angle to be 109.5°, the same as the H–C–H bond angle in methane. However, unlike methane, water has two nonbonding pairs of electrons. The repulsion between the two nonbonding pairs causes the bond angle to be less than 109.5°. Experiments indicate a value of approximately 104.5°.
Your Turn 3.8
Predicting Molecular Shapes, Part 1
Using the strategies just described, sketch the shape for each of these molecules. a. CCl4 (carbon tetrachloride) b. CCl2F2 (Freon-12; dichlorodifluoromethane) c. H2S (hydrogen sulfide) Answer a. Total outer electrons: 4 1 4(7) 5 32. Eight of these electrons form 4 single bonds around the central C atom, one to each Cl atom. The other 24 are in 12 nonbonding pairs on the 4 Cl atoms. The bonding electron pairs on C arrange themselves to maximize the separation, and the molecule is tetrahedral.
Cl Cl C Cl Cl
Cl
or
Cl
C Cl
Cl Cl
or
Cl C Cl 109.5⬚Cl
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H
H O H
O
H
O
119
H
H 104.5⬚
(a)
(b)
Figure 3.13 Representations of H2O. (a) Lewis structures and structural formula; (b) Space-filling model.
We already looked at the structures of several molecules important for understanding the chemistry of climate change. What about the structure of the carbon dioxide molecule? With 16 outer electrons, the C atom contributes 4 electrons and 6 come from each of the 2 oxygen atoms. If only single bonds were involved, each atom would not have an octet. But the octet rule still can be obeyed if the central carbon atom forms a double bond with each of the 2 oxygen atoms, thus sharing 4 electrons. What is the shape of the CO2 molecule? Again, groups of electrons repel one another, and the most stable configuration provides the furthest separation of the negative charges. In this case, the groups of electrons are the double bonds, and these are furthest apart with an O–C–O bond angle of 180°. The model predicts that all three atoms in a CO2 molecule will be in a straight line and that the molecule will be linear. This is, in fact, the case as shown in Figure 3.14.
O C O
O
C
O
O
C
Revisit Section 2.3 for more about drawing Lewis structures for molecules with double bonds.
O
180⬚
(a)
(b)
Figure 3.14 Representations of CO2. (a) Lewis structures and structural formula; (b) Space-filling model.
We applied the idea of electron pair repulsion to molecules in which there are four groups of electrons (CH4, NH3, and H2O) and two groups of electrons (CO2). Electron pair repulsion also applies reasonably well to molecules that include three, five, or six groups of electrons. In most molecules, the electrons and atoms are still arranged to keep the separation of the electrons at a maximum. This logic accounts for the bent shape we associated with the ozone molecule. The Lewis structure for the ozone (O3) molecule with its 18 outer electrons contains a single bond and a double bond, and the central oxygen atom carries a nonbonding lone pair of electrons. Thus, the central O atom has three groups of electrons: the pair that makes up the single bond, the two pairs that constitute the double bond, and the lone pair. These three groups of electrons repel one another, and the minimum energy of the molecule corresponds to their furthest separation. This occurs when the electron groups are all in the same plane and at an angle of about 120° from one another. We predict, therefore, that the O3 molecule should be bent, and the angle made by the three atoms should be approximately 120°. Experiments show the bond angle to be 117°, just slightly smaller than the prediction (Figure 3.15). The nonbonding electron pair on the central oxygen atom occupies an effectively greater volume than bonding pairs of electrons, causing a greater repulsion force responsible for the slightly smaller bond angle.
Your Turn 3.9
Predicting Molecular Shapes, Part 2
Using the strategies just described, predict and sketch the shapes of SO2 (sulfur dioxide) and SO3 (sulfur trioxide). Hint: Because S and O are in the same group on the periodic table, the structures for SO2 and O3 will be closely related.
The O3 molecule is best represented by two equivalent resonance structures. Again see Section 2.3.
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O O O
O
O
O
O
O
O
117⬚
(a)
(b)
Figure 3.15 Representations of O3. (a) Lewis structures and structural formula for one resonance form; (b) Space-filling model.
As promised, in this section we helped you see that molecules have different shapes, ones that can be predicted. In the next, we return to our story of greenhouse gases, putting your knowledge of shapes to work to help you understand why not all gases are greenhouse gases.
3.4
|
Vibrating Molecules and the Greenhouse Effect
How do greenhouse gases trap heat, keeping our planet at more or less comfortable temperatures? In part, the answer lies in how molecules respond to photons of energy. This topic is complex, but even so, we can give you enough basics so you can understand how the greenhouse gases in our atmosphere function. At the same time, we’ll reveal why some gases do not trap heat. We begin this topic by revisiting the interaction of ultraviolet (UV) light with molecules, something we discussed earlier in Chapter 2 in connection with the ozone layer. You saw that a photon in the UV region of the electromagnetic spectrum had sufficient energy to break some covalent bonds. In particular, you saw that UV-C could break the bonds in O2 and that photons of lower energy (UV-B) could break the bonds in O3. Put another way, both the ozone and the oxygen molecule can absorb UV radiation. When this absorption occurs, an oxygen-to-oxygen bond is broken. Fortunately, IR photons do not contain enough energy to cause chemical bonds to break. Instead, a photon of IR radiation can add energy to the vibrations in a molecule. Depending on the molecular structure, only certain vibrations are possible. The energy of the incoming photon must correspond exactly to the vibrational energy of the molecule for the photon to be absorbed. This means that different molecules absorb IR radiation at different wavelengths and thus vibrate at different energies. We illustrate these ideas with the CO2 molecule, representing the atoms as balls and the covalent bonds as springs. Every CO2 molecule is constantly vibrating in the four ways pictured in Figure 3.16. The arrows indicate the direction of motion of each atom during each vibration. The atoms move forward and backward along the arrows. Vibrations a and b are stretching vibrations. In vibration a, the central carbon atom is
(a)
(b)
(c)
(d)
Figure 3.16 Molecular vibrations in CO2. Each spring represents a C5O double bond. Vibrations a and b are stretching vibrations; c and d are bending vibrations.
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stationary and the oxygen atoms move back and forth (stretch) in opposite directions away from the central atom. Alternatively, the oxygen atoms can move in the same direction and the carbon atom in the opposite direction (vibration b). Vibrations c and d look very much alike. In both cases, the molecule bends from its normal linear shape. The bending counts as two vibrations because it occurs in either of two possible planes. In vibration c the molecule is shown bending up and down in the plane of the paper on which the diagram is printed, whereas in vibration d the molecule is shown bending out of the plane of the paper. If you ever examined a spring, you probably noticed that more energy is required to stretch it than to bend it. Similarly, more energy is required to stretch a CO2 molecule than to bend it. This means that more energetic photons, those with shorter wavelengths, are needed to add energy to stretching vibrations a or b than to add energy to bending vibrations c or d. For example, absorption of IR radiation with a wavelength of 15.0 micrometers (μm) adds energy to the bending vibrations (c and d). When that occurs, the atoms move farther from their equilibrium positions and move faster (on average) than they do normally. For the same thing to happen with vibration b, higher energy radiation having a wavelength of 4.3 μm is required. Together, vibrations b, c, and d account for the greenhouse properties of carbon dioxide. In contrast, direct absorption of IR radiation does not add energy to vibration a. In a CO2 molecule, the average concentration of electrons is greater on the oxygen atoms than on the carbon atom. This means that the oxygen atoms carry a partial negative charge relative to the carbon atom. As the bonds stretch, the positions of the electrons change, thereby changing the charge distribution in the molecule. Because of the linear shape and symmetry of CO2, the changes in charge distribution during vibration a cancel and no infrared absorption occurs. The infrared (heat) energy that molecules absorb can be measured with an instrument called an infrared spectrometer. IR radiation from a glowing filament is passed through a sample of the compound to be studied, in this case gaseous CO2. A detector measures the amount of radiation, at various wavelengths, transmitted by the sample. High transmission means low absorbance, and vice versa. This information is displayed graphically, where the relative intensity of the transmitted radiation is plotted versus wavelength. The result is the infrared spectrum of the compound. Figure 3.17 shows the infrared spectrum of CO2. The infrared spectrum shown in Figure 3.17 was acquired using a laboratory sample of CO2, but the same absorption takes place in the atmosphere. Molecules of CO2 that absorb specific wavelengths of infrared energy experience different fates. Some hold that extra energy for a brief time, and then reemit it in all directions as heat. Others collide with atmospheric molecules like N2 and O2 and can transfer some
Percent transmittance
100 80 60 40 20 (b) stretching
0 3.33
4.00
5.00 6.67 Wavelength (µm)
(c) and (d) bending 10.0
20.0
Figure 3.17 Infrared spectrum of carbon dioxide. The letters (b), (c), and (d) refer to the molecular vibrations shown in Figure 3.16.
121
A micrometer is equal to onemillionth of a meter: 1 μm 5 1 3 1026 m 5 1000 nm.
The property of electronegativity, a measure of an atom’s ability to attract bonded electrons, is discussed in Section 5.5.
Spectroscopy is the field of study that examines matter by passing electromagnetic energy through a sample.
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Percent transmittance
100 99 98 97 96 95 94 93 2.5
2.86
3.33
4.00 5.00 Wavelength (µm)
6.67
10.0
20.0
Figure 3.18 Infrared spectrum of water vapor.
You already learned about the role of CFCs in ozone depletion in Chapter 2. Nitrous oxide, N2O, is also called dinitrogen monoxide. You will encounter this gas again in Chapter 6.
of the absorbed energy to those molecules, also as heat. Through both of these processes, CO2 “traps” some of the infrared radiation emitted by the Earth, keeping our planet comfortably warm. This is what makes CO2 a greenhouse gas. Any molecule that can absorb photons of IR radiation can behave as a greenhouse gas. There are many such molecules. Water is by far the most important gas in maintaining Earth’s temperature, followed by carbon dioxide. Figure 3.18 shows the IR spectrum of H2O molecules absorbing IR radiation. However, methane, nitrous oxide, ozone, and chlorofluorocarbons (such as CCl3F) are among the other substances that help retain planetary heat.
Consider This 3.10
Bending and Stretching Water Molecules
a. Use Figure 3.18 to estimate the wavelengths corresponding to the strongest IR absorbencies for water vapor. b. Which wavelength do you predict represents bending vibrations and which represents stretching? Explain the basis of your predictions. Hint: Compare the IR spectrum of H2O with that of CO2.
Diatomic gases, such as N2 and O2, are not greenhouse gases. Although molecules consisting of two identical atoms do vibrate, the overall electric charge distribution does not change during these vibrations. Hence, these molecules cannot be greenhouse gases. Earlier we discussed this lack of overall electric charge distribution as the reason why stretching vibration a in Figure 3.17 was not responsible for the greenhouse gas behavior of CO2. So far, you have encountered two ways that molecules respond to radiation. Highly energetic photons with high frequencies and short wavelengths (such as UV radiation) can break bonds within molecules. The less energetic photons (such as IR radiation) cause an increase in molecular vibrations. Both processes are depicted in Figure 3.19, but the figure also includes another response of molecules to radiant energy that is probably a good deal more familiar to you. Longer wavelengths than those in the IR range have only enough energy to cause molecules to rotate faster. For example, microwave ovens generate electromagnetic radiation that causes water molecules to spin faster. The radiation generated in such a device is of relatively long wavelength, about a centimeter. Thus the energy per photon is quite low. As the H2O molecules absorb the photons and spin more rapidly, the resulting friction cooks
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Ultraviolet
Wavelength Visible
molecule dissociates
Infrared
molecule vibrates
123
longer
Microwave
molecule rotates
Figure 3.19 Molecular response to types of radiation.
your food, warms up the leftovers, or heats your coffee. The same region of the spectrum is used for radar. Beams of microwave radiation are sent out from a generator. When the beams strike an object such as an airplane, the microwaves bounce back and are detected by a sensor.
3.5
|
The Carbon Cycle: Contributions from Nature and Humans
In his book The Periodic Table, the late chemist, author, and World War II concentration camp survivor Primo Levi wrote eloquently about CO2. “This gas which constitutes the raw material of life, the permanent store upon which all that grows draws, and the ultimate destiny of all flesh, is not one of the principal components of air but rather a ridiculous remnant, an ‘impurity’ thirty times less abundant than argon, which nobody even notices. . . . [F]rom this ever renewed impurity of the air we come, we animals and we plants, and we the human species, with our four billion discordant opinions, our millenniums of history, our wars and shames, nobility and pride.” In the essay from which this quotation is taken, Levi traces a brief portion of the life history of a carbon atom from a piece of limestone (calcium carbonate, CaCO3), where it lies “congealed in an eternal present,” to a CO2 molecule, to a molecule of glucose in a leaf, and ultimately to the brain of the author. And yet that is not the final destination. “The death of atoms, unlike our own,” writes Levi, “is never irrevocable.” That carbon atom, already billions of years old, will continue to persist into the unimagined future. This marvelous continuity of matter, a consequence of its conservation, is beautifully illustrated by the carbon cycle (Figure 3.20). Even without Primo Levi’s poetic description, the cycle is important to our understanding of the human effects on the global ecosystem. It is certain that without the proper functioning of the carbon cycle, every aspect of life on Earth could undergo dramatic change.
Your Turn 3.11 a. b. c. d.
Understanding the Carbon Cycle
Which processes add carbon (in the form of CO2) to the atmosphere? Which processes remove carbon from the atmosphere? What are the two largest reservoirs of carbon? Which parts of the carbon cycle are most influenced by human activities?
The carbon cycle is a dynamic system, with all the processes illustrated happening simultaneously but at far different rates. Michael B. McElroy of Harvard University estimated, “The average carbon atom has made the cycle from sediments through the
Levi’s book, The Periodic Table was written in 1975. Over 6.9 billion people now inhabit Earth.
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Chapter 3 atmosphere (750 Gt) deforestation (1.5 Gt/yr) respiration (60 Gt/yr)
from ocean (90 Gt/yr) Atmospheric CO2
reforestation photosynthesis (61 Gt/yr) (0.5 Gt/yr) forest (610 Gt) sand and silt 3 (1.2 ⫻ 10 Gt)
soils 3 (1.6 ⫻ 10 Gt) carbonate minerals in rocks 7 (1.8 ⫻ 10 Gt) fossil fuels 7 (2.5 ⫻ 10 Gt)
to ocean (92 Gt/yr) burning fossil fuel (6 Gt/yr) surface water 3 (1.0 ⫻ 10 Gt)
deep ocean 4 (3.8 ⫻ 10 Gt)
Figure 3.20 A gigatonne (Gt) is a billion metric tons, or about 2200 billion pounds. For comparison, a fully loaded 747 jet weighs about 800,000 lb. It would take nearly 3 million 747s to have a total mass of 1 Gt.
The global carbon cycle. The numbers show the quantity of carbon, expressed in gigatonnes (Gt), that is stored in various carbon reservoirs (black numbers) or moving through the system per year (red numbers).
more mobile compartments of the Earth back to sediments some 20 times over the course of Earth’s history.” CO2 in the air today may have been released from campfires burning more than a thousand years ago. Notice the presence of both natural emission and removal mechanisms. Respiration adds carbon dioxide to the atmosphere, and photosynthesis removes it. Similarly, the oceans both absorb and emit carbon dioxide. As members of the animal kingdom, we Homo sapiens participate in the carbon cycle along with our fellow creatures. As is true for any animal, we inhale and exhale, ingest and excrete, live and die. In addition though, human civilization relies on processes that put much more carbon into the atmosphere than they remove (Figure 3.21). Widespread burning of coal for electricity production, of petroleum products for transportation, and of natural gas for home heating all transfer carbon from the largest underground carbon reservoir into the atmosphere. Another human influence on CO2 emissions is deforestation by burning, a practice that releases about 1.5 Gt of carbon to the atmosphere each year. It is estimated
Transportation 17%
Residential/ commercial 9% Deforestation 22%
Power/heating stations 28% Industry 24%
Figure 3.21 Global carbon dioxide emissions by end use. Source: IPCC Fourth Assessment Report, Working Group III, 2007.
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that forested land the size of two football fields is lost every second of every day from the rain forests of the world. Although firm numbers are rather elusive, Brazil continues as the country with the greatest annual loss of rain forest acreage; over 5.4 million acres of Amazon rain forest is vanishing each year. Trees, those very efficient absorbers of carbon dioxide, are removed from the cycle through deforestation. If the wood is burned, vast quantities of CO2 are generated; if it is left to decay, that process also releases carbon dioxide, but more slowly. Even if the lumber is harvested for construction purposes and the land is replanted in cultivated crops, the loss in CO2-absorbing capacity may approach 80%. The total quantity of carbon released by the human activities of deforestation and burning fossil fuels is about 7.5 Gt per year. About half of this is eventually recycled into the oceans and the biosphere, which serve as carbon sinks, natural processes that remove CO2 from the atmosphere. These processes do not always remove CO2 with the speed required by the ever-increasing concentrations of CO2. Much of the CO2 emitted stays in the atmosphere, adding between 3.1 and 3.5 Gt of carbon per year to the existing base of 750 Gt noted in Figure 3.20. We are concerned primarily with the relatively rapid increase in atmospheric carbon dioxide, because the excess CO2 is implicated in global warming. Therefore, it would be useful to know the mass (Gt) of CO2 added to the atmosphere each year. In other words, what mass of CO2 contains 3.3 Gt of carbon, the midpoint between 3.1 Gt and 3.5 Gt? Answering this question requires that we return to some more quantitative aspects of chemistry.
3.6 |
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Remember that the natural “greenhouse effect” makes life on Earth possible. Problems occur when the amount of greenhouse gases increases faster than the sinks can accommodate the increases. The result is the enhanced greenhouse effect.
Quantitative Concepts: Mass
To solve the problem just posed, we need to know how the mass of C is related to the mass of CO2. Regardless of the source of CO2, its chemical formula is stubbornly the same. The mass percent of C in CO2 is also unwavering, and therefore we must calculate the mass percent of C in CO2, based on the formula of the compound. As you work through this and the next section, keep in mind that we are seeking a value for that percentage. The approach requires the use of the masses of the elements involved. But this raises an important question: How much does an individual atom weigh? The mass of an atom is mainly attributable to the neutrons and protons in the nucleus. Thus, elements differ in mass because their atoms differ in composition. Rather than using absolute masses of individual atoms, chemists have found it convenient to employ relative masses—in other words, to relate the masses of all atoms to some convenient standard. The internationally accepted mass standard is carbon-12, the isotope that makes up 98.90% of all carbon atoms. C-12 has a mass number of 12 because each atom has a nucleus consisting of 6 protons and 6 neutrons plus 6 electrons outside the nucleus. The periodic table in the text shows that the atomic mass of carbon is 12.01, not 12.00. This is not an error; it reflects the fact that carbon exists naturally as three isotopes. Although C-12 predominates, 1.10% of carbon is C-13, the isotope with six protons and seven neutrons. In addition, natural carbon contains a trace of C-14, the isotope with six protons and eight neutrons. The tabulated mass value of 12.01 is often called by the name atomic weight, an average that takes into consideration the masses and percent natural abundance of all naturally occurring isotopes of carbon. This isotopic distribution and average mass of 12.01 characterize carbon obtained from any chemical source—a graphite (“lead”) pencil, a tank of gasoline, a loaf of bread, a lump of limestone, or your body. The radioactive isotope carbon-14, although present only in trace amounts, provides direct evidence that the combustions of fossil fuels is the predominant cause of the rise in atmospheric CO2 concentrations over the past 150 years. In all living things, only 1 out of 1012 carbon atoms is a C-14 atom. A plant or animal constantly exchanges CO2 with the environment, and this maintains a constant C-14 concentration in the organism. However, when the organism dies, the biochemical processes that exchange
Isotopes and the relative masses of subatomic particles were discussed in Section 2.2.
The term atomic weight is a familiar (but not technically correct) term used for the relative scale of atomic masses.
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You will learn to write equations for nuclear reactions in Section 7.2.
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carbon stop functioning and the C-14 is no longer replenished. This means that after the death of the organism, the concentration of C-14 decreases with time because it undergoes radioactive decay to form N-14. Coal, oil, and natural gas are remnants of plant life that died hundreds of millions of years ago. Hence, in fossil fuels, and in the carbon dioxide released when fossil fuels burn, the level of C-14 is essentially zero. Careful measurement show that the concentration of C-14 in atmospheric CO2 has recently decreased. This strongly suggests that the origin of the added CO2 is indeed the burning of fossil fuels, a decidedly human activity.
Your Turn 3.12
Isotopes of Nitrogen
Nitrogen (N) is an important element in the atmosphere and in biological systems. It has two naturally occurring isotopes: N-14 and N-15. a. b. c. d.
Figure 3.22 Six tennis balls have a greater mass than six golf balls.
Use the periodic table to find the atomic number and atomic mass of nitrogen. What is the number of protons, neutrons, and electrons in a neutral atom of N-14? Compare your answers for part b with those for a neutral atom of N-15. Given the atomic mass of nitrogen, which isotope has the greatest natural abundance?
Having reviewed the meaning of isotopes, we return to the matter at hand—the masses of atoms and particularly the atoms in CO2. Not surprisingly, it is impossible to weigh a single atom because of its extremely small mass. A typical laboratory balance can detect a minimum mass of 0.1 mg; this corresponds to 5 3 1018 carbon atoms, or 5,000,000,000,000,000,000 carbon atoms. An atomic mass unit is far too small to measure in a conventional chemistry laboratory. Rather, the gram is the chemist’s mass unit of choice. Therefore, scientists use exactly 12 g of carbon-12 as the reference for the atomic masses of all the elements. We define atomic mass as the mass (in grams) of the same number of atoms that are found in exactly 12 g of carbon-12. This number of atoms is, of course, very large. This important chemical number is named after an Italian scientist with the impressive name of Count Lorenzo Romano Amadeo Carlo Avogadro di Quaregna e di Ceretto (1776–1856). (His friends called him Amadeo.) Avogadro’s number is the number of atoms in exactly 12 g of C-12. Avogadro’s number, if written out, is 602,000,000,000,000,000,000,000. It is more compactly written in scientific notation as 6.02 3 1023. This is the incredible number of atoms in 12 g of carbon, no more than a tablespoonful of soot! Avogadro’s number counts a large collection of atoms, much like the term dozen counts a collection of eggs. It does not matter if the eggs are large or small, brown or white, “organic” or not. No matter, for if there are 12 eggs, they are still counted as a dozen. A dozen ostrich eggs has a greater mass than a dozen quail eggs. Figure 3.22 illustrates this point with a half-dozen tennis and a half-dozen golf balls. Like atoms of different elements, the masses of a tennis ball and a golf ball differ. The number of balls is the same—six in each bag, a half dozen.
Skeptical Chemist 3.13
Marshmallows and Pennies
Avogadro’s number is so large that about the only way to hope to comprehend it is through analogies. For example, one Avogadro’s number of regular-sized marshmallows, 6.02 3 1023 of them, would cover the surface of the United States to a depth of 650 miles. Or, if you are more impressed by money than marshmallows, assume 6.02 3 1023 pennies were distributed evenly among the approximately 7 billion inhabitants of the Earth. Every man, woman, and child could spend $1 million every hour, day and night, and half of the pennies would still be left unspent at death. Can these fantastic claims be correct? Check one or both, showing your reasoning. Come up with an analogy of your own.
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Knowledge of Avogadro’s number and the atomic mass of any element permit us to calculate the average mass of an individual atom of that element. Thus, the mass of 6.02 3 1023 oxygen atoms is 16.00 g, the atomic mass from the periodic table. To find the average mass of just one oxygen atom, we must divide the mass of the large collection of atoms by the size of the collection. In chemist’s terms, this means dividing the atomic mass by Avogadro’s number. Fortunately, calculators help make this job quick and easy. 16.00 g oxygen 6.02 3 10
23
oxygen atoms
5 2.66 3 10223 g oxygen / oxygen atom
This very small mass confi rms once again why chemists do not generally work with small numbers of atoms. We manipulate trillions at a time. Therefore, practitioners of this art need to measure matter with a sort of chemist’s dozen—a very large one, indeed. To learn about it, read on . . . but only after stopping to practice your new skill.
Your Turn 3.14
Calculating Mass of Atoms Calculation Tip
a. Calculate the average mass in grams of an individual atom of nitrogen. b. Calculate the mass in grams of 5 trillion nitrogen atoms. c. Calculate the mass in grams of 6 3 1015 nitrogen atoms.
Predict: Will the answer be a large or a small number?
Answer 14.01 g nitrogen a. 5 2.34 3 10223 g nitrogen /nitrogen atom 6.02 3 1023 nitrogen atoms
3.7
|
Check: Does your answer match your prediction and is it reasonable?
Quantitative Concepts: Molecules and Moles
Chemists have another way of communicating the number of atoms, molecules, or other small particles present. This is to use the term mole (mol), defined as containing an Avogadro’s number of objects. The term is derived from the Latin word to “heap,” or “pile up.” Thus, 1 mol of carbon atoms consists of 6.02 3 1023 C atoms, 1 mol of oxygen gas is made up of 6.02 3 1023 oxygen molecules, and 1 mol of carbon dioxide molecules corresponds to 6.02 3 1023 carbon dioxide molecules. As you already know from previous chapters, chemical formulas and equations are written in terms of atoms and molecules. For example, reconsider the equation for the complete combustion of carbon in oxygen. C ⫹ O2
CO2
When used together with a number, mol is an abbreviation for mole.
[3.1]
This equation tells us that one atom of carbon combines with one molecule of oxygen to yield one molecule of carbon dioxide. Thus it reflects the ratio in which the particles interact. It is equally correct to say that 10 C atoms react with 10 O2 molecules (20 O atoms) to form 10 CO2 molecules. Or, putting the reaction on a grander scale, we can say 6.02 3 1023 C atoms combine with 6.02 3 1023 O2 molecules (12.0 3 1023 O atoms) to yield 6.02 3 1023 CO2 molecules. The last statement is equivalent to saying: “one mole of carbon plus one mole of oxygen yields one mole of carbon dioxide.” The point is that the numbers of atoms and molecules taking part in a reaction are proportional to the numbers of moles of the same substances. The ratio of two oxygen atoms to one carbon atom remains the same regardless of the number of carbon dioxide molecules, as summarized in Table 3.1.
There are 2 mol of oxygen atoms, O, in every mole of oxygen molecules, O2.
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Table 3.1
Ways to Interpret a Chemical Equation
C
O2
1
→
CO2
1 atom
1 molecule
1 molecule
6.02 3 1023 atoms
6.02 3 1023 molecules
6.02 3 1023 molecules
1 mol
1 mol
1 mol
In the laboratory and the factory, the quantity of matter required for a reaction is often measured by mass. The mole is a practical way to relate number of particles to the more easily measured mass. The molar mass is the mass of one Avogadro’s number, or mole, of whatever particles are specified. For example, the mass of a mole of carbon atoms, rounded to the nearest tenth of a gram, is 12.0 g. A mole of oxygen atoms has a mass of 16.0 g. But we can also speak of a mole of O2 molecules. Because there are two oxygen atoms in each oxygen molecule, there are two moles of oxygen atoms in each mole of molecular oxygen, O2. Consequently, the molar mass of O2 is 32.0 g, twice the molar mass of O. Some refer to this as the molecular mass or molecular weight of O2, emphasizing its similarity to atomic mass or atomic weight. The same logic for the molar mass of the element O2 applies to compounds, which brings us, at last, to the composition of carbon dioxide. The formula, CO2, reveals that each molecule contains one carbon atom and two oxygen atoms. Scaling up by 6.02 3 1023, we can say that each mole of CO2 consists of 1 mol of C and 2 mol of O atoms (see Table 3.1). But remember that we are interested in the mass composition of carbon dioxide—the number of grams of carbon per gram of CO2. This requires the molar mass of carbon dioxide, which we obtain by adding the molar mass of carbon to twice the molar mass of oxygen: 1 mol CO2 ⫽ 1 mol C ⫹ 2 mol O ⫽ 1 mol C ⫻
12.0 g C 16.0 g O ⫹ 2 mol O ⫻ 1 mol C 1 mol O
⫽ 12.0 g C ⫹ 32.0 g O 1 mol CO2 ⫽ 44.0 g CO2 This procedure is routinely used in chemical calculations, where molar mass is an important property. Some examples are included in the next activity. In every case, you multiply the number of moles of each element by the corresponding atomic mass in grams and add the result.
Your Turn 3.15
Molecular Molar Mass
Calculate the molar mass of each of these greenhouse gases. a. O3 (ozone) b. N2O (dinitrogen monoxide or nitrous oxide) c. CCl3F (Freon-11; trichlorofluoromethane) Answer a. 1 mol O3 ⫽ 3 mol O
⫽ 3 mol O ⫽ 48.0 g O3
⫻
16.0 g O 1 mol O
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We started out on this mathematical excursion so that we could calculate the mass of CO2 produced from burning 3.3 Gt of carbon. We now have all the pieces assembled. Out of every 44.0 g of CO2, 12.0 g is C. This mass ratio holds for all samples of CO2, and we can use it to calculate the mass of C in any known mass of CO2. More to the point, we can use it to calculate the mass of CO2 released by any known mass of carbon. It only depends on how we arrange the ratio. The C-to-CO2 12.0 g C 44.0 g CO2 ratio is , but it is equally true that the CO2-to-C ratio is . 44.0 g CO2 12.0 g C For example, we could compute the number of grams of C in 100.0 g CO2 by setting up the relationship in this manner. 100.0 g CO2 ⫻
12.0 g C ⫽ 27.3 g C 44.0 g CO2
The fact that there is 27.3 g of carbon in 100.0 g of carbon dioxide is equivalent to saying that the mass percent of C in CO2 is 27.3%. Note that carrying along the units “g CO2” and “g C” helps you do the calculation correctly. The unit “g CO2” can be canceled, and you are left with “g C.” Keeping track of the units and canceling where appropriate are useful strategies in solving many problems. This method is sometimes called “unit analysis.”
Your Turn 3.16
Mass Ratios and Percents
Calculation Tip Predict:
a. Calculate the mass ratio of S in SO2. b. Find the mass percent of S in SO2. c. Calculate the mass ratio and the mass percent of N in N2O. Answers a. The mass ratio is found by comparing the molar mass of S with the molar mass of SO2. 32.1 g S 0.501 g S 5 64.1 g SO2 1.00 g SO2
Will the answer be larger or smaller than the given value? What are the units?
Check: Does the answer match your prediction? Have units canceled, leaving the one needed for the answer?
b. To find the mass percent of S in SO2, multiply the mass ratio by 100. 0.501 g S 3 100 5 50.1% S in SO2 1.00 g SO2
To find the mass of CO2 that contains 3.3 gigatons (Gt) of C, we use a similar approach. We could convert 3.3 Gt to grams, but it is not necessary. As long as we use the same mass unit for C and CO2, the same numerical ratio holds. Compared with our last calculation, this problem has one important difference in how we use the ratio. We are solving for the mass of CO2, not the mass of C. Look carefully at the units this time. 3.3 Gt C ⫻
44.0 Gt CO2 ⫽ 12 Gt CO2 12.0 Gt C
Once again the units cancel and we are left with Gt of CO2. Our burning question, “What is the mass of CO2 added to the atmosphere each year from the combustion of fossil fuels?” has finally been answered: 12 gigatons. Of course, we also managed to demonstrate the problem-solving power of chemistry and to introduce five of its most important ideas: atomic mass, molecular mass, Avogadro’s number, mole, and molar mass. The next few activities provide opportunities to practice your skill with these concepts.
The question to be answered is: What mass of CO2 contains 3.3 Gt of carbon?
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Your Turn 3.17
SO2 from Volcanoes
a. It is estimated that volcanoes globally release about 19 3 10 6 t (19 million metric tons) of SO2 per year. Calculate the mass of sulfur in this amount of SO2 . b. If 142 3 106 t of SO2 is released per year by fossil-fuel combustion, calculate the mass of sulfur in this amount of SO2. Answer a. The mass ratio of S to SO2 is known from Your Turn 3.16. 6
19 ⫻ 10 t SO2 ⫻
6
32.1 ⫻ 10 t S 6
64.1 ⫻ 10 t SO2
6
⫽ 9.5 ⫻ 10 t S
If you know how to apply these ideas, you have gained the ability to critically evaluate media reports about releases of C or CO2 (and other substances as well) and judge their accuracy. One can either take such statements on faith or check their accuracy by applying mathematics to the relevant chemical concepts. Obviously, there is insufficient time to check every assertion, but we hope that readers develop questioning and critical attitudes toward all statements about chemistry and society, including those found in this book.
Skeptical Chemist 3.18
Checking Carbon from Cars
A clean-burning automobile engine emits about 5 pounds of C in the form of CO2 for every gallon of gasoline it consumes. The average American car is driven about 12,000 miles per year. Using this information, check the statement that the average American car releases its own weight in carbon into the atmosphere each year. List the assumptions you make in solving this problem. Compare your list and your answer with those of your classmates.
3.8
|
Methane and Other Greenhouse Gases
Concerns about an enhanced greenhouse effect are based primarily on increases in concentrations of atmospheric CO2. However, other gases also play a role. Methane, nitrous oxide, chlorofluorocarbons, and even our friend ozone all take part in trapping heat in the atmosphere. Our level of concern regarding each of these gases is related to their concentration in the atmosphere but also to other important characteristics. The global atmospheric lifetime characterizes the time required for a gas added to the atmosphere to be removed. It is also referred to as the “turnover time.” Greenhouse gases also vary in their effectiveness in absorbing infrared radiation. This is quantified by the global warming potential (GWP), a number that represents the relative contribution of a molecule of the atmospheric gas to global warming. The GWP of carbon dioxide is assigned the reference value of 1; all other greenhouse gases are indexed with respect to it. Gases with relatively short lifetimes, such as water vapor, tropospheric ozone, tropospheric aerosols, and other ambient air pollutants, are distributed unevenly around the world. It is difficult to quantify their effect, and therefore GWP values are not usually assigned. Table 3.2 lists four greenhouse gases, their main sources, and their important properties in the climate change conversation.
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Table 3.2 Name and Chemical Formula
131
Examples of Greenhouse Gases Preindustrial Concentration (1750)
Concentration in 2008
Atmospheric Lifetime (years)
270 ppm
methane CH4
700 ppb
1760 ppb
12
Rice paddies, waste dumps, livestock
21
nitrous oxide N2O
275 ppb
322 ppb
120
Fertilizers, industrial production, combustion
310
0.56 ppb
102
Liquid coolants, foams
50-200*
Fossil fuel combustion, deforestation, cement production
Global Warming Potential
carbon dioxide CO2
CFC-12 CCl2F2
388 ppm
Anthropogenic Sources
1
8100
*A single value for the atmospheric lifetime of CO2 is not possible. Removal mechanisms take place at different rates. The range given is an estimate based on several removal mechanisms.
Your Turn 3.19
Greenhouse Gases on the Rise
Using the data in Table 3.2, calculate the percentage increases for CO2, CH4, and N2O since the beginning of the Industrial Revolution. Rank the three in order of their percentage increase.
The current atmospheric concentration of CH4 is about 50 times lower than that of CO2, but as an infrared absorber, methane is about 20 times more efficient than carbon dioxide. Fortunately, CH4 is quite readily converted to other chemical species by interaction with tropospheric free radicals, and therefore has a relatively short lifetime. By comparison, carbon dioxide is much less reactive. The primary removal mechanisms for CO2 are dissolution in oceans, and the much longer process of mineralization into carbonate rocks. Methane emissions arise from both natural and human sources. About 40% of total CH4 emissions come from natural sources, of which emanations from wetlands are by far the largest contributor. These marshy habitats are perfectly suited for anaerobic bacteria, those that can function without the use of molecular oxygen. As they decompose organic matter, many types of anaerobic bacteria produce methane, which then escapes into the atmosphere. In Alaska, Canada, and Siberia, however, much of the methane produced from thousands of years of decomposition has remained trapped underground by the permafrost. There is concern that melting of the surface in the Northern latitudes might trigger a massive release of methane into the atmosphere. There is geological evidence that such a release has occurred in the past, and led to higher global temperatures. Methane is also released from the oceans, where a substantial amount of it appears to be trapped in “cages” made of water molecules. Such deposits are referred to as methane hydrates. Australia’s Commonwealth Scientific and Industrial Research Organization (CSIRO) has been taking a series of ocean core drillings to gather evidence about methane hydrates and their role in global warming (Figure 3.23). There is concern that if some of these hydrates become unstable then large amounts of methane might rapidly be released to the atmosphere. Termites are another natural source of methane. These ubiquitous insects have special bacteria in their guts that allow them to metabolize cellulose, the main component of wood. But instead of making water and CO2, termites produce methane and CO2. Not only can they inflict direct damage to homes, but they also add to greenhouse gas concentrations. The sheer number of termites is staggering, estimated to be more than half a metric ton for every man, woman, and child on the planet!
Global atmospheric lifetime values, although useful for comparison, are best thought of as approximations.
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(a)
(b)
Figure 3.23 (a) A floating drilling platform used by the CSIRO. (b) A sample of methane hydrate from the continental shelf off the coast of Florida.
For more information on using biogas as a fuel, check Section 4.10.
The major human source of CH4 is agriculture, with the biggest culprits being rice cultivation and the raising of livestock. Rice is grown with its roots under water, where, again, anaerobic bacteria produce methane. Most of the methane is released to the atmosphere. Additional agricultural CH4 comes from an increasing number of cattle and sheep. The digestive systems of these ruminants (animals that chew their cud) contain bacteria that break down cellulose. In the process, methane is formed and released through belching and flatulence—about 500 liters of CH4 per cow per day! The ruminants of the Earth release a staggering 73 million metric tons of CH4 each year. Landfills add another large quantity of methane to the atmosphere. The chemistry occurring within our buried garbage is controlled by the same anaerobic bacteria found in wetlands and produces the same result. Some of this methane is captured (biogas) and burned as a fuel, but the vast majority is released into the atmosphere. The other main anthropogenic source of methane originates from our extraction of fossil fuels. Methane is often found with oil and coal deposits, and drilling and mining procedures release most of that methane to the atmosphere while recovering the liquid or solid products. There are also significant losses from transporting, purifying, and using natural gas.
Consider This 3.20
Methane Concentrations Stabilizing?
Concentrations of atmospheric methane have stabilized in recent years. Use the resources of the web to find some hypotheses about why the concentrations have leveled out.
The role of N2O in destroying stratospheric ozone was discussed in Section 2.8.
Another gas that contributes to global warming is nitrous oxide, also known as “laughing gas.” It has been used as an inhaled anesthetic for dental and medical purposes. Its sources and sinks are not as well established as are those for carbon dioxide and methane. The majority of N2O molecules in the atmosphere come from the bacterial removal of nitrate ion (NO32) from soils, followed by removal of oxygen. Agricultural practices, again linked to population pressures, can speed up the removal of reactive compounds of nitrogen from soils. Other sources include ocean upwelling, and stratospheric interactions of nitrogen compounds with high-energy oxygen atoms. Major anthropogenic sources of N2O are automobile catalytic converters, ammonia fertilizers, biomass burning, and certain industrial processes (nylon and nitric acid production). In the atmosphere, a typical N2O molecule persists for about 120 years,
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Table 3.3
Climate Change and Ozone Depletion: A Comparison Climate Change
Ozone Depletion
region of atmosphere
primarily the troposphere
the stratosphere
major players
H2O, CO2, CH4, and N2O
O3, O2, and CFCs
interaction with radiation
Molecules absorb IR radiation. This causes them to vibrate and return heat energy to the Earth.
Molecules absorb UV radiation. This causes one or more bonds in the molecule to break.
nature of problem
Greenhouse gases are increasing in concentration. In turn this is trapping more heat, causing an increase in the average global temperatures.
CFCs are causing a decrease in concentrations of O3 in the stratosphere. In turn, this is causing an increase in the UV radiation at the surface of the Earth.
absorbing and emitting infrared radiation. Over the past decade, global atmospheric concentrations of N2O have shown a slow but steady rise. A few comments need to be made about ozone, a gas we encountered in Chapter 2. Often there is confusion between the phenomena of climate change and ozone depletion. Both are often in the news, both involve complex atmospheric processes, and both have anthropogenic as well as natural sources. In fact, ozone itself can act like a greenhouse gas, but its efficiency depends very much on its altitude. It appears to have its maximum warming effect in the upper troposphere, around 10 km above the Earth. Therefore, depletion of ozone has a slight cooling effect in the stratosphere, and it may also promote slight cooling at Earth’s surface. Other differences are summarized in Table 3.3. Depletion of the stratospheric ozone layer is clearly not a principal cause of climate change. However, stratospheric ozone depletion and climate change are linked in an important way, through ozone-destroying substances. CFCs, HCFCs, and halons, all implicated in the destruction of stratospheric ozone, also absorb infrared radiation and are all greenhouse gases. Emissions of these synthetic gases have risen by 58% from 1990–2005, although their concentrations are still very low.
Your Turn 3.21
Comparing Greenhouse Gas Effectiveness
Multiplying GWP by tropospheric abundance provides a number that can be used to compare the warming effectiveness of a greenhouse gas. a. Compare the effectiveness of CFC-12 (CCl2F2) as a greenhouse gas relative to that of CO2. b. HFC-134a (CF3CH2F lifetime 5 13.8 years) is a replacement for Freon-12 and has a GWP value of 1300 and a tropospheric abundance of 7.5 parts per trillion (1998 data). Compare its effectiveness as a greenhouse gas to that of CO2. Hint: Use the same unit of tropospheric abundance for both gases. c. How do their global atmospheric lifetimes affect their overall ability to function as greenhouse gases?
3.9 |
133
How Warm Will the Planet Get?
“Prediction is very difficult, especially about the future.” Niels Bohr, one of the foremost contributors to our modern view of the atom, spoke these words years ago. His words still hold true today!
HCFCs were discussed in Section 2.12.
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The IPCC received the 2007 Nobel Peace Prize (shared with former U.S. Vice President Al Gore) for its work in understanding global warming.
The unique properties of water, including its unusually large specific heat, will be described in Chapter 5.
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Chapter 3
Although admittedly a difficult task, we still need to make predictions. To this end, in 1988, the United Nations Environment Programme and the World Meteorological Organization teamed up to establish the UN Intergovernmental Panel on Climate Change (IPCC). The IPCC was charged with assessing the data on climate change, not just the scientific data, but the socioeconomic information as well. Thousands of international scientists were involved in this review. In their fourth and most recent report published in 2007, the vast majority of scientists agreed on several key points: (1) The Earth is getting warmer; (2) Human activities (primarily the combustion of fossil fuels and deforestation) are responsible for much of the recent warming; and (3) If the rate of greenhouse gas emissions is not curtailed, our water resources, food supply, and even our health will suffer. The challenge, however, is to understand current climate change well enough to predict future changes and by doing so, to determine the decrease in emissions required to minimize harmful changes. To make predictions, scientists work with models. They design computer models of the oceans and the atmosphere that take into account the ability of each to absorb heat as well as to circulate and transport matter (Figure 3.24). If that weren’t difficult enough, the models must also include astronomical, meteorological, geological, and biological factors, ones that are often incompletely understood. Human influences, such as population, industrialization levels, and pollution emissions must also be included. Dr. Michael Schlesinger, who directs climate research at the University of Illinois, remarked: “If you were going to pick a planet to model, this is the last planet you would choose.” Climate scientists call the factors (both natural and anthropogenic) that influence the balance of Earth’s incoming and outgoing radiation by the term radiative forcings. Negative forcings have a cooling effect; positive forcings a warming effect. The primary forcings used in climate models are solar irradiance, greenhouse gas concentrations, land use, and aerosols. The effects of these forcings on the Earth’s energy balance are summarized in Figure 3.25. Red, orange, and yellow bars represent positive forcings, and blue bars indicate negative ones. Each forcing has an error bar associated with it; the larger the error bar, the more uncertain the value.
Figure 3.24
CO2
Climate scientists use computer simulations to understand future climate change.
Long-lived greenhouse gases
N2O
Anthropogenic
CH4 Ozone Surface albedo
Stratospheric
Halocarbons Tropospheric
Land use
Black carbon on snow
Natural
Direct effect Total aerosol Cloud albedo effect Solar irradiance –2
–1
1
2
Radiative forcing (W m–2)
Figure 3.25 Selected radiative forcings of climate from 1750 to 2005. The units are in watts per square meter (W m22 ), the light energy hitting a square meter of the Earth’s surface every second. Source: Adapted from Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change.
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135
Solar Irradiance (“solar brightness”) We can directly observe the natural seasonal variations in sunlight intensity. In the higher latitudes, temperatures are warmer in the summer. Compared to winter months, the Sun is higher in the sky and stays up longer. Across the globe, these variations essentially cancel, because when it is winter in the Northern Hemisphere it is summer in the Southern Hemisphere. Subtle periodic changes occur in the brightness of the Sun. The Earth’s orbit oscillates slightly over a 100,000-year period, changing its shape. In addition, the magnitude of the tilt of the Earth’s axis and the direction of that tilt both change over the course of several tens of thousands of years, affecting the amount of solar radiation hitting the Earth. Neither of these occurs on a time scale short enough to explain the recent warming, however. Additionally, sunspots occur in large numbers about every 11 years. You might think that dark spots on the Sun would mean a smaller amount of radiation hitting the Earth, but exactly the opposite is true. Sunspots occur when there is increased magnetic activity in the outer layers of the Sun, and the stronger magnetic fields stir up a larger amount of charged particles that emit radiation. Notably, the 17th and 18th centuries, sometimes called the “Little Ice Age” because of the below average temperatures in Europe, were preceded by a period of almost no sunspot activity. However, the solar brightness over those 11-year cycles varies only by about 0.1%. As you can see from Figure 3.25, this natural variability is the smallest of any positive forcing listed.
Your Turn 3.22
Periodic orbital eccentricities are a possible cause of the ice age oscillations shown in Figure 3.9.
During periods of high sunspot activity, the aurora borealis (“northern lights”) is more spectacular because of the greater number of charged particles striking the Earth’s atmosphere.
Radiation from the Sun
Sunlight strikes the Earth continually. Which types of light are emitted by the Sun? Which one makes up the largest percentage of sunlight? Hint: Refer back to Figure 2.7.
Greenhouse Gases These are the dominant anthropogenic forcings. Largest among these is CO2, constituting about two thirds of the warming from all greenhouse gases. However, as we explained in the previous section, methane, nitrous oxide, and other gases do contribute. Notice the relatively small contribution from “halocarbons” (CFCs and HCFCs) as shown in Figure 3.25. It has been estimated that without the ban on CFC production imposed by the Montreal Protocol, by 1990 the forcings from CFCs would have outweighed those from CO2. In sum, the positive forcings from greenhouse gases are more than 30 times greater than the natural changes in solar irradiance.
The Montreal Protocol was discussed in Section 2.11.
Land Use Changes in land use drive climate change because these changes alter the amount of incoming solar radiation that is absorbed by the surface of the Earth. The ratio of electromagnetic radiation reflected from a surface relative to the amount of radiation incident on it is called the albedo. In short, albedo is a measure of the reflectivity of a surface. The albedo of the Earth’s surface varies between about 0.1 and 0.9, as you can see from the values listed in Table 3.4. The higher the number, the more reflective the surface. As the seasons change, so does the albedo of the Earth. When a snow-covered area melts, the albedo decreases and more sunlight is absorbed, creating a positive feedback loop and additional warming. This effect helps to explain the greater increases in average temperature observed in the Arctic, where the amount of sea ice and permanent snow cover is decreasing. Similarly, when glaciers retreat and expose darker rock, the albedo decreases, causing further warming. Human activity also can change the Earth’s albedo, most notably through deforestation in the tropics. The crops we plant reflect more incoming light than does the
Earth has an average albedo of 0.39. In contrast, that of the Moon is about 0.12.
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Table 3.4
Albedo Values for Different Ground Covers
Surface fresh snow
Range of Albedo 0.80–0.90
old/melting snow
0.40–0.80
desert sand
0.40
grassland
0.25
deciduous trees
0.15–0.18
coniferous forest
0.08–0.15
tundra
0.20
ocean
0.07–0.10
dark green foliage of the rainforests, causing an increase in the albedo and hence resulting in cooling. In addition, sunlight is more consistent in the tropics, so changes in land use at low latitudes produce greater effects than changes in the polar regions. The conversion of tropical rainforest to crop and pastureland has more than offset the decrease in the amount of sea ice and snow cover near the poles. Therefore, the changes in the Earth’s albedo have caused a net cooling effect.
Consider This 3.23
White Roofs, Green Roofs
a. In 2009, U.S. Energy Secretary Steven Chu suggested that painting roofs white would be one way to combat global warming. Explain the reasoning behind this course of action. b. The idea of “green roofs” is also attracting attention. Planting gardens on rooftops has benefits in addition to those of white roofs. But such gardens also have limitations. Explain.
The term aerosol was defined in Section 1.11. The role of aerosols in acid rain will be discussed in Section 6.6.
Mt. Pinatubo eruption, 1991.
Aerosols A complex class of materials, aerosols have a correspondingly complex effect on climate. Many natural sources of aerosols exist, including dust storms, ocean spray, forest fires, and volcanic eruptions. Human activity can also release aerosols into the environment in the form of smoke, soot, and sulfate aerosols from coal combustion. The effect of aerosols on climate is probably the least well understood of the forcings listed in Figure 3.25. Tiny aerosol particles (,4 μm) are efficient at scattering incoming solar radiation. Other aerosols absorb incoming radiation, and still other particles both scatter and absorb. Both processes decrease the amount of radiation available for greenhouse gases to absorb and therefore have a cooling effect (negative forcing). In a dramatic example, the 1991 eruption of Mt. Pinatubo in the Philippines spewed over 20 million tons of SO2 into the atmosphere. In addition to providing spectacular sunsets for several months, the sulfur dioxide caused temperatures around the world to drop slightly. The results provided the climate modelers a mini-control experiment. The most reliable models were able to reproduce the cooling effect caused by the eruption. In addition to that direct cooling effect, aerosol particles can serve as nuclei for the condensation of water droplets and hence promote cloud formation. Clouds reflect incoming solar radiation, although the effects of increased cloud cover are more complex than this. Therefore, in both direct and indirect ways, aerosols counter the warming effects of greenhouse gases.
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1.0 0.5 0.0 1900
1950 Year
2000
Global land 1.0 0.5 0.0 1900
1950 Year
2000
Temperature anomaly (⬚C)
Global
Temperature anomaly (⬚C)
Temperature anomaly (⬚C)
The Chemistry of Global Climate Change
137 Global ocean
1.0 0.5 0.0 1900
1950 Year
2000
Figure 3.26 Climate model predictions of annual global mean surface temperatures for the 20th century. Black lines display temperature data relative to the average temperature for the years 1901–1950. The blue bands indicate the predicted temperature range using natural forcings only. The pink bands indicate the predicted temperature range using both natural and anthropogenic forcings. Source: Adapted from Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change.
Given the complexity inherent in all the forcings that we have just described, you can appreciate that assembling these forcings into a climate model is no easy task. Furthermore, once a model has been built, scientists have difficulty assessing its validity. However, scientists do have one trick in their back pockets. They can test climate models with known data sets as a means to tease apart the contributions of different forcings. For example, we know the temperature data of the 20th century. In Figure 3.26, the black lines represent the known data. Next examine the blue bands. These represent temperature ranges that were predicted by the climate model using only natural forcings. As you can see, the natural forcings do not map well onto the actual temperatures. Finally, examine the pink bands to see that when anthropogenic forcings are included, the temperature increases of the 20th century can be accurately reproduced. So although the last 30 years of warming were influenced by natural factors, the actual temperatures cannot be accounted for without including the effects of human activities.
Your Turn 3.24
Assessing Climate Models
Between 1950 and 2000, the climate models that used natural forcings only (blue bands in Figure 3.26) showed an overall cooling effect and thus did not match the observed temperatures. a. Name the forcings included in the models that only included natural forcings. b. List two additional forcings included in the models that more accurately recreate the temperatures of the 20th century (pink bands in Figure 3.26). Answer a. Aerosols (such as those from volcanic eruptions), solar irradiance.
The magnitude of future emissions, and hence the magnitude of future warming, depends on many factors. As you might expect, one is population. As of 2010, the global population stood at about 6.9 billion. Assuming that there will be more feet on the planet in the future, we humans are likely to have a larger carbon footprint, an estimate of the amount of CO2 and other greenhouse gas emissions in a given time frame, usually a year. Having more people to feed, clothe, house, and transport will require the consumption of more energy. In turn, this translates to more CO2 emissions, at least if using current energy sources. In addition, scientists who create climate models have to include values for two factors: (1) the rate of economic growth and (2) the rate of development of “green” (less carbon-intensive) energy sources. Again as you might expect, both are difficult to predict.
Chapter 0 introduced the concept of an ecological footprint. Carbon footprints are a subset of the more general term.
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Chapter 3 Global surface warming (⬚C)
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4.0 3.0 2.0 1.0 0.0 –1.0 1900
2000 Year
2100
Figure 3.27 Four model projections for temperature scenarios in the 21st century based on different socioeconomic assumptions. The black line is the data for the 20th century with the gray regions indicating the uncertainty in those values. The four dark lines represent projected 21st-century temperatures, with the wider lighter colored bands representing the uncertainty range for each scenario. Source: Adapted from Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change.
So what, if anything, can computer models tell us about the Earth’s future climate? Given the uncertainties that we have listed, hundreds of different projected temperature scenarios for the 21st century are possible. Figure 3.27 shows four of these, together with the actual temperature data for the 20th century. The four scenarios for 21st-century temperatures are based on different assumptions. The orange line assumes that emissions levels are kept at 2000 levels, admittedly an unrealistic target given the increases that already have occurred since 2000. Even with this most optimistic scenario, some additional warming will take place due to the persistence of CO2 in the atmosphere for years to come. Both the blue and green lines assume that the global population will increase to 9 billion by 2050 but then gradually decrease. However, the blue line includes the more rapid development of energy-efficient technologies, leading to lower CO2 emissions. The red line assumes a continually increasing population combined with a slower and less globally integrated transition to new, cleaner technologies. All of the lines point in the same direction—up. With some amount of future warming virtually ensured, we now turn our discussion to the consequences of climate change.
3.10
| The Consequences of Climate Change
Considering even the most extreme predictions of warming described in the last section, you may be thinking, “So what?” After all, the temperature changes predicted in Figure 3.27 are only a few degrees. At any single spot on the planet, the temperature fluctuates several times that amount daily. An important distinction needs to be made between the terms climate and weather. Weather includes the daily highs and lows, the drizzles and downpours, the blizzards and heat waves, and the fall breezes and hot summer winds, all of which have relatively short durations. In contrast, climate describes regional temperatures, humidity, winds, rain, and snowfall over decades, not days. And while the weather varies on a daily basis, our climate has stayed relatively uniform over the last 10,000 years. The values quoted for the “average global temperature” are but one measure of climate phenomena. The key point is that relatively small changes in average global temperature can have huge effects on many aspects of our climate. In addition to modeling various future temperature scenarios (see Figure 3.27), the 2007 IPCC report estimated the likelihood of various consequences. The report employed descriptive terms (“judgmental estimates of confidence”) to help both policy
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Table 3.5 Term
139
Judgmental Estimates of Confidence Probability That a Result Is True (%)
virtually certain
.99
very likely
90–99
likely
66–90
medium likelihood
33–66
unlikely
10–33
very unlikely
1–10
Source: Adapted from Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change.
makers and the general public better understand the inherent uncertainty of the data. These terms are listed in Table 3.5 together with the definitions that will continue to be used in all subsequent updates to the IPCC report. Conclusions from the 2007 IPCC report are listed in Table 3.6. For example, it was judged very unlikely that all of the observed global warming was due to natural climate variability. Rather, the scientific evidence strongly supports the position that human activity is a significant factor causing the increase in average global temperature observed over the last century. Furthermore, from the scientific evidence for global warming, it was judged virtually certain that human activities were the main drivers of recent warming. Check Table 3.6 for other conclusions relevant to any discussion of global climate change. Many scientific organizations, including the American Association for the Advancement of Science and the American Chemical Society, also have recognized the threats posed by climate change. In an open letter to United States senators, the organizations cited sea level rises, more extreme weather events, increased water scarcity, and disturbances of local ecosystems as likely eventualities of a warmer planet. To conclude this section, we describe these and other outcomes we can expect, including sea ice disappearance, more extreme weather, changes in ocean chemistry, loss of biodiversity, and harm to human health.
Table 3.6
IPCC Conclusions, 2007
Very Likely ■ ■ ■
■
Human-caused emissions are the main factor causing warming since 1950. Higher maximum temperatures are observed over nearly all land areas. Snow cover decreased about 10% since the 1960s (satellite data); lake and river ice cover in the middle and high latitudes of the Northern Hemisphere was reduced by 2 weeks per year in the 20th century (independent ground-based observations). In most of the areas in the Northern Hemisphere, precipitation has increased.
Likely ■ ■ ■ ■
Temperatures in the Northern Hemisphere during the 20th century have been the highest of any century during the past 1000 years. Arctic sea ice thickness declined about 40% during late summer to early autumn in recent decades. An increase in rainfall, similar to that in the Northern Hemisphere, has been observed in tropical land areas falling between 108° North and 108° South. Summer droughts have increased.
Very Unlikely ■
The observed warming over the past 100 years is due to climate variability alone, providing new and even stronger evidence that changes must be made to stem the influence of human activities.
Each of the potential consequences can be considered in the context of the tragedy of the commons, which we encountered in Chapters 1 and 2.
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Figure 3.28 The extent of Arctic ice in 1979 (left) and in 2003 (right), composite images. Source: Earth Observatory, NASA.
Sea Ice Disappearance As shown in Figure 3.8, the temperatures in the Arctic are rising faster than anywhere else on Earth. One result is that sea ice is shrinking (Figure 3.28). In 2007, the ice cover at the end of the Arctic summer was 23% lower than the previous record. A new analysis that uses both computer models and data from actual conditions in the Arctic region forecasts that most of the Arctic sea ice will be gone in 30 years. Not only would significant populations of wildlife be endangered, but the accompanying decrease in albedo would lead to even more warming.
Sea-Level Rise Warmer temperatures result in an increase in sea level. This increase occurs primarily because as water warms, it expands. A smaller effect is caused by the influx of freshwater into the ocean from glacier runoff. According to a 2008 study published in the journal Nature, the increase was about 1.5 millimeters each year between 1961 and 2003. However, the increases are not seen uniformly across the globe. In addition, they are influenced by regional weather patterns. Even so, these small increases in sea levels can cause erosion in coastal areas and the stronger storm surges associated with hurricanes and cyclones.
Consider This 3.25
External Costs
The consequences described earlier and later on are examples of what are known as external costs. These costs are not reflected in the price of a commodity, such as the price of a gallon of gasoline or a ton of coal, but nevertheless take a toll on the environment. The external costs of burning fossil fuels often are shared by those who emit very little carbon dioxide, such as the people of the island nation of Maldives. Although a rise of sea level of just a few millimeters may not seem like much, the effects could be catastrophic for nations that lie close to sea level. Use the resources of the web to investigate how the people of Maldives are preparing for rising sea levels.
More Extreme Weather An increase in the average global temperature could cause more extreme weather, including storms, floods, and droughts. In the Northern Hemisphere, the summers are predicted to be drier and the winters wetter. Over the past several decades, more frequent wildfires and floods have occurred on every continent. The severity (although
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not the frequency) of cyclones and hurricanes also may be increasing. These tropical storms extract their energy from the oceans; a warmer ocean provides more energy to feed the storms.
Changes in Ocean Chemistry “Over the past 200 years, the oceans have absorbed approximately 550 billion tons of CO2 from the atmosphere, or about a third of the total amount of anthropogenic emissions over that period,” reports Richard A. Feely, a senior scientist with the National Pacific Marine Environmental Laboratory in Seattle. Scientists estimate that a million tons of CO2 is absorbed into the oceans every hour of every day! In their role as carbon sinks, the world’s oceans have mitigated some of the warming that carbon dioxide would have caused had it remained in the atmosphere. However, this absorption has come with a cost. Critical changes are already occurring in the oceans, as we will further explore in Chapter 6. For example, carbon dioxide is slightly soluble in water and dissolves to form carbonic acid. In turn, this is affecting marine organisms that rely on a constant level of acidity in the ocean to maintain the integrity of their shells and skeletons. The increase in carbon dioxide concentrations in the atmosphere (and hence the corresponding concentration in the oceans) is putting entire marine ecosystems at risk.
Consider This 3.26
Look for more about carbon dioxide and ocean acidification in Chapter 6.
Plankton and You
Plankton are microscopic plant- and animal-like creatures found in both salt and freshwater systems. Many plankton species have shells made of calcium carbonate that could be weakened by more acidic environments. Although humans do not eat plankton, many other marine organisms do. Construct a food chain to show the link between plankton and humans.
Loss of Biodiversity Climate change already is affecting plant, insect, and animal species around the world. Species as diverse as the California starfish, Alpine herbs, and checkerspot butterflies all have exhibited changes in either their ranges or their habits. Dr. Richard P. Alley, a Pennsylvania State University expert on past climate shifts, sees particular significance in the fact that animals and plants that rely on each other will not necessarily change ranges or habits at the same rate. Referring to affected species, he said, “You’ll have to change what you eat, or rely on fewer things to eat, or travel farther to eat, all of which have costs.” In extreme cases, those costs can cause the extinction of species. Currently, the rate of extinction worldwide is nearly 1000 times greater than at any time during the last 65 million years! A 2004 report in the journal Nature projects that about 20% of the plants and animals considered will face extinction by 2050, even under the most optimistic climate forecasts.
Many different species of checkerspot butterflies exist. This one is found in parts of Wisconsin.
Vulnerability of Freshwater Resources Like polar and sea ice, glaciers in many parts of the world are shrinking (Figure 3.29). Billions of people rely on glacier runoff for both drinking water and crop irrigation. The 2007 report of the IPCC predicts that a 1 °C increase in global temperature corresponds to more than half a billion people experiencing water shortages that they have not known before. The redistribution of freshwater also has implications in food production. Drought and high temperatures could reduce crop yields in the American Midwest, but the growing range might extend farther into Canada. It is also possible that some desert regions could get sufficient rain to become arable. One region’s loss may well become another locale’s gain, but it is too early to tell.
For more on the chemistry of water availability and use, see Section 5.3.
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Figure 3.29 A view of the Exit Glacier in Kenji Fjords National Park, Alaska in 2008. The sign in the foreground marks the extent of the ice flow in 1978.
Human Health We may all be losers in a warmer world. In 2000, the WHO attributed over 150,000 premature deaths worldwide to the effects of climate change. Those effects included more frequent and severe heat waves, increased droughts in already water-stressed regions, and infectious diseases in regions where they had not occurred before. Further increases in average temperatures are expected to expand the geographical range of mosquitoes, tsetse flies, and other disease-carrying insects. The result could be a significant upturn in illnesses such as malaria, yellow and dengue fevers, and sleeping sickness in new areas, including Asia, Europe, and the United States.
3.11
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What Can (or Should) We Do About Climate Change?
The debate over climate change has shifted in the last 15 years. Today’s scientific data leave little room for doubt about whether it is occurring. For example, measurements of higher surface and ocean temperatures, retreating glaciers and sea ice, and rising sea levels are unequivocal. In addition, the carbon isotopic ratio found in atmospheric CO2 (discussed in Section 3.4) leaves little doubt that human activity is responsible for much of the observed warming. However, at issue is what we can do and what we should do about the changes that are occurring.
Consider This 3.27
Carbon Footprint Calculations
Investigate three websites that calculate your carbon footprint. To save you time, the textbook’s website provides a list. a. For each site, list the name, the sponsor, and the information requested in order to calculate the carbon footprint. b. Does the information requested differ from site to site? If so, report the differences. c. List two advantages and two disadvantages of doing a carbon footprint calculation.
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Energy is essential for every human endeavor. Personally, you obtain the energy you need by eating and then metabolizing food. As a community or nation, we meet our energy needs in a variety of ways, including by burning coal, petroleum, and natural gas. The combustion of these carbon-based fuels produces several waste products, including carbon dioxide. The countries with large populations and those that are highly industrialized tend to burn the largest quantities of fuels and as a result emit the most CO2. According to the Carbon Dioxide Information Analysis Center (CDIAC) of Oak Ridge National Laboratory, in 2006, the top CO2 emitters were China, the United States, the Russian Federation, India, and Japan. Which other nations rank high on the list? The next activity shows you how to find out.
Consider This 3.28
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For more on food metabolism as a source of energy, see Section 11.9.
Carbon Emissions by Nation
CDIAC publishes a list of the top 20 nations for CO2 emissions. a. From what you already know, predict any five of the nations (in addition to those listed in the previous paragraph) that are on this list. Then use the link provided at the textbook’s website to check how accurate your predictions were. b. How would these rankings change if they were listed per capita?
Recall the quotation that opened this chapter. In that same 2008 address, John Holdren summarized our options in dealing with climate change with three words: mitigation, adaptation, and suffering. “Basically, if we do less mitigation and adaptation, we’re going to do a lot more suffering,” he concluded. But who will be responsible for the mitigation? Who will be forced to adapt? Who will bear the brunt of the suffering? It is likely that significant disagreements will arise regarding answers to these questions. But we can agree that any practical solution must be global in nature and include a complicated mix of risk perception, societal values, politics, and economics. Climate mitigation is any action taken to permanently eliminate or reduce the long-term risk and hazards of climate change to human life, property, or the environment. The most obvious strategy for minimizing anthropogenic climate change is to reduce the amount of CO2 emitted into the atmosphere in the first place. Take a look back at Figure 3.21. It is difficult to imagine curtailing any of these “necessities” to any great extent. Therefore decreasing our energy consumption will not be easy, at least in the short term. The simplest and least expensive approach is to improve energy efficiency. Due to the inefficiencies associated with energy production, saving energy on the consumer end multiplies its effect on the production end three to five times. However, relying on the individual consumers worldwide to buy the right goods and do the right things will not be sufficient to hold CO2 emissions below dangerous levels. A developing technology aimed at slowing the rate of carbon dioxide emissions is to capture and isolate the gas after combustion. Carbon capture and storage (or CCS) involves separating CO2 from other combustion products and storing (sequestering) it in a variety of geologic locations. If the CO2 is properly immobilized, it cannot reach the atmosphere and contribute to global warming. In addition to the large technological challenges posed by CCS, high start-up costs, usually in excess of $1 billion per plant, so far are limiting this approach as a mitigation strategy. Although at least two dozen projects are in development worldwide, as of 2009, only four industrial-scale CCS projects were in operation. Three remove CO2 from natural gas reservoirs and store it in various underground geologic formations (Figure 3.30). The fourth and largest project, located in Saskatchewan, Canada, takes CO2 captured from a coal-fired power plant in North Dakota and injects it into a depleted oil field. By doing so, additional oil is forced up through the existing wells for recovery. The benefits of enhancing oil recovery combined with CO2 sequestration could become a model for other types of projects. Combined, these CCS efforts store about 5 million metric tons of carbon dioxide annually.
Chapter 4 focuses on energy from fossil fuels, Chapter 7 on nuclear energy, and Chapter 8 on some alternative energy sources such as wind and solar power.
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Dispersed CO2
Capture and separation Pond with bacteria Coal bed methane formations
CO2 Pipelines Geological formations
Depleted oil & gas reservoirs Deep aquifer
Figure 3.30 Methods for carbon dioxide sequestration.
Your Turn 3.29
Carbon Capture Limitations
Refer back to the global carbon cycle in Figure 3.20. What percent of global carbon dioxide emissions from fossil fuel burning is captured by current CCS technology?
Critics of CCS technology dispute its ultimate efficacy for slowing atmospheric CO2 buildup, citing high costs as well as the long time frame for commercial implementation. Others contend that pursuing CCS simply delays and distracts attention from developing carbon-free energy sources. Finally, there is the sheer magnitude of the problem. According to the International Energy Agency in order for CCS to make a meaningful contribution to mitigation efforts by 2050, it would require nearly 6000 installations each injecting a million metric tons of CO2 per year into the ground. A low-tech sequestration strategy is to reverse the extensive deforestation activities that are occurring predominantly in the world’s tropical rainforests. Started in 2006 by the United Nations Development Programme and the World Agroforestry Center, the “Billion Trees Campaign” seeks to slow climate change by planting trees in depleted forests. During the first 18 months of the program, over 2 billion trees were planted, mostly in Africa. The organizers have now expanded their goals to 7 billion plantings by 2010. If 7 billion trees seem like a lot, remember the scale of deforestation; in 2005, forest area equal to about 35,000 football fields was cleared every day!
Your Turn 3.30
Trees as Carbon Sinks
An average-sized tree absorbs 25 to 50 pounds of carbon dioxide each year. In the United States, the average annual per capital CO2 emission is 19 tons. a. How many trees would be required to absorb the annual CO2 emissions for an average U.S. citizen? b. What percentage of annual global emissions could be absorbed by 7 billion trees? Hint: Refer back to Figure 3.20.
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Regardless of any potential decreases in future emissions, some effects of climate change are unavoidable. As mentioned previously, many of the CO2 molecules emitted today will remain in the atmosphere for centuries. Climate adaptation refers to the ability of a system to adjust to climate change (including climate variability and extremes) to moderate potential damage, to take advantage of opportunities, or to cope with the consequences. Some adaptive methods include developing new crop varieties and shoring up or constructing new coastline defense systems for low-lying countries and islands. The further spread of infectious diseases could be minimized by enhanced public health systems. Many of these strategies are win–win situations that would benefit societies even in the absence of climate change challenges. Compared with the scientific consensus on understanding the role greenhouse gases play in the Earth’s climate, there is much less agreement among governments regarding what actions should be taken to limit greenhouse gas emissions. One outcome from the Earth Summit held in 1992 in Rio de Janeiro was the Framework Convention on Climate Change. The goal of this international treaty was “to achieve stabilization of greenhouse gas concentrations in the atmosphere at a low enough level to prevent dangerous anthropogenic interference with the climate system.” Not only was this treaty nonbinding, but also there was no agreement about what “dangerous anthropogenic interference” meant, or what level of greenhouse gas emissions would be necessary to avoid it. In 1997, the first international treaty imposing legally binding limits on greenhouse gas emissions was written by nearly 10,000 participants from 161 countries gathered in Kyoto, Japan. The result has come to be known as the Kyoto Protocol. Binding emission targets based on 1990 levels were set for 38 developed nations to reduce their emissions of six greenhouse gases. The gases regulated include carbon dioxide, methane, nitrous oxide, hydrofluorocarbons (HFCs), perfluorocarbons (PFCs), and sulfur hexafluoride. The United States was expected to reduce emissions to 7% below its 1990 levels, the European Union (EU) nations 8%, and Canada and Japan 6% by 2012.
Consider This 3.31
The British Experience
The British Labour Party in 1997, under the leadership of Tony Blair, boldly committed to cut British greenhouse gas emissions 20% by 2010. This is significantly more than the 12.5% required by the Kyoto treaty. Did Britain meet its goal? Research this question and write a short report on the British experience in reducing greenhouse gases. Have other countries been able to reduce their emissions significantly since 1997?
Although the treaty went into effect in 2005 (when ratified by the Russian Federation), the United States never opted to participate. One reason was the belief that meeting the reduction requirements set by the protocol would cause serious harm to the U.S. economy. Another reason for not ratifying the protocol was concern about the lack of emissions limitations on developing nations, mainly China and India; those countries are expected to show the most dramatic increases in carbon dioxide emissions in the coming years. The administration of President George W. Bush argued that such unequal burdens between developed and developing countries would be economically disastrous to the United States. The United States has also resisted domestic legislation to restrict CO2 emissions on similar economic grounds. Voluntary reduction programs implemented during the early 2000s proved insufficient to reduce emissions for a variety of reasons. One “problem” is that fossil fuels are too cheap. A second problem is that any mitigation measures entail significant up-front costs, and just as importantly, the cost of mitigation is not known with certainty, making it difficult for corporations to plan effectively. The world’s current energy infrastructure cost $15 trillion to develop and distribute, and reducing carbon dioxide emissions will mean replacing much of that
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Genetically engineered foods, specifically corn, are described in Section 12.6.
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infrastructure. A final problem lies in the fact that the benefits of emissions reductions will not be felt for decades because of the long residence time of CO2 molecules in the atmosphere. Now, 15 years after the Earth Summit, scientific consensus is beginning to focus on determining what levels of CO2 are considered “dangerous.” At the United Nations Climate Conference in 2007, participating scientists concluded that greenhouse gas emissions need to peak by about 2020, and then be reduced to well below half of current levels by 2050. In absolute terms, that means that annual global emissions must be decreased by about 9 billion tons. To give you a scale of the magnitude of this goal, reducing emissions by 1 billion tons requires one of the following changes. ■ ■ ■ ■
Section 6.11 describes in detail the damage caused by the oxides of nitrogen and sulfur.
Cutting energy usage in the world’s buildings by 20–25% below business-as-usual. Having all cars get 60 mpg instead of 30 mpg. Capturing and sequestering carbon dioxide at 800 coal-burning power plants. Replacing 700 large coal-burning power plants with nuclear, wind, or solar power.
Clearly, implementation of any one of those (and the projected goal is 9 billion tons) will not be accomplished on a purely voluntary basis. In the United States and elsewhere, there is a burgeoning realization that laws and regulations are needed to reduce greenhouse gas emissions. One example is a “cap-and-trade” system, such as the one that has been successful in reducing the emission of oxides of both sulfur and nitrogen in the United States. The “trade” part of the cap-and-trade system works through a system of allowances. Companies are assigned allowances that authorize the emission of a certain quantity of CO2, either during the current year or any year thereafter. At the end of a year, each company must have sufficient allowances to cover its actual emissions. If it has extra allowances, it can trade or sell them to another company that might have exceeded their emissions limit. If a company has insufficient allowances, it must purchase them. The “cap” is enforced by creating only a certain number of allowances each year. Here’s an example of how cap-and-trade works. Without emission restrictions, Plant A emits 600 tons of CO2 and Plant B emits 400 tons. To get under the imposed cap, they are required to reduce their combined emissions by 300 tons (30%). One way to accomplish this is for each to reduce their own emissions by 30%, each accruing the associated costs. It is likely, however, that one of the plants (Plant B in Figure 3.31)
Traditional Approach: 30% Mandatory Reduction Plant A
Plant B
Flexible Cap-and-Trade Approach Plant A
Plant B
Before: 600 tons
Before: 600 tons After: 420 tons Before: 400 tons
Before: 400 tons
After: 500 tons
After: 200 tons
permits
After: 280 tons
payment
180 tons reduced
120 tons reduced
Total Emissions Reduced: 300 tons Cost to Reduce: $12,000
100 tons reduced
200 tons reduced
Total Emissions Reduced: 300 tons Cost to Reduce: $9,000
Figure 3.31 The emissions cap-and-trade concept. Source: EPA, Clearing the Air, The Facts About Capping and Trading Emissions, 2002, page 3.
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would be more efficient in their emissions reductions, and lower their emissions below the proscribed 30%. In that case, Plant A can purchase some unused emissions permits from Plant B, at a cost less than that required for Plant A to comply with the 30% emissions reduction. The overall emissions reductions are then arrived at in the most financially beneficial way for both plants. The cap-and-trade system has some possible disadvantages, including a potentially volatile market for the emissions permits. Energy providers might experience wide, often unpredictable swings in their energy costs. Those swings would result in large fluctuations in consumer costs. As an alternative to cap-and-trade, some advocate a carbon tax instead of a cap-and-trade program. Instead of limiting emissions and letting the market decide how “best” to comply, a carbon tax simply increases the cost of burning fossil fuels. Placing an additional cost based on the amount of carbon contained in a certain quality of fuel is intended to make alternative energy sources more competitive in the near term. Of course, levying a tax on carbon fuels or emissions will mean higher prices for consumers as well.
Consider This 3.32
Climate Change Insurance?
Mitigation of climate change can be seen as a risk–benefit scenario. As such, uncertainty about future effects may discourage governments from taking financially costly actions. Another way of tackling climate change is to view it as a risk-management problem, analogous to the reasons we buy insurance. Having car insurance doesn’t reduce the likelihood of being involved in an accident, but it can limit the costs if an accident should occur. How might the insurance analogy fit in with climate change actions and policies?
Although the U.S. federal government has been slow to produce binding climate change legislation, individual states have taken matters into their own hands. The 10 northeastern states that make up the Regional Greenhouse Gas Initiative (RGGI) signed the first U.S. cap-and-trade program for carbon dioxide. The program began by capping emissions at current levels in 2009 and then reducing emissions 10% by 2019. The Midwestern Regional Greenhouse Gas Reduction Accord states developed a multisector cap-and-trade system to help meet a long-term target of 60–80% below current emissions levels. Western Climate Initiative states, as well as British Columbia and Manitoba (the first participating jurisdictions outside of the United States), agreed to mandatory emissions reporting, as well as regional efforts to accelerate development of renewable energy technologies. More locally, the U.S. Mayors Climate Protection Agreement included 227 cities committed to cutting emissions to meet the targets of the Kyoto Protocol. The cities represented include some of the largest in the Northeast, the Great Lakes region, and West Coast, and their mayors represent some 44 million people.
Skeptical Chemist 3.33
Drop in the Bucket?
Critics suggest that actions made by individual states or countries, even if successful, cannot possibly have a significant effect on global emissions of greenhouse gases. Proponents for immediate action, such as NASA climate scientist James Hansen, take a different approach. “China and India have the most to lose from uncontrolled climate change because they have huge populations living near sea level. Conversely then, they also have the most to gain from reduced local air pollution. They must be a part of the solution to global warming, and I believe they will be if developed nations such as the United States take the appropriate first steps.” After studying this chapter, which side do you fall on? Explain.
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Conclusion Climate change is not the first environmental challenge created by our highly technological world, but two aspects make it the most daunting. The first is the timescale. We have recognized only recently the harm that air pollution, stratospheric ozone depletion, and acid rain can bring. Yet when we realized that human activities were responsible for the creation of those crises, we did not shrink from the challenges they presented. Combining regulation, new technologies, and adaptation strategies, we quickly began to restore air quality, retard the growth of the Antarctic ozone hole, and reduce emissions that caused acid rain. Our ability to address and correct those problems relied on the relatively short lifetimes of the pollutants involved; we have no such advantage with carbon dioxide and climate change. In 2008, United Nations Resident Coordinator Khalid Malik spoke to this very issue: “What we do today about climate change has consequences that will last a century or more. The part of that change that is due to greenhouse gas emissions is not reversible in the foreseeable future. The heat trapping gases we send into the atmosphere in 2008 will stay there until 2108 and beyond. We are therefore making choices today that will affect our own lives, but even more so the lives of our children and grandchildren. This makes climate change different and more difficult than other policy challenges.” The words bear striking resemblance to the definition of sustainability in Chapter 0. If we wait until our children experience the effects of climate change first-hand, it will likely be too late to avoid disastrous consequences. Yet our current political institutions are ill-suited to prepare for and respond to such long-term threats, rendering the possibility of corrective global action slim at best. The second daunting aspect is the immense scale of global greenhouse gas emissions. In 2009, over 80% of the energy produced worldwide was supplied by burning fossil fuels and was responsible for the majority of greenhouse gas emissions. The next chapter explores in detail how that energy is derived from combustion, mainly of fossil fuels, but also of biofuels. Given the projected doubling of energy demand by 2050, the need for developing large-scale, carbon-free energy sources becomes clear. Toward that end, the advantages and disadvantages associated with harnessing the energy of splitting atomic nuclei is the subject of Chapter 7. Technological challenges of converting radiation from our most abundant energy source, the Sun, directly to electricity are discussed in Chapter 8. Like it or not, we are in the midst of conducting a planet-wide experiment, one that will test our ability to sustain both our economic development and our environment. It is vital that we as individuals and as a society respond with wisdom, compassion, commitment, and courage. We may get only one chance.
Chapter Summary Having completed this chapter you should be able to: ■ Understand the different processes that take part in Earth’s energy balance (3.1) ■ Compare and contrast the Earth’s natural greenhouse effect and the enhanced greenhouse effect (3.1) ■ Understand the major role that certain atmospheric gases play in the greenhouse effect (3.1–3.2) ■ Explain the methods used to gather past evidence of greenhouse gas concentrations and global temperatures (3.2)
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Use Lewis structures to determine molecular geometry and bond angles (3.3) Relate molecular geometry to absorption of infrared radiation (3.4) List the major greenhouse gases and explain why each has the appropriate molecular geometry to be a greenhouse gas (3.4) Explain the roles that natural processes play in the carbon cycle and climate change (3.5)
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Evaluate how human activities contribute to the carbon cycle and climate change (3.5) Understand how molar mass is defined and used (3.6) Calculate the average mass of an atom using Avogadro’s number (3.6) Demonstrate the usefulness of the chemical mole (3.7) Assess the sources, relative emission quantities, and effectiveness of greenhouse gases other than CO2 (3.8) Evaluate the roles of natural and anthropogenic climate forcings (3.9) Recognize the successes and limitations of computerbased models in predicting climate change (3.9)
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Correlate some of the major consequences of climate change with their likelihood (3.10) Evaluate the advantages and disadvantages of proposed greenhouse gas regulations (3.11) Provide examples of climate mitigation and climate adaptation strategies (3.11) Analyze, interpret, evaluate, and critique news stories on climate change (3.1–3.12) Take an informed position with respect to issues surrounding climate change (3.1–3.12)
Questions Emphasizing Essentials 1. The chapter opened with a quote from John Holdren: “Global warming is a misnomer, because it implies something that is gradual, something that is uniform, something that is quite possibly benign. What we are experiencing with climate change is none of those things.” Use examples to: a. explain why climate change is not uniform. b. explain why it is not gradual, at least in comparison to how quickly social and environmental systems can adjust. c. explain why it probably will not be benign. 2. The surface temperatures of both Venus and Earth are warmer than would be expected on the basis of their respective distances from the Sun. Explain. 3. Using the analogy of a greenhouse to understand the energy radiated by Earth, of what are the “windows” of Earth’s greenhouse made? In what ways is the analogy not precisely correct? 4. Consider the photosynthetic conversion of CO2 and H2O to form glucose, C6H12O6, and O2. a. Write the balanced equation. b. Is the number of each type of atom on either side of the equation the same? c. Is the number of molecules on either side of the equation the same? Explain. 5. Describe the difference between climate and weather. 6. a. It is estimated that 29 megajoules per square meter (MJ/m2) of energy comes to the top of our atmosphere from the Sun each day, but only 17 MJ/m2 reaches the surface. What happens to the rest? b. Under steady-state conditions, how much energy would leave the top of the atmosphere? 7. Consider Figure 3.9. a. How does the present concentration of CO2 in the atmosphere compare with its concentration
8.
9.
10.
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20,000 years ago? With its concentration 120,000 years ago? b. How does the present temperature of the atmosphere compare with the 1950–1980 mean temperature? With the temperature 20,000 years ago? How does each of these values compare with the average temperature 120,000 years ago? c. Do your answers to parts a and b indicate causation, correlation, or no relation? Explain. Understanding Earth’s energy balance is essential to understanding the issue of global warming. For example, the solar energy striking the Earth’s surface averages 168 watts per square meter (W/m2), but the energy leaving Earth’s surface averages 390 W/m2. Why isn’t the Earth cooling rapidly? Explain each of these observations. a. A car parked in a sunny location may become hot enough to endanger the lives of pets or small children left in it. b. Clear winter nights tend to be colder than cloudy ones. c. A desert shows much wider daily temperature variation than a moist environment. d. People wearing dark clothing in the summertime put themselves at a greater risk of heatstroke than those wearing white clothing. Construct a methane molecule (CH4) from a molecular model kit (or use Styrofoam balls or gumdrops to represent the atoms and toothpicks to represent the bonds). Demonstrate that the hydrogen atoms would be farther from one another in a tetrahedral arrangement than if they all were in the same plane (square planar arrangement). Draw the Lewis structure and name the molecular geometry for each molecule. a. H2S b. OCl2 (oxygen is the central atom) c. N2O (nitrogen is the central atom)
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12. Draw the Lewis structure and name the molecular geometry for these molecules. a. PF3 b. HCN (carbon is the central atom) c. CF2Cl2 13. a. Draw the Lewis structure for methanol (wood alcohol), H3COH. b. Based on this structure, predict the H–C–H bond angle. Explain your reasoning. c. Based on this structure, predict the H–O–C bond angle. Explain your reasoning. 14. a. Draw the Lewis structure for ethene (ethylene), H2CCH2, a small hydrocarbon with a C5C double bond. b. Based on this structure, predict the H–C–H bond angle. Explain your reasoning. c. Sketch the molecule showing the predicted bond angles. 15. Three different modes of vibration of a water molecule are shown. Which of these modes of vibration contributes to the greenhouse effect? Explain.
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16. If a carbon dioxide molecule interacts with certain photons in the IR region, the vibrational motions of the atoms are increased. For CO2, the major wavelengths of absorption occur at 4.26 μm and 15.00 μm. a. What is the energy corresponding to each of these IR photons? b. What happens to the energy in the vibrating CO2 species? 17. Water vapor and carbon dioxide are greenhouse gases, but N2 and O2 are not. Explain. 18. Explain how each of these relates to global climate change. a. volcanic eruptions b. CFCs in the stratosphere 19. Termites possess enzymes that allow them to break down cellulose into glucose, C6H12O6, and then metabolize the glucose into CO2 and CH4. a. Write a balanced equation for the metabolism of glucose into CO2 and CH4.
26.
27.
b. What mass of CO2, in grams, could one termite produce in one year if it metabolized 1.0 mg glucose in one day? Consider Figure 3.21. a. Which sector has the highest CO2 emission from fossil-fuel combustion? b. What alternatives exist for each of the major sectors of CO2 emissions? Silver has an atomic number of 47. a. Give the number of protons, neutrons, and electrons in a neutral atom of the most common isotope, Ag-107. b. How do the numbers of protons, neutrons, and electrons in a neutral atom of Ag-109 compare with those of Ag-107? Silver only has two naturally occurring isotopes: Ag-107 and Ag-109. Why isn’t the average atomic mass of silver given on the periodic table simply 108? a. Calculate the average mass in grams of an individual atom of silver. b. Calculate the mass in grams of 10 trillion silver atoms. c. Calculate the mass in grams of 5.00 3 1045 silver atoms. Calculate the molar mass of these compounds. Each plays a role in atmospheric chemistry. a. H2O b. CCl2F2 (Freon-12) c. N2O a. Calculate the mass percent of chlorine in CCl3F (Freon-11). b. Calculate the mass percent of chlorine in CCl2F2 (Freon-12). c. What is the maximum mass of chlorine that could be released in the stratosphere by 100 g of each compound? d. How many atoms of chlorine correspond to the masses calculated in part c? The total mass of carbon in living systems is estimated to be 7.5 3 1017 g. Given that the total mass of carbon on Earth is estimated to be 7.5 3 1022 g, what is the ratio of carbon atoms in living systems to the total carbon atoms on Earth? Report your answer in percent and in ppm. Consider the information presented in Table 3.2. a. Calculate the percent increase in CO2 when comparing 1998 concentrations with preindustrial concentrations. b. Considering CO2, CH4, and N2O, which has shown the greatest percentage increase when comparing 1998 concentrations with preindustrial concentrations?
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28. Other than atmospheric concentration, what two other properties are included in the calculation of the global warming potential for a substance? 29. Total greenhouse gas emissions in the United States rose 16% from 1990 to 2005, growing at a rate of 1.3% a year since 2000. How is this possible when CO2 emissions grew by 20% in the same time period? Hint: See Table 3.2. Concentrating on Concepts 30. John Holdren, quoted at the opening of the chapter, suggests that we use the term global climatic disruption rather than global warming. After studying this chapter, do you agree with his suggestion? Explain. 31. The Arctic has been called “our canary in the coal mine for climate impacts that will affect us all.” a. What does the phrase “canary in the coal mine” mean? b. Explain why the Arctic serves as a canary in a coal mine. c. The melting of the tundra accelerates changes elsewhere. Give one reason. 32. Do you think the comment made in the cartoon is justified? Explain.
37.
38.
39.
40.
Pepper . . . and Salt 41. 42.
43. 44. “This winter has lowered my concerns about global warming . . . ” Source: From The Wall Street Journal. Permission by Cartoon Features Syndicate.
33. Given that direct measurements of Earth’s atmospheric temperature over the last several thousands of years are not available, how can scientists estimate past fluctuations in the temperature? 34. A friend tells you about a newspaper story that stated, “The greenhouse effect poses a serious threat to humanity.” What is your reaction to that statement? What would you tell your friend? 35. Over the last 20 years, about 120 billion tons of CO2 has been emitted from the burning of fossil fuels, yet the amount of CO2 has risen only by about 80 billion tons. Explain. 36. Carbon dioxide gas and water vapor both absorb IR radiation. Do they also absorb visible radiation? Offer
45.
46.
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some evidence based on your everyday experiences to help explain your answer. How would the energy required to cause IR-absorbing vibrations in CO2 change if the carbon and oxygen atoms were connected with single rather than with double bonds? Explain why water in a glass cup is quickly warmed in a microwave oven, but the glass cup warms much more slowly, if at all. Ethanol, C2H5OH, can be produced from sugars and starches in crops such as corn or sugarcane. The ethanol is used as a gasoline additive and when burned, it combines with O2 to form H2O and CO2. a. Write a balanced equation for the complete combustion of C2H5OH. b. How many moles of CO2 are produced from each mole of C2H5OH completely burned? c. How many moles of O2 are required to burn 10 mol of C2H5OH? Explain whether each of the radiative forcings described in Section 3.9 is positive or negative and rank them in terms of importance to overall climate change predictions. Why is the atmospheric lifetime of a greenhouse gas important? Compare and contrast stratospheric ozone depletion andclimate change in terms of the chemical species involved, the type of radiation involved, and the predicted environmental consequences. Explain the term radiative forcings to someone unfamiliar with climate modeling. It is estimated that Earth’s ruminants, such as cattle and sheep, produce 73 million metric tons of CH4 each year. How many metric tons of carbon are present in this mass of CH4? The 10 warmest years since 1880 all occurred between 1997 and 2009. Does this prove that the enhanced greenhouse effect (global warming) is taking place? Explain. A possible replacement for CFCs is HFC-152a, with a lifetime of 1.4 years and a GWP of 120. Another is HFC-23, with a lifetime of 260 years and a GWP of 12,000. Both of these possible replacements have a significant effect as greenhouse gases and are regulated under the Kyoto Protocol. a. Based on the given information, which appears to be the better replacement? Consider only the potential for global warming. b. What other considerations are there in choosing a replacement?
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47. This figure shows global emissions of CO2 in metric tons per person per year for the selected countries. 6.0 5.6 5.2
United States Canada Australia Republic of Korea Russia Japan Germany United Kingdom Italy China India
53.
3.5 3.3 3.0 3.0 2.8 2.3 1.3 0.3 0
1
2 3 4 5 Metric tons per capita
6
a. The values in this figure are reported as metric tons of CO2, not as metric tons of C in CO2. How are these values related? Hint: Think about the mass relationships developed in Section 3.6. b. Find the value for U.S. CO2 emissions from a source other than shown in this figure. Do the values agree? Explain. 48. Compare and contrast a cap-and-trade system with a carbon tax. 49. When Arrhenius first theorized the role of atmospheric greenhouses, he calculated that doubling the concentration of CO2 would result in an increase of 5–6 °C in the average global temperature. How far off was he from the current IPCC modeling? 50. Now that you have studied air quality (Chapter 1), stratospheric ozone depletion (Chapter 2), and global warming (Chapter 3), which do you believe poses the most serious problem for you in the short run? In the long run? Discuss your reasons with others and draft a short report on this question.
54.
55.
56.
57. Exploring Extensions 51. Former Vice President Al Gore writes in his 2006 book and film, An Inconvenient Truth: “We can no longer afford to view global warming as a political issue— rather, it is the biggest moral challenge facing our global civilization.” a. Do you believe that global warming is a moral issue? If so, why? b. Do you believe that global warming is a political issue? If so, why? 52. China’s growing economy is fueled largely by its dependence on coal, described as China’s “doubleedged sword.” Coal is both the new economy’s “black gold” and the “fragile environment’s dark cloud.” a. What are some of the consequences of dependence on high-sulfur coal?
b. Sulfur pollution from China may slow global warming, but only temporarily. Explain. c. What other country is rapidly stepping up its construction of coal-fired power plants and is expected to have a larger population than China by the year 2030? The quino checkerspot butterfly is an endangered species with a small range in northern Mexico and southern California. Evidence reported in 2003 indicates that the range of this species is even smaller than previously thought. a. Propose an explanation why this species is being pushed north, out of Mexico. b. Propose an explanation why this species is being pushed south, out of southern California. c. Propose a plan to prevent further harm to this endangered species. Data taken over time reveal an increase in CO2 in the atmosphere. The large increase in the combustion of hydrocarbons since the Industrial Revolution is often cited as a reason for the increasing levels of CO2. However, an increase in water vapor has not been observed during the same period. Remembering the general equation for the combustion of a hydrocarbon, does the difference in these two trends disprove any connection between human activities and global warming? Explain your reasoning. In the energy industry, 1 standard cubic foot (SCF) of natural gas contains 1196 mol of methane (CH4) at 15.6 °C (60 °F). Hint: See Appendix 1 for conversion factors. a. How many moles of CO2 could be produced by the complete combustion of 1 SCF of natural gas? b. How many kilograms of CO2 could be produced? An international conference on climate change was held in Copenhagen in December 2009. Write a brief summary of the outcomes of this conference. A solar oven is a low-tech, low-cost device for focusing sunlight to cook food. How might solar ovens help mitigate global warming? Which regions of the world would benefit most from using this technology?
58. In 2005, the European Union adopted a cap-and-trade policy for carbon dioxide. Write a short report on the outcomes of this policy, both in terms of the economic result and the effect it has had on European greenhouse gas emissions. 59. The world community responded differently to the atmospheric problems described in Chapters 2 and 3. The evidence of ozone depletion was met with the Montreal Protocol, a schedule for decreasing the production of ozone-depleting chemicals. The evidence of global warming was met with the Kyoto Protocol, a plan calling for targeted reduction of greenhouse gases.
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a. Suggest reasons why the world community dealt with the issue of ozone depletion before that of global warming. b. Compare the current status of the two responses. When was the latest amendment to the Montreal Protocol? How many nations have ratified it? Has the
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level of chlorine in the stratosphere dropped as a result of the Montreal Protocol? How many nations have ratified the Kyoto Protocol? What has happened since it went into effect? Have any other initiatives been proposed? Have levels of greenhouse gases dropped as a result of the Kyoto Protocol?
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Energy from Combustion
Since the beginning of recorded history, fire has been a source of heat, light, and security.
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Ever since fire was harnessed by our early ancestors, combustion has been central to society. Our modern fuels—the substances we burn—come in many different forms. We use coal in power plants to generate electricity. We use gasoline to run our cars. We use natural gas or heating oil to warm our homes. We use propane, charcoal, or wood to cook our food at a summer barbeque. We might even use wax to provide light for a romantic candlelight dinner. In each of these cases, using fuels means burning them. The process of combustion releases the energy stored in the molecules that these substances contain. However, the rate at which we are burning fuels is not sustainable. Perhaps you are somewhat skeptical of this claim. The supply of coal, petroleum, and natural gas may appear to be adequate, because new deposits are always being found and extraction technologies are continually improving. But even if the supply of fossil fuels were infinite (it is not), sustainability involves more than just availability. In Chapter 0, we mentioned the need to consider how our actions today will affect those who live tomorrow. A Lakota Sioux proverb emphasizes the same idea: “We don’t inherit this land from our ancestors, we borrow it from our children.” The effects of our current fossil fuel use will be felt for many decades to come.
Consider This 4.1
Fuels in the News!
a. Locate two recent news articles concerning a fuel of your choice. Cite the title, author, date, and source. According to what you read, what is the intended use of the fuel? b. Interpret each article in terms of the Lakota Sioux proverb. What, if anything, are we borrowing from our children?
Burning fossil fuels for energy fails to meet the criteria of sustainability in two ways. First, the fuels themselves—taking hundreds of millions of years to produce—are nonrenewable. Once gone, they cannot be replaced. Although our modern economies, which are based on fossil fuels, have been functioning for nearly two centuries, we need new long-term solutions. Second, the waste products of combustion have adverse effects on our environment. The previous chapter described how atmospheric carbon dioxide concentrations have risen dramatically since the beginning of the Industrial Revolution. Burning coal also releases soot, carbon monoxide, mercury, and the oxides of sulfur and nitrogen. These emissions have undeniable links to serious environmental concerns such as global warming, acid rain, and the deterioration of air and water quality. In this chapter, we will describe fuels and their characteristics. We begin with what happens inside a power plant. In the context of energy transformation, we introduce a law that tells us that energy is never created or destroyed; rather, it just changes forms. We also consider the efficiency (actually the inefficiency) of energy transformations, a factor in our ability to harness energy in convenient forms. But fuels differ, and so we need to describe how all fuels are not created equal; that is, how they have different heat contents and release different amounts of carbon dioxide. To do this, we take a closer look at coal and petroleum, describing their chemical composition, physical properties, the structures of the molecules they contain, and the ways we manipulate them for use. We learn how these molecules store energy and how to write the chemical reactions that describe energy release. We then move to biofuels, exploring the advantages and disadvantages of these renewable resources. This chapter closes with a discussion of the scale of global energy production and usage, revisiting the challenges of meeting our future energy needs.
Chapter 1 described the connection between combustion and poor air quality. Chapter 3 dealt with carbon dioxide as a greenhouse gas. Look for acid rain in Chapter 6.
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Chapter 4
4.1
Operating at full capacity, a large power plant can burn up to 10,000 tons of coal a day!
| Fossil Fuels and Electricity
About 70% of the electricity generated in the United States comes from burning fossil fuels—primarily coal. How do electrical power plants “produce” electricity and what really goes on inside them? Our task in this section is to take a closer look at the energy transformations in a power plant. In Section 4.3, we discuss the chemistry of coal. The first step in producing electricity from coal is to burn it. Examine the photographs in Figure 4.1. You can almost feel the heat from the burning coal! In the coal beds of the boilers, the temperature can reach 650 °C. To generate this heat, this small power plant burns a train car load of coal every few hours. As we pointed out in Chapter 1, combustion is the chemical process of burning; that is, the rapid combination of fuel with oxygen to release energy in the form of heat and light. Note that the two most common combustion products, CO2 and H2O, both contain oxygen. The second step in producing electricity is to use the heat released from combustion to boil water, usually in a closed, high-pressure system (Figure 4.2). The elevated pressure serves two purposes: it raises the boiling point of the water and it compresses the resulting water vapor. The hot high-pressure steam is then directed at a steam turbine. The third and final step generates electricity. As the steam expands and cools, it rushes past the turbine, causing it to spin. The shaft of the turbine is connected to a large coil of wire that rotates within a magnetic field. The turning of this coil generates an electric current. Meanwhile, the water vapor leaves the turbine and continues to cycle through the system. It passes through a condenser where a stream of cooling water carries away the remainder of the heat energy originally acquired from the fuel. The condensed water then reenters the boiler, ready to resume the energy transfer cycle. To help you better understand these different steps, we define two types of energy. Potential energy, as the name suggests, is stored energy or the energy of position. For example, energy can be stored in the position of a book lifted against the force of gravity. The heavier the book and the higher you lift it, the more potential energy it has. The potential energy of a reactant or product is often referred to as its “chemical
(a)
(b)
(c)
(d)
Figure 4.1 Photos from a small coal-fired electric power plant. (a) Piles of coal outside the plant. (b) A row of boilers into which the coal is fed. (c) Behind the blue door in photograph (b). (d) A close-up of coal burning on the boiler bed.
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Steam Turbine Generator Electricity Boiler Warm water
Condenser
Cooling water Condensate Burner
Body of water
Pump
Pump
Figure 4.2 Diagram of an electric power plant illustrating the conversion of energy from the combustion of fuels to electricity.
energy.” We discuss the chemical energy stored in fuel molecules in Section 4.5. In contrast, kinetic energy is the energy of motion. The heavier an object is and the faster it is moving, the more kinetic energy it possesses. Would you rather be hit by a baseball traveling at 90 mph or a Ping-Pong ball traveling at 90 mph? The baseball has considerably more kinetic energy because of its larger mass. Molecules that have high potential energy make good fuels. The process of combustion converts some of the potential energy of the fuel molecules into heat, which in turn is absorbed by the water in the boiler. As the water molecules absorb the heat, they move faster and faster in all directions; their kinetic energy increases. The temperature we observe is simply a measure of the average speed of that molecular motion. Hence, the temperature increases as the amount of kinetic energy of the molecules increases. When the water is vaporized to steam, the water molecules acquire a tremendous amount of kinetic energy. That energy is transformed into the mechanical energy of the spinning turbine that then turns the generator converting the mechanical energy into electrical energy. These energy transformation steps are summarized in Figure 4.3.
Potential energy (fuel molecules)
Burner
Kinetic energy
Turbine
Mechanical energy
Figure 4.3 Energy transformations in a fossil fuel electric power plant.
Consider This 4.2
Energy Conversion
Although power plants require several steps to transform potential energy into electrical energy, other devices do this more simply. For example, a battery converts chemical energy to electrical energy in one step. List three other devices that convert energy from one form to another. For each one, name the types of energy involved.
Generator
Electrical energy
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Section 3.2 introduced three fossil fuels: coal, petroleum, and natural gas.
The Sun’s energy is nuclear in origin; more specifically, nuclear fusion.
Plants contain nitrogen as well. Section 6.9 tells more about the nitrogen cycle.
Although coal is usable as a fuel, its combustion products are not. We stated that good fuels have high potential energy, but from what source does this energy come? A clue lies in the name “fossil fuel” itself. The creation of fossil fuels began when sunlight was captured by green plants that flourished on our primeval planet. Also, as mentioned earlier, photosynthesis is the process by which green plants capture the energy of sunlight to produce glucose and oxygen from carbon dioxide and water. In essence, the energy from sunlight is converted into the potential energy of glucose and oxygen. 6 CO2 6 H2O
chlorophyll
C6H12O6 6 O2
[4.1]
glucose
When living organisms die and decay, they release energy and reverse this process, producing CO2 and H2O. Under certain conditions, however, the carbon-containing compounds that make up the organism only partially decompose. This happened in the prehistoric past when vast quantities of plant and animal life became buried beneath layers of sediment in swamps or at the bottom of the oceans. Oxygen failed to reach the decaying material, thus retarding the decomposition process. The temperature and pressure increased as additional layers of mud and rock covered the buried remnants, causing additional chemical reactions to occur. Over time, the plants that once captured the Sun’s rays were transformed into the substances we call coal, petroleum, and natural gas. In a very real sense, these fossils are ancient solar energy (sunshine) stored in the solid, liquid, and gaseous state. Yes, today’s plants will become tomorrow’s fossil fuels. But this will not occur in a time frame useful to humans. It is staggering to realize that we will consume in a few centuries what it took nature hundreds of millions of years to produce. We discuss the details of fuels at the molecular level in Section 4.6.
Your Turn 4.3
Steamy Compost
Want to recycle and reuse plant and animal material? Start a compost pile. Under the right weather conditions, steam can be seen rising from a pile of compost. Explain this observation.
plant material O2 Potential energy
The law of conservation of matter and mass was introduced in Section 1.9.
Revisit the processes of combustion and photosynthesis. Energy is released in combustion, but is required for photosynthesis. The relationship between the two hints at a cycle, as shown in Figure 4.4. The first law of thermodynamics, also called the law of conservation of energy, states that energy is neither created nor destroyed. It implies that although the forms of energy change, the total amount of energy before and after any transformation remains the same. The solar energy that is stored as potential energy during photosynthesis is released as heat and light during combustion.
Photosynthesis
Combustion
heat light CO2 H2O
Figure 4.4 The energy relationship between photosynthesis and combustion.
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Energy from Combustion
4.2 |
Efficiency of Energy Transformation
By the first law of thermodynamics, we are assured that the total energy of the universe is conserved. If this is true, how can we ever experience an energy crisis? To be sure, no new energy is created during combustion, but none is destroyed, either. Although we may not be able to win, can we at least break even? The question is not as facetious as it might sound. In fact, we cannot break even. In burning coal, natural gas, and petroleum, we always convert at least some of the energy in the fuels into forms that we cannot easily use. You probably have seen the desktop toy known as a Newton’s cradle (Figure 4.5). This device transforms energy but is much simpler than a power plant! Here’s how it works. ■ ■
■
■
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A ball is lifted at one end. This gives potential energy to this ball. The ball is released and falls back toward its starting point. The potential energy (energy of position) is converted into kinetic energy (energy of motion). The ball hits the row of stationary balls. Kinetic energy is transferred along the row to the ball at the other end. The ball at the other end swings up. Kinetic energy is gradually converted into potential energy as the ball rises and slows.
The second ball begins to fall and the process repeats. With each successive cycle, however, each ball does not rise quite as high as the previous one. Eventually, the balls all come to rest at their original positions. Why do they stop moving? Where did their energy go? Is this a violation of the first law of thermodynamics? Fortunately, it isn’t. In each collision, some of the energy is used to make sound and some is used to generate heat. If we could measure precisely enough, we would observe the balls heating up slightly. This heat is then transferred to the surrounding atoms and molecules in the air, thus increasing their kinetic energy. In keeping with the law of conservation of energy, neither kinetic energy nor potential energy is conserved independently, but the sum of the two is. Therefore, when all the balls finally come to rest, the energy of the universe has been conserved. All the energy that you initially put into the system has been dissipated as random motion of the atoms and molecules in the surrounding air. In essence, the device is a fun way of dissipating a little bit of potential energy into heat (the kinetic energy of the atoms and molecules in the air). These same principles can be used to explain why no electric power plant, no matter how well designed, can completely convert one type of energy into another. In spite of the best engineers and the most competent green chemists, inefficiency is inevitable. It is caused by the transformation of energy into useless heat. Overall, the net efficiency is given by the ratio of the electrical energy produced to the energy supplied by the fuel. electrical energy produced Net efficiency 5 3 100 [4.2] heat from fuel Newer boiler systems and advanced turbine technologies have pushed the efficiencies of each step in Figure 4.3 to 90% or better. Efficiencies are multiplicative, so you might be surprised to learn that the net efficiencies of most fossil fuel power plants are between 35 and 50%. Why so low? The problem is that not all of the heat energy from the fuel combustion in the boilers can be converted into electricity. Consider, for example, the high-temperature steam that initially spins the turbines. As the steam transfers energy to the turbines, the kinetic energy of the steam decreases, it cools, and its pressure drops. It isn’t long before the steam does not have enough energy to spin the turbines anymore. Yet, the production of this “unused” steam still required a significant amount of energy; energy that is not converted into electricity. Power plants using very high temperature steam (600 °C) have efficiencies at the high end of the range. In fact, the efficiency goes up as the difference between the steam temperature and the temperature outside the plant increases. Of course there is a limit. Higher temperature steam means higher pressures and improved construction materials that need to be able to withstand such extreme conditions.
Figure 4.5 A Newton’s cradle.
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One potato (148 g) has about 100 food Calories (100 kilocalories).
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Chapter 4
Before we discuss a specific example, we need to say a word or two about energy units. The calorie was introduced with the metric system in the late 18th century and was defined as the amount of heat necessary to raise the temperature of one gram of water by one degree Celsius. When Calorie is capitalized, it generally means kilocalorie. The values tabulated on package labels and in cookbooks are, in fact, kilocalories. 1 kilocalorie (kcal) 5 1000 calories (cal) 5 1 Calorie (Cal) The modern system of units uses the joule (J), a unit of energy equal to 0.239 cal. One joule (1 J) is approximately equal to the energy required to raise a 1-kg book 10 cm against the force of gravity. On a more personal basis, each beat of the human heart requires about 1 J of energy. Now consider the case of electrical home heating, sometimes advertised as being clean and efficient. Assume that electricity from a coal-burning power plant (efficiency of 37%) is used to heat a house. If the house requires 3.5 3 107 kJ of energy for heat each day, a typical value for a northern city in January, how much coal would be burned? To answer this question, we need a value for the energy content of the coal. Let’s assume that the combustion of 1 gram of this particular coal releases about 29 kJ. Remember that only 37% of the energy released by burning the coal is available to heat the house. We now can calculate the total quantity of heat that we need to generate by burning coal at the power plant. energy generated at plant 3 efficiency 5 energy required to heat house energy generated at plant 3 0.37 5 3.5 3 107 kJ 5 energy required to heat house 3.5 3 107 kJ 5 9.5 3 107 kJ 5 energy required to heat house 0.37 In these calculations, note that we expressed the % efficiency in decimal form. We now take into account that each gram of coal burned yields 29 kJ. 1 g coal 9.5 3 107 kJ 3 29 kJ 5 3.3 3 106 g coal 5 coal required to heat house This shows that 3.3 3 106 g of coal must be burned in order to furnish the needed 3.5 3 107 kJ of energy to heat the house. This calculation assumed an efficiency of 37% at the coal plant. Higher efficiencies would mean that less fuel would have to be burned to generate the same amount of energy and that less carbon dioxide and other pollutants would be emitted. The next activity explores these connections. energy generated at plant 5
Your Turn 4.4
Comparing Power Plants
Consider two coal-fired power plants that generate 5.0 3 1012 J of electricity daily. Plant A has an overall net efficiency of 38%. Plant B, a proposed replacement, would operate at higher temperatures with an overall net efficiency of 46%. The grade of coal used releases 30 kJ of heat per gram. Assume that coal is pure carbon. a. If 1000 kg of coal costs $30, what is the difference in daily fuel costs for the two plants? b. How many fewer grams of CO2 are emitted daily by Plant B, assuming complete combustion? Answer a. Coal costs for Plant A 5 $13,150/day. Coal costs for Plant B 5 $10,900/day.
Cars and trucks also convert energy from one form to another. The internal combustion engine uses the gaseous combustion products (CO2 and H2O) to push a series of pistons, thus converting the potential energy of the gasoline or diesel fuel into mechanical energy. Other mechanisms transform that mechanical energy eventually into the kinetic energy of the vehicle’s motion. Internal combustion engines are even less efficient than coal-fired power plants. Only about 15% of the energy released by the combustion
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Energy from Combustion
of the gasoline actually is used to move the vehicle. Much of the energy is dissipated as waste heat, including about 60% lost from the internal combustion engine alone.
Consider This 4.5
Transportation Inefficiency
a. List some of the energy losses that take place when driving a car. Use the web to verify and expand your list if necessary. The textbook’s website provides a direct link. b. Given the assumption that only 15% of the energy from fuel combustion is used to move the vehicle, estimate the percent used to move the passengers.
To bring this section to a close, we ask you to revisit the Newton’s cradle. You would never expect the balls at rest to start knocking into one another on their own, right? For this to occur, all the heat energy dissipated when the balls were colliding would have to be gathered back together. The inability of a Newton’s cradle to start up on its own relates to another concept—entropy. Entropy is a measure of how much energy gets dispersed in a given process. The second law of thermodynamics has many versions, the most general of which is that the entropy of the universe is constantly increasing. The Newton’s cradle provides an example of the second law of thermodynamics. When we lift one of the balls of the Newton’s cradle, we add potential energy. After the balls knock for awhile and come to rest, this potential energy has become transformed into the chaotic (and hence more random and dispersed) motion of heat energy and never the other way around. The entropy of the universe has increased. Do you find it difficult to visualize how energy can disperse? If so, here is an analogy that might help. Imagine that you were sitting in the middle of a large auditorium and someone down in the front broke a bottle of perfume. You don’t smell anything at first, because it takes time for the molecules of the perfume to diffuse to where you’re sitting. This process of diffusion is predicted by the second law of thermodynamics. When the perfume molecules disperse into a larger volume (from the smaller volume of the bottle), the energy of the molecules gets dispersed as well. As with the Newton’s cradle, the end result is an increase in the entropy of the universe. It is extremely unlikely that all of the perfume molecules would suddenly gather in one corner of the room. Rather, once dispersed they stay dispersed unless energy is expended to recollect them. In the same way, it is essentially impossible for the Newton’s cradle to begin to move on its own after the energy originally added was dissipated as heat. Though it may not be as obvious, the second law of thermodynamics also explains the inability of a power plant or an auto engine to convert energy from one type to another with 100% efficiency.
Consider This 4.6
Can Entropy Decrease?
Processes that result in a decrease in “local” entropy require an input of energy. For example, it requires energy to arrange the socks in a bureau drawer. Identify another process in which entropy appears to decrease but is actually coupled with an increase in entropy elsewhere in the universe. Hint: Consider the entropy associated with burning coal.
4.3 |
The Chemistry of Coal
About two centuries ago, the Industrial Revolution began the great exploitation of fossil fuels that continues today. In the early 1800s, wood was the major energy source in the United States. Coal turned out to be an even better energy source than wood, because it yielded more heat per gram. Coal continued to provide more than 50% of the nation’s energy until about 1940.
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Sections 8.4 and 8.6 discuss more efficient hybrid and fuel cell vehicles.
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By the 1960s, most coal was used for generating electricity, and today the electrical power sector accounts for 92% of all U.S. coal consumption. Figure 4.6 displays the history of U.S. energy consumption.
Consider This 4.7
Evolution of Fuels
a. In the period from 1950 to 2005, which sources of energy have shown significant growth and which have not? Propose reasons for the observed trends. Hint: See Figure 4.6. b. Estimate the percent total energy consumption now supplied by coal.
In Your Turn 4.4, we assumed that coal was pure carbon. In fact, coal contains small amounts of other elements as well. Although not a single compound, coal can be approximated by the chemical formula C135H96O9NS. This formula corresponds to a carbon content of about 85% by mass. The smaller amounts of hydrogen, oxygen, nitrogen, and sulfur come from the ancient plant material and other substances present when the plants were buried. In addition, some samples of coal typically contain trace amounts of silicon, sodium, calcium, aluminum, nickel, copper, zinc, arsenic, lead, and mercury.
Your Turn 4.8
Coal Calculations
a. Assuming the composition of coal can be approximated by the formula C135H96O9NS, calculate the mass of carbon (in tons) in 1.5 million tons of coal. This quantity of coal might be burned by a typical power plant in 1 year. b. Compute the amount of energy (in kilojoules) released by burning this mass of coal. Assume the process releases 30 kJ/g of coal. Recall that 1 ton 5 2000 lb and that 1 pound 5 454 g. c. What mass of CO2 would be formed by the complete combustion of 1.5 million tons of this coal? d. How many molecules of CO2 would be formed by the complete combustion of 1.5 million tons of this coal? Answers a. Calculate the approximate molar mass of coal. The subscripts for each element give the number of moles: 135 mol C
12.0 g C 1620 g C 1 mol C
96 mol H
1.0 g H 96 g H 1 mol H
9 mol O
16.0 g O 144 g O 1 mol O
1 mol N
14.0 g N 14.0 g N 1 mol N
1 mol S
32.1 g S 32.1 g S 1 mol S
The sum of these elemental contributions for C135H96O9NS is 1906 g/mol. Therefore, every 1906 g of coal contains 1620 g C. Similarly, 1906 tons of coal contains 1620 tons of carbon. 1620 tons C Mass of carbon 5 1.5 3 106 tons C135H96O9NS 3 1906 tons C135H96NS b. 4.1 3 1013 kJ
c. 4.8 million tons
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Energy consumption, EJ
50 Petroleum Wood Nuclear electric power
Alternative energy sources, such as solar, wind, geothermal, and biofuels are barely visible in Figure 4.6 because they appeared only in the most recent decade.
Coal Natural gas Hydroelectric power
25
0 1800
1825
1850
1875
1900 Year
1925
1950
1975
2000
Figure 4.6 History of U.S. energy consumption by source, 1800–2008. Note: 1 EJ 5 1018 J. Source: Annual Energy Review, 2008. DOE/EIA-0384, June 2009.
Coal occurs in varying grades, but all grades are better fuels than wood because they contain a higher percentage of carbon and a lower percentage of oxygen. Generally speaking, the more oxygen a fuel contains, the less energy per gram it releases on combustion. In other words, oxygen-containing fuels lie lower on the potential energy scale. For example, burning 1 mole of C to produce CO2 yields about 40% more energy than is obtained from burning 1 mole of CO to produce CO2. Soft lignite, or brown coal, is the lowest grade (Figure 4.7). The plant matter from which it originated underwent the least amount of change, and its chemical composition is similar to that of wood or peat (Table 4.1). Consequently, the amount of energy released when lignite is burned is only slightly greater than that of wood. The higher grades of coal, bituminous and anthracite, have been exposed to higher pressures and temperatures for longer periods of time in the earth. During that process, they lost more oxygen and moisture and became a good deal harder, more mineral than vegetable (see Figure 4.7). These grades of coal contain a higher percentage of carbon than lignite. Anthracite has a relatively high carbon content and a low sulfur content, both of which make it the most desirable grade of coal. Unfortunately, the deposits of anthracite are relatively small, and the supply of it in the United States is almost exhausted.
Figure 4.7 Samples of anthracite (left) and lignite coal (right).
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Table 4.1
In 2009, a U.S. district court in West Virginia issued an injunction against new mountaintop removal projects in the southern part of that state. Further legislation in other parts of the country is likely.
Mercury, a contaminant in soils and in drinking water, will be discussed in Section 5.5.
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Energy Content of U.S. Coals
Type of Coal
State of Origin
anthracite
Pennsylvania
Energy Content (kJ/g) 30.5
bituminous
Maryland
30.7
sub-bituminous
Washington
24.0
lignite (brown coal)
North Dakota
16.2
peat
Mississippi
13.0
wood
various
10.4–14.1
Although coal is available across the globe and remains a widely used fuel, it has serious drawbacks, the first of which relates to underground mining, which is both dangerous and expensive. Although mine safety has dramatically improved in the United States, since 1900 more than 100,000 workers have been killed by accidents, cave-ins, fires, explosions, and poisonous gases. Many more have been incapacitated by respiratory diseases. Worldwide, the picture is far worse. A second drawback is the environmental harm caused by coal mining. Many streams and rivers in Appalachia suffer from the effects of decades of mining operations. When groundwater floods abandoned mine shafts, or comes in contact with sulfur-rich rock often associated with coal deposits, it becomes acidified. This acid mine drainage also dissolves excessive amounts of iron and aluminum, making the water uninhabitable for many fish species and placing drinking water sources at risk for many communities. When coal deposits lie sufficiently close to the surface, mining techniques safer for miners are possible, but they still have environmental costs. One technique, called mountaintop mining, is most common in West Virginia and eastern Kentucky. The process calls for scraping away the overlying vegetation and then blasting off the top several hundred feet of a mountain to reveal the underlying coal seam. Mountaintop mining creates massive quantities of rubble (“overburden”) that often is disposed of by dumping the debris into nearby river valleys. In 2005, the U.S. EPA estimated that over 700 miles of Appalachian streams were completely buried as a result of mountaintop mining between 1985 and 2001. Furthermore, increased sediments and mineral content in the surrounding water systems has adversely affected many aquatic ecosystems. A third drawback is that coal is a dirty fuel. It is, of course, physically dirty, but the issue here is the dirty combustion products. Soot from countless coal fires in cities in the 19th and early 20th century blackened both buildings and lungs. The oxides of nitrogen and sulfur are less visible but equally damaging. Although coal contains only minor amounts of mercury (50–200 ppb), mercury is concentrated in the fly ash that escapes as particulate matter into the atmosphere. In the United States, coal-fired power plants emit roughly 48 metric tons of mercury to the environment each year. The “bottom” ash left on site also presents a storage hazard. For example, Figure 4.8 shows the devastation caused by millions of gallons of fly ash sludge that spilled down a valley when the retaining walls of a storage pond failed.
Your Turn 4.9
Coal Emissions
In the United States, coal-burning power plants are responsible for two thirds of the sulfur dioxide emissions and one fifth of the nitrogen monoxide emissions. a. Why does burning coal produce SO2? Name another source of SO2 in the atmosphere. b. Why does burning coal produce nitrogen oxides? Name two other sources of NO. Hint: Revisit Chapter 1.
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Figure 4.8 In December 2008, 300 million gallons of coal sludge buried homes near Knoxville, Tennessee.
A fourth drawback may ultimately be the most serious, that burning coal produces carbon dioxide, a greenhouse gas. Coal combustion produces more CO2 per kilojoule of heat released than either petroleum or natural gas. As of 2009, coal combustion accounted for about 40% of global CO2 emissions. Because of these drawbacks, and given that coal reserves are relatively plentiful in the United States, significant research efforts are underway to develop new coal technologies. Though it may sound like an oxymoron, “clean coal” is promoted by its supporters as an important step toward decreasing our reliance on petroleum imports and reducing air pollution. The term “clean coal technology” actually encompasses a variety of methods that aim to increase the efficiency of coal-fired power plants while decreasing harmful emissions. Here we list several technologies already implemented in selected power plants. ■
■
■
“Coal washing” to remove sulfur and other mineral impurities from the coal before it is burned. “Gasification” to convert coal to a mixture of carbon monoxide and hydrogen (equation 4.10). The resulting gas burns at a lower temperature, thus reducing the generation of nitrogen oxides. “Wet scrubbing” to chemically remove SO2 before it goes up the smokestack. This is accomplished by reacting the SO2 with a mixture of ground limestone and water.
Of course, none of these address greenhouse gas emissions. This requires the most ambitious clean coal technology: carbon capture and storage. Serious questions remain about the viability of the technology involved. What does the future hold for the dirtiest of the fossil fuels? The answer may depend on where you live. Figure 4.9 compares coal consumption in different regions of the globe for 1999 and 2009. Though most regions showed modest increases, the use of coal in Asia is skyrocketing. On one hand, this makes sense, as China has enormous coal reserves. But on the other, coal burning (by any nation) clearly does not meet the criteria for sustainability. In contrast, Europe and Eurasia have actually decreased their coal consumption. The United States owns more than one quarter of the world’s coal reserves, and coal combustion accounts for more than half of all U.S. electricity generation. According to the U.S. Department of Energy, 214 new coal-fired power plants were proposed in the United States between 2000 and 2009. This upturn was met with significant public, political, and financial opposition; one third have been refused construction
Carbon dioxide capture and storage, also known as sequestration, was discussed in Section 3.11.
China has about 12% of the world’s coal reserves. Only the United States (28%) and Russia (21%) have more.
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2400
1999
2200
2009
2000 Million Metric Tons in Oil Equivalent
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1800 1600 1400 1200 1000 800 600 400 200 0
North America
S. & Cent. America
Europe & Middle East Eurasia & Africa
Asia Pacific
Figure 4.9 Global coal consumption by region for 1999 ( green) and 2009 ( gold ). The unit is million metric tons oil equivalent, the approximate energy released in burning a million metric tons of oil. Source: British Petroleum International, 2010.
licenses or are not being pursued. Another one third face on-going court battles, and many banks are refusing to finance traditional coal-fired power plants because of potential new emissions limits.
Consider This 4.10
How Clean Is Coal?
The merits of clean coal technology are widely debated. In 2004, the Sierra Club’s Dan Becker proclaimed, ”There is no such thing as ’clean coal’ and there never will be. It’s an oxymoron.” That same year, David Hawkins, director of the Natural Resources Defense Council’s Climate Center took a different stand. ”Coal is an inevitable and substantial part of the global energy mix [for the foreseeable future], so to the extent that we are going to use it, we believe coal-based generation should be . . . with carbon capture and storage.” Use the resources available to you to develop your own conclusions about clean coal. Then write an entry for an energy blog that states your position.
4.4 |
Petroleum
Although you may never have set eye on a lump of coal, undoubtedly you have seen gasoline. Around 1950, petroleum surpassed coal as the major energy source in the United States. The reasons are relatively easy to understand. Petroleum, like coal, is partially decomposed organic matter. However, it has the distinct
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advantage of being a liquid, making it easily pumped to the surface from its natural, underground reservoirs, transported via pipelines, refi ned, and fed to its point of use. Moreover, petroleum yields approximately 40–60% more energy per gram than does coal. Typical figures are 48 kJ/g for petroleum versus 30 kJ/g for a high grade of coal. Petroleum (“crude oil”) is a mixture of thousands of different compounds. The great majority are hydrocarbons, compounds that consist only of the elements hydrogen and carbon. The hydrocarbons in petroleum can contain from 1 to as many as 60 carbon atoms per molecule. Many are alkanes, hydrocarbons with only single bonds between carbon atoms (Table 4.2). Compared with coal, the concentrations of sulfur and other contaminants in petroleum generally are quite low, which minimizes the emissions of pollutants such as SO2. The oil refinery is the icon of the petroleum industry (Figure 4.10). During the initial step in the refining process, the crude oil is separated into fractions that consist of compounds with similar properties. Distillation is a separation process in which a solution is heated to its boiling point and the vapors are condensed and collected. The distillation of crude oil takes place in this manner.
Table 4.2
167
The simplest hydrocarbon, methane, contains only one carbon atom, and therefore no carbon–carbon bonds.
Petroleum is classified as “sweet crude” when it contains less than 0.5% sulfur. “Sour crude” has more than 1% sulfur.
Selected Alkanes
Name and Chemical Formula
Condensed Structural Formula
Structural Formula H H
methane CH4
H
CH4
H
H
ethane C 2H 6
H
H
C
C
H
H
propane C 3H 8
H
n-butane C4H10
H
H
n-hexane C6H14
H
H
CH3CH3
H H
H
C
C
C
H
H
H
H
H
H
H
H
C
C
C
C
H
H
CH3CH2CH3
H
H
H
H
H
C
C
C
C
C
H
H
H
H
H
H
H
H
H
H
H
H
C
C
C
C
C
C
H
H
CH3CH2CH2CH3
H
H
H
n-heptane C7H16
H
H
H
n-pentane C5H12
n-octane C8H18
C
H
H
CH3CH2CH2CH2CH3
H
CH3CH2CH2CH2CH2CH3
H
H
H
H
H
H
H
H
C
C
C
C
C
C
C
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
C
C
C
C
C
C
C
C
H
H
H
H
H
H
H
H
CH3CH2CH2CH2CH2CH2CH3
H
CH3CH2CH2CH2CH2CH2CH2CH3
Note: Butane, pentane, hexane, heptane, and octane all have other isomers. We explain names such as n -butane and n -pentane in Section 4.7.
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Figure 4.10 An oil refinery showing the tall distillation towers.
As of 2008, 143 oil refineries were operating in the United States. No new refineries have been built since 1976.
■ ■
■
■
■
The crude oil is pumped into a large vessel and heated. As the temperature increases, the compounds with the lowest boiling points vaporize. As the temperature further increases, compounds with higher boiling points begin to vaporize. Once in the gas phase, compounds travel up the tall distillation or fractionation tower. The compounds with smaller molar masses travel higher; the compounds with larger molar masses travel shorter distances up the column. Fractions are condensed back to the liquid state at different levels in the tower.
Figure 4.11 illustrates a distillation tower and lists some of the fractions obtained. These include refinery gases such as methane and propane, liquids such as gasoline, kerosene, and jet fuel, as well as waxy solids and asphalt. Note that boiling points generally increase with increasing number of carbon atoms in the molecule and hence with increasing molar mass and size. The fractions distilled from crude oil differ in the hydrocarbon molecules that constitute them. The most volatile components, “refinery gases,” have 1–4 carbon atoms per molecule. Refinery gases often are used to fuel the distillation towers. They also can be liquefied and sold for home use. Chemical manufacturers use refinery gases as a starting material for many different compounds. A major fraction is gasoline, a mixture of hydrocarbons with 5–12 carbon atoms per molecule. Although produced since the mid-1800s, gasoline became valuable in the early 20th century with the advent of the automobile and internal combustion engine. Also valuable are the larger molecules (12–16 carbon atoms) that constitute the diesel and jet fuel fractions. The highest boiling fractions, containing molecules with over 20 carbon atoms, are useful as industrial oils, lubricants, and asphalt.
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Refinery gases
C1–C4
Reformer
C5–C12
C12–C16
Jet fuel
C14–C16
Diesel
C15–C18 Cracker Crude oil
Coker >C20
Industrial fuel Lubricants Asphalt
Distillation tower
Figure 4.11 Diagram of a distillation tower for crude oil that shows the fractions and typical uses.
Figure 4.12 shows the range of products from a barrel (42 gallons) of crude oil. Most are burned for heating and transportation. The remaining 7 or so gallons are used primarily for nonfuel purposes, including the gallon or two that serve as starting materials, or “feedstocks,” to produce a myriad of plastics, pharmaceuticals, fabrics, and other carbon-based products. As these hydrocarbon feedstocks are nonrenewable resources, one easily can predict how one day petroleum products might become too valuable to burn.
Consider This 4.11
Refinery gases Gasoline Gasoline Jet fuel Diesel
C16–C20
Boiler
Gasoline
Products from a Barrel of Crude
Through research, chemists have increased the amount of gasoline that can be derived from a barrel of crude oil. For example, in 1904 a barrel of crude oil produced 4.3 gal of gasoline, 20 gal of kerosene, 5.5 gal of fuel oil, 4.9 gal of lubricants, and 7.1 gal of miscellaneous products. By 1954, the products were 18.4 gal of gasoline, 2.0 gal of kerosene, 16.6 gal of fuel oil, 0.9 gal of lubricants, and 4.1 gal of miscellaneous products. a. Compare these values with those shown in Figure 4.12. b. Offer two reasons why the distribution of products has changed over time.
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Other products 7.3 gal Liquified refinery gas (LRG) 1.7 gal Heavy fuel oil 1.7 gal Jet fuel 3.8 gal Diesel and home heating oil 10.0 gal
Gasoline 19.2 gal
Figure 4.12 Products (in gallons) from the refining of a barrel of crude oil. Note: A barrel is 42 gallons. The refinery products occupy a volume of about 44 gallons. Source: U.S. Energy Information Administration, 2009.
The United States has about 770 cars for every 1000 people. In China, there are only 10 cars per 1000 people.
Oil production from Alaskan oil fields peaked in 1988.
Other estimates of the amount of recoverable petroleum are more pessimistic, with some experts claiming that peak oil may have already occurred.
In his 2006 State of the Union address, former President George W. Bush observed that the United States is “addicted to oil.” It is, in fact, a global addiction. Due to increases in population, urbanization, and industrialization, the demand for oil has reached unprecedented levels. The largest increases are in developing countries. For example, China put nearly 10 million new cars on the road in 2008 and currently is the second largest car market in the world. The 2004 update to the United Nations World Energy Assessment Overview claims that at current rates of consumption, proven oil reserves will be depleted in 50–100 years and coal reserves in about 250 years. Nontraditional sources and new extraction technologies may extend these limits, but regardless of which scenario we choose to believe, the supply of fossil fuels is finite. The real issue here is not the quantity of fossil fuels remaining on Earth, but rather the rate at which we are able to extract them (and, of course, whether it makes sense to burn what we extract). In the mid-1950s, yearly global oil consumption was 4 billion barrels, and over 30 billion barrels of new deposits were found annually. Today, those numbers are nearly reversed. Experts therefore predict that sometime in the near future, oil production will peak and then decline as we deplete the most easily recoverable deposits. In fact, this has already happened in the United States, where oil production from the lower 48 states has slowly but steadily declined since 1970. Oil company executives voice doubts that production could ever keep pace with future demand. In 2008, Royal Dutch Shell’s CEO, Jeroen van der Veer, commented that “after 2015, supplies of easy-to-access oil and gas will no longer keep up with demand.” Assuming petroleum consumption increases 2% per year, the U.S. Energy Information Agency has developed three “peak oil” scenarios (Figure 4.13); one low, one average, and one high. The differences lie in the degree to which oil reserves worldwide are recoverable, and even the most optimistic scenario gives a maximum before 2050. In none of the scenarios do we abruptly “run out” of oil. Rather, dramatically higher prices and increasing scarcity will characterize the era when oil production peaks. Thus at some point in this century, we will no longer be able to sustain our addiction to this “black gold.” We now turn our attention to natural gas, also a fossil fuel. Although the natural gas that enters your home is practically pure methane, the “raw” gas from oil and gas wells usually includes ethane (2–6%), other small hydrocarbons, and varying quantities of water vapor, carbon dioxide, hydrogen sulfide, and helium. Before natural gas can be transported by pipeline, it must be processed to remove these impurities.
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70 2047 60 Billion barrels per year
2037 50 2026
40 30 20 10 0 1900
1925
1950
1975
2000
2025
2050
2075
2100
Year
Figure 4.13 Worldwide oil production since 1900 (black) and three peak oil scenarios. The blue, red, and green lines correspond to low, average, and high estimates of recoverable oil, respectively. Source: Energy Information Agency, U.S. Department of Energy.
Natural gas provides heat for about two thirds of the single-family homes and apartment buildings in the United States. Increasingly, though, natural gas is burned to generate electricity and to power vehicles. The reason is that natural gas burns more cleanly than other fossil fuels. Natural gas releases essentially no sulfur dioxide when burned because the sulfur-containing compounds are removed at the refinery. In addition, the levels of unburned volatile hydrocarbons, particulates, carbon monoxide, and nitrogen oxides are relatively low. No ash residue containing toxic metals remains after combustion. Although burning natural gas does produce the greenhouse gas carbon dioxide, the amount is less per unit of energy released than for the other fossil fuels. Check the numbers yourself in the next activity.
Your Turn 4.12
Coal Versus Natural Gas
The combustion of one gram of natural gas releases 50.1 kJ of heat. a. Calculate the mass of CO2 released when natural gas is burned to produce 1500 kJ of heat. Assume that natural gas is pure methane, CH4. b. Select one of the grades of coal from Table 4.1. Compare the mass of CO2 produced when enough of this coal is burned to produce the same 1500 kJ of heat. Hint: Assume coal is C135H96O9NS. You calculated its molar mass in Your Turn 4.8.
4.5 |
Measuring Energy Changes
The ability of a substance to release energy makes it a good fuel. As you have seen, both calories and joules can be used to express the energy contained in a food or fuel. In this section, you learn to quantify energy changes in chemical reactions. The next activity gives you practice with energy units.
2125
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Your Turn 4.13 Donuts contain lipids and carbohydrates, which are discussed in Sections 11.3 and 11.5, respectively.
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Energy Calculations
a. When a donut is metabolized, 425 kcal (425 Cal) are released. Express this value in kilojoules. b. Calculate the number of 1-kg books you could lift to a shelf 2 m off the floor with the amount of energy from metabolizing one donut. Answers a. Recall that 1 kcal is equivalent to 4.184 kJ. 425 kcal
4.184 kJ 1.78 103 k J 1 kcal
b. Earlier, we stated that one joule is approximately equal to the energy required to raise a 1-kg book a distance of 10 cm against Earth’s gravity. We can use this information to calculate the number of 1-kg books that could be lifted 2 meters. First note that 2 m is equivalent to 200 cm. Next, calculate the energy ( joules) required to lift one 1-kg book the entire 2 m. 200 cm 3
1J 5 20 J 10 cm
Then, express this value in kilojoules. 1 kJ 20 J 3 3 5 0.020 kJ 10 J Use this value to make the final calculation. 1.78 103 k J
1 book 8.9 104 books 0.020 k J
To work off one donut requires lifting almost 90,000 books!
Skeptical Chemist 4.14
Checking Assumptions
A simplifying (and erroneous) assumption was made in doing the calculations in part b of the preceding activity. What was the assumption and is it reasonable? Based on this assumption, is your answer too high or too low? Explain your reasoning.
The connections between oceans and climate are discussed further in Sections 3.10 and 6.5.
The metabolism of foods, including minimally healthy ones such as donuts, helps keep our bodies at a constant temperature. Temperature is a measure of the average kinetic energy of the atoms and/or molecules present in a substance. Everything around us is at some temperature—hot, cold, or lukewarm. When we perceive a particular object as “cold,” this means its atoms and molecules are moving more slowly on average relative to an object we perceive as “hot.” Therefore, for the temperature of an object to increase, the kinetic energy of its atoms and molecules must increase. Where does that energy come from? Heat is the kinetic energy that flows from a hotter object to a colder one. When two bodies are in contact, heat always flows from the object at the higher temperature to one at a lower temperature. Although the concepts of temperature and heat are related, they are not identical. Your bottle of water and the Pacific Ocean may be at the same temperature, but the ocean contains and can transfer far more heat than the bottle of water. Indeed, bodies of water can affect the climate of an entire region as a consequence of their ability to absorb and transfer heat. The calorimeter is a device used to experimentally measure the quantity of heat energy released in a combustion reaction. Figure 4.14 shows a schematic representation of a calorimeter. To use it, you introduce a known mass of fuel and an excess of oxygen into the heavy-walled stainless steel container. The container is then sealed and submerged in a bucket of water. The reaction is initiated with a spark. The heat evolved by the reaction flows from the container to the water and the rest of the apparatus. As a consequence, the temperature of the entire calorimeter system
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Electrical leads for igniting sample Stirrer
Thermometer
Oxygen inlet Water Insulated container Fuse wire in contact with sample Cup holding sample Bomb (reaction chamber)
Figure 4.14 Schematic drawing of a calorimeter.
increases. The quantity of heat given off by the reaction can be calculated from this temperature rise and the known heat-absorbing properties of the calorimeter and the water it contains. The greater the temperature increase, the greater the quantity of energy evolved from the reaction. Experimental measurements of this sort are the source of most of the tabulated values of heats of combustion. As the name suggests, the heat of combustion is the quantity of heat energy given off when a specified amount of a substance burns in oxygen. Heats of combustion are typically reported as positive values in kilojoules per mole (kJ/mol), kilojoules per gram (kJ/g), kilocalories per mole (kcal/mol), or kilocalories per gram (kcal/g). For example, the experimentally determined heat of combustion of methane is 802.3 kJ/mol. This means that 802.3 kJ of heat is given off when 1 mole of CH4(g) reacts with 2 moles of O2(g) to form 1 mole of CO2(g) and 2moles of H2O(g). CH4(g) ⫹ 2 O2(g)
CO2(g) ⫹ 2 H2O(g) ⫹ 802.3 kJ
[4.3]
We can use this value to calculate the number of kilojoules released for a gram, rather than for a mole. The molar mass of CH4, calculated from the atomic masses of carbon and hydrogen, is 16.0 g/mol. We then can calculate the heat of combustion per gram of methane. 802.3 kJ 1 mol CH4 ⫻ ⫽ 50.1 kJ/g CH4 1 mol CH4 16.0 g CH4 As fuels go, this is a high heat of combustion! Look ahead to Figure 4.16 to see how this value compares with those for other fuels. Burning methane is analogous to water tumbling down from the top of a waterfall. Initially in a state of higher potential energy, the water drops down to one of lower potential energy. The potential energy is converted into kinetic energy, which is then released when the water hits the rocks below. Similarly, when methane is burned, energy is released when the atoms in the reactants “fall” to a state of lower potential energy as the products are formed. Figure 4.15 is a schematic representation of this process. The downward arrow indicates that the energy associated with 1 mole of CO2(g) and 2 moles of H2O(g) is less than the energy associated with 1 mole of CH4(g)
Heats of combustion, by convention, are tabulated as positive values even though all combustion reactions release heat.
The mole was defined in Section 3.7.
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CH4 (g) 2 O2 (g) (1 mol) (2 mol)
Energy of reactants
Energy difference 802.3 kJ (heat released)
CO2 (g) 2 H2O (g) (1 mol) (2 mol)
Energy of products
Figure 4.15 Energy difference in the combustion of methane—an exothermic reaction.
For an exothermic reaction, E products 2 E reactants , 0.
and 2 moles of O2(g). The combustion of methane is exothermic, a term applied to any chemical or physical change accompanied by the release of heat. In this reaction, the energy difference is 2802.3 kJ. The negative sign attached to the energy change for all exothermic reactions signifies the decrease in potential energy going from reactants to products. Not surprisingly, the amount of energy released depends on the amount of fuel burned. By now, it is probably clear that good fuels have high potential energies. The higher the potential energy of a fuel, the more heat it releases when it is burned to produce CO2 and H2O. Figure 4.16 compares the energy difference (in kJ/g) of several different fuels. We can make some interesting generalizations based on the chemical formulas of the fuels. First, the fuels with the highest heats of combustion are hydrocarbons. Second, as the ratio of hydrogen-to-carbon decreases, the heat of combustion decreases. And third, as the amount of oxygen in the fuel molecule increases, the heat of combustion decreases.
Consider This 4.15
Coal Versus Ethanol
On the basis of their chemical composition, explain why ethanol and coal have very different chemical formulas but similar heats of combustion.
methane CH4 2 O2
octane C8H18 25/2 O2
coal C O2
ethanol C2H5OH 3 O2
Energy difference glucose Energy difference 50.1 kJ/g 44.4 kJ/g Energy difference Energy difference C6H12O6 6 O2 32.8 kJ/g 28.9 kJ/g Energy difference 14.2 kJ/g CO2 2 H2O
8 CO2 9 H2O
CO2
2 CO2 3 H2O
6 CO2 6 H2O
Figure 4.16 Energy differences (in kJ/g) for the combustion of methane (CH4), n -octane (C8H18), coal (assumed to be pure carbon), ethanol (C2H5OH), and wood (assumed to be glucose).
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Endothermic Versus Exothermic Reactions
Table 4.3 Endothermic Reaction
Exothermic Reaction
Energyproducts . Energyreactants
Energyproducts , Energyreactants
Net energy change is positive.
Net energy change is negative.
Energy is absorbed.
Energy is released.
Many naturally occurring reactions are not exothermic but rather absorb energy as they occur. We discussed two important examples in earlier chapters. One is the decomposition of O3 to yield O2 and O, and the other is the combination of N2 and O2 to yield two molecules of NO. Both reactions require energy in the form of an electrical discharge, a high-energy photon, or a high temperature. These reactions are endothermic, the term applied to any chemical or physical change that absorbs energy. A chemical reaction is endothermic when the potential energy of the products is higher than that of the reactants. The energy change for an endothermic reaction is always positive. Table 4.3 compares the energy changes in endothermic and exothermic reactions. Photosynthesis also is endothermic. This process requires the absorption of 2800 kJ of sunlight per mole of C6H12O6, or 15.5 kJ per gram of glucose formed. The complete process involves many steps, but the overall reaction can be described with this equation. 2800 kJ 6 CO2 (g) 6 H2O(l)
chlorophyll
C6H12O6 (s) 6 O2 (g)
Section 2.6 described the decomposition of O3 with the absorption of UV light. Section 1.9 described the formation of NO at high temperature.
[4.4]
glucose
The reaction requires the participation of the green pigment chlorophyll. The chlorophyll molecule absorbs energy from the photons of visible sunlight and uses this energy to drive the photosynthetic process, an energetically uphill reaction. Photosynthesis plays an essential role in the carbon cycle, as it removes CO2 from the atmosphere. The potential energy of any specific chemical species, and therefore the amount of energy released on combustion, is related to the chemical bonds in the molecules of the fuel. In the following section, we illustrate how knowledge of molecular structure can be used to calculate heats of combustion and can allow us to pinpoint the differences between fuels.
4.6 |
Energy Changes at the Molecular Level
Chemical reactions involve the breaking and forming of chemical bonds. Energy is required to break bonds, just as energy is required to break chains or to tear paper. In contrast, forming chemical bonds releases energy. The overall energy change associated with a chemical reaction depends on the net effect of the bond breaking and bond forming. If the energy required to break the bonds in the reactants is greater than the energy released when the products form, the overall reaction is endothermic; energy is absorbed. If, on the other hand, the bond-making energy of the products is greater than the bond breaking in the reactants, then the net energy change is exothermic; energy is released by the reaction. For example, consider the combustion of hydrogen. Hydrogen is desirable as a fuel because, compared with other fuels, it releases a large amount of energy per gram when it burns. 2 H2 (g) O2 (g)
2 H2O(g) energy
[4.5]
To calculate the energy change associated with the combustion of hydrogen to form water vapor, let us assume that all the bonds in the reactant molecules are broken, and then the individual atoms are reassembled to form the products. In fact, the reaction does not occur this way, but we are interested in only the overall (net) change, not
Look for more about hydrogen as a fuel in Sections 8.5 and 8.6.
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Chapter 4 Covalent Bond Energies (in kJ/mol)
Table 4.4 H
C
N
O
S
F
Cl
Br
I
Single Bonds H
436
C
416
356
N
391
285
160
O
467
336
201
S
347
272
}
}
226
F
566
485
272
190
326
158
146
Cl
431
327
193
205
255
255
242
Br
366
285
}
234
213
}
217
193
I
299
213
}
201
}
}
209
180
C5N
616
C5O*
803
C;O
1073
151
Multiple Bonds C5C
598
C;C
813
C;N
866
N5N
418
O5O
498
N;N
946
*in CO2.
the details. Therefore, we proceed with our convenient plan and see how well our calculated result agrees with the experimental value. The covalent bond energies given in Table 4.4 provide the numbers needed for our computation. Bond energy is the amount of energy that must be absorbed to break a specific chemical bond. Thus, because energy must be absorbed, breaking bonds is an endothermic process, and all the bond energies in Table 4.4 are positive. The values are expressed in kilojoules per mole of bonds broken. Note that the atoms appear both across the top and down the left side of the table. The number at the intersection of any row and column is the energy (in kilojoules) needed to break a mole of the bonds between the two atoms. The amount of energy required depends on the number of bonds broken: more bonds take more energy. Each value in Table 4.4 is for one mole of bonds. For example, the energy required to break 1 mole of H–H bonds, as in the H2 molecule, is 436 kJ. Similarly, the energy required to break 1 mole of O5O double bonds, as in the O2 molecule, is 498 kJ. We need to keep track of whether energy is absorbed or released. To do this, we indicate the energy absorbed with a positive sign. This is the energy absorbed when the bond is broken. Forming a bond releases energy, and the sign is negative. For example, the bond energy for the O5O double bond is 498 kJ/mol. Accordingly, when 1 mole of O5O double bonds is broken, the energy change is 1498 kJ, and when 1 mole of O5O double bonds is formed, the energy change is 2498 kJ. Now we are finally ready to apply these concepts and conventions to the burning of hydrogen gas, H2. The next equation shows the Lewis structures of the species involved so that we can count the bonds that need to be broken and formed: 2H
H O
O
2
H
O
H
[4.6]
Remember that chemical equations can be read in terms of moles. Both equation 4.5 and 4.6 indicate “2 moles of H2 plus 1 mole of O2 yields 2 moles of H2O.” To use bond energies, we need to count the number of moles of bonds involved. Here is a summary.
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Energy from Combustion
Molecule
Bonds per Molecule
Moles in Reaction
Moles of Bonds
Bond Process
Energy per Bond
Total Energy
H–H
1
2
13252
breaking
1436 kJ
2 3 (1436) 5 1872 kJ
O5O
1
1
13151
breaking
1498 kJ
1 3 (1498) 5 1498 kJ
H–O–H
2
2
23254
forming
2467 kJ
4 3 (2467) 5 21868 kJ
From the last column, we can see that the overall energy change in breaking bonds (872 kJ 1 498 kJ 5 1370 kJ) and forming new ones (21868 kJ) results in a net energy change of 2498 kJ. This calculation is diagrammed in Figure 4.17. The energy of the reactants, 2 H2 and O2, is set at zero, an arbitrary but convenient value. The green arrows pointing upward signify energy absorbed to break the bonds in the reactant molecules and form 4 H atoms and 2 O atoms. The red arrow on the right pointing downward represents energy released as these atoms bond to form the product molecules: 2 H2O. The shorter red arrow corresponds to the net energy change of 2498 kJ signifying that the overall combustion reaction is strongly exothermic. The products are lower in energy than the reactants, so the energy change is negative. The net result is the release of energy, mostly in the form of heat. Another way to look at such exothermic reactions is as a conversion of reactants involving weaker bonds to products involving stronger ones. In general, the products are more stable (lower potential energy) and less reactive than the starting substances. The energy change we just calculated from bond energies, 2498 kJ for burning 2 mol of hydrogen, compares favorably with the experimentally determined value when all of the species are gases. This agreement justifies our rather unrealistic assumption that all the bonds in the reactant molecules are first broken and then all the bonds in
1500
Energy (kJ)
1000
500
500
Breaking 1 mol of O O bonds 498 kJ Breaking 2 mol of H H bonds 2 (436 kJ) 872 kJ 2 H2 O2 (reactants)
Net energy change 498 kJ
Forming 4 mol of H O bonds 4 (467 kJ) 1868 kJ
2 H2O (product)
1000
Figure 4.17 The energy changes during the combustion of hydrogen to form water vapor.
Figures Alive! Visit the textbook’s website to learn more about the energy changes of this reaction.
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Experimental values differ somewhat from those calculated using bond energies.
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the product molecules are formed. The energy change that accompanies a chemical reaction depends on the energy difference between the products and the reactants, not on the particular process, mechanism, or individual steps that connect the two. This is an extremely powerful idea when doing calculations related to energy changes in reactions. Not all calculations come out as easily and well as this one did. For one thing, the bond energies in Table 4.4 apply only to gases, so calculations using these values agree with experiment only if all the reactants and products are in the gaseous state. Moreover, tabulated bond energies are average values. The strength of a bond depends on the overall structure of the molecule in which it is found; in other words, on what else the atoms are bonded to. Thus, the strength of an O–H bond is slightly different in HOH, HOOH, and CH 3OH. Nevertheless, the procedure illustrated here is a useful way of estimating energy changes in a wide range of reactions. The approach also helps illustrate the relationship between bond strength and chemical energy. This analysis also helps clarify why products of combustion reactions (such as H2O or CO2) cannot be used as fuels. There are no substances into which these compounds can be converted that have stronger bonds and that are lower in energy. Bottom line: You cannot run a car on its exhaust fumes!
Your Turn 4.16
Heat of Combustion for Ethyne
Use the bond energies in Table 4.4 to calculate the heat of combustion for ethyne, C2H2, also called acetylene. Report your answer both in kilojoules per mole (kJ/mol) C2H2 and kilojoules per gram (kJ/g) C2H2. Here is the balanced chemical equation. 2H
C
C
H 5O
O
4O
C
O 2
H
O
H
Answer Energy change 5 21256 kJ/mol C2H2, or 248.3 kJ/g C2H2 Heat of combustion 5 1256 kJ/mol C2H2, or 48.3 kJ/g C2H2
Your Turn 4.17
O2 Versus O3
As noted in Chapter 2, ozone absorbs UV radiation having wavelengths less than 320 nm, and oxygen absorbs electromagnetic radiation with wavelengths less than 242 nm. Use the bond energies in Table 4.4 plus information about the resonance structures of O3 from Chapter 2 to explain why.
4.7 |
The Chemistry of Gasoline
Equipped with our understanding of the molecular nature of fuels and the energy changes associated with combustion, we now return to petroleum. The distribution of compounds obtained by distilling crude oil does not correspond to the prevailing pattern of commercial use. For example, the demand for gasoline is considerably greater than that for higher boiling fractions. Chemists employ several processes to change the natural distribution and to obtain more gasoline of higher quality. These include cracking, combining, and reforming (see Figure 4.11). Thermal cracking, a process that breaks large hydrocarbon molecules into smaller ones by heating them to a high temperature, was developed first. In this procedure, the heaviest crude oil fractions are heated between 400 and 450 °C. This heat “cracks” the heaviest tarry crude oil molecules into smaller ones useful for gasoline
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and diesel fuel. For example, at high temperature, one molecule of C16H34 can be cracked into two nearly identical molecules. C16H34
C8H18 C8H16
[4.7]
Thermal cracking also can produce different-sized molecules. C16H34
C11H22 C5H12
[4.8a]
In either case, the total number of carbon and hydrogen atoms is unchanged from reactants to products. The larger reactant molecules simply have been fragmented into smaller, more economically important molecules. We can use space-filling models to show the size difference more clearly.
C16H34
C11H22
The space-filling model of C11H22 shows a “bend” because of the geometry of the atoms at the C5C double bond.
C5H12 [4.8b]
Your Turn 4.18
More Practice with Cracking
a. Draw structural formulas for one pair of products formed when C16H34 is thermally cracked (see equation 4.8a). Hint: Draw the atoms of the product molecules in an unbranched chain and include one double bond (only in one product). b. Revisit the alkanes shown in Table 4.2. Look closely to find a pattern for the number of H atoms per C atom. Use this pattern to write the generic chemical formula. c. Write the generic chemical formula for a hydrocarbon with one C5C double bond. Answer b. CnH2n12 (where n is an integer)
The problem with thermal cracking is the energy required to produce the high temperature. Catalytic cracking is a process in which catalysts are used to crack larger hydrocarbon molecules into smaller ones at relatively low temperatures, thus reducing energy use. Chemists at all major oil companies have developed important cracking catalysts and continue to find more selective and inexpensive processes. We discuss how catalysts affect the rates of chemical reactions in Section 4.8. Sometimes chemists want to combine molecules, rather than split them apart. To produce more of the intermediate-sized molecules needed for gasoline, catalytic combination can be used. In this process, smaller molecules are joined. 4 C2H4
catalyst
C8H16
[4.9]
Another important chemical process is catalytic reforming. Here, the atoms within a molecule are rearranged, usually starting with linear molecules and producing ones with more branches. As we will see, the more highly branched molecules burn more smoothly in automobile engines. It turns out that molecules with the same molecular formula are not necessarily identical. For example, octane has the formula C8H18. Careful analysis reveals 18 different compounds with this formula. Molecules with the same molecular formula but with different chemical structures and different properties are called isomers. In n-octane (normal octane) the carbon atoms are all in a continuous chain (Figure 4.18a). In isooctane, the carbon chain has several branch points (Figure 4.18b). The chemical and physical properties of these two isomers are similar, but they are not identical. For example, the boiling point of n-octane is 125 °C, compared with 99 °C for iso- octane. Although the heats of combustion for n-octane and iso-octane are nearly identical, they burn differently in an auto engine. The more compact shape of the latter compound imparts a “smoother” burn. In a well-tuned car engine, gasoline vapor and
Using catalysts to produce very large molecules (polymers) from smaller ones (monomers) is discussed in Section 9.3.
Combustion of branched hydrocarbons releases 2–4% more energy than combustion of their straight-chain isomers.
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CH3 CH3CCH2CHCH3
CH3 CH3
CH3CH2CH2CH2CH2CH2CH2CH3
(a)
(b)
Figure 4.18 Condensed structural formula and space-filling model for (a) n -octane and (b) iso-octane.
Younger people today may never have heard the sound of an engine “knocking,” because it rarely occurs with current engine technology and gasoline blends.
air are drawn into a cylinder, compressed by a piston, and ignited by a spark. Normal combustion occurs when the spark plug ignites the fuel–air mixture and the flame front travels rapidly across the combustion chamber, consuming the fuel. Sometimes, however, compression alone is enough to ignite the fuel before the spark occurs. This premature firing is called preignition. It results in lower engine efficiency and higher fuel consumption because the piston is not in its optimal location when the burned gases expand. “Knocking,” a violent and uncontrolled reaction, occurs after the spark ignites the fuel, causing the unburned mixture to burn at supersonic speed with an abnormal rise in pressure. Knocking produces an objectionable metallic sound, loss of power, overheating, and engine damage when severe. In the 1920s, knocking was shown to depend on the chemical composition of the gasoline. The “octane rating” was developed to designate a particular gasoline’s resistance to knocking. Iso-octane performs exceptionally well in automobile engines and arbitrarily has been assigned an octane rating of 100. Like n-octane, n-heptane is a straight-chain hydrocarbon, but with one fewer –CH2 group. It also has a high tendency to knock and is assigned an octane rating of 0 (Table 4.5). When you go to the gasoline pump and fill up with 87 octane, you are buying gasoline that has the same knocking characteristics as a mixture of 87% iso-octane (octane number 100) and 13% n-heptane (octane number 0). Higher grades of gasoline also are available: 89 octane (regular plus) and 92 octane (premium). These blends contain a higher percent of compounds with higher octane ratings (Figure 4.19). Although n-octane has a poor octane rating, it is possible to catalytically re-form n-octane to iso-octane, thus greatly improving its performance. This rearrangement is accomplished by passing n-octane over a catalyst consisting of rare and expensive elements such as platinum (Pt), palladium (Pd), rhodium (Rh), and iridium (Ir). Re-forming isomers to improve their octane rating became important starting in the late 1970s because of the nationwide efforts to ban the use of tetraethyl lead (TEL) as an antiknock additive.
Figure 4.19 Gasoline is available in a variety of octane ratings.
Table 4.5
Octane Ratings of Several Compounds Compound n-octane
Octane Rating 220
n-heptane
iso-octane
100
methanol
107
ethanol
108
MTBE
116
MTBE, methyl tertiary-butyl ether.
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Consider This 4.19
Getting the Lead Out
The United States completed the ban on leaded gasoline in 1996 because of the hazards associated with lead exposure. But other sources of lead still exist. Be a detective on the web to identify:
181
In the United States, TEL was phased out during the 1970s because of the toxic effects of lead, especially to children. See Section 5.10 for more information on lead in the environment.
a. an occupational source of lead exposure. b. a hobby that is a source of lead exposure. c. a source of lead exposure that particularly affects children. The textbook’s website provides helpful links to aid your search.
Elimination of TEL as an octane enhancer necessitated finding substitutes that were inexpensive, easy to produce, and environmentally benign. Several were tried, including ethanol and MTBE (methyl tertiary-butyl ether), each with an octane rating greater than 100 (see Table 4.5). As we will see, however, MTBE did not turn out as well as expected.
H
H
H
C
C
H
H ethanol
H O
H
H
C H
CH3 O
C
CH3
CH3 MTBE
Fuels containing these additives are referred to as oxygenated gasolines, blends of petroleum-derived hydrocarbons with added oxygen-containing compounds such as MTBE, ethanol, or methanol (CH3OH). Because they contain oxygen, these gasoline blends burn more cleanly and produce less carbon monoxide than their nonoxygenated counterparts. Since 1995, about 90 cities and metropolitan areas with the highest groundlevel ozone levels have adopted the Year-Round Reformulated Gasoline Program mandated by the Clean Air Act Amendments of 1990. This program requires the use of reformulated gasolines (RFGs), oxygenated gasolines that also contain a lower percentage of certain more volatile hydrocarbons found in nonoxygenated conventional gasoline. RFGs cannot contain more than 1% benzene (C6H6) and must be at least 2% oxygen. Because of their composition, reformulated gasolines evaporate less readily than conventional gasolines and produce less carbon monoxide emissions. As pointed out earlier in Chapter 1, the volatile organic compounds (VOCs) in conventional gasoline play a role in tropospheric ozone formation, especially in hightraffic areas. When RFGs were introduced in the 1990s, MTBE was the oxygenate of choice. However, concerns over its toxicity and its ability to leach from gasoline storage tanks into the groundwater have led many states to ban MTBE and switch to ethanol. As an additive and a fuel in its own right, ethanol is described more fully in Section 4.9.
4.8 |
New Uses for an Old Fuel
World supplies of coal are predicted to last for hundreds of years, much longer than current estimates of remaining available oil reserves. Unfortunately, the fact that coal is a solid makes it inconvenient for many applications, especially as a fuel for vehicles. Therefore, research and development projects are underway aimed at converting solid coal into fuels that possess characteristics similar to petroleum products. Before large supplies of natural gas were discovered and exploited, cities were lighted with water gas, a mixture of carbon monoxide and hydrogen. Water gas is
Benzene, C6H6, is a known carcinogen. Its molecular structure is introduced in Section 9.6 and further discussed in Section 10.2.
Volatile organic compounds and their role in ozone formation were explained in Section 1.11.
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formed by blowing steam over hot coke, the impure carbon that remains after volatile components have been distilled from coal. C(s) ⫹ H2O(g)
CO(g) ⫹ H2(g)
coke
[4.10]
water gas
This same reaction is the starting point for the Fischer–Tropsch process for producing synthetic gasoline from coal. German chemists Emil Fischer (1852–1919) and Hans Tropsch (1889–1935) developed the process during the 1920s. At that time, Germany had abundant coal reserves, but little petroleum. The Fischer–Tropsch process can be described by this general equation. n CO(g) ⫹ (2n ⫹ 1) H2(g)
CnH2n⫹2(g,l) ⫹ n H2O(g)
[4.11]
The hydrocarbon products can range from small molecules like methane, CH4 (n 5 1), to the medium-sized molecules (n 5 5–8) typically found in gasoline. This chemical reaction proceeds when the carbon monoxide and hydrogen are passed over a catalyst containing iron or cobalt. To better understand the role of the catalyst, consider a typical exothermic reaction, as shown in Figure 4.20. The potential energy of the reactants (left side) is higher than the potential energy of the products (right side) because it is an exothermic reaction. Now examine the pathways that connect the reactants and products. The green line indicates the energy changes during a reaction in the absence of a catalyst. Overall, this reaction gives off energy, but the energy initially increases because some bonds break (or start to break) first. The energy necessary to initiate a chemical reaction is called its activation energy and is indicated by the green arrow. Although energy must be expended to get the reaction started, energy is given off as the process proceeds to a lower potential energy state. Generally, reactions that occur rapidly have low activation energies; slower reactions have higher activation energies. However, there is no direct relationship between the height of the activation barrier and the net energy change in the reaction. In other words, a highly exothermic reaction can have a large or a small activation energy. Increasing the temperature often results in increased reaction rates; when molecules have extra energy, a greater fraction of collisions can overcome the required activation energy. Sometimes, however, increasing the temperature isn’t a practical solution. The blue line shows how a catalyst can provide an alternative reaction pathway and thus a lower activation energy (represented by the blue arrow), without raising the temperature. In the Fischer–Tropsch process, strong C;O triple bonds must be broken for the reaction to proceed. Breaking this bond corresponds to an activation energy so large the reaction simply does not proceed. This is the point at which the metal catalyst
Uncatalyzed reaction Catalyzed reaction Reactants
Energy
Catalysts were introduced in the context of automobile catalytic converters in Section 1.11.
catalyst
Products Reaction pathway
Figure 4.20 Energy–reaction pathway for the same reaction with (blue line) and without ( green line) a catalyst. The green and blue arrows represent the activation energies. The red arrow represents the overall energy change for either pathway.
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enters the reaction. Molecules of CO can form bonds with the metal surface, and when this happens, the C;O bonds weaken. The hydrogen molecules also attach to the metal surface, completely breaking the H–H single bonds. The rest of the reaction proceeds quickly, producing the higher molecular weight hydrocarbons. Thee beauty Th beau be auty tyy ooff a ca cata catalyst taly ly yst is is th that at it it is nnot ot ccon consumed onsu sume medd and and thus thus only oonl nlyy sm smal small alll amounts amou am ount ntss off it it must must be be used. ussed ed. d Green Gree Gree Gr eenn chemists chem ch hem emiis ists vvalue ists alluee catalytic ccat atallyt ytic ic reactions rrea eact ea ctiio ct ions not ions not only onl nly ly because becaus beca be use se small smal all ll amounts amou am ount ou ntss of cat nt ccatalysts atal at alys al ysts ys ts are are emp eemployed, mplo mp loye lo yedd, but ye but als aalso lsoo be ls beca because caus ca usee th us thee re reac reaction acti ac tion ti on oft ooften ften ft en can can be be carried carr ca rrie rr iedd ie out at lower temperatures. Historically, commercialization of Fischer–Tropsch technology has been limited. South Africa, a coal-rich and oil-poor nation, is the only country that synthesizes a majority of its gasoline and diesel fuel from coal. Any spike in oil prices, coupled with a plentiful domestic coal supply, may spark increased use of the Fischer–Tropsch process in other energy-hungry countries. As of 2008, China is constructing a coal-to-liquid fuels plant in Inner Mongolia. In the United States, an Australian energy corporation announced plans to build a $7 billion coal-to-liquids plant in Big Horn County, Montana, home to the Crow Tribe. The deposits there are estimated to contain almost 9 billion tons of coal. Coal, whether solid or converted to liquid fuels, still burns to produce CO2. Recent work by the National Renewable Energy Laboratory indicates that greenhouse gas emissions over the entire fuel cycle for producing coal-based liquid fuels are nearly twice as high as their petroleum-based equivalent. Clearly, we need to search for fuels to replace coal.
4.9 |
Biofuels I—Ethanol
As we mentioned in the opening of this chapter, the current rate of fossil fuel combustion is not sustainable. In Chapter 0, we also quoted Donella Meadows, a biophysicist and the founder of the Sustainability Institute: “A sustainable society is one that is far-seeing enough, flexible enough, and wise enough not to undermine either its physical or social systems of support.” Clearly, burning fossil fuels today is undermining our support systems. What are our options? Many people today believe that moving to a more sustainable energy future may be possible with the increased use of biofuels, the generic term for renewable fuels derived from plant matter such as trees, grasses, agricultural crops, or other biological material. Compared with fossil fuels, burning biofuels should release less net CO2, because the plants from which biofuels are grown had absorbed CO2 from the atmosphere through photosynthesis. This statement assumes, however, that the fuels burned to grow the crop did not cancel out this net benefit. The most common biofuel, wood, is in insufficient supply to meet our energy demands. Cutting down trees for fuel also destroys effective absorbers of carbon dioxide. Instead of relying on direct combustion of wood or other biomass, scientists currently are eyeing liquid fuels, including ethanol made via fermentation of grains, and biodiesel made from different plant oils (Figure 4.21). Since ancient times, people have known how to produce ethanol from the fermentation of the starches and sugars found in grains. Admittedly, though, this fermentation was to brew alcoholic beverages rather than to fuel auto engines. In early human history, corn was not the grain of choice. As we will see in Chapter 12 on genetic engineering, the people of the New World bred corn from a wild strain. Those living in other continents brewed ethanol from other grains, including rice and barley. Today, however, we ferment corn in the United States to produce fuel grade ethanol. The first step involves making a “soup” of corn kernels and water. The second step uses enzymes to catalyze the breakdown of the large starch molecules into individual glucose molecules. In the third step, yeast cells take over by releasing different enzymes that catalyze the conversion of glucose to ethanol. C6H12O6 glucose
2 C2H5OH 2 CO2 ethanol
[4.12]
Donella Meadows 1941–2001
In many regions of the planet, entire ecosystems are being lost because people burn wood for cooking and heat. Section 3.5 discussed how deforestation releases carbon dioxide, a greenhouse gas.
Biofuels for automobiles are not new. Henry Ford envisioned ethanol as the fuel of choice for the Model T, and Rudolph Diesel ran the first of his engines on peanut oil.
Enzymes (biological catalysts) are introduced in Section 10.4. Further examples are given in Chapters 11 and 12.
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Figure 4.21 An advertisement for renewable fuels.
The distillation of ethanol works on the same principles as that of crude oil. Both distillations separate components by their boiling points. The more volatile ethanol boils first, allowing it to be separated.
The result is essentially “beer minus the hops,” as the yeast cells die when the alcohol content reaches about 10%. The final step is to distill the mixture to separate, concentrate, and purify the ethanol. Most of the ethanol produced in the United States is blended with gasoline to make “gasohol.” Usually the mixture contains 10% ethanol, one that can be used in standard automobile engines (Figure 4.22a). As mentioned earlier, the oxygenated blends reduce emissions that produce ozone in urban areas. Significant interest (and investment) also exists in producing cars and trucks that run on a higher percent of ethanol. For example, of the 13 million vehicles in Brazil, more than 4 million use the pure ethanol produced from fermented sugarcane. The remainder operate on a mixture of ethanol and gasoline. As of 2007, more than 6 million flexible fuel vehicles (FFVs) on the road in the United States can use E85 (85% ethanol and 15% gasoline), gasoline, or any mixture of gasoline and E85 (Figure 4.22b).
(a)
(b)
Figure 4.22 Ethanol can be blended with gasoline to make (a) Gasohol, 10% ethanol, 90% gasoline or (b) E85, 85% ethanol, 15% gasoline.
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Now let’s do the energy analysis. First, recall that ethanol releases less energy per gram when burned than does a hydrocarbon mixture such as gasoline: 29.7 kJ/g (C2H5OH) versus 47.8 kJ/g (C8H18). Ethanol releases less energy because it contains oxygen (see Figure 4.16). C2H5OH(l) 3 O2(g)
2 CO2(g) 3 H2O(l) 1367 kJ
[4.13]
Don’t confuse gas mileage with the octane rating. The octane rating relates to how smoothly the fuel burns in the engine rather than to the energy content of the fuel. The octane rating of gasohol is higher than that of gasoline; the gas mileage is typically lower with an alcohol blend.
Your Turn 4.20
Ethanol and Your Gas Mileage
Even though E85 (85% ethanol and 15% gasoline) has a significantly higher octane rating than regular or even premium gasoline, the energy content per gallon is less. This translates into fewer miles per gallon. Use the values in Figure 4.16 to estimate the percent decrease in gas mileage when using E85 instead of “regular” gasoline.
Consider This 4.21
Ethanol in Brazil
Brazil derives almost all of its automotive fuel from fermented sugarcane. Use the resources of the web to learn why ethanol is so much more cost-effective in Brazil than in other countries.
Growing our fuel is far from a magic bullet. Scientists and citizens alike are questioning the sustainability of ethanol production, both short term and long. Recall the Triple Bottom Line: Healthy economies, healthy communities, and healthy ecosystems. To give you a better sense of how sustainable ethanol production might be, we now dive into the details. We begin with the economic bottom line. With the crude oil prices of 2009, it was still more expensive to produce a gallon of ethanol than a gallon of gasoline. So why the booming corn ethanol market? The answer varies by region. For example, in 2007, the U.S. government provided more than $3 billion dollars in tax credits to ethanol producers. This is one factor that encouraged the use of this fuel. In terms of other energy costs, the good news is that the Sun freely provides the energy for plants to grow. The bad news, however, is that growing corn requires additional energy inputs. Planting, cultivating, and harvesting all require energy. The same is true for producing and applying the fertilizers, manufacturing and maintaining the necessary farm equipment, and distilling the alcohol from the fermented grains. Currently, this energy is supplied by burning fossil fuels at a significant monetary cost and with significant carbon dioxide emissions. The overall energy cost for corn ethanol is difficult to quantify. Some studies estimate that for every joule put into ethanol production, 1.2 J is recovered. Others contend that the combined energy inputs outweigh the energy content of the ethanol produced. Next, we consider the social bottom line. Many rural communities in the Midwest have benefited greatly from booming ethanol demand. Construction of a distillery not only provides jobs to local workers, but also a buyer for locally grown corn. Several communities, hit hard by the depressed viability of family farms, have experienced a resurgence thanks to demand for ethanol. There are certain drawbacks as well. Increased demand for corn leads to increased prices on many other products (especially foods); prices that everyone in that community (and others around the world) must pay.
The Triple Bottom Line was first introduced in Chapter 0.
Can genetic engineering reduce the pesticides used to grow corn? Section 12.1 explores this question.
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Consider This 4.22
Passing the Buck
Many blame the skyrocketing cost of corn and corn products in the United States on the ethanol industry. However, the industry claims that high fossil fuel prices are to blame. Identify the ways in which fossil fuels are used in the production of corn ethanol.
Lastly, we turn to the environmental bottom line. Growing corn relies on the heavy use of fertilizers, herbicides, and insecticides. The manufacturing and transport of these chemicals requires the burning of fossil fuels, which in turn releases carbon dioxide. Furthermore, these chemicals, once introduced to the field, degrade the soil and water quality. Although corn growers can and do follow responsible practices, they certainly face challenges when called on to produce more corn. Increased demand for corn has also directly affected deforestation rates. In some countries, entire sections of rainforests have been cut to provide increased acreage on which to grow corn for fuel. There are also ethical considerations, including whether valuable farmland, normally used to grow crops that feed people and animals, should be used to produce ethanol. Historically, the United States has produced a significant surplus of corn and other grains for export. That surplus is dwindling as corn is diverted to fuel production. For example, in 2007, over 25% of the U.S. corn crop was used for ethanol. Increasing nonfood demand for corn is part of the cause of the recent spike in corn prices. In turn, food prices are up worldwide, and several countries are facing possible food shortages. People around the world who count on imports of American grain are paying the price for our increasing use of ethanol. Even with these issues, it appears likely that ethanol production will increase in the coming years. A record 6.5 billion gallons of ethanol was produced in the United States in 2007, more than double the output in 2002. In 2007, then President George W. Bush signed the Energy Independence and Security Act. A major part of this legislation was the Renewable Fuels Standard, a law that requires the use of 36 billion gallons of renewable fuels by 2022, to be made up predominantly of ethanol and biodiesel. The Department of Energy estimated that only about 12 billion gallons (10% of U.S. gasoline demand) could be met by corn ethanol.
Consider This 4.23
The Politics of Ethanol
States have responded in different ways to the federal mandate on renewable fuels. Some seek to increase their use of renewable fuels beyond what the Renewable Fuels Standard requires, but others have asked for partial waivers of the requirements. Use the resources of the web to find out what, if anything, your state is proposing. Draft a letter to your state legislator informing him or her of your position on the issue.
The topic of genetic modification is explored in Section 12.6. Genetically modified bacteria may be able to efficiently convert agricultural wastes into ethanol.
To address the problems with corn ethanol, researchers are looking for ways to produce ethanol from agricultural waste products, and nonfood-based crops, and the use of marginal lands to produce these crops. Cellulosic ethanol is the ethanol produced from corn stalks, switchgrass, wood chips, and other materials that are nonedible by humans. Agricultural wastes are more chemically complex than the starch from corn kernels and are more difficult to break down into glucose. As of 2009, there were no commercial cellulosic ethanol plants in the United States, but the Department of Energy has commissioned six commercial-scale plants to test the technological viability of new conversion processes. For cellulosic ethanol, the supply of raw materials is significantly greater. Without taking additional food from anyone’s plate, we have enough raw materials to make 60 billion gallons of ethanol a year. This would satisfy about 30% of the current U.S. gasoline needs. As we pointed out at the start of this section, ethanol is not the only biofuel in town. The next section is devoted to one of the newcomers: biodiesel.
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Crops for Fuel
List some desirable characteristics of nonfood crops designed for ethanol production. Consideration of the Triple Bottom Line might help in your deliberations.
4.10
|
Biofuels II—Biodiesel, Garbage, and Biogas
The production of biodiesel has grown dramatically during the last few years. It is unique among transportation fuels in that it can be produced economically in small batches by individual consumers (Figure 4.23). Biodiesel is primarily made from vegetable oils, but animal fats work as well. Both vegetable oils and animal fats contain compounds called triglycerides. Because of their high molar masses, triglycerides are not suitable for direct use in an auto engine. In contrast, biodiesel is. Reacting a triglyceride molecule with methanol and sodium hydroxide (as a catalyst) produces three molecules of biodiesel and one molecule of glycerol (C3H8O3). O a triglyceride 3 CH3OH
NaOH
3 CH3CH2CH2CH2CH2CH2CH2CH2COCH3 C3H8O3 a biodiesel molecule
glycerol
[4.14]
Consider This 4.25
Heat of Combustion for Biodiesel
Examine the structural formula of a biodiesel molecule. Then estimate the heat of combustion per gram of biodiesel in comparison to that for gasoline and to that for ethanol. Explain your reasoning.
Notably, biodiesel releases much more energy when burned than it costs to produce. In 1998, the U.S. Department of Energy and the U.S. Department of Agriculture undertook a life cycle study of the energy balance (energy in vs. energy out) of biodiesel.
(a)
(b)
Figure 4.23 (a) Biodiesel formulated from recycled restaurant vegetable oil by a regional vendor. (b) B20 is a mixture of 80% petroleum diesel and 20% biodiesel.
Section 11.3 discusses triglycerides in the context of the fats and oils found in foods.
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The study concluded that for each unit of fossil fuel energy used in the entire biodiesel production cycle, 3.2 units of energy are gained when the fuel is burned; a positive energy balance of 320%.
Consider This 4.26
Origin of Biodiesel
The U.S. Department of Energy has an extensive website dedicated to biofuels. Use the link provided at the textbook’s website to answer these questions. a. How many biodiesel plants are located in your state? b. The feedstocks for the synthesis of biodiesel differ by region. What might account for this?
Like fossil fuels, ethanol and biodiesel contain carbon and thus produce carbon dioxide when burned. But because they are derived from plants, they are more carbon-neutral fuels. By this we mean that the carbon released in combustion is at least partially offset by the carbon absorbed as the plants grew. Some of the oxygen produced by the plants is then used to burn the fuel. The production of the fuels from the raw plants does require energy (usually from fossil fuels). Thus there is a net release of CO2 even with biofuels, though less than that of coal, petroleum, or natural gas. Onee in On inte interesting tere te rest re stin st ingg facet in face fa cett of biodiesel ce bbio iodi io dies di esel es el synthesis ssyn y th yn thes esis es is is is th thee gl gglycerol y er yc erol ol bby-product. byy-pr p od oduc uctt. G uc Glycerol Gly lyyce cero roll ro is an is an impo iimportant mpo p rt rtan tantt compound comp co mp poundd used use usedd in in many man manyy different diff di ffer ff eren entt consumer cons co nsum umer er products. pro roddu duct cts ts. How H However, owev ever er, for for every ever ev eryy 9 po er poun pounds unds un ds ooff bi biod biodiesel, odie od iese ie sell, 1 ppou se pound ound ou nd ooff gl glyc glycerol ycer yc erol er ol is is pr prod produced, oduc od uced uc ed, wh ed whic which ichh ha ic hass re resu resulted sult su lted lt ed in in a glut of glycerol on the market. In 2006, Galen Suppes and coworkers at the University of Missouri earned a Presidential Green Chemistry Challenge Award for a process to convert glycerol to propylene glycol. H H
Section 1.13 cited propylene glycol as an antifreeze used in water-based interior paints. However, since the compound is volatile, the new low VOC paints no longer contain this.
C
H C
H C
H
copper catalyst
H
H
H
H
C
C
C
OH OH OH
OH OH H
glycerol
propylene glycol
H
[4.15]
The FDA has approved propylene glycol as a food additive in alcoholic beverages, ice cream, seasonings, and flavorings. The compound also finds uses in cosmetics, pharmaceuticals, paints, and detergents. Conversion of the glycerol into value-added products lowers the cost of biodiesel production, making it more competitive with petroleum-derived diesel fuel. Here, the production of propylene glycol is from a renewable resource, whereas other methods of production require petroleum as a feedstock. Biomass will become an increasingly important source of organic compounds as the supply of crude oil begins to decline. Another energy source that is cheap, always present in abundant supply, and always being renewed is garbage. Other than in a movie, no one is likely to design a car that will run on orange peels and coffee grounds. However, approximately 90 waste-to-energy power plants in the United States generate electricity from garbage. One of these, pictured in Figure 4.24, is the Hennepin Energy Resource Recovery Facility in Minneapolis, Minnesota. Hennepin County produces about 1 million tons of solid waste each year. One truckload of garbage, or about 27,000 pounds, generates the same quantity of energy as 21 barrels of oil. In 2008, the plant converted 365,000 tons of garbage into electricity to provide power to the equivalent of about 25,000 homes. In addition, over 11,000 tons of iron-containing metals are recovered from the garbage and recycled. Elk River Resource Recovery Facility, the second in Hennepin County, converts another 235,000 tons of garbage to electricity. This “resource recovery” approach, as it is sometimes called, simultaneously addresses two major problems: our growing need for energy and our growing mountains
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Figure 4.24 Hennepin County Resource Recovery Facility, a garbage-burning power plant.
of garbage. When used as a power plant fuel, the great majority of the trash is converted to carbon dioxide and water; no supplementary fuel is needed. The unburned residue is disposed of in landfills, but it represents only about 10% of the volume of the original refuse. Although some citizens have expressed concern about gaseous emissions from garbage incinerators, the incinerator’s stack effluent is carefully monitored and must be maintained within established limits. The energy needs of over 500 million people worldwide, mostly in Europe and Japan, are supplied by power plants running on garbage. In the United States, waste-to-energy plants supply power to about 2.3 million households.
Consider This 4.27
Building a Waste-Burning Plant
Imagine you were the administrator of a city of a million residents charged with drafting a proposal to your city council outlining the pros and cons of a waste-burning plant. Use the link provided at the textbook’s website for the Hennepin Energy Recovery Center in Minnesota as a model to collect information and examples. Issues such as resource recovery, pollution, and the concerns of local residents might be helpful starting points.
Biogas generators provide another good example of using waste as an energy source. In the absence of oxygen, certain strains of bacteria are able to decompose organic matter. The bacteria produce a fuel composed of about 60% CH4, 35% CO2 (by volume), and small amounts of water, hydrogen sulfide, and carbon monoxide. The biogas can be used for cooking, heating, lighting, refrigeration, and generating electricity. Both sewage and manure can be used as the source material. The leftover waste also can be used as excellent compost. The technology lends itself very well to smallscale applications. The daily manure from one or two cows can generate enough biogas to meet most of the cooking and lighting needs of a farm family. For example, more than 25 million rural households in China were using biogas in 2008. Within a decade, this number is expected to more than double.
Section 3.8 discussed methane as a potent greenhouse gas. Methane capture from landfills equates to less methane in the atmosphere.
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4.11
| The Way Forward
One thing is certain. Future generations are going to need more energy than we do today. Figure 4.25 displays the trends in energy consumption for the past several decades together with projections for the next 15 years. The data on this figure are split into three different economies. The mature market economies include the United States, Canada, and Western Europe. The emerging economies include Central and South America, India, China, and Africa. And the transitional economies include Eastern Europe and the former Soviet Union. One piece of the energy puzzle not shown in Figure 4.25 is how energy use varies across the globe. Roughly 25% of the world’s energy supply is consumed by the 5% of the world’s population that lives in North America. Why do all three future projections slope upward? One reason is increasing population. More people will require more energy to go about their daily lives. Another reason is that energy is a primary driver of industrial and economic progress. Although the future energy use in each type of economy is projected to increase, the 4.5% rate of increase in the emerging economies—dominated by the tremendous growth in China and India—far exceeds that of either the established (1.2%) or transitional economies (1.7%). Energy consumption and gross domestic product (GDP) strongly correlate. However, the GDP does not tell the full story of the quality of life for people. A more complete picture is given by the Human Development Index (HDI), one that includes GDP but also takes into account two other factors. One is life expectancy, a general measure of the health of a population. The other relates to education and includes the adult literacy rate as well as amount of schooling. For convenience, the HDI values are scaled to fall between 0 and 1.0. Norway and Iceland both fall close to the upper bound of 1.0 and thus have very high qualities of life. Examine Figure 4.26 to see the HDI is the y-axis value. Note that nations are color-coded: ■ ■
■
very low quality of life (Haiti, Pakistan, and many African nations) average quality of life (Egypt, most of Central America, states of the former Soviet Union) very high quality of life (much of Europe, Japan, Israel, Saudi Arabia, Australia)
These categories roughly correlate with a nation’s being underdeveloped, developing, or developed. In Figure 4.26, the x-axis is energy consumption. The unit is the kilograms of oil equivalent (kgoe) per year, analogous to the million metric tons of oil equivalent that
300 250 Mature market economies Exajoules, 1018 J
GDP is a measure of the total values of all goods as services produced by a country during a specific period of time.
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200 150 100 50 0
Emerging economies
Transitional economies
1970 1975 1980 1985 1990 1995 2000 2005 2010 2015 2020 2025 Year
Figure 4.25 The history (data points) and projected future (solid lines) of energy consumption worldwide. Source: Annual Energy Review 2005, Energy Information Agency, U.S. Department of Energy.
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Norway
Hong Kong Denmark
0.9 Human Development Index
United Kingdom 0.8
Canada
Iceland
United States
China 0.7 0.6
India
0.5 Very low quality of life Average quality of life Very high quality of life
0.4 0.3 0
2000
4000 6000 8000 10000 Energy consumption (kgoe/person)
12000
14000
Figure 4.26 The relationship between Human Development Index and energy consumption. Energy consumption is reported in kilograms of oil equivalent (kgoe) per person. 1 kgoe 5 4.4 3 1010 J or the approximate energy released in burning a kilogram of oil. Source: United Nations Human Development Report 2007/2008 .
we encountered earlier with Figure 4.9. As you might expect, energy consumption correlates to some degree with the quality of life. All of the countries whose people have a low quality of life and many of those with an average quality of life have low per capita energy consumptions, many less than 1000 kgoe per person. However, the story of energy consumption does not end here. Look again at Figure 4.26 and you will see that countries with high standards of living vary greatly in the amount of energy they consume per person. For example, Norway has the highest HDI (tied with Iceland), but a person in Norway consumes considerably less energy than does one in the United States. What does this mean? Most simply, it means that some countries are much more energy-efficient than others. A related question concerns the minimum amount of energy per capita required to produce a high standard of living. Although there is no definitive value, from Figure 4.26 we can ballpark it in the range of 2000–2500 kgoe per person. Hong Kong currently falls in this range. Energy consumption beyond the value of 2500 kgoe per person results in relatively small and possibly insignificant changes in the quality of living. To those living in countries with higher per capita energy consumption, this is good news. It brings some measure of hope that many in the world can maintain a certain quality of life while decreasing energy use.
Consider This 4.28
Not So Fast!
It is one thing to say that we could simply produce more energy, but it is another to say that we could do so sustainably. In this chapter, you learned about the combustion of fossil fuels. a. Why is burning coal not sustainable? List three reasons. b. In terms of sustainability, do natural gas and petroleum stack up any better?
Do we have enough energy available so that everybody on the planet can achieve a high standard of living, say, to the 0.9 HDI level? This question may be on your mind; actually, many people are asking it. If you do the math, our world energy production falls short of the math. Furthermore, as the world population increases, so does the total amount of energy needed.
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Section 8.6 describes photovoltaic cells, a means of capturing the energy from the Sun.
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Given that global energy resources are limited, nations face a dual challenge. First, they must make choices to use their energy resources wisely; that is, they must conserve. Secondly, they must develop new energy sources, especially ones that come closer to meeting the criteria for sustainability. For example, China’s construction of the immense Three Gorges Dam and hydroelectric power station helps to meet skyrocketing electrical energy demands. Denmark leads the way in wind power. Other countries are exploring other clean, renewable sources including geothermal, tidal, and solar energy. Conservation is not as easy as it may sound. Remember from our earlier discussions in Section 4.2 that, even with efficient technologies, about two thirds of the energy produced is lost due to the inherent inefficiencies of energy transformations. Let’s consider just one example where we could conserve: transportation. In the United States, motor vehicles account for more than 70% of the oil consumption and for about one fifth of carbon dioxide emissions. One legislative response to the energy crisis in the early 1970s was the Energy Policy Conservation Act of 1975. One lasting result of that act was the establishment of Corporate Average Fuel Economy (CAFE) standards that set fuel economy limits for each manufacturer’s fleet of cars and light trucks. Though the United States is the leader in many aspects of energy, fuel economy is not one of them. The average European car gets about 40 miles per gallon (mpg), and those in Japan get about 45 mpg. In contrast, automobiles in the United States average around 25 mpg.
Skeptical Chemist 4.29
Cars and CO2
A news reporter asserts that the average car in the United States emits 6 to 9 tons of CO2 yearly. What do you think, is this a reasonable estimate? Do a calculation to evaluate the reporter’s statement.
The era of stagnant CAFE standards for U.S. cars and trucks is coming to an end. In addition to mandating significant increases in biofuel production (as noted earlier), the Energy Independence and Security Act of 2007 also raises CAFE standards to 35.5 mpg by 2020. In his first four months in office, President Obama fast-tracked the transition by requiring the new standards be in place nationwide by 2016. Critics contend that to meet the stringent standards, vehicles will cost more and will need to be significantly lighter and therefore more dangerous in accidents. Improvements in safety technology, as well as development of lighter alloys and composites, will likely mean that the vehicles of the next decade will be both economical and safe. The fact remains that the automobile is an energy-intensive means of transportation and an inefficient one at that. A mass transit system is far more economical, provided it is heavily used. In Japan, 47% of travel is by public transportation, compared with only 6% in the United States. Of course, Japan is a compact country with a high population density. Although the entirety of North America is not ideally suited to mass transit, such systems could be employed in population-dense regions. However, one also must reckon with the long love affair between Americans and their automobiles. Improving end-use efficiency is another way to save energy and energy resources. In a Scientific American article published in September 2005, Amory B. Lovins asserted, “With the help of efficiency improvements and competitive renewable energy sources, the U.S. can phase out oil use by 2050.” Driving this push for efficiency is simple economics, as individuals and corporations realize that it is much cheaper to conserve fossil fuels than it is to burn them. To be sure, the future is filled with more questions than answers. Without doubt, making sustainable development a reality will be one of the greatest challenges we humans have ever faced. Often we feel pessimistic about our ability to affect change, especially on a global scale; recycling an aluminum can here and there will not lead to a sustainable future. The fact is, our actions don’t just leave our footprints in our neighborhood anymore, they leave global impressions. A few slightly smaller shoes won’t make a difference, but a few billion most certainly will. And as always, with
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challenge comes opportunity. To end the chapter on a note of optimism, we invite you to complete the following activity.
Consider This 4.30
A Sustainable Future
In 2002, then Secretary General of the United Nations Kofi Annan called sustainability, “. . . an exceptional opportunity—economically to build markets, socially to bring people in from the margins, and politically to reduce tensions over resources that could give every man and woman a voice and a choice in deciding their own future.” Expand on the Secretary General’s remarks, giving some specifics for each area mentioned.
Conclusion Fire! To early humans, fire was a source of security to ward off animals and brought the ability to cook food and minimize the spread of certain diseases. But it did more than that. Fire allowed people to venture into colder regions of the planet. The nightly campfire also became an important social vehicle, a place to come together and share stories as a community. Today, combustion still is central to our human community. We use it daily to cook; heat or cool our dwellings; transport goods and crops; and to travel the roads, rails, and skies of our planet. As we have seen in this chapter, the process of combustion converts matter and energy into less useful forms. For example, when we burn a hydrocarbon, we dissipate the potential energy it contains in the form of heat. The products of combustion—carbon dioxide and water—are not usable as fuels. One of these products, CO2, is a greenhouse gas linked to global warming. Few chemical reactions have as far-sweeping consequences as does our combustion of fuels. Given the realities of combustion, we have no option. We must expand and diversify our energy sources. Biofuels like ethanol and biodiesel likely will be a part of a sustainable energy future, but they will not be enough. We discuss nuclear energy in Chapter 7 and solar energy in Chapter 8 as other possible ways of satisfying our everincreasing appetite for energy. The next energy crisis, if not already here, will be fundamentally different from those of the past. There is little doubt that overcoming the challenges it will bring will require fundamental changes. As individuals and as a society, we must decide what sacrifices we are willing to make in speed, comfort, and convenience for the sake of our dwindling fuel supplies and the good of the planet. Recall the definition of sustainability from Chapter 0, “meeting the needs of the present without compromising the ability of future generations to meet their own needs.” We must continue to ask if we are borrowing too much from future generations. One thing is clear: the sooner we honestly examine our options, our priorities, and our will, the better. Sustainable energy, chemistry, and society are complexly intertwined, and this chapter has been an attempt to untangle them.
Chapter Summary Having studied this chapter, you should be able to: ■ Name the fossil fuels, describe the characteristics of each, and compare them in terms of how cleanly they burn and how much energy they produce (4.1–4.7) ■ Evaluate fossil fuels as a sustainable source of energy (4.1–4.7) ■ Correlate the process of electricity generation from fossil fuels with the steps in energy transformation (4.1)
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Compare and contrast kinetic energy and potential energy, both on the macroscopic and molecular level (4.1) Apply the concept of entropy to explain the second law of thermodynamics (4.2) Describe “clean coal technologies” and comment on their viability, long-term and short (4.3) Explain how and why petroleum is refined (4.4)
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List the different fractions obtained by distilling petroleum. Compare and contrast these in terms of their chemical composition, chemical properties, boiling points, and end uses (4.4) Apply the terms endothermic and exothermic to chemical reactions based on calculations or chemical intuition (4.5) Calculate energy changes in reactions using bond energies (4.6) Assess how gasoline additives affect automobile efficiencies, tailpipe emissions, human health, and the environment (4.7) Understand activation energy and how it relates to the rates of reaction (4.8)
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Compare and contrast the production and uses of ethanol and biodiesel as fuels (4.9–4.10) Compare and contrast biofuels with gasoline in terms of chemical composition, energy released on combustion, and energy required to produce (4.9–4.10) Correlate energy use to population, environmental pollution, and economic expansion (4.10) Take an informed stand on various energy conservation measures, including to what extent they are likely to produce energy savings (4.11) Evaluate news articles on energy sustainability measures and judge their accuracy (4.11)
Questions Emphasizing Essentials 1. a. List three fossil fuels. b. What is the origin of fossil fuels? c. Are fossil fuels renewable resources? 2. Explain each energy transformation step that takes place when coal is burned in a power plant. 3. Compare the processes of combustion and photosynthesis. 4. Describe how grades of coal differ and the significance of these differences. 5. A coal-burning power plant generates electrical power at a rate of 500 megawatts (MW) or 5.00 3 108 J/s. The plant has an overall efficiency of 37.5% (0.375) for the conversion of heat to electricity. a. Calculate the electrical energy (in joules) generated in 1 year of operation and the heat energy used for that purpose. b. Assuming the power plant burns coal that releases 30 kJ/g, calculate the mass of coal (in grams and metric tons) that is burned in 1 year of operation. Hint: 1 metric ton 5 1 3 103 kg 5 1 3 106 g. 6. The complete combustion of methane is given in equation 4.3. a. By analogy, write a chemical equation for the combustion of ethane, C2H6. b. Rewrite this equation using Lewis structures. c. The heat of combustion for ethane, C2H6, is 47.8 kJ/g. How much heat is produced if 1.0 mol of ethane undergoes complete combustion? 7. List three major drawbacks of using coal as a fuel. 8. Mercury (Hg) is a contaminant of coal, ranging from 50–200 ppb. Consider the amount of coal burned by the power plant in Your Turn 4.8. Calculate tons of mercury in the coal based on the lower (50 ppb) and higher (200 ppb) concentrations. 9. An energy consumption of 650,000 kcal per person per day is equivalent to an annual consumption of 65 barrels of oil or 16 tons of coal. Calculate
the amount of energy available in kilocalories for each of these. a. one barrel of oil b. 1 gallon of oil (42 gallons per barrel) c. 1 ton of coal d. 1 pound of coal (2000 pounds per ton) 10. Use the information in the previous question to find the ratio of the quantity of energy available in 1 pound of coal to that in 1 pound of oil. Hint: One pound of oil has a volume of 0.56 quart. 11. Consider the data for these three hydrocarbons. Compound, Formula pentane, C5H12 triacontane, C30H62 propane, C3H8
Melting Point (°C)
Boiling Point (°C)
2130 66
36 450
2188
242
Predict the physical state (solid, liquid, or gas) of each at room temperature. 12. a. Write the chemical equation for the complete combustion of n-heptane, C7H16. b. The heat of combustion for n-heptane is 4817 kJ/mol. How much heat is released if 250 kg of n-heptane burns completely? 13. Figure 4.17 shows energy differences for the combustion of hydrogen, an exothermic chemical reaction. The combination of nitrogen gas and oxygen gas to form nitrogen monoxide is an example of an endothermic reaction: N2(g) O2(g)
2 NO(g)
The bond energy in NO is 607 kJ/mol. Sketch an energy diagram for this reaction, and calculate the overall energy change. 14. A single serving bag of Granny Goose Hawaiian Style Potato Chips has 70 Cal. Assuming that all of the energy
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from eating these chips goes toward keeping your heart beating, how long can these chips sustain a heartbeat of 80 beats per minute? Note: 1 kcal 5 4.184 kJ, and each human heart beat requires approximately 1 J of energy. 15. A 12-oz serving of a soft drink has an energy equivalent of 92 kcal. a. In kilojoules, what is the energy released when metabolizing this beverage? b. Assume that you use this energy to lift concrete blocks that weigh 10 kg each. How many blocks could you lift to a height of 3.0 m with the energy calculated in part a? 16. One way to produce ethanol for use as a gasoline additive is the reaction of water vapor with ethylene: CH2CH2(g) H2O(g)
17. From personal experience, state whether these processes are endothermic or exothermic. a. A charcoal briquette burns. b. Water evaporates from your skin. c. Ice melts. 18. Use the bond energies in Table 4.4 to explain why: a. chlorofluorocarbons, CFCs, are so stable. b. it takes less energy to release Cl atoms than F atoms from CFCs. 19. Use the bond energies in Table 4.4 to calculate the energy changes associated with each of these reactions. Label each reaction as endothermic or exothermic. Hint: Draw Lewis structures of the reactants and products to determine the number and kinds of bonds. 2 NH3(g) a. N2(g) 3 H2(g) b. 2 C5H12(g) 11 O2(g) 10 CO(g) 12 H2O(l) c. H2(g) Cl2(g) 2 HCl(g) 20. Use the bond energies in Table 4.4 to calculate the energy changes associated with each of these reactions. Label each reaction as endothermic or exothermic. b. H2(g) O2(g)
c. The structural formulas shown are two-dimensional. Use the bond angle information in Chapter 3 to predict the C–C–C and H–C–H bond angles in n-decane. 23. Consider this equation representing the process of cracking. C16H34
C5H12 C11H22
a. Which bonds are broken and which bonds are formed in this reaction? Use Lewis structures to help answer this question. b. Use the information from part a and Table 4.4 to calculate the energy change during this cracking reaction. 24. Here is a ball-and-stick representation for one isomer of butane (C4H10).
CH3CH2OH(l)
a. Rewrite this equation using Lewis structures. b. Use the bond energies in Table 4.4 to calculate the energy change for this reaction. Is the reaction endothermic or exothermic?
a. 2 H2(g) CO(g)
195
CH3OH(g) H2O2(g)
Br2(g) Cl2(g) c. 2 BrCl(g) 21. Use Figure 4.6 to compare the sources of U.S. energy consumption. Arrange the sources in order of decreasing percentage and comment on the relative rankings. 22. The structural formulas of straight-chain (normal) alkanes containing 1 to 8 carbon atoms are given in Table 4.2. a. Draw the structural formula for n-decane, C10H22. b. Predict the chemical formula for n-nonane (9 carbon atoms) and for n-dodecane (12 carbon atoms).
a. Draw the Lewis structure for this isomer. b. Draw Lewis structures for all other isomers. Hint: Watch for duplications! 25. A premium gasoline available at most stations has an octane rating of 92. What does that tell you about: a. the knocking characteristics of this gasoline? b. whether the fuel contains oxygenates? 26. Figure 4.16 gives the energy content of several fuels in kilojoules per gram (kJ/g). Calculate the energy content in kilojoules per mole (kJ/mol) for each. How does the chemical composition of a fuel relate to its energy content? Visit Figures Alive! at the textbook’s website for related activities. Concentrating on Concepts 27. How might you explain the difference between temperature and heat to a friend? Use some practical, everyday examples. 28. Write a response to this statement: “Because of the first law of thermodynamics, there can never be an energy crisis.” 29. A friend tells you that hydrocarbon fuels containing larger molecules liberate more heat than those with smaller ones. a. Use these data, together with appropriate calculations, to discuss the merits of this statement. Hydrocarbon
Heat of Combustion
octane, C8H18 butane, C4H10
5070 kJ/mol 2658 kJ/mol
b. Based on your answer to part a, do you expect the heat of combustion per gram of candle wax, C25H52, to be more or less than that of octane? Do you expect the molar heat of combustion of candle wax to be more or less than that of octane? Justify your predictions.
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30. Halons are synthetic chemicals similar to CFCs but include bromine. Although halons are excellent materials for fire fighting, they more effectively deplete ozone than CFCs. Here is the Lewis structure for halon-1211. Br F
C
F
Cl a. Which bond in this compound is broken most easily? How is that related to the ability of this compound to deplete ozone? b. The compound C2HClF4 is being considered as a replacement for halons in fire extinguishers. Draw its Lewis structure and identify the bond broken most easily. 31. The Fischer–Tropsch conversion of hydrogen and carbon monoxide into hydrocarbons and water was given in equation 4.11: n CO (2n 1) H2
CnH2n2 n H2O
a. Determine the heat evolved by this reaction when n 5 1. b. Without doing a calculation, do you think that more or less energy is given off per mole in the formation of larger hydrocarbons (n . 1)? Explain your reasoning. 32. During petroleum distillation, kerosene and hydrocarbons with 12–18 carbons used for diesel fuel condense at position C marked on this diagram. A B
36. Octane ratings of several substances are listed in Table 4.5. a. What evidence can you give that the octane rating is or is not a measure of the energy content of a gasoline? b. Octane ratings are measures of a fuel’s ability to minimize or prevent engine knocking. Why is the prevention of knocking important? c. Why are higher octane gasolines more expensive than lower octane gasoline? 37. Section 4.8 states that both n-octane and iso-octane have essentially the same heat of combustion. How is that possible if they have different structures? 38. At present, the United States is dependent on foreign oil. One possible consequence is periodic gasoline shortages. These shortages affect more than individual motorists. List two ways in which a gasoline shortage could affect your life. 39. It was stated in the text that emissions of some pollutants are lower using biodiesel than using petroleum diesel. Based on the methods of production for each fuel, explain the lower amounts of a. sulfur dioxide emissions. b. CO emissions. 40. These three structures have the chemical formula C8H18. The hydrogen atoms and C–H bonds have been omitted for simplicity. C C C
C C C
C
C
C
C
C C
C C C
C
C
C
C
C
C
C
D
structure 1
C
C
structure 2
structure 3
C
a. Separating hydrocarbons by distillation depends on differences in a specific physical property. Which one? b. How does the number of carbon atoms in the hydrocarbon molecules separated at A, B, and D compare with those separated at position C? Explain your prediction. c. How do the uses of the hydrocarbons separated at A, B, and D differ from those separated at position C? Explain your reasoning. 33. Explain why cracking is necessary in the refinement of crude oil. 34. Consider equation 4.15. Are glycerol and propylene glycol isomers? Explain. 35. Catalysts speed up cracking reactions in oil refining and allow them to be carried out at lower temperatures. What other examples of catalysts were given in the first three chapters of this text?
a. Redraw the structures to show the missing hydrogen atoms. Hint: Check that all structures have 18 H atoms. b. Which (if any) of these structures are identical? c. Obtain a model kit and build one of these molecules. What are the C–C–C bond angles? d. Draw the structural formulas of two additional isomers of C8H18. e. If you were to build models of these two isomers, would the C–C–C bond angles be the same as that in part c? Explain. 41. Here is a ball-and-stick model of ethanol, C2H6O. Another compound, dimethyl ether, has this same chemical formula. Draw the Lewis structure of dimethyl ether. Hint: Remember to follow the octet rule.
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42. Describe how the growth in oxygenated gasolines relates to each of these. a. restrictions on the use of lead in gasoline b. federal and state air quality regulations 43. Compare the energy released on combustion of 1 gallon of ethanol and 1 gallon of gasoline. Assume gasoline is pure octane (C8H18). Explain the difference. 44. Your neighbor is shopping for a new family vehicle. The salesperson identified a van of interest as a flexible fuel vehicle (FFV). a. Explain what is meant by FFV to your neighbor. b. What is E85 fuel? c. Would your neighbor and his family be particularly interested in using E85 fuel depending on what region of the country they live? 45. The concept of entropy and probability is used in games like poker. Describe how the rank of hands (from a simple high card to a royal flush) is related to entropy and probability. 46. Bond energies such as those in Table 4.4 are sometimes found by “working backward” from heats of reaction. A reaction is carried out, and the heat absorbed or evolved is measured. From this value and known bond energies, other bond energies can be calculated. For example, the energy change associated with the combustion of formaldehyde (H2CO) is 2465 kJ/mol. H2CO(g) O2(g)
51.
52.
53.
54.
55.
CO2(g) H2O(g)
Use this information and the values found in Table 4.4 to calculate the energy of the C5O double bond in formaldehyde. Compare your answer with the C5O bond energy in CO2 and speculate on why there is a difference. Exploring Extensions 47. Revisit the Six Principles of Green Chemistry found on the inside of the front cover. Which of these are met by the synthesis by Suppes of propylene glycol from glycerol? Hint: See equation 4.15. 48. Another claim in the Scientific American article by Lovins referenced in Section 4.11 was that replacing an incandescent bulb (75 W) with a compact fluorescent bulb (18 W) would save about 75% in the cost of electricity. Electricity is generally priced per kilowatthour (kWh). Using the price of electricity where you live, calculate how much money you would save over the life of one compact fluorescent bulb (about 10,000 hr). 49. Section 4.7 states that RFGs burn more cleanly by producing less carbon monoxide than nonoxygenated fuels. At the molecular level, what evidence supports this statement? 50. Another type of catalyst used in the combustion of fossil fuels is the catalytic converter that was discussed in Chapter 1. One of the reactions that these catalysts speed up is the conversion of NO(g) to N2(g) and O2(g).
56.
57.
58.
59.
197
a. Draw a diagram of the energy of this reaction similar to the one shown in Figure 4.20. b. Why is this reaction important? Hint: See Sections 1.9 and 1.11. Chemical explosions are very exothermic reactions. Describe the relative bond strengths in the reactants and products that would make for a good explosion. Because the United States has large natural gas reserves, there is significant interest in developing uses for this fuel. List two advantages and two disadvantages ofusing natural gas to fuel vehicles. You may have seen some General Motors advertisements using the slogan “Live Green by Going Yellow” for their FlexFuel vehicles that can use E85 gasoline. To what do the colors in this slogan refer? China’s large population has increased energy consumption as the standard of living increases. a. Report on China’s increasing number of automobiles over the last 10 years. b. What evidence suggests that the increase in the number of vehicles has affected air quality? What interventions, if any, does the Chinese government have underway? Quality of life and energy consumption are related as shown in Figure 4.26. What ethical considerations (if any) about their lifestyle do citizens of a country having a per capita consumption of 8000 kgoe (kilograms oil equivalent) have to the rest of the world? What are the advantages and disadvantages of replacing gasoline with renewable fuels such as ethanol? Indicate your personal position on the issue and state your reasoning. According to the EPA, driving a car is “a typical citizen’s most polluting daily activity.” a. Do you agree? Explain. b. What pollutants do cars emit? Hint: Information on automobile emissions provided by the EPA (together with the information in this text) can help you fully answer this question. c. RFGs play a role in reducing emissions. Where in the country are RFGs required? Check the current list published on the web by the EPA. d. Explain which emissions RFGs are supposed to lower. Research the Three Gorges Dam in China. Investigate some of the major issues concerning this dam. Present your findings in a format of your choice. C. P. Snow, a noted scientist and author, wrote an influential book called The Two Cultures, in which he stated: “The question, ‘Do you know the second law of thermodynamics?’ is the cultural equivalent of ‘Have you read a work by Shakespeare?’ ” How do you react to this comparison? Discuss his remark in light of your own educational experiences.
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Water for Life
“Water has never lost its mystery. After at least two and a half millennia of philosophical and scientific inquiry, the most vital of the world’s substances remains surrounded by deep uncertainties. Without too much poetic license, we can reduce these questions to a single bare essential: What exactly is water?” PhilipBall, in Life’s Matrix: A Biography of Water, University of California Press, Berkeley, CA, 2001, p. 115.
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Neeru, shouei, maima, aqua. In any language, water is the most abundant compound on the surface of the Earth. Look at any map and you will see that oceans, rivers, lakes, and ice cover more than 70% of our planet’s surface. Recognizing the importance of water, the United Nations General Assembly proclaimed 2005–2015 as a decade for international action on water. This “Water for Life” decade addresses many themes, including the scarcity of water, sanitation, food, agriculture, and water pollution. Indeed, water is for life. It plays a key role in the cycling of nutrients on our planet. The water cycle drives weather and climate and helps to shape the contours of land masses. Water is incredibly versatile, dissolving many substances and suspending others. It is the essential medium for the biochemical reactions in the cells of all living species, including humans. Your body can go weeks without food but only days without water. If the water content in your body were reduced by 2%, you would get thirsty. With a 5% water loss, you would feel fatigue and have a headache. At a 10–15% loss your muscles would become spastic and you would feel delirious. And with greater than 15% dehydration, you would die. Water has unique properties that make life on Earth possible. It is the only common substance that you can find as a solid, a liquid, and a gas. Most solids are denser than their liquid counterparts, but ice is an exception. It is less dense than liquid water, and so it floats, allowing ecosystems in lakes to survive under the ice during winter. And because water absorbs more heat per gram than many other substances, bodies of water act as heat reservoirs. Oceans and lakes moderate extreme temperature swings. Because water has properties that are unique in supporting life, when scientists search for life on other planets, they search for water. Take a sip from a tap, bottle, or can. Steam some vegetables. Wash some laundry. Flush a toilet. Water is part of your daily routine. You also depend on water in ways that may be less obvious. For example, it takes water to irrigate crops and to prepare them for consumption. Industrial processes also require water to produce our vast array of consumer goods. Individually and collectively, we get water, use it for purposes that most likely dirty it, and then dispose of it without thinking where that dirty water ends up. This has been called the “flush and forget” syndrome. Natural cycles can clean water, but these processes occur over long periods of time. As a resource, fresh water is neither unlimited nor renewable fast enough to meet our burgeoning needs. We are creating dirty water faster than nature can clean it. Consider also that water is distributed unevenly on our planet. Just as we fight wars over oil, in the future we could be fighting wars over a much more basic necessity— water. Water is a strategic resource, and its scarcity brews conflicts and raises questions of who has the right to access and use it. In this chapter, we explore many facets of water, including how we use it, the issues related to its use, and how we might arrive at local and global water solutions. As we do so, we will keep a close eye on water and its properties. The behavior of the water molecule drives many of the phenomena we observe on this, our wet planet.
Consider This 5.1
WAT E R F O R L I F E 2005 – 2015
The amount of water you need daily depends on your size, age, health, and physical activity. To stay hydrated, a rule of thumb is to drink when you feel thirsty.
A renewable natural resource can be replenished over short periods of time.
About 20% of the people in the world lack access to safe drinking water.
“Many of the wars of the 20th century were about oil, but wars of the 21st century will be over water.” Ismail Serageldin, former Vice-President for Environmentally and Socially Sustainable Development at the World Bank.
Keep a Water Log
Pick a 12-hour waking segment of your day. Log all of your activities that involve water by time and activity. Also log: a. The role the water played in your life. For example, are you consuming it? Are you using it in some process? Is it part of your outdoor experience? b. The source of the water, the quantity involved, and where it went afterward. c. The extent to which you got the water dirty.
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Consider This 5.2
Beyond Toilets
Flushing a toilet is just one part your daily water routine. Learn more about your daily indoor water use. A link is provided at the textbook’s website. a. What surprised you about your water use? b. How does this information relate to your water log from the previous activity?
5.1
For covalent substances, as the molar mass increases, the boiling point generally increases as well.
Revisit Sections 2.3 and 3.3 for more information about the water molecule.
Electronegativity values were developed by the quantum chemist and biochemist Linus Pauling (1901–1994).
Clearly, water is essential to our lives. What may not be as apparent is that water has a number of unusual properties. In fact, these properties are quite peculiar and we are very fortunate that they are. If water were a more conventional compound, life as we know it could not exist. Let us begin with its physical state. Water is a liquid at room temperature (about 25 °C, or 77 °F) and normal atmospheric pressure. This is surprising, because almost all other compounds with similar molar masses are gases under those conditions. Consider these three gases found in air: N2, O2, and CO2. Their molar masses are 28, 32, and 44 g/mol, respectively, all greater than that of water (18 g/mol). Yet none of these are liquids! Not only is water a liquid under these conditions, but also it has an anomalously high boiling point of 100 °C (212 °F). When water freezes, it exhibits another somewhat bizarre property—it expands. Most liquids contract when they solidify. These and other unusual properties derive from the molecular structure of water. First, recall the chemical formula of water, H2O. This is probably the world’s most widely known bit of chemical trivia. Next, recall that water is a covalently bonded molecule with a bent shape. Figure 5.1 shows the same representations of the water molecule that we used in Chapter 3. New to our discussion in this chapter is the fact that the electrons are not shared equally in the O–H covalent bond. Experimental evidence indicates that the O atom attracts the shared electron pair more strongly than does the H atom. In chemical language, oxygen is said to have a higher electronegativity than hydrogen. Electronegativity is a measure of the attraction of an atom for an electron in a chemical bond. The scale runs from about 0.7 to 4.0. The values have no units and are set relative to each other. The greater the electronegativity, the more an atom attracts the electrons in a chemical bond toward itself. Table 5.1 shows electronegativity values for the first 18 elements. Examine it to see that: ■ ■ ■
If the electronegativity difference between two atoms is more than 1.0, the bond is considered polar. If it is greater than 2.0, the bond is considered ionic. Use this information as a guideline rather than as a rule.
| The Unique Properties of Water
Fluorine and oxygen have the highest values. Metals such as lithium and sodium have low values. Values increase from left to right in a row of the periodic table (from metals to nonmetals) and decrease going down a group.
The greater the difference in electronegativity between two bonded atoms, the more polar the bond is. Accordingly, we can use electronegativity values to estimate bond polarities. For example, the electronegativity difference between oxygen and
H O H
H
O
H
O H
H 104.5
(a)
(b)
Figure 5.1 Representations of H2O. (a) Lewis structures and structural formula; (b) Space-filling model.
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Table 5.1 1A
201 Electronegativity value (EN) 3.5 2.1
Electronegativity Values for Selected Elements
2A
3A
4A
5A
6A
7A
H 2.1
8A
Li 1.0
Be 1.5
B 2.0
C 2.5
N 3.0
O 3.5
F 4.0
Ne *
Na 0.9
Mg 1.2
Al 1.5
Si 1.8
P 2.1
S 2.5
Cl 3.0
Ar *
O
He *
H
EN difference 1.4
Figure 5.2 Representation of the polar covalent bond between a hydrogen and oxygen atom. The electrons are pulled toward the more electronegative oxygen atom.
*Noble gases rarely (if ever) bond to other elements
hydrogen is 1.4. The electrons in an O–H bond are pulled closer to the more electronegative oxygen atom. This unequal sharing results in a partial negative charge (d2) on the O atom and a partial positive charge (d1) on the H atom, as shown in Figure 5.2. An arrow is used to indicate the direction in which the electron pair is displaced. The result is a polar covalent bond, a covalent bond in which the electrons are not equally shared but rather are closer to the more electronegative atom. A polar covalent bond is an example of an intramolecular force, a force that exists within a molecule.
Compare: ■
Your Turn 5.3
Polar Bonds
■
Intramolecular forces are within molecules. Intramural sports are played within a college.
For each pair, which is the more polar bond? In the bond you select, the electron pair is more strongly attracted to one of the atoms. Which one? Hint: Use Table 5.1. a. H–F or H–Cl b. N–H or O–H c. N–O or O–S Answer a. The H–F bond is more polar. The electron pair is more strongly attracted to the F atom.
We have made the case that bonds can be polar, some more than others. What about molecules? To help you predict if a molecule is polar, we offer two useful generalizations: ■
■
A molecule that contains only nonpolar bonds must be nonpolar. For example, the Cl2 and H2 molecules are nonpolar. A molecule that contains polar covalent bonds may or may not be polar. The polarity depends on the geometry of the molecule.
For example, the water molecule contains two polar bonds and the molecule is polar (Figure 5.3). Each H atom carries a partial positive charge (d1), and the oxygen atom carries a partial negative charge (d2). Because the molecule is bent, overall it is polar. Many of the unique properties of water are a consequence of its polarity. But before we continue the story of water, take a moment to complete this activity.
␦⫺
O ␦⫹
H
H
␦⫹
Figure 5.3 H2O, a polar covalent molecule with polar covalent bonds.
Consider This 5.4
The Carbon Dioxide Molecule
Revisit the carbon dioxide molecule. You can find its Lewis structure in Figure 3.14. a. Are the covalent bonds in CO2 polar or nonpolar? Use Table 5.1. b. Analogous to Figure 5.3, draw a representation for CO2. c. In contrast to the H2O molecule, the CO2 molecule is not polar. Explain.
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5.2 |
H
H
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O
O
H
H
Hydrogen O bonds H H O H H Covalent bonds
H
O
H
Figure 5.4 Hydrogen bonding in water (distances not to scale).
Figures Alive! Visit the
The Role of Hydrogen Bonding
Consider what happens when two water molecules approach each other. Because opposite charges attract, a H atom (d1) on one of the water molecules is attracted the O atom (d2) on the neighboring water molecule. This is an example of an intermolecular force, that is, a force that occurs between molecules. But with more than two water molecules, the story gets more complicated. Examine each H2O molecule in Figure 5.4 and note the two H atoms and two nonbonding pairs of electrons on the O atom. These allow for multiple intermolecular attractions. This phenomenon of attracting between molecules is called “hydrogen bonding.” A hydrogen bond is an electrostatic attraction between a H atom bonded to a highly electronegative atom (O, N, or F) and a neighboring O, N, or F atom, either in another molecule or in a different part of the same molecule. Hydrogen bonds typically are only about one tenth as strong as the covalent bonds connecting atoms within molecules. Also, the atoms involved in hydrogen bonding are farther apart than they are in covalent bonds. In liquid water there may be three or four hydrogen bonds per water molecule, as shown in Figure 5.4.
textbook’s website to learn more about hydrogen bonding.
Your Turn 5.5 Compare: ■ ■
Intermolecular forces are between molecules. Intercollegiate sports are played between colleges.
Sulfur is less electronegative than oxygen and nitrogen. Although H atoms bonded to N or O atoms can form hydrogen bonds, H atoms bonded to S atoms cannot.
Bonding in Water
a. Explain the dashed lines between water molecules in Figure 5.4. b. In the same figure, label the atoms on two adjacent water molecules with d1 or d2. How do these partial charges help to explain the orientation of the molecules? c. Are hydrogen bonds intermolecular or intramolecular forces? Explain.
Although hydrogen bonds are not as strong as covalent bonds, hydrogen bonds still are quite strong compared with other types of intermolecular forces. The boiling point of water gives us evidence for this assertion. For example, consider H2S, a molecule that is analogous to water but does not hydrogen bond. H2S boils at about 260 °C and so is a gas at room temperature. In contrast, water boils at 100 °C. Because of hydrogen bonding, water is a liquid at room temperature as well as at body temperature (about 37 °C). Life’s very existence on our planet depends on this fact.
Consider This 5.6
Bonds Within and Between Water Molecules
Are any covalent bonds broken when water boils? Explain with drawings. Hint: Start with molecules of water in the liquid state as shown in Figure 5.4. Make a second drawing to show water in the vapor phase.
Hydrogen bonding also can help you understand why ice cubes and icebergs float. Ice is a regular array of water molecules in which every H2O molecule is hydrogenbonded to four others. The pattern is shown in Figure 5.5. Note the empty space in the form of hexagonal channels. When ice melts, the pattern is lost, and individual H2O molecules can enter the open channels. As a result, the molecules in the liquid state are more closely packed than in the solid state. Thus, a volume of one cubic centimeter (1 cm3) of liquid water contains more molecules than 1 cm3 of ice. Consequently, liquid water has a greater mass per cubic centimeter than ice. This is simply another way of saying that the density, the mass per unit volume, of liquid water is greater than that of ice. People often confuse density with mass. For example, popcorn has a low density, and people say that a bag of popcorn feels “light.” Similarly, you may hear someone say that lead is “heavy.” Large pieces of lead are indeed often quite heavy, but it is more accurate to say that lead has a high density (11.3 g/cm3).
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O H covalent bond hydrogen bond
Figure 5.5 The hydrogen-bonded lattice structure of the common form of ice. Note the open channels between “layers” of water molecules that cause ice to be less dense than water.
We usually express the mass of water in grams. Expressing its volume is a bit trickier. We use either cubic centimeters or milliliters (mL)—the two units are equivalent. The density of water is 1.00 g/cm3 at 4 °C and varies only slightly with temperature. So for convenience, we sometimes say that 1 cm 3 of water has a mass of 1 g. On the other hand, 1.00 cm 3 of ice has a mass of 0.92 g, so its density is 0.92 g/cm 3. The bottom line? The ice cubes in your favorite beverage float rather than sink. Unlike water, most substances are denser as solids. The fact that water shows the reverse behavior means that in the winter, ice floats on lakes rather than sinking. This topsy-turvy behavior means that surface ice, often covered by snow, can act as an insulator and keep the lake water beneath from freezing solid. Aquatic plants and fish thus can live in a freshwater lake during winter. And when the ice melts in spring, the water formed sinks, helping to mix the nutrients in the freshwater ecosystem. Needless to say, water’s unique behavior has implications both for the biological sciences and for life itself. The phenomenon of hydrogen bonding is not restricted to water. It can occur in other molecules that contain covalent O–H or N–H bonds. The H bonds help stabilize the shape of large biological molecules, such as proteins and nucleic acids. For example, the double-helix structure of the DNA molecule is stabilized by hydrogen bonds between the two DNA strands. When DNA undergoes transcription, it “unzips” as the hydrogen bonds across the two strands break. Again, hydrogen bonding plays an essential role in the processes of life. We end this section by examining one last unusual property of water, its uncommonly high capacity to absorb and release heat. Specific heat is the quantity of heat energy that must be absorbed to increase the temperature of 1 gram of a substance by 1 °C. The specific heat of water is 4.18 J/g ? °C. This means that 4.18 J of energy is needed to raise the temperature of 1 g of liquid water by 1 °C. Conversely, 4.18 J of heat must be removed in order to cool 1 g of water by 1 °C. Water has one of the highest specific heats of any substance and is said to have a high heat capacity. Because of this, it is an exceptional coolant. When water evaporates, it can be used to carry away the excess heat in a car radiator, in a power plant, or in the human body.
For any liquid at any temperature: 1 cm3 5 1 mL
To reiterate, water is most dense at 4 °C. At 0 °C, it is slightly less dense.
DNA molecules form hydrogen bonds between different strands of DNA. In contrast, proteins can form hydrogen bonds within different regions within the same molecule. Look for more about the structures of proteins and DNA in Chapter 12.
The joule and the calorie, units of energy, were defined in Section 4.2. The specific heat of water can also be expressed (using calories) as 1.00 cal/g ? °C.
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Consider This 5.7
A Barefoot Excursion
Have you ever walked barefoot across a carpeted floor and then onto a tile or stone floor? If not, try it and see what you notice. Based on your observation, does carpet or tile have the higher heat capacity?
Because of water’s high specific heat, large bodies of water influence regional climate. When water evaporates from seas, rivers, and lakes, heat is absorbed. By absorbing vast quantities of heat, the oceans and the droplets of water in clouds help mediate global temperatures. Since water has a higher capacity to “store” heat than the ground does, when the weather turns cold, the ground cools more quickly. Water retains more heat and is able to provide more warmth for a longer time to the areas bordering it. Such properties should be familiar to anyone who has ever lived near a large body of water. We have just examined some of the critical properties of water that influence life on our planet. Before we explore its ability to dissolve many different substances, we seek a broader picture of what water is used for and what issues are related to its use.
5.3 |
Water Use
Just as we need clean, unpolluted air to breathe, we also need potable water; that is, water that is safe to drink and to cook with. Nonpotable water may contain toxic metals such as arsenic, or it may be contaminated with bacteria such as those that cause cholera. Nonetheless, water that is not safe to drink still has its uses. For example, untreated water from nearby rivers and lakes can be transported for street washing, keeping down dust, or irrigation, as shown in Figure 5.6.
Figure 5.6 Water truck at the University of Alaska, Fairbanks, with a warning that the water is not fit to drink.
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Our need for potable water is but a small part of the larger picture of water use. Over 390 liters (, 100 gallons) of water per day are required to support the lifestyle of the average U.S. citizen. In addition, it can take 4 L to process one can of fruit or vegetables, and a whopping 240,000 L is needed to produce 1 ton of steel. Globally, we consume about 10% of water for domestic or household use, 20% for industrial needs, and 70% for agriculture. For a given region of the globe, these percentages vary. For example, water availability depends on climate, and so more water is needed for agriculture in the arid parts of our planet. Water use also can change over time. If a country shifts its economic base, its water use may increase or decrease. For example, as China shifts from agriculture to industry, its need for water will change. Agriculture takes the biggest gulp of water. Worldwide, we grow wheat, rice, corn, soybeans, and other crops that we have come to depend on. In feedlots, water is needed to raise the beef, pork, and chicken sold in grocery stores. Table 5.2 shows the relative amounts of water needed to produce some of the foods we eat. Although it takes a smaller sip, industry also uses its share. Industrial water use includes chemical processing (such as dying textiles), cooling (running power plants), and washing (such as cleaning fibers). Our discussion of water use falls in a larger context. A water footprint is an estimate (for an individual or a nation) of the amount of water required to sustain the consumption of goods and services. The total global water footprint is about 7 3 1015 liters per year (L/yr). Doing the math, this translates to 1 3 10 6 L/yr for every person on Earth. Alternatively, this is about half the water required to fill an Olympic-sized swimming pool. The numbers listed in Table 5.2 are examples of water footprints.
Consider This 5.8
Your Own Water Footprint
Thanks to several organizations, you can now calculate your water footprint. A link is provided at the textbook’s website. a. Calculate your personal water footprint. What did you discover about your water usage? b. Do you feel that the survey was fair in assessing the water you use? Explain. c. Name ways you now might use water differently, knowing what you know.
Figure 5.7 ranks nations according to their water footprints. The United States has the largest per capita water footprint at 2 3 106 L/person/yr. In part, this is due to the high consumption of meat and industrial goods. Although India has the largest total consumption of any nation at 1 3 1015 L/yr, this value stems from its large population. As you can see from Figure 5.7, each U.S. citizen uses nearly two and a half times more water than does a citizen in India.
Table 5.2 Food (1 kg)
Fresh Water Needed to Produce Food Water (L)
beef
15,500
pork chicken sheep
Food (1 kg)
Water (L)
rice
3,400
4,800
soybeans
1,800
3,900
wheat
1,300
6,100
corn
Source: Water Footprint Network. www.waterfootprint.org.
900
205
A liter (L) contains 1000 milliliters (mL). One gallon is about 3.8 L.
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Water footprint (liters per person per year)
3000 Domestic water consumption Industrial goods Agricultural goods
2500 2000 1500 1000 500
U.S.
Italy
Thailand
Nigeria
Russia
Mexico
Brazil
Indonesia
Pakistan
Japan
India
China
Figure 5.7 The national water footprint per capita and the contribution of different consumption categories for some selected countries. Source: Data from A. Y. Hoekstra and A. K. Chapagain. Water footprints of nations, Water Resource Management, (2007) 21:35–48.
On a more personal level, our water use can be viewed through products we may encounter each day. Consider a 200-mL glass of milk. The volume of water used to produce this glass of milk is 2000 L, or ten thousand times the volume in one glass of milk! This includes the water to care for the cow and the water used to grow the food that it eats. It also includes the water used at a dairy farm to collect the milk and clean the equipment. You can check out the water footprints for some of your favorite beverages and consumer goods in Table 5.3. Water footprint values are controversial and far from exact. Our intent in providing them is not to label items as either good or bad for consumption. Rather, these values are meant to increase your awareness of how we use water to produce goods and to provide you with a more-inclusive picture of water use. For example, on first inspection of Table 5.3, you might be tempted to forego cotton T-shirts. Cotton is indeed a thirsty crop and has been grown in arid climates in which farmers have had to import water, often using a canal system. Huge amounts of water are required to process cotton fibers into the shirts that we wear. This large water footprint for cotton can encourage us to irrigate more efficiently and to design industrial practices that conserve water. A green chemistry solution applied to the processing of cotton is discussed in Section 5.12.
Table 5.3
Product Water Footprints
Product 1 cup of coffee (125 mL)
Water Footprint (L) 140
1 apple (100 g)
70
1 orange (100 g)
50
1 glass of orange juice (200 mL)
170
1 egg (60 g)
200
1 hamburger (150 g)
2400
1 cotton T-shirt (250 g)
2700
1 computer chip (2 g)
32
Source: Adapted from A. Y. Hoekstra and A. K. Chapagain, Water footprints of nations, Water Resource Management (2007) 21: 35–48.
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Water Issues
Of all the water on Earth, amazingly only 3% is fresh water, as shown in Figure 5.8. About 68% of this fresh water is locked up in glaciers, ice caps, and snowfields, although with global warming, the ice is decreasing. About 30% of our fresh water is underground and must be pumped for our use. Less than 1% is tied up in our atmosphere, and finally about 0.3% is the most easily accessible on Earth’s surfaces in lakes, swamps, and rivers. If all the water on Earth were represented by the contents of a 2-L bottle, then 60 mL of this would be fresh water, and the amount of this fresh water that is accessible is about four drops!
Skeptical Chemist 5.9
A Drop to Drink
The previous section states that four drops in 2 L corresponds to the amount of fresh water available for our use. Is this accurate? Make a determination of your own. Hint: Use both the relationships shown in Figure 5.8 and assume 20 drops/mL.
The previous activity makes water look precious, and arguably it is, depending on where you live. We strive to use the most convenient source of fresh water for human activities, which in many cases comes from surface water, the fresh water found in lakes, rivers, and streams (Figure 5.9). Less convenient to access is groundwater, fresh water found in underground reservoirs also known as aquifers. However, by pumping groundwater from wells drilled into these underground reservoirs, people worldwide come to depend on this source of fresh water. A third plentiful source, sea water, only works if we remove its salt through a process called desalination. In the United States, the average household spends about $2.00 for 1000 gallons (3800 L) of water for home use. Compared with other commodities such as gasoline, this is relatively inexpensive. Most citizens in the United States obtain their water from a faucet or drinking fountain (Figure 5.10a). Municipal tap water may be consumed with or without further home filtering. Some people draw water from their own wells. If out on a hiking trail, water may be purified from nearby streams. Water also can be purchased in plastic, aluminum, or glass containers (Figure 5.10b). Each of these options most certainly has advantages and disadvantages. However, the point is that an Lakes, rivers, swamps, atmosphere, soil moisture
Surface water 0.3%
Fresh water 3% 3%
Oceans 97% (salt water)
Figure 5.8 Distribution of fresh water on Earth.
Look for more about desalination in Section 5.12.
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Figure 5.9 Lakes and reservoirs provide much of our drinking water. This one, Hetch Hetchy, provides water to San Francisco, California.
(a)
infrastructure supports the availability of water. Those who live in economically advantaged nations with a strong infrastructure can make decisions about how frequently they drink, how much they drink, and about the source of the water they drink. What if you live where you cannot turn on a tap or buy bottled water? You might need to walk for miles to reach a water source, fill a container, and carry it home to your family (Figure 5.11a). You might need to regularly depend on a water truck to stop by your house and deliver water (Figure 5.11b). On a larger scale, you might have to depend on governments working with engineers to design mega-structures to help move water from one region of the country to another. Aqueducts in the United States move water from the Colorado River to the Southwest. In China, a massive south-to-north
(b)
Figure 5.10 (a) Some people (but not all) take the safety of drinking water for granted. (b) Across the planet, people drink bottled water for many different reasons.
(a)
(b)
Figure 5.11 (a) Young girls walking home with water buckets. (b) A truck provides water for a community.
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water transfer project costing billions of dollars promises to funnel more than 45 trillion liters of water per year from the Yangtze River basin in southern China to the water-thirsty north. Such major diversions of water usually come with a considerable cost, as we will explore in a later section. Unfortunately, a mismatch exists between where water is found on our planet and where people need to use it. The phrase “location, location, location” applies well to water. Several issues, including global climate change, overconsumption and inefficient use of water, and contamination further complicate the availability of water. We now discuss each in turn.
Global Climate Change Violent storms and floods bring water in ferocious abundance, as witnessed by recent flooding in the midwestern United States, Europe, and China. At the other extreme, drought or desertification creates crippling shortages. Either way, climate affects the supply and demand for water. When regions become hotter and drier, the demand increases for water, especially water for irrigation. For example, consider the Great Lakes Basin Water Resource Compact that was first signed in 2005 by the governors of the eight states and the leaders of the two Canadian provinces that border the Great Lakes. In part, this pact was driven by fears that drought-stricken regions might try to raid the water supply of the Great Lakes (Figure 5.12).
Figure 5.12 Political cartoon of the evolution of the Great Lakes Compact. Source: Joe Heller
Your Turn 5.10
Great Lakes, Great News!
A long-debated agreement to regulate the withdrawal of water from the Great Lakes passed the U.S. House of Representatives on September 23, 2008. A local newspaper in Wisconsin offered its readers these quiz questions. a. How long does it take a single drop of rain to cycle through Lake Superior? b. How much water is permanently lost from the Great Lakes drainage basin every year? c. How many national parks and lakeshores are there on the Great Lakes? Note: As a reference point, the Great Lakes contain about 220 quadrillion liters of water. Answers a. About 200 years b. About 1 trillion liters, or more than 15,000 L for each of the basin’s 37 million residents c. Ten national parks plus hundreds of state and provincial parks with more than 70 million visitors annually
Weather describes the physical conditions of the atmosphere, such as its temperature, pressure, moisture, and wind. In contrast, climate describes the long-term weather patterns in a region.
Lake Superior holds over 10% of the world’s fresh water supply.
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In Chapter 3, you learned about the carbon cycle.
This newspaper quiz brings up an important point: A drop of water that enters Lake Superior will one day leave it. Just as our planet has a carbon cycle, it also has a water cycle. The water that falls on land either evaporates or eventually finds its way to the ocean. The water cycle not only includes lakes and rivers, but also their frozen cousins—glaciers and sea ice. Climate plays an important role in the timing of the water cycle. For example, glaciers stabilize water flow over time as they accumulate snow pack during winter months and then release a regular stream of water during summer months. The great glaciers of the Himalayas feed seven of the largest rivers in Asia, ensuring a reliable water supply for 2 billion people—almost one third of the world’s population. If climate is altered and these glaciers are not replenished on an annual cycle to a point where they cannot melt and sustain the rivers in the region, the effect on the people in Asia who have come to rely on glaciers as water reservoirs could be devastating. Climate change also affects the timing of events in ecosystems. For example, insects, birds, and plants need to appear in the right order so that the birds can feed, the insects can pollinate, and the plants can grow. If birds migrate earlier in the spring, they may arrive before enough insects have hatched for food. Conversely, if too many insects hatch before the birds are present to eat them, the insects may devastate crops. Either way, water is a key variable.
Overconsumption and Inefficient Use
In the context of air quality, the tragedy of the commons was first mentioned in Section 1.12.
Figure 5.13 One of the world’s largest aquifers, the High Plains Aquifer, is show in dark blue on this map.
In many places, water is being pumped out of the ground faster than it can be replenished by the natural water cycle. For example, much of the bountiful grain harvest from the central United States is a result of using water from the Ogallala Aquifer (also called the High Plains Aquifer). This vast aquifer trapped water from the last ice age and runs from South Dakota to Texas (Figure 5.13). Clearly, it is not a sustainable practice to continuously pump water from all aquifers. Some aquifers do recharge more quickly by precipitation and runoff. Others may take hundreds or even thousands of years to recharge naturally. Continuous pumping can bring other harmful outcomes as well. For example, if water is removed from a geologically unstable area near the coast, salt water may intrude into a freshwater aquifer. This situation presents us with another example of the tragedy of the commons. The water from aquifers is a resource used in common, yet no one in particular is responsible for this resource. If water is overdrawn for agriculture or some other purpose, this act can be to the detriment of all. An overdraw of surface water can create other problems. Consider Kazakhstan and Uzbekistan, countries that border the Aral Sea. Until recently, this sea was the world’s fourth largest inland body of fresh water. In the 1960s, workers in the former Soviet Union built a network of canals that diverted this water from the rivers that fed the Aral Sea in order to grow cotton in the arid climate. Consequently, the Aral Sea dried up, as shown in Figure 5.14. Although the ecosystem once was rich as a fishery, today only three very salty pools of water remain. The United Nations has called this the greatest environmental disaster of the 20th century. Dust that is laden with toxins, pesticides, and salt now blows in the region, causing health problems and contributing to poverty. Not only were the rivers feeding the Aral Sea diverted, but the river water taken was also used inefficiently. For example, the water used to irrigate cotton was transported in open canals. Given the arid climate, much of the water was lost through evaporation. Other wasteful practices in areas where water is scarce include using water sprinklers to irrigate fields, cultivating lush green lawns in residential areas, and not fixing leaky pipes in aging water distribution systems. Many factors influence this unsustainable use of water, including a lack of knowledge of other irrigation options, subsidies to keep the cost of water low, and the high cost of repairing a water distribution system.
Contamination We expect our water to be safe; that is, devoid of harmful chemicals and microbes. Access to clean water varies worldwide, as shown in Figure 5.15. More than a billion people (1 in 6), principally in developing nations, lack access to safe drinking water.
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Aral Sea 1973
Aral Sea 1987
Aral Sea 1999
Aral Sea 2009
Figure 5.14 The Aral Sea has lost more than 80% of its water over a period of 30 years. The rivers that fed it were diverted to irrigate crops.
Percentage of Safe Drinking Water Access by Total Population Over 90%
75–90%
60–75%
45–60%
Figure 5.15 Access to safe drinking water varies widely across the world. Source: © Compare Infobase.
30–45%
Under 30%
No Data
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Arsenic and fluoride ions occur in water in the form of a cation and an anion, respectively. Look for more about ions in Section 5.6.
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Another 2.6 billion lack basic sanitation; that is, garbage disposal and treatment of industrial and other wastewater. Each day, more than 5000 deaths worldwide occur in infants and young children because of infectious disease agents borne by unsafe water. In response to the global water needs in the 1980s, the United Nations Children’s Fund (UNICEF) funded pumps and the sinking of wells to tap into underground sources of water. Although water from aquifers is generally potable, tragically, the water pumped to the surface in India and Bangladesh was not. It contained arsenic and fluoride ions, both naturally occurring in minerals found in the bedrock of the aquifers. Arsenic and fluoride ions are cumulative poisons, and it may take years before the amount ingested reaches a harmful level. Thousands of people have been irreversibly poisoned, because the arsenic and fluoride ions were detected too late in their drinking water.
Consider This 5.11
A Decade for Water
As pointed out earlier in this chapter, the United Nations has designated 2005–2015 as the Decade for Action regarding water. The textbook’s website provides a link.
WAT E R F O R L I F E 2005 – 2015
a. Find two significant issues related to water that were not discussed in this section. Make a case for the importance of each. b. Summarize the theme for this year’s World Water Day.
Let us now turn to topics that help us better understand why water is able to dissolve and mix with so many substances, including essential nutrients and contaminants.
5.5 |
Aqueous Solutions
Water dissolves a remarkable variety of substances. As we will see, some of them, including salt, sugar, ethanol, and the air pollutant SO2, are very soluble in water. In comparison, limestone rock, oxygen, and carbon dioxide dissolve only in tiny amounts. To build your understanding about water quality, you need to know what dissolves in water, why it dissolves, and how to specify the concentration of the resulting solution when it does dissolve. This section tackles solution concentrations; the section that follows addresses solubility. Let’s begin with some useful chemical terminology. Water is a solvent, a substance, often a liquid, that is capable of dissolving one or more pure substances. The solid, liquid, or gas that dissolves in a solvent is called the solute. The result is called a solution, a homogeneous (of uniform composition) mixture of a solvent and one or more solutes. In this section, we are particularly interested in aqueous solutions, solutions in which water is the solvent. Because water is such a good solvent, it practically never is “100% pure.” Rather, it contains impurities. For example, when water flows over the rocks and minerals of our planet, it dissolves tiny amounts of the substances that they contain. Although this usually causes no harm to our drinking water, occasionally the ions dissolved in water are toxic. For example, as we noted in the previous section, if the water contacts minerals that contain arsenic or fluoride ions, the water may be rendered nonpotable. The water on our planet also comes in contact with air. When it does, it dissolves tiny amounts of the gases in the air, most notably oxygen and carbon dioxide. Some air pollutants are very soluble in water. So when it rains, the water actually cleans some of the pollutants out of the air, including SO2 and NO2. As we will see in Chapter 6, the acidic solutions that form can have serious consequences for the environment. Humans also contribute to the number of substances dissolved in water. When we wash clothes, we add not only the spent detergent, but also whatever made our clothing dirty in the first place. When we flush a toilet, we add liquid and solid wastes. Our urban streets add solutes to rain water during the process of storm run-off. And our agricultural practices add fertilizers and other soluble compounds to water.
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Water for Life
What does water’s being a good solvent mean for our drinking water? In order to assess water quality, you need to know several things. One is a way to specify how much of a substance has dissolved, so that you can compare the value with a known standard. In other words, you need to understand the concept of concentration. This was first introduced in Chapter 1 in relation to the composition of air. For example, O2 and N2 are about 21% and 78% of dry air, respectively. We revisited concentration again in Chapters 2 and 3, exploring the concentrations of chlorine compounds in the stratosphere and greenhouse gases in the troposphere. For example, carbon dioxide has a concentration of about 390 ppm in the air. Now we examine this concept in terms of substances dissolved in water. As we will see, percent and parts per million are valid ways of expressing concentrations for aqueous solutions as well. To get started with solution concentrations, let’s use a familiar analogy—sweetening a cup of tea. If 1 teaspoon of sugar is dissolved in a cup of tea, the resulting solution has a concentration of 1 teaspoon per cup. Note that you would have this same concentration if you were to dissolve 3 teaspoons of sugar in 3 cups of tea, or half a teaspoon in half a cup of tea. If your recipe is tripled or halved, the sugar and tea are adjusted proportionally. Therefore, the concentration, the ratio of the amount of solute to the amount of solution—or in this case the ratio of sugar dissolved to make the solution—is the same in each case. Solute concentrations in aqueous solution follow the same pattern but are expressed with different units. We use four ways to express concentration: percent, parts per million, parts per billion, and molarity. Three of these should already be familiar to you. The fourth, molarity, uses the mole concept introduced in Chapter 3. Percent (%) means parts per hundred. For example, an aqueous solution containing 0.9 grams of sodium chloride (NaCl) in 100 grams of solution is a 0.9% solution by mass. This concentration of sodium chloride is referred to as “normal saline” in medical settings when given intravenously. You may find the antiseptic hydrogen peroxide (H2O2) in your medicine cabinet as a 3% aqueous solution by volume. It contains 3 milliliters of H2O2 in every 100 milliliters of aqueous solution. Percent is used to express the concentration of a wide range of solutions. But when the concentration is very low, as is the case for many substances dissolved in drinking water, parts per million (ppm) is more commonly used. For example, water that contains 1 ppm of calcium ion contains the equivalent of 1 gram of calcium (in the form of the calcium ion) dissolved in 1 million grams of water. The water we drink contains substances naturally present in the parts per million range. For example, the acceptable limit for nitrate ion, found in well water in some agricultural areas, is 10 ppm; the limit for the fluoride ion is 4 ppm. Although parts per million is a useful concentration unit, measuring 1 million grams of water is not very convenient. We can do things more easily by switching to the unit of a liter. One ppm of any substance in water is equivalent to 1 mg of that substance dissolved in a liter of solution. Here is the math: 1 ppm
1000 mg solute 1000 g water 1 mg solute 1 g solute 1 g solute 1 L water 1 L water 1 106 g water
Municipal water utilities typically use the unit milligrams per liter (mg/L) to report the minerals and other substances dissolved in tap water. For example, Table 5.4 shows a tap water analysis from an aquifer that supplies a midwestern community in the United States.
Table 5.4 Cation calcium ion
Tap Water Mineral Report mg/L 97
Anion
mg/L
sulfate ion
45 75
magnesium ion
51
chloride ion
sodium ion
27
nitrate ion
4
fluoride ion
1
213
For solutions at low concentration, the mass of the solution is approximately the mass of the solvent.
These limits for fluoride and nitrate ions reflect the U.S. standards. See Section 5.10.
1000 grams (1 3 103 g) of H2O can be taken to have a volume of 1 liter. Strictly speaking, this is true only at 4 °C.
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Chapter 5
Mercury in water is present in a soluble form (Hg21) rather than as elemental Hg (“quicksilver”).
In aqueous solutions, 1 ppb 5 1 mg/L 1 ppm 5 1 mg/L
Some contaminants are of concern at concentrations much lower than parts per million and are reported as parts per billion (ppb). Assuming that 1 ppm corresponds to 1 second in nearly 12 days, then 1 ppb corresponds to 1 second in 33 years. Another way of looking at this is that one part per billion corresponds to a few centimeters on the circumference of the Earth! One contaminant found in the range of parts per billion is mercury. For humans, the primary source of exposure to mercury is food, mainly fish and fish products. Even so, the concentration of mercury in water needs to be monitored. One part per billion of mercury (Hg) in water is equivalent to 1 gram of Hg dissolved in 1 billion grams of water. In more convenient terms, this means 1 microgram (1 3 1026 g, or 1 mg) of Hg dissolved in 1 liter of water. The acceptable limit for mercury in drinking water is 2 ppb: 2 ppb Hg
2 g Hg 1 106 g Hg 1000 g H2O 2 g Hg 9 1 10 g H2O 1 g Hg 1 L H2O 1 L H2O
Convince yourself that the units cancel, as in the previous example.
Your Turn 5.12
Mercury Ion Concentrations
a. A 5-L sample of water contains 80 mg of dissolved mercury ion. Express the concentration of the solution in ppm and ppb. b. Would your answer in part a. be in compliance with a federal maximum mercury concentration of 2 ppb? Explain.
Molarity (M), another useful concentration unit, is defined as a unit of concentration represented by the number of moles of solute present in 1 liter of solution. Molarity (M) 5
The molar mass of NaCl (58.5 g) is calculated by adding the molar mass of sodium (23.0 g) plus the molar mass of chlorine (35.5 g). Section 3.7 explains molar mass calculations.
(aq) is short for aqueous, indicating that the solvent is water.
The great advantage of molarity is that solutions of the same molarity contain exactly the same number of moles of solute and hence the same number of molecules (ions or atoms) of solute. The mass of a solute varies depending on its identity. For example, 1 mole of sugar has a different mass than 1 mole of sodium chloride. But if you take the same volume, all 1 M solutions (read as “one molar”) contain the same number of moles. As an example, consider a solution of NaCl in water. The molar mass of NaCl is 58.5 g; therefore, 1 mol of NaCl has a mass of 58.5 g. By dissolving 58.5 g of NaCl in some water and then adding enough water to make exactly 1.00 L of solution, we would have a 1.00 M NaCl aqueous solution (Figure 5.16). We have prepared a onemolar solution of sodium chloride. Note the use of a volumetric flask, a type of glassware that contains a precise amount of solution when filled to the mark on its neck. But because concentrations are simply ratios of solute to solvent, there are many ways to make a 1.00 M NaCl(aq) solution. Another possibility is to use 0.500 mol NaCl (29.2 g) in 0.500 L of solution. This requires the use of a 500-mL volumetric flask, rather than the 1-L flask shown in Figure 5.16. 1 M NaCl(aq) 5
Remember that 1 ppm 5 1 mg/L and that the molar mass of Hg is 200.6 g/mol.
moles of solute liter of solution
1 mol NaCl 0.500 mol NaCl or , etc. 1 L solution 0.500 L solution
Let’s say you have a water sample with 150 ppm of dissolved mercury. What is this concentration expressed in molarity? You might do the calculation this way:
150 ppm Hg
1 g Hg 1 mol Hg 7.5 104 mol Hg 150 mg Hg 7.5 104 M Hg 1000 mg Hg 200.6 g Hg 1 L H2O 1 L H2O
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1. Add 1.00 mol (58.5 g) NaCl to empty 1.000 L flask. 2. Add water until flask is about half full. Swirl to mix water and NaCl. 3. Add water until liquid level is even with 1000-mL mark.
1000 mL
4. Stopper and mix well. 1.00 M NaCl solution
Figure 5.16 Preparing a 1.00 M NaCl aqueous solution.
Thus, a sample of water containing 150 ppm of mercury also can be expressed as 7.5 3 1024 M Hg.
Your Turn 5.13
Moles and Molarity
a. Express a concentration of 16 ppb Hg in units of molarity. b. For 1.5 M and 0.15 M NaCl(aq), how many moles of solute are present in 500 mL of each? c. A solution is prepared by dissolving 0.50 mol NaCl in enough water to form 250 mL of solution. A second solution is prepared by dissolving 0.60 mol NaCl to form 200 mL of solution. Which solution is more concentrated? Explain. d. A student was asked to prepare 1.0 L of a 2.0 M CuSO4(aq) solution. The student placed 40.0 g of CuSO4 crystals in a volumetric flask and filled it with water to the 1000-mL mark. Was the resulting solution 2.0 M? Explain.
In this section, we made the case that water is an excellent solvent for a wide variety of substances and that we can express the concentration of these substances numerically. As promised, the next section helps you to build an understanding of how and why substances dissolve in water.
5.6 |
A Closer Look at Solutes
Sugar and salt both dissolve in water. However, the solutions of these two solutes are quite different in nature—the former conducts electricity and the latter does not. Experimentally, we demonstrate this difference using a conductivity meter, an apparatus that produces a signal to indicate that electricity is being conducted. The conductivity meter shown in Figure 5.17 is built from wires, a battery, and a light bulb. As long as the electrical circuit is not completed, the bulb does not glow. For example, if the two wires are placed into distilled water or a sugar solution in distilled water, the bulb does not light. However, if the separate wires are placed into an aqueous solution of salt, the bulb turns on. Perhaps the light also has gone on in the mind of the experimenter! Distilled water does not conduct electricity. The same is true for an aqueous solution of sugar. Sugar is a nonelectrolyte, a solute that is nonconducting in aqueous solutions. But an aqueous solution of common table salt, NaCl, conducts electricity and the light bulb glows. Sodium chloride is classified as an electrolyte, a solute that conducts electricity in aqueous solution.
Distillation, a process to purify water, is discussed in Section 5.12.
The term electrolyte is used in connection with sports drinks. Some taste slightly salty because they contain sodium salts.
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Chapter 5
(b)
(a)
(c)
Figure 5.17 Conductivity experiments. (a) Distilled water (nonconducting). (b) Sugar dissolved in distilled water (nonconducting). (c) Salt dissolved in distilled water (conducting).
Remember to indicate ions in aqueous solution using (aq).
Cl
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Figure 5.18 The arrangement of Na1 and Cl2 in a crystal of sodium chloride.
Valence (outer) electrons were described in Section 2.2.
What makes salt solutions behave differently from sugar solutions or from pure water? The observed flow of electric current through a solution involves the transport of electric charge. The ability of aqueous NaCl solutions to conduct electricity suggests that they contain some charged species capable of moving electrons through the solution. When solid NaCl dissolves in water, it separates into Na1(aq) and Cl2(aq). An ion is an atom or group of atoms that has acquired a net electric charge as a result of gaining or losing one or more electrons. The term is derived from the Greek for “wanderer.” Na1 is an example of a cation, a positively charged ion. Similarly, Cl2 is an example of an anion, a negatively charged ion. No such separation occurs with covalently bonded sugar or water molecules, making these liquids unable to carry electric charge. It may surprise you to learn that Na1 and Cl2 exist both in crystals of salt (such as those in a salt shaker) and in an aqueous solution of NaCl. Solid sodium chloride is a three-dimensional cubic arrangement of sodium and chloride ions. An ionic bond is the chemical bond formed when oppositely charged ions attract. In the case of NaCl, ionic bonds hold the crystal together; there are no covalently bonded atoms, only positively charged cations and negatively charged anions held together by electrical attractions. An ionic compound is composed of ions that are present in fixed proportions and arranged in a regular, geometric structure. In the case of NaCl, each Na1 is surrounded by six oppositely charged Cl2 ions. Likewise, each Cl2 is surrounded by six positively charged Na1 ions. A single, tiny crystal of sodium chloride consists of many trillions of sodium ions and chloride ions in the arrangement shown in Figure 5.18. We’ve described ionic compounds, but we still need to explain why certain atoms lose or gain electrons to form ions. Not surprisingly, the answer involves the distribution of electrons within atoms. For example, recall that a neutral sodium atom has 11 electrons and 11 protons. Sodium, like all metals in Group 1A, has one valence electron. This electron is rather loosely attracted to the nucleus and can be easily lost. When this happens, the Na atom forms Na1, a cation. Na
Na e
[5.1]
1
The Na has a 11 charge because it contains 11 protons but only 10 electrons. It also has a complete octet, just like the neon atom. Table 5.5 shows the comparison between Na, Na1, and Ne.
Table 5.5
Electronic Bookkeeping for Cation Formation
Sodium Atom
Sodium Ion
Neon Atom
Na
Na1
Ne
11 protons
11 protons
10 protons
11 electrons
10 electrons
10 electrons
net charge of 0
net charge of 11
net charge of 0
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Table 5.6
217
Electronic Bookkeeping for Anion Formation
Chlorine Atom
Chloride Ion
Argon Atom
Cl
Cl2
Ar
17 protons
17 protons
18 protons
17 electrons
18 electrons
18 electrons
net charge of 0
net charge of 12
net charge of 0
Unlike sodium, chlorine is a nonmetal. Recall that a neutral chlorine atom has 17 electrons and 17 protons. Chlorine, like all nonmetals in Group 7A, has seven valence electrons. Because of the stability associated with eight outer electrons, it is energetically favorable for chlorine to gain one electron. Cl e
Cl
[5.2]
2
The chloride ion (Cl ) has 18 electrons and 17 protons; thus the net charge is 12 (Table 5.6). Sodium metal and chlorine gas react vigorously when they come in contact. The result is the aggregate of Na1 and Cl2, known as sodium chloride. In the formation of an ionic compound such as sodium chloride, the electrons are actually transferred from one atom to another, not simply shared as they would be in a covalent compound. Is there evidence for electrically charged ions in pure sodium chloride? Experimental tests show that crystals of sodium chloride do not conduct electricity. This makes sense, because in the crystal, the ions are fixed in place and so are unable to move and transport charge. However, when these crystals are melted, the ions are free to move, and the hot liquid conducts electricity. This provides evidence that ions are present. Like other ionic compounds, NaCl crystals are hard yet brittle. When hit sharply, they shatter rather than being flattened. This suggests the existence of strong forces that extend throughout the ionic crystal. Strictly speaking, there is no such thing as a specific, localized “ionic bond” analogous to a covalent bond in a molecule. Rather, ionic bonding holds together a large assembly of ions; in this case, Na1 and Cl2. Generally speaking, electron transfer to form cations and anions occurs between metallic elements and nonmetallic elements. Sodium, lithium, magnesium, and other metals have a strong tendency to give up electrons and form positive ions. They have very low electronegativity values. On the other hand (or the other side of the periodic table), chlorine, fluorine, oxygen, and other nonmetals have a strong attraction for electrons and readily gain them to form negative ions. Nonmetals have relatively high electronegativity values. Potassium chloride (KCl) and sodium iodide (NaI) are two of many such compounds.
Your Turn 5.14
Predicting Ionic Charge
a. Predict whether these atoms will form an anion or a cation based on their electronegativity values.
Li
S
K
N
b. Predict the ion that each of these will form. Then draw a Lewis structure for the atom and the ion, clearly labeling the charge on the ion.
Br
Mg
O
Al
Hint: Use the periodic table as a guide to the number of outer electrons. Then determine how many electrons must be lost or gained to achieve stability with an octet of electrons. Answer b. Bromine (Group 7A) gains one electron. The resulting ion has a charge of 12, just as was the case for chlorine. Here are the Lewis structures. Br and Br
⫺
Revisit Section 1.6 for more about metals and nonmetals.
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Chapter 5
This section opened with a discussion of sugar and salt. The names salt and sugar are both in common use, and you knew what we were talking about. Salt is the stuff you sprinkle on French fries. And sugar is the stuff that some people use to sweeten coffee. In fact, ordinary table salt (NaCl) is such an important example of an ionic compound that chemists frequently refer to others simply as “salts,” meaning crystalline ionic solids. As you will see in Chapter 11, sugars are another important class of compounds, and what we call “sugar” is really the compound sucrose. To delve further into the issues of water quality, you need to know the names of other salts; that is, ionic compounds. As you might guess, chemists name them using a long and careful set of rules. Luckily, we won’t discuss all of these rules here. Rather, we follow the “need-to-know” philosophy, helping you learn the ones that you need for understanding water quality.
5.7
|
Names and Formulas of Ionic Compounds
In this section, we work on the “vocabulary” you need in order to work with ionic compounds. As we pointed out in Chapter 1, chemical symbols are the alphabet of chemistry and chemical formulas are the words. Earlier, we helped you to “speak chemistry” by correctly using chemical formulas and names for the substances in the air you breathe. Now we do the same for the substances in the water you drink. Let’s begin with the ionic compound formed from the elements calcium and chlorine: CaCl 2. The explanation for the 1:2 ratio of Ca to Cl lies in the charges of the two ions. Calcium, a member of Group 2A, loses its two outer electrons to form Ca 21. Ca
Halogens were described in Sections 1.6 and 2.9.
Ca2 2 e
[5.3]
Chlorine, as we saw in equation 5.2, gains an outer electron to form Cl2. In an ionic compound, the sum of the positive charges equals the sum of the negative charges. Hence, the formula for this compound is CaCl2. The logic is the same with MgO and Al2O3, two other ionic compounds. These both contain oxygen, but in different ratios. Recall that oxygen, Group 6A, has six outer electrons. Thus a neutral oxygen atom can gain two electrons to form O22. The magnesium atom loses two electrons to form Mg21. These two ions must then combine in a 1:1 ratio so the overall charge is zero; the chemical formula is MgO. Note that although the charge always must be written on an individual ion, we omit the charges in the chemical formulas of ionic compounds. Thus, it is not correct to write the chemical formula as Mg21O22. The charges are implied by the chemical formula. Here is another example. Armed with the knowledge that aluminum tends to lose three electrons to form Al31, you can write the chemical formula of the ionic compound formed from Al31 and O22 as Al2O3. Here, a 2:3 ratio of ions is needed so that the overall electric charge on the compound is zero. Again, it is not correct to write the chemical formula as Al231O322. Earlier in the chapter, we referred to several ionic compounds by their names, including sodium chloride, sodium iodide, and potassium chloride. Observe the pattern: name the cation first, then the anion, modified to end in the suffix -ide. Thus, CaCl2 is calcium chloride, with each ion named for its element and with chlorine modified to read chloride. Similarly, NaI is sodium iodide and KCl is potassium chloride. The elements presented thus far formed only one type of ion. Group 1A and 2A elements only form 11 and 21 ions, respectively. The halogens form only 12 ions. Lithium bromide is LiBr. The ratio of 1:1 is understood because lithium only forms Li1 and bromine only forms Br2. The prefixes mono-, di-, tri-, and tetra- are not used when naming ionic compounds such as these. There is no need to call it monolithium monobromide. MgBr2 is magnesium bromide, not magnesium dibromide. Magnesium only forms Mg21, and the ratio of 1:2 is understood and so does not need to be stated.
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1 1A
18 8A 2 2A
13 3A
14 4A
Li Na Mg2
3 3B
4 4B
5 5B
6 6B
7 7B
8
9 8B
10
11 1B
12 2B
15 5A
16 6A
17 7A
N3 O2
F
S2
Cl
Al3
K Ca2
Cr2 Mn2 Fe2 Co2 Ni2 Cu Zn2 Cr3 Mn4 Fe3 Co3 Ni3 Cu2
Rb Sr2
Ag Cd2
Sn2 Sn4
Cs Ba2
Hg22 Hg2
Pb2 Pb4
Br I
Figure 5.19 Common ions formed from their elements. Ions in green (cations) or blue (anions) have only one charge. Ions in red (cations) have more than one possible ionic charge.
But some elements do form more than one ion, as you can see in Figure 5.19. Prefixes still are not used, but rather the charge on the ion must be specified using a Roman numeral. Take copper for example. If your instructor asks you to head down to the stockroom and grab some copper oxide, what do you do? You ask if what is wanted is copper(I) oxide or copper(II) oxide, right? Similarly, iron can form different oxides. Two forms are FeO (formed from Fe21) and Fe2O3 (commonly called rust and formed from Fe31). The names for FeO and Fe2O3 are iron(II) oxide and iron(III) oxide, respectively. Note the space after but not before the parenthesis enclosing the Roman numeral. Again compare. The name CuCl2 is copper(II) chloride, but the name of CaCl2 is calcium chloride. Calcium only forms one ion (Ca21), whereas copper can form two ions: Cu1 and Cu21.
Your Turn 5.15
Prefixes such as di- and tri- generally are not used in naming ionic compounds. Roman numerals are used with the name of the cation if it has more than one possible charge.
Ionic Compounds
Each pair of elements forms one or more ionic compounds. For each, write the chemical formulas and names. a. Ca and S d. Cl and Al
b. F and K e. Co and Br
c. Mn and O
Answer e. From Figure 5.19, Co can form Co21 and Co31. Br forms only Br2. The possible chemical formulas are CoBr2, cobalt(II) bromide, and CoBr3, cobalt(III) bromide.
One or both of the ions in an ionic compound can be a polyatomic ion, two or more atoms covalently bound together that have an overall positive or negative charge. An example is the hydroxide ion, OH2, with an oxygen atom covalently bonded to a hydrogen atom. The Lewis structure shown in Figure 5.20 reveals that there are 8 electrons, 1 more than the 6 valence electrons provided by one O atom and one H atom. The “extra” electron gives the hydroxide ion a charge of 12. Table 5.7 lists common polyatomic ions. Most are anions, but polyatomic cations also are possible, as in the case of the ammonium ion, NH41. Note that some elements (carbon, sulfur, and nitrogen) form more than one polyatomic anion with oxygen. The rules for naming ionic compounds containing polyatomic ions are similar to those for ionic compounds of two elements. Consider, for example, aluminum sulfate,
Brackets in a Lewis structure, such as with the hydroxide ion, call attention to an ion that fulfilled its octet by losing or gaining one or more electrons.
O
H
Figure 5.20 The Lewis structure for the hydroxide ion, OH2.
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Chapter 5 Common Polyatomic Ions
Table 5.7 Name
Formula
Name
Formula
acetate
C2H3O22
nitrite
NO22
bicarbonate*
HCO32
phosphate
PO432
carbonate
CO322
sulfate
SO422
hydroxide
OH
sulfite
SO322
hypochlorite
ClO2
ammonium
NH41
nitrate
NO32
2
*Also called the hydrogen carbonate ion.
an ionic compound that is used in many water purification plants. The compound is formed from Al31 and SO422. When you see Al2(SO4)3, mentally read this chemical formula as a compound that contains two ions: aluminum and sulfate. These ions are in a 2:3 ratio. As is true for all ionic compounds, the name of the cation is given first. The parentheses in Al2(SO4)3 are meant to help you. The subscript 3 applies to the entire SO422 ion that is enclosed in parentheses. Accordingly, “read” this as three sulfate ions, not as one larger unit composed of three sulfate ions. Similarly, in the ionic compound ammonium sulfide (see Table 5.8), the NH41 is enclosed in parentheses. The subscript of 2 indicates that there are two ammonium ions for each sulfide ion. Note that the charges are not shown in the chemical formula; they are assumed to be there. In some cases, though, the polyatomic ion is not enclosed in parentheses. Table 5.8 shows two examples. The PO432 ion in aluminum phosphate has no parentheses; similarly, the NH41 ion in ammonium chloride has no parentheses. Parentheses are omitted when the subscript of the polyatomic ion is 1. Nonetheless, you still have to “read” the chemical formula of AlPO4 as containing the phosphate ion, and you have to “read” NH4Cl as containing the ammonium ion. These activities will help you practice using polyatomic ions.
Your Turn 5.16
Polyatomic Ions I
Write the chemical formula for the ionic compound formed from each pair of ions. a. Na1 and SO422
b. Mg21 and OH2
Answers a. Na2SO4
c. Al31 and C2H3O22
d. CO322 and K1
b. Mg(OH)2
Your Turn 5.17
Polyatomic Ions II
Name each of these compounds. a. KNO3
b. (NH4)2SO4
Answers a. potassium nitrate
Table 5.8
c. NaHCO3
d. CaCO3
e. Mg3(PO4)2
b. ammonium sulfate
Ionic Compounds Containing Polyatomic Ions
Chemical formula
Al2(SO4)3
(NH4)2S
AlPO4
NH4Cl
Cation(s)
Al31 Al31
NH41 NH41
Al31
NH41
Anion(s)
SO422 SO422 SO422
S22
PO432
Cl2
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Your Turn 5.18
Polyatomic Ions III
Write the chemical formula for each of these compounds. a. b. c. d.
sodium hypochlorite (used to disinfect water) magnesium carbonate (found in some limestone rocks, makes water “hard”) ammonium nitrate (fertilizer, runoff can contaminate groundwater) calcium hydroxide (an agent used to remove impurities from water)
Answer d. Ca(OH)2. Two hydroxide ions (OH2) are needed for each calcium ion (Ca21).
5.8
|
The Ocean—An Aqueous Solution with Many Ions
Salt water! As we pointed out earlier, about 97% of the water on our planet is found in the oceans. This source of water contains much more than simple table salt (NaCl) dissolved in water. You are now in a position to understand why so many other ionic compounds can be found dissolved in our oceans. Recall from Section 5.1 that water molecules are polar. When you take salt crystals and dissolve them in water, the polar H2O molecules are attracted to the Na1 and Cl2 ions contained in these crystals. The partial negative charge (d2) on the O atom of a water molecule is attracted to the positively charged Na1 cations of the salt crystal. At the same time, the H atoms in H2O, with their partial positive charges (d1), are attracted to the negatively charged Cl2 anions. Over time, the ions are separated and then surrounded by water molecules. Equation 5.4 and Figure 5.21 represent the process of forming an aqueous solution. NaCl(s)
H2O
Na(aq) Cl(aq)
[5.4]
Cl
Na
Cl
Na
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Na
Cl
Cl
Figure 5.21 Sodium chloride dissolving in water.
Figures Alive! Visit the textbook’s website to learn more about sodium chloride dissolving in water.
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The process is similar for forming solutions of compounds containing polyatomic ions. For example, when solid sodium sulfate dissolves in water, the sodium ions and sulfate ions simply separate. Note that the sulfate ion stays together as a unit. Na2SO4(s)
H2O
2 Na(aq) SO42(aq)
[5.5]
Many ionic compounds dissolve in this manner. This explains why almost all naturally occurring water samples contain various amounts of ions. The same is also true for our bodily fluids, as these also contain significant concentrations of electrolytes.
Consider This 5.19
Electricity and Water Don’t Mix
Small electric appliances such as hair dryers and curling irons carry prominent warning labels advising the consumer not to use the appliance near water. Why is water a problem since it does not conduct electricity? What is the best course of action if a plugged-in hair dryer accidentally falls into a sink full of water?
If the principles we just described applied to all ionic compounds, our planet would be in trouble. When it rained, ionic compounds such as calcium carbonate (limestone) would dissolve and end up in the ocean! Fortunately, many ionic compounds are only slightly soluble or have extremely low solubilities. The differences arise because of the sizes and charges of the ions, how strongly they attract one another, and how strongly the ions are attracted to water molecules. Table 5.9 is your guide to solubility. For example, calcium nitrate, Ca(NO3)2, is soluble in water, as are all compounds containing the nitrate ion. Calcium carbonate, CaCO3, is insoluble, as are most carbonates. By similar reasoning, copper(II) hydroxide, Cu(OH)2, is insoluble, but copper(II) sulfate, CuSO4, is soluble.
Your Turn 5.20
Solubility of Ionic Compounds
Which of these compounds are soluble in water? Use Table 5.9 as your guide. a. b. c. d.
ammonium nitrate, NH4NO3, a component of fertilizers sodium sulfate, Na2SO4, an additive in laundry detergents mercury(II) sulfide, HgS, known as the mineral cinnabar aluminum hydroxide, Al(OH)3, used in water purification processes
Answer a. Soluble. All ammonium compounds and all nitrate compounds are soluble.
The landmasses on Earth are largely composed of minerals, that is, ionic compounds. Most have extremely low solubility in water as we mentioned earlier. Table 5.10 summarizes some environmental consequences of solubility.
Table 5.9
Water Solubility of Ionic Compounds
Ions
Solubility of Compounds
Solubility Exceptions
Examples
sodium, potassium, and ammonium
all soluble
none
NaNO3 and KBr. Both are soluble.
nitrates
all soluble
none
LiNO3 and Mg(NO3)2. Both are soluble.
chlorides
most soluble
silver and some mercury chlorides
MgCl2 is soluble. AgCl is insoluble.
sulfates
most soluble
strontium, barium, and lead sulfate
K 2SO4 is soluble. BaSO4 is insoluble.
NH41
carbonates
mostly insoluble*
Group 1A and soluble.
hydroxides and sulfides
mostly insoluble*
Group 1A and NH41 hydroxides and sulfides are soluble.
carbonates are
Na2CO3 is soluble. CaCO3 is insoluble. KOH is soluble. Al(OH)3 is insoluble.
*Insoluble means that the compounds have extremely low solubilities in water (less than 0.01 M). All compounds have at least a very small solubility in water.
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Table 5.10
223
Environmental Consequences of Solubility
Source
Ions
Solubility and Consequences
salt deposits
sodium and potassium halides*
These salts are soluble. Over time, they dissolve and wash into the sea. Thus, oceans are salty and sea water cannot be used for drinking without expensive purification.
agricultural fertilizers
nitrates
All nitrates are soluble. The runoff from fertilized fields carries nitrates into surface and groundwater. Nitrates can be toxic, especially for infants.
metal ores
sulfides and oxides
Most sulfides and oxides are insoluble. Minerals containing iron, copper, and zinc are often sulfides and oxides. If these minerals had been soluble in water, they would have washed out to sea long ago.
mining waste
mercury, lead
Most mercury and lead compounds are insoluble. However, they may leach slowly from mining waste piles and contaminate water supplies.
*Halides are the anions in Group 7A, such as Cl2 and I2.
5.9
|
Covalent Compounds and Their Solutions
From the previous discussion, you might have gotten the impression that only ionic compounds dissolve in water. But remember that sugar dissolves in water as well. The white granules of “table sugar” that you use to sweeten your coffee or tea are sucrose, a polar covalent compound with the chemical formula C12H22O11 (Figure 5.22). When sucrose dissolves in water, the sucrose molecules disperse uniformly among the H2O molecules. The sucrose molecules remain intact and do not separate into ions. Evidence for this includes the fact that aqueous sucrose solutions do not conduct electricity (see Figure 5.17b). However, the sugar molecules do interact with the water molecules, as they are both polar and are attracted to one another. Furthermore, the sucrose molecule contains eight –OH groups and three additional O atoms that can participate in hydrogen bonding (see Figure 5.22). Solubility is always promoted when an attraction exists between the solvent molecules and the solute molecules or ions. This suggests a general solubility rule: Like dissolves like. Let’s also consider two other familiar polar covalent compounds, both of which are highly soluble in water. One is ethylene glycol, the main ingredient in antifreeze; and the other is ethanol, or ethyl alcohol, found in beer and wine. These molecules both contain the polar –OH group and are classified as alcohols (Figure 5.23).
H C OH
CH2OH C O H H OH C C H OH
H C
HOCH2 O C H HO O C C OH H
H C CH2OH
Figure 5.22 Structural formula of sucrose. The –OH groups are shown in red.
Your Turn 5.21
Alcohols and Hydrogen Bonds
Some of the H and O atoms in the ethanol and the ethylene glycol molecule (see Figure 5.23) bear partial charges. Label these d1 and d2, respectively. Hint: If the electronegativity difference between two atoms is more than 1.0, the bond is considered polar.
Look for more about sucrose and other sugars in Chapter 11.
“Like dissolves like” is a useful generalization for solubility.
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H
H
H
C
C
H
H
O
ethanol
H
H
O
H
H
C
C
H
H
O
H
ethylene glycol
Figure 5.23 Lewis structures of ethanol and ethylene glycol. The –OH groups are shown in red.
In connection with indoor air quality, Chapter 1 mentioned propylene glycol, a “glycol” used as an antifreeze in paints.
The H in the –OH group can hydrogen bond, just as was the case for water (Figure 5.24). This is why water and ethanol have a great affinity for each other. Any bartender can tell you that alcohol and water form solutions in all proportions. Again, both molecules are polar and like dissolves like. Ethylene glycol is another example of an alcohol, sometimes called a “glycol.” Ethylene glycol is added to water, such as the water in the radiator of your car, to keep it from freezing. It also is one of the VOCs that some water-based paints emit when drying, an additive to keep the paint from freezing. Examine its structural formula in Figure 5.23 to see that it has two –OH groups available for hydrogen bonding. These intermolecular attractions give high water-solubility to ethylene glycol, a necessary property for any antifreeze. It often has been observed that “oil and water don’t mix.” Water molecules are polar, and the hydrocarbon molecules in oil are nonpolar. When in contact, water molecules tend to attract to other water molecules; in contrast, hydrocarbon molecules stick with their own. Since oil is less dense than water, oil slicks float on top (Figure 5.25).
Your Turn 5.22
More About Hydrocarbons
Hydrocarbon molecules such as pentane and hexane contain C–H and C–C bonds. Use the electronegativity values in Table 5.1 to determine whether these bonds are polar or nonpolar.
Since water is a poor solvent for grease and oil, we cannot use water to wash these off. Instead, we wash our hands (and clothes) with the aid of soaps and detergents. These compounds are surfactants, compounds that help polar and nonpolar compounds to mix, sometimes called “wetting agents.” The molecules of surfactants contain both polar and nonpolar groups. The polar groups allow the surfactant to dissolve in water while the nonpolar ones are able to dissolve in the grease.
H H H O
H
C C O
H
H H
H
O
H O
H
H
H
covalent bond hydrogen bond
Figure 5.24 Hydrogen bonding between an ethanol molecule and three water molecules.
Figure 5.25 Oil and water do not dissolve in each other.
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Another way to dissolve nonpolar molecules is to use nonpolar solvents. Like dissolves like! Worldwide, the production of nonpolar solvents (sometimes called “organic solvents”) is estimated to be 15 billion kilograms. These solvents are widely used, including in the production of drugs, plastics, paints, cosmetics, and cleaning agents. For example, dry cleaning solvents typically are chlorinated hydrocarbons. One example, “perc,” is a cousin of ethene. Take ethene (sometimes called ethylene), a compound with a C5C double bond, and replace all the H atoms with Cl atoms. The result is tetrachloroethylene, also called perchloroethylene. Its nickname is perc. H
H C
Cl
C
H
ethylene
Cl
Look for more about ethylene (ethene) in Chapter 9 on polymers.
Cl C
H
225
C Cl
tetrachloroethylene (‘‘perc’’)
Perc and other chlorinated hydrocarbons like it are carcinogens or suspected carcinogens. They have serious health consequences whether we are exposed to them in the workplace or as contaminants of our air, water, or soil. Green Gree Gr eenn chemists chem ch emis ists ts aim aim to to re rede redesign desi sign g processes gn pro roce cess sses es so so that that they the theyy don’t don’ do n t require requ re q ir qu iree solvents. solv so lven ents ts. But But if thi this thi hiss is is nnot ott ppossible, ossi sibl ibl blee, th tthey hey try hey try to to rep rreplace ep place lace harmful har harmffull solvents sol olve lvent nts ts like likke li ke perc per ercc with wiith ones one oness that thatt are are friendly frie fr iend ie ndly nd ly to to th thee en envi environment. viro vi ronm ro nmen nm entt. One en One possibility pos pos ossi sibi si bili bi lity li ty is is liquid liqu li quid qu id car ccarbon arbo ar bonn dioxide. bo diox di oxid ox idee. Under id Und Und nder er conditions con con ondi diti di tion ti onss on of high pressure, the gas you know as CO2 can condense to form a liquid! Compared with organic solvents, CO2(l) offers many advantages. It is nontoxic, nonflammable, chemically benign, non-ozone-depleting, and it does not contribute to the formation of smog. Although you may be concerned with the fact that it is a greenhouse gas, carbon dioxide that is used as a solvent is a recovered waste product from industrial processes and it is generally recycled. Adapting liquid CO2 to dry cleaning posed a challenge, as it is not very good at dissolving oils, waxes, and greases found in soiled fabrics. To make carbon dioxide a better solvent, Dr. Joe DeSimone, a chemist and chemical engineer at the University of North Carolina–Chapel Hill, developed a surfactant to use with CO2(l). For his work, DeSimone received a 1997 Presidential Green Chemistry Challenge Award. His breakthrough process paves the way for designing environmentally benign, inexpensive, and easily recyclable replacements for conventional organic and water solvents currently in use. DeSimone was instrumental in the beginnings of Hangers Cleaners, a dry cleaning chain that uses the process that he developed.
Consider This 5.23
Liquid CO2 as a Solvent
a. Which of the six principles of Green Chemistry (see inside the front cover) are met by the use of liquid carbon dioxide as a solvent to replace organic solvents? Explain. b. Comment on this statement: “Using carbon dioxide as a replacement for organic solvents simply replaces one set of environmental problems with another.” c. If a local dry cleaning business switched from “perc” to carbon dioxide, how might this business report a different Triple Bottom Line?
The tendency of nonpolar compounds to dissolve in other nonpolar substances explains how fish and animals accumulate nonpolar substances such as PCBs (polychlorinated biphenyls) or the pesticide DDT (dichlorodiphenyltrichloroethane) in their fatty tissues. When fish ingest these, the molecules are stored in body fat (nonpolar) rather than in the blood (polar). PCBs can interfere with the normal growth and development of a variety of animals, including humans, in some cases at concentrations of trillionths of a gram per liter. The higher you go on the food chain, the greater concentrations of harmful nonpolar compounds like DDT you fi nd. This is called biomagnification, the increase in concentration of certain persistent chemicals in successively higher levels
PCBs (mixtures of highly chlorinated compounds) were widely used as coolants in electrical transformers until banned in 1977. Like CFCs, they do not burn easily. They were released into the environment during manufacture, use, and disposal.
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DDT in fish-eating birds 25 ppm
DDT in large fish 2 ppm
DDT in small fish 0.5 ppm
DDT in zooplankton 0.04 ppm
DDT in water 0.000003 ppm or 0.003 ppb
Figure 5.26 Organisms in the water take up and store DDT. They are eaten by larger creatures that in turn are eaten by still larger ones. Creatures highest on the food chain have the highest concentration of DDT. Source: From William and Mary Ann Cunningham. Environmental Science: A Global Concern, 10th ed., 2008. Reprinted with permission of the McGraw-Hill Companies, Inc.
of a food chain. Figure 5.26 shows a biomagnification process that was studied extensively in the 1960s. At that time, DDT was shown to interfere with the reproduction of peregrine falcons and other predatory birds at the top of their food chain.
5.10
The SDWA does not apply to the 10% of the people in the United States whose water comes from private wells.
|
Protecting Our Drinking Water: Federal Legislation
One way or another, many different substances get added to fresh water. Is this water safe to drink? The answer depends on what is present in the water, how much of it is present, and how much of it you drink in a day. In this section, we address issues relating to water quality. The need to keep public water supplies safe to drink has long been recognized. In 1974, the U.S. Congress passed the Safe Drinking Water Act (SDWA) in response to public concern about harmful substances in the water supply. The aim of this act, as amended in 1996, was to ensure potable water to those who depend on community water supplies. As required by the SDWA, contaminants that may be health risks are regulated by the
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Environmental Protection Agency (EPA). The EPA sets legal limits for contaminants according to their toxicities (Table 5.11). These limits also take into account the practical realities that water utilities face in trying to remove the contaminants with available technology. For each water-soluble contaminant, the EPA established a maximum contaminant level goal (MCLG), the maximum level of a contaminant in drinking water at which no known or anticipated adverse effect on human health would occur. Expressed in parts per million or parts per billion, this level has no ill effects for a person weighing 70 kilograms (154 pounds) who consumes 2 liters of water every day for 70 years. Each MCLG allows for uncertainties in data collection and for how different people might react to each contaminant. An MCLG is not a legal limit to which water systems must comply; rather, it is a goal based on human health considerations. For known carcinogens, the EPA has set the health goal at zero under the assumption that any exposure presents a cancer risk. Before regulatory action can be taken against a water utility, the concentration of an impurity must exceed the maximum contaminant level (MCL), the legal limit for the concentration of a contaminant expressed in parts per million or parts per billion. The EPA sets legal limits for each impurity as close to the MCLG as possible, keeping in mind any practical realities that may make it difficult to achieve the goals. Except for contaminants regulated as carcinogens (for which the MCLG is set at zero), most legal limits and health goals are the same. Even when less strict than the MCLGs, the MCLs still provide substantial public health protection.
Consider This 5.24
What’s in Drinking Water?
Table 5.11 is merely a starting point for information available about contaminants in drinking water. The EPA Office of Ground Water and Drinking Water offers a consumer fact sheet on dozens of contaminants. Both general summaries and technical fact sheets are available; the latter is recommended. a. Select a contaminant listed in Table 5.11. How does it get into the water supply? b. How would you know if it were in your drinking water? Is your state one of the top states that releases it? The textbook’s website contains useful links as well as hints on how to locate the data.
Consider This 5.25
Understanding MCLGs and MCLs
Most people are unfamiliar with the terms MCLG and MCL from the Safe Drinking Water Act. How would you explain these abbreviations to the general public? Prepare an outline for a presentation. Be prepared to answer questions from the audience, including why MCLs are not set to zero for all carcinogens.
Water legislation continually needs to be updated. In part, this need arises because water chemists keep improving their ability to detect what is in the water. But the need
Table 5.11
MCLGs and MCLs for Drinking Water
Contaminant
MCLG (ppm)
cadmium (Cd21) 31
chromium (Cr ,
CrO422)
21
lead (Pb ) 21
mercury (Hg ) nitrates (NO32)
MCL (ppm)
0.005
0.005
0.1
0.1
0.015
0.002 10
0.002 10
benzene (C6H6)
0.005
trihalomethanes (CHCl3 and others)
0.080
Look for more examples of trihalomethanes in the next section.
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also arises because our knowledge base is growing. MCL limits should be raised or lowered as we learn more about toxicity. Currently, more than 80 contaminants are regulated: ■ ■ ■
■ ■
The chemical symbol Pb comes from the Latin name for lead, plumbum. The word plumbing comes from this Latin word as well and harkens back to the time when most water pipes were made of lead.
Cold water is recommended because some lead compounds, notably PbCl2, are more soluble in hot water than in cold.
metal ions such as Cd21, Cr31, CrO422, Hg21, Cu21, and Pb21 nonmetal ions such as NO32, F2, and various arsenic-containing ions miscellaneous compounds, including pesticides, industrial solvents, and compounds associated with plastics manufacturing radioisotopes, including radon and uranium biological agents, including Cryptosporidium and intestinal viruses
Depending on the particular contaminant, MCLs range from about 10 ppm to less than 1 ppb. Some contaminants interfere with liver or kidney function. Others can affect the nervous system if ingested over a long period at levels consistently above the legal limit (MCL). For example, unlike many contaminants, lead is a cumulative poison. Lead pipes and solder were once commonly used in water distribution systems. When ingested by humans and animals, lead accumulates in bones and the brain, causing severe and permanent neurological problems. Severe exposure in adults causes symptoms such as irritability, sleeplessness, and irrational behavior. Lead is a particular problem for children because Pb21 can be incorporated rapidly into bone along with Ca21. Since children have less bone mass than adults, some Pb21 may remain in the blood where it can damage cells, especially in the brain. Children may suffer mental retardation and hyperactivity as a result of lead exposure, even at relatively low concentrations. Fortunately, very little lead is present in most public water supplies. Amounts exceeding allowable limits are estimated to be present in less than 1% of public water supply systems, and these serve less than 3% of the U.S. population. Most of this lead comes from corrosion of plumbing systems, not from the source water itself. When lead is reported, consumers are advised to take simple steps to minimize exposure, such as letting water run before using and using only cold water for cooking. Both actions minimize the chances of ingesting dissolved Pb21. Until recently, the MCL for Pb21 in drinking water was 15 ppb. In 1992, the EPA converted this to an “action level,” meaning that the EPA will take legal action if 10% of tap water samples exceed 15 ppb. The hazard from lead is so great that the EPA has established an MCLG of 0, even though lead is not a carcinogen.
Your Turn 5.26
Comparing Lead Content
Two samples of drinking water contained lead ion. One had a concentration of 20 ppb; the other a concentration of 0.003 mg/L. a. Which sample has higher concentration of lead ion? Explain. b. How does each sample compare with the current acceptable limit?
In the United States, about 1000 nitrate violations are reported yearly. These violations usually affect a small number of people and only for a short period of time. In contrast, the number of microbial violations is up to 10 times this value.
Whereas contaminants such as lead cause chronic long-term health problems, other substances in drinking water present more immediate and acute effects. For example, in infants, nitrate ion (NO32) may be converted into nitrite ion (NO22), a substance that limits blood’s ability to carry oxygen. Infants who drink formula made from water containing high levels of nitrate ion may experience difficulty breathing and possibly permanent brain damage from lack of oxygen. Although a maximum contaminant level (MCL) is set for nitrate ion in drinking water, this level may be exceeded for a variety of reasons, including fertilizer or manure runoff that gets into well water. Figure 5.27 shows water quality data for nitrate ion in California. As you can see, some water sources exceeded the MCL of 10 ppm. Because nitrate is toxic to infants, monitoring nitrate levels and informing communities of any violations are important. Water can also be contaminated by biological agents such as bacteria, viruses, and protozoa. Examples include Cryptosporidium and Giardia. News media warnings announcing a “boil-water emergency” are typically the result of a “total coliform” violation. Coliforms are a broad class of bacteria that live in the digestive tracts of humans and other animals. Most are harmless. The presence of a high coliform concentration in water
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Nitrate Concentration National Water Quality Assessment Data/ US Geological Survey Data from the National Water Information System
/ / /
Less than one order of magnitude of Maximum Contaminant Level* Less than, but within one order of magnitude of MCL Greater than MCL
Total Acres Irrigated for Agriculture by County In Million Gallons/Day 0–45 45–195 195–315 315–630 Greater than 630 Principal Aquifer Central Valley Groundwater Basins * Nitrate MCL = 10 milligrams/Liter
N W
E S
Figure 5.27 Map showing nitrate concentrations from California domestic groundwater wells and agricultural irrigation. Source: Environmental Waikato, 2000.
usually indicates that the water-treatment or distribution system is allowing fecal contamination to enter the drinking water supply. Diarrhea, cramps, nausea, and vomiting—the symptoms of microbial-related illness—are generally not serious for a healthy adult but can be life-threatening for the very young, the elderly, or those with weakened immune systems.
Consider This 5.27
A Cryptic Microbe
The EPA surface water-treatment rules require systems using surface water or groundwater under the direct influence of surface water to remove or deactivate 99% of Cryptosporidium. a. What is Cryptosporidium and how does it get into drinking water? What are its potential health effects? b. Regarding Cryptosporidium, what is the LT2 Rule (Long Term 2 Surface Water Treatment Rule) and when did it take effect?
In addition to the Safe Drinking Water Act, other federal legislation controls pollution of lakes, rivers, and coastal areas. The Clean Water Act (CWA), passed in 1974 by Congress and amended several times, provided the foundation for reducing surface water pollution. The CWA established limits on the amounts of pollutants that industry can discharge, removing over a billion pounds of toxic pollutants from U.S. waters every year. In keeping with the new trend toward green chemistry, industries are finding ways both to convert waste materials into useful products and to design processes that neither use toxic substances nor harm water quality. Improvements in surface water quality have at least two major beneficial effects. First, they reduce the amount of cleanup needed for public drinking water supplies. And second, they result in a more
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healthful natural environment for aquatic organisms. In turn, more healthy aquatic ecosystems have many indirect benefits for humans.
| Water Treatment
5.11
This section explores the chemistry that takes place at a local drinking water treatment plant. We assume that the plant treats lake water, which is true in many municipalities. For example, if you live in Chicago, nearby Lake Michigan is the source. And if you live in San Francisco, the water comes from a reservoir in the Hetch Hetchy valley, over a hundred miles away (see Figure 5.9). In a typical water treatment plant (Figure 5.28), the first step is to pass the water through a screen that physically removes items such as weeds, sticks, and beverage bottles. The next step is to add aluminum sulfate and calcium hydroxide. Take a moment to review these two chemicals.
Your Turn 5.28
Water Treatment Chemicals
a. Write chemical formulas for these ions: sulfate, hydroxide, calcium, and aluminum. b. What compounds can be formed from these four ions? Write their chemical formulas. c. The hypochlorite ion plays a role in water purification. Write chemical formulas for sodium hypochlorite and calcium hypochlorite.
Aluminum sulfate and calcium hydroxide are flocculating agents; that is, they react in water to form a sticky floc (gel) of aluminum hydroxide. This gel collects suspended clay and dirt particles on its surface. Al2(SO4)3(aq) 3 Ca(OH)2(s)
Chlorine only can kill the microorganisms with which it comes in contact. Chlorine does not kill bacteria or viruses that hide inside particles of silt or clay. This is one reason why particles need to be removed before the chlorination step.
2 Al(OH)3(s) 3 CaSO4(aq)
[5.7]
As the Al(OH)3 gel slowly settles, it carries particles with it that were suspended in the water (see Figure 5.28). Any remaining particles are removed as the water is filtered through charcoal or gravel and then sand. The crucial step comes next—disinfecting the water to kill disease-causing microbes. In the United States, this is most commonly done with chlorine-containing compounds. Chlorination is accomplished by adding chlorine gas (Cl2), sodium hypochlorite Storage
Fluoridation Flocculating agents Chlorination Paddles
Intake pipe
Screens Pump
Flocculator
Settling tank Charcoal, sand filter
Lake
Figure 5.28 Typical municipal water treatment facility.
Pump
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(NaClO), or calcium hypochlorite (Ca(ClO)2). All of these generate the antibacterial agent hypochlorous acid, HClO. A very low concentration of HClO, 0.075 to 0.600 ppm, remains to protect the water against further bacterial contamination as it passes through pipes to the user. Residual chlorine refers to the chlorine-containing chemicals that remain in the water after the chlorination step. These include hypochlorous acid (HClO), the hypochlorite ion (ClO2), and dissolved elemental chlorine (Cl2). Before chlorination, thousands died in epidemics spread via polluted water. In a classic study, John Snow, an English physician, was able to trace a mid-1800s cholera epidemic in London to water contaminated with the excrement of cholera victims. A more recent example occurred in 2007 in war-torn Iraq. After extremists had put chlorine tanks on suicide truck bombs earlier that year, authorities kept tight controls on chlorine. The chlorine killed two dozen people in several attacks, sending up noxious clouds that left hundreds of people panicked and gasping for breath. At one point, a shipment of 100,000 tons of chlorine was held up for a week at the Jordanian border amid fears for its safe passage through Iraq. With the water infrastructure disrupted and the quality of water and sanitation poor, levels of fecal coliform bacteria increased dramatically, resulting in thousands of Iraqis contracting cholera. Even in peacetime when the transportation of chlorine is relatively safe, chlorination has its drawbacks. The taste and odor of residual chlorine can be objectionable and is commonly cited as a reason why people drink bottled water or use filters to remove residual chlorine. A more serious drawback is the reaction of residual chlorine with other substances in the water to form by-products at concentrations that may be toxic. The most widely publicized, trihalomethanes (THMs), are compounds such as CHCl3 (chloroform), CHBr3 (bromoform), CHBrCl2 (bromodichloromethane), and CHBr2Cl (dibromochloromethane) that form from the reaction of chlorine or bromine with organic matter in drinking water. Like HClO, hypobromous acid (HBrO) used to disinfect spa tubs can generate trihalomethanes.
Your Turn 5.29
231
NaClO is present in Clorox and other brands of laundry bleach. Ca(ClO)2 is commonly used to disinfect swimming pools.
HClO is sometimes written as HOCl. The second shows the order in which the atoms are bonded.
THMs at a Glance
a. Draw Lewis structures for any two THM molecules. b. THMs differ from CFCs in their chemical composition. How? c. THMs differ from CFCs in their chemical properties. How? Answer c. CFCs are chemically inert and nontoxic. In contrast, THMs are chemically reactive and quite toxic.
Many European and a few U.S. cities use ozone to disinfect their water supplies. In Chapter 1, we discussed tropospheric O3 as a serious air pollutant. But in water treatment, the toxicity of O3 serves a beneficial purpose. One advantage is that a lower concentration of ozone than chlorine is required to kill bacteria. Furthermore, ozone is more effective than chlorine against water-borne viruses. But ozonation also comes with disadvantages. One is cost. Ozonation only becomes economical for large water-treatment plants. Another is that ozone decomposes quickly and hence does not protect water from possible contamination as it is piped through the municipal distribution system. Consequently, a low dose of chlorine must be added to ozonated water as it leaves the treatment plant. Disinfecting water using ultraviolet (UV) light is gaining in popularity. By UV, we mean UV-C, the high-energy UV radiation that can break down DNA in microorganisms, including bacteria. Disinfection with UV-C is fast, leaves no residual byproducts, and is economical for small installations, including rural homes with unsafe well water. Like ozone, however, UV-C does not protect the water after it leaves the treatment site. Again, a low dose of chlorine must be added. Depending on local needs, one or more additional purification steps may be taken after disinfection at the water-treatment facility. Sometimes the water is sprayed into the air to remove volatile chemicals that create objectionable odors and taste. If little natural
See Chapter 2 for more about UV light. See Chapter 12 for more about DNA.
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In water, sodium fluoride dissolves to form Na1(aq ) and F2(aq ).
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fluoride ion is present in the water supply, some municipalities add fluoride ion (~1 ppm NaF) to protect against tooth decay. Learn more about fluoridation in the next activity.
Consider This 5.30
The Fluoride Ion Controversy
Community water fluoridation is cited as one of 10 great public health achievements of the 20th century by the U.S. Centers for Disease Control and Prevention. In 2005, the American Dental Association celebrated “20 years of community water fluoridation.” But in 2006, the U.S. National Academies published a report claiming that this practice can damage bones and teeth, and that Federal standards may put children at risk. List the arguments on both sides. Your class might use this topic for a lively debate.
Typical BOD values: A pristine river, 1mg/L; effluent from a municipal sewage treatment plant, 20 mg/L; and untreated sewage, 200 mg/L.
Flood waters and agricultural runoff containing nitrates and phosphates disrupt ecosystems in the Mississippi Delta. Chapter 11 (Figure 11.15) shows a stunning photo.
We just jus ustt described desc de scri ribe bedd how how wa wate water terr is tre ttreated reat ated ed before bef befor oree it is is ready read re adyy to drink dri drink nk out out of of the the tap. tapp. ta Buut once B But once w wee tu turn rn oonn th thi this his is ttap tap, ap p, we sta sstart tart ta rt tth the he pprocess he roce ro cess ce ss of of ge ggetting ttiin tt ing th ing the he water wate wa terr dirty te dirt di rtyy again. rt ag gai ain in. We We addd wa ad wast waste stee to the st the water wat wat ater er eac eeach achh time ac time it it leaves leav le aves av es our our bathrooms bat bat athr hroo hr ooms oo ms in in a to toil toilet ilet il et flu flush ush, ush sh, runs runs down dow dow ownn the drain after a soapy shower, or goes down the sink after we wash the dishes. Clearly it makes sense to use as little water as possible because if we dirty it, it has to be cleaned again before it is released back to the environment. Remember Green Chemistry! It is better to prevent waste than to treat or clean up waste after it is formed. How do we clean sewage and other wastes from water? If the drains in your home are connected to a municipal sewage system, then the wastewater flows to a sewage treatment plant. Once there, it undergoes similar cleaning processes to those for water treatment, with the exception of end-stage chlorination before it is released back to the environment. Cleaning sewage is more complicated, though, because it contains waste in the form of organic compounds and nitrate ions. To many aquatic organisms, this waste is a source of food! As these organisms feed, they deplete oxygen from surface waters. Biological oxygen demand (BOD) is a measure of the amount of dissolved oxygen microorganisms use up as they decompose organic wastes found in water. A low BOD is one indicator of good quality water, whereas a high BOD indicates polluted water. Nitrates and phosphates both contribute to BOD, as these ions are important nutrients for aquatic life. An overabundance of either can disrupt the normal flow of nutrients and lead to algal blooms that clog waterways and deplete oxygen from the water. In turn, this reduced oxygen can lead to massive fish kills. The problem of reduced oxygen in water is compounded by the fact that the solubility of oxygen in water is so very low in the first place. Some treatment plants are using wetland areas to capture nutrients such as nitrates and phosphates before the water is returned to the surface water or recharges the groundwater. Plants and soil microorganisms in these wetland areas (marshes and bogs) facilitate nutrient recycling, thus reducing the nutrient load in the water. If treated sewage water is clean enough, why not just use it as a source of drinking water? Los Angeles, California, has been considering doing this very thing. Los Angeles has an arid climate, and the Colorado River cannot supply enough water to meet the current needs of those who live there. Water conservation helps, but Los Angeles still is challenged by not having enough water. The next activity can help you to explore the issues that the citizens of metropolitan areas face.
Consider This 5.31
Uses for Treated Sewage Water
Use resources of the textbook’s website to answer these questions. a. Name the pros and cons of using treated sewage in agricultural practices. b. If the quality of the water produced from the sewage treatment process matched the quality of the water in our current drinking water system, would you accept treated sewage water as drinking water? Comment either way.
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| Water Solutions for Global Challenges
As we saw at the opening of this chapter, 2005 began the United Nations International Decade for Action called Water for Life. These words frame the U.N. call to action: “Water is crucial for sustainable development, including the preservation of our natural environment and the alleviation of poverty and hunger. Water is indispensable for human health and well-being.” In this final section, we showcase three efforts that demonstrate how water connects to sustainable development. The first relates to the production of fresh water from salt water. The second describes how individuals in developing nations can purify their own drinking water. And the third relates to industry—a green chemistry solution for cotton production.
Fresh Water from Salt Water “Water, water everywhere, nor any drop to drink.” These words from The Rime of the Ancient Mariner are as true today as they were in 1798 when written by Samuel Taylor Coleridge. The high salt content (3.5%) of sea water makes it unfit for human consumption. Neither the ancient mariner nor any who live by the seacoast today can survive drinking sea water. Today, in certain parts of the world where fresh water is not available, we are able to tap the sea as a source of water for both agriculture and drinking. Desalination is any process that removes ions from salt water, thus producing potable water. According to the International Desalination Association, in 2008, over 13,000 desalination plants worldwide together produced over 60 billion L/day of water. The world’s largest single desalination plant is in the United Arab Emirates and produces over 450 million L/day (Figure 5.29). Besides the Middle East, desalination facilities have been built in Europe, Australia, and North America. One means of desalination is distillation, a separation process in which a liquid solution is heated and the vapors are condensed and collected. Impure water is heated. As the water vaporizes, it leaves behind most of its dissolved impurities. Distillation requires energy! Figure 5.30 shows this energy being provided by a Bunsen burner in one case and by the Sun in the other. Recall from Section 5.2 that water has an unusually high specific heat and requires an unusually large amount of energy to convert to a vapor. Both result from the extensive hydrogen bonding in water. The so-called distilled
Figure 5.29 Desalination plant (right) at Jebel Ali in the United Arab Emirates.
Recall from Section 5.3 that the average U.S. citizen uses about 390 L water/day.
A process similar to distillation takes place in the natural water cycle. Water evaporates, condenses, and then falls as rain or snow.
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Water out Thermometer
Condenser Distillation flask
Salt water Cooling water in Distilled water
(a)
(b)
Figure 5.30 (a) Laboratory distillation apparatus. (b) Table-top solar still.
water that you use in introductory laboratory experiments most likely was “deionized” rather than distilled. Again, distillation requires too much energy. Large-scale distillation operations used today often involve either multiple-effect distillation (MED) or multistage flash (MSF) evaporation. These newer technologies have increased energy efficiency over the basic distillation process shown in Figure 5.30a. However, because fossil fuels currently supply the majority of energy necessary to distill fresh water, these new technologies are not sustainable. An alternative to using fossil fuels to heat water is to use smaller solar distillation units (Figure 5.30b).
Consider This 5.32
Other Forms of Distillation
Both MED and MSF evaporation separate salt from water by changing the pressure of heated water as it passes through different compartments. List two ways in which these newer technologies are improvements over the basic process of distillation. The textbook’s website provides some helpful links.
Reverse osmosis is another desalination technique. Osmosis is the passage of water through a semipermeable membrane from a solution that is less concentrated to a solution that is more concentrated. The water diffuses through the membrane and the soluble does not. This is why the membrane is called “semipermeable.” However, with an input of energy, osmosis can be reversed. Reverse osmosis uses pressure to force the movement of water through a semipermeable membrane from a solution that is more concentrated to a solution that is less concentrated. When using this process to purify water, pressure is applied to the saltwater side, forcing water through the membrane to leave the ions behind (Figure 5.31). Reverse osmosis desalination requires a tremendous amount of energy input to supply the pressure needed to make this process work. Despite this high requirement of energy, almost 60% of desalination plants use reverse osmosis technology due to the high quality of water produced. It can be used to produce some bottled water. Portable units are suitable for use on sailboats (Figure 5.32).
Point-of-Use LifeStraws Imagine that your water source was contaminated with bacteria. A billion people are sickened or die each year due to cholera, diarrhea, typhoid, and other ailments caused by microbes in untreated water. A European company, Vestergaard Frandsen,
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Salt water in High-pressure pump Reverse osmosis chamber Concentrated brine out Membrane Pure water out
Figure 5.31 Water purification by reverse osmosis.
Figure 5.32 has developed the LifeStraw, a type of pipe filter that removes virtually all bacteria, viruses, and parasites from water. Aptly named, this device is used to suck water through filters and a disinfection unit (Figure 5.33). This unit can be used to drink water from a stream, river, or lake. It costs about $3, can purify about 700 liters of water, and lasts about a year. For household use, the LifeStraw Family unit filters a minimum of 18,000 liters of water per family for up to 3 years. The units cost about $25 each and are distributed with the help of charitable organizations. The LifeStraw has limitations, however. It is not a long-term solution to the lack of potable water. In addition, it filters neither heavy metals such as arsenic and lead nor some of the protozoan parasites responsible for diarrhea. However, it does provide an interim stopgap in regions where there is plenty of fresh water contaminated with microbes. Although it is not known how many people have used the LifeStraw, charitable organizations such as the Carter Center purchased 23 million units to be distributed in developing nations.
Skeptical Chemist 5.33
LifeStraw
The company that produces LifeStraw has a set of FAQs on the web. One reads: “Does LifeStraw filter heavy metals like arsenic, iron, and fluoride?” What might the Skeptical Chemist say in response to such a question? The textbook’s website provides a link to the answer provided (which we hope isn’t what you would provide).
(a)
(b)
Figure 5.33 (a) Boys using personal LifeStraw to drink. (b) The portable LifeStraw unit is easy to carry.
A small reverse osmosis apparatus for converting sea water to potable water.
The LifeStraw disinfects using iodine and possibly other halogens as well.
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The Cotton Industry Recall the water requirement for a cotton T-shirt, listed in Table 5.2.
A cotton T-shirt, anybody? Or perhaps pants, socks, or a baseball cap? Another solution has targeted a problem associated with cotton clothing. Globally, over 40 billion pounds of this very popular material is produced each year. Even with the growing trend toward synthetic fibers like nylon and polyester, cotton still accounts for over half the market share for apparel in the United States. Currently, the United States is the world’s largest exporter of cotton and exports close to 50% of the world’s supply. Cotton serves not only as the fiber for clothing, but also for many other consumer products. Explore these in the next activity.
Consider This 5.34
Water Footprint Content of a T-Shirt
In an earlier section of this chapter, we pointed out that the water footprint of a 250-g cotton T-shirt was 2700 L (see Table 5.3). a. List four ways that water is used in the production of a T-shirt. b. Other than for clothing, cotton has many other uses. Name three. c. For these uses, estimate whether the water footprint is larger, smaller, or roughly the same as for cotton clothing.
Figure 5.34 Raw cotton has a waxy cuticle (outer layer) that must be removed.
Answers b. Cotton uses include curtains, upholstery, netting, gauze bandages, and swabs. c. These all can be expected to have a large water footprint because of the water used both to grow and clean the cotton.
Even though cotton is a natural fiber, its production leaves a significant footprint on the environment. In Section 5.3, we mentioned that cotton was a thirsty crop that drank large amounts of water as it grew. Once produced, raw cotton must be treated before it can be bleached and dyed (Figure 5.34). In particular, the cuticle, or outermost layer of cotton must be removed. The process (“scouring”) requires copious quantities of caustic chemicals, water, and energy. The wastewater produced has a BOD equal to that of raw sewage! In addition, the wastewater is contaminated with significant amounts of caustic chemicals, ones that weaken the cotton fibers. In 2001, Novozymes won a Presidential Green Chemistry Challenge Award for developing an alternative process to remove the waxy cotton cuticle. The milder process (“biopreparation”) uses an enzyme that breaks down the cuticle. As a result, cotton’s environmental footprint improved. The BOD of wastewater dropped by more than 20%, the caustic chemicals were eliminated, and the amount of water, energy, and time required was significantly reduced. Recall one of the principles of Green Chemistry: It is better to use and generate substances that are not toxic.
Your Turn 5.35
Green Chemistry in Action
We just named one of the six principles of Green Chemistry (see inside front cover) that is met by through the Novozyme “biopreparation” process. Which others apply as well?
This is another example of how green chemistry improves the Triple Bottom Line (economic, societal, environmental) and helps to move us toward more sustainable practices. Reduced and less toxic waste benefits the environment. Using less energy and materials reduces the cost. And society gains a more robust cotton fiber and conserves water.
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Consider This 5.36
The Future of Water
a. Select two of the Green Chemistry principles listed on the inside front cover. For each, brainstorm an idea that might help us to use water more efficiently. b. Identify an important global water issue. Suggest two factors that make it important. Name two ways in which people currently are addressing this issue.
Conclusion Like the air we breathe, water is essential to our lives. It bathes our cells, transports nutrients through our bodies, provides most of our body mass, and cools us when it evaporates. Water also is central to our way of life. We drink it, cook with it, clean things in it, use it to irrigate our crops, and manufacture goods with it. As we do these things, we add waste to the water. Although fresh water purifies itself through a cycle of evaporation and condensation, we humans are dirtying water faster than nature can regenerate clean water. Remember the first principle of Green Chemistry: It is better to prevent waste than to treat or clean up waste after it is formed. So catch the rain water and use it on a garden, rather than letting it run off and join the streams of runoff that pick up pollutants. Instead of using the garbage disposal to grind up food wastes, put the scraps in a compost pile and save the tap water. Turn off the faucet when you are brushing your teeth, limit your time in the shower, and fix that dripping faucet and running toilet! You may feel like your efforts are a mere drop in a much larger bucket. Indeed they are. But remember that like the raindrop shown in the winning Earth Day poster below, your efforts are part of the bigger water picture on this planet (Figure 5.35). Although fresh water is a renewable resource, the demands of population growth, rising affluence, and other global issues are amplifying shortages of this essential commodity. If we are to achieve sustainability we must think water! Think water! Your life and the lives of other creatures depend on it.
Figure 5.35 An Earth Day Haiku Poster Winner, 2008.
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Chapter Summary Having studied this chapter you should be able to: ■ Describe how water is linked to life on this planet (introduction) ■ Connect the electronegativity of atoms with the polarity of the bonds formed from these atoms (5.1) ■ Describe hydrogen bonding and relate it to the properties of water (5.2) ■ Compare the densities of ice and water and be able to account for the difference (5.2) ■ Relate the specific heat of water to the roles played by water on the planet (5.2) ■ Discuss the relationship between the properties of water and its molecular structure (5.2) ■ Describe the major ways people use water on our planet (5.3) ■ Discuss how the concept of water footprint shapes our view of water use (5.3) ■ Connect global climate change with the supply and demand of water (5.4) ■ Use concentration units: percent, ppm, ppb, and molarity (5.5) ■ Discuss why water is such an excellent solvent for many (but not all) ionic and covalent compounds (5.5) ■ Relate these terms: cation, anion, and ionic compound (5.6)
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Write the names and chemical formulas for ionic compounds, including those with common polyatomic ions (5.7)
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Describe what occurs when an ionic compound dissolves in water (5.8) Explain why some solutions conduct electricity and others do not (5.9) Describe the role of surfactants as solubility agents (5.9) Explain the saying “like dissolves like” and relate this to biomagnification (5.9) Understand the role of federal legislation in protecting safe drinking water (5.10) Contrast the maximum contaminant level goal (MCLG) and the maximum contaminant level (MCL) established by the EPA to ensure water quality (5.10) Discuss how drinking water can be made safe to drink (5.11) Understand the processes of distillation and reverse osmosis for producing potable water (5.12) Describe how green chemistry and its applications can contribute to clean water (5.9, 5.12) Summarize at least two possible solutions to our global water challenges (5.12)
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Questions Emphasizing Essentials 1. The chapter opens with these words: “Neeru, shouei, maima, aqua. In any language, water is the most abundant compound on the surface of the Earth.” a. Explain the term compound and also why water is not an element. b. Draw the Lewis structure for water and explain why its shape is bent. 2. Today we are creating dirty water faster than nature can clean it for us. a. Name five daily activities that dirty the water. b. Name two ways in which polluting substances naturally are removed from water. c. Name five steps you could take to keep water cleaner in the first place. 3. Life on our planet depends on water. Explain each of these. a. Bodies of water act as heat reservoirs, moderating climate. b. Ice protects ecosystems in lakes because it floats rather than sinks. 4. Why might a water pipe break if left full of water during extended frigid weather?
5. Here are four pairs of atoms. Consult Table 5.1 to answer these questions. N and C S and O
N and H S and F
a. What is the electronegativity difference between the atoms? b. Assume that a single covalent bond forms between each pair of atoms. Which atom attracts the electron pair in the bond more strongly? c. Arrange the bonds in order of increasing polarity. 6. Consider a molecule of ammonia, NH3. a. Draw its Lewis structure. b. Does the NH3 molecule contain polar bonds? Explain. c. Is the NH3 molecule polar? Hint: Consider its geometry. d. Would you predict NH3 to be soluble in water? Explain. 7. In some cases, the boiling point of a substance increases with its molar mass. a. Does this hold true for hydrocarbons? Explain with examples. Hint: See Section 4.4.
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b. Based on the molar masses of H2O, N2, O2, and CO2, which would you expect to have the lowest boiling point? c. Unlike N2, O2, and CO2, water is a liquid at room temperature. Explain. 8. Both methane (CH4) and water are compounds of hydrogen and another nonmetal. a. Give four examples of nonmetals. In general, how do the electronegativity values of nonmetals compare with those of metals? b. How do the electronegativity values of carbon, oxygen, and hydrogen compare? c. Which bond is more polar, the C–H bond or the O–H bond? d. Methane is a gas at room temperature, but water is a liquid. Explain. 9. This diagram represents two water molecules in a liquid state. What kind of bonding force does the arrow indicate? Is this an intermolecular or intramolecular force?
hydrogen atom oxygen atom
10. The density of water at 0 °C is 0.9987 g/cm3; the density of ice at this same temperature is 0.917 g/cm3. a. Calculate the volume occupied at 0 °C by 100.0 g of liquid water and by 100.0 g of ice. b. Calculate the percentage increase in volume when 100.0 g of water freezes at 0 °C. 11. Consider these liquids. Liquid dishwashing detergent maple syrup vegetable oil
Density, g/mL 1.03 1.37 0.91
a. If you pour equal volumes of these three liquids into a 250-mL graduated cylinder, in what order should you add the liquids to create three separate layers? Explain. b. Predict what would happen if a volume of water equal to the other liquids were poured into the cylinder in part a and the contents then were mixed vigorously. 12. Let’s say the water in a 500-L drum represents the world’s total supply. How many liters would be suitable for drinking? Hint: See Figure 5.8.
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13. Based on your experience, how soluble is each of these substances in water? Use terms such as very soluble, partially soluble, or not soluble. Cite supporting evidence. a. orange juice concentrate b. household ammonia c. chicken fat d. liquid laundry detergent e. chicken broth 14. a. Bottled water consumption was reported to be 29 gallons per person in the United States in 2007. The 2000 U.S. census reported the population as 3.0 3 108 people. Given this, estimate the total bottled water consumption. b. Convert your answer in part a to liters. 15. NaCl is an ionic compound, but SiCl4 is a covalent compound. a. Use Table 5.1 to determine the electronegativity difference between chlorine and sodium, and between chlorine and silicon. b. What correlations can be drawn about the difference in electronegativity between bonded atoms and their tendency to form ionic or covalent bonds? c. How can you explain on the molecular level the conclusion reached in part b? 16. For each of these atoms, draw a Lewis structure. Also draw the Lewis structure for the corresponding ion. Hint: Consult Tables 5.5 and 5.6. a. Cl b. S c. Ne d. Ba e. Li 17. Give the chemical formula and name of the ionic compound that can be formed from each pair of elements. a. Na and Br b. Cd and S c. Ba and Cl d. Al and O e. Rb and I 18. Write the chemical formula for each compound. a. calcium bicarbonate b. calcium carbonate c. magnesium chloride d. magnesium sulfate 19. Name each compound. a. KC2H3O2 b. LiOH c. CoO d. ZnS e. Ca(ClO)2 f. Na2SO4 g. MnCl2 h. K2O 20. Explain why CoCl2 is named cobalt(II) chloride, whereas CaCl2 is named calcium chloride. 21. The MCL for mercury in drinking water is 0.002 mg/L. a. Does this correspond to 2 ppm or 2 ppb mercury? b. Is this mercury in the form of elemental mercury (“quicksilver”) or the mercury ion (Hg21)?
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22. The acceptable limit for nitrate, often found in well water in agricultural areas, is 10 ppm. If a water sample is found to contain 350 mg/L, does it meet the acceptable limit? 23. A student weighs out 5.85 g of NaCl to make a 0.10 M solution. What size volumetric flask does he or she need? Hint: See Figure 5.16. 24. Solutions can be tested for conductivity using this type of apparatus. Bulb Plugged into wall outlet
32. Is there any such thing as “pure” drinking water? Discuss what is implied by this term, and how the meaning of this term might change in different parts of the world. 33. Some vitamins are water-soluble, whereas others are fat-soluble. Would you expect either or both to be polar compounds? Explain. 34. A new sign is posted at the edge of a favorite fishing hole that says “Caution: Fish from this lake may contain over 1.5 ppb Hg.” Explain to a fishing buddy what this unit of concentration means, and why the caution sign should be heeded. 35. This periodic table contains four elements identified by numbers.
Wires 1
Solution being tested
25. 26.
27. 28.
29.
30.
Predict what will happen when each of these dilute solutions is tested for conductivity. Explain your predictions briefly. a. CaCl2(aq) b. C2H5OH(aq) c. H2SO4(aq) An aqueous solution of KCl conducts electricity, but a solution of sucrose does not. Explain. Based on the generalizations in Table 5.9, which compounds are likely to be water-soluble? a. KC2H3O2 b. LiOH c. Ca(NO3)2 d. Na2SO4 For a 2.5 M solution of Mg(NO3)2, what is the concentration of each ion present? Explain how you would prepare these solutions using powdered reagents and any necessary glassware. a. Two liters of 1.50 M KOH b. One liter of 0.050 M NaBr c. 0.10 L of 1.2 M Mg(OH)2 a. A 5-minute shower requires about 90 L of water. How much water would you save for each minute that you shorten your shower? b. Running the water while you brush your teeth can consume another liter. How much water can you save in a week by turning it off? Rank these in order of water footprint: producing the meat for a hamburger, growing an orange, growing a potato, and producing a pint of beer. Explain your reasoning.
Concentrating on Concepts 31. Explain why water is often called the universal solvent.
2 4
3
a. Based on trends within the periodic table, which of the four elements would you expect to have the highest electronegativity value? Explain. b. Based on trends within the periodic table, rank the other three elements in order of decreasing electronegativity values. Explain your ranking. 36. A diatomic molecule XY that contains a polar bond must be a polar molecule. However, a triatomic molecule XY2 that contains a polar bond does not necessarily form a polar molecule. Use some examples of real molecules to help explain this difference. 37. Imagine you are at the molecular level, watching water vapor condense.
a. Sketch four water molecules using a space-filling representation similar to this one. Sketch them in the gaseous state and then in the liquid state. How does the collection of molecules change when water vapor condenses to a liquid? b. What happens at the molecular level when water changes from a liquid to a solid? 38. Propose an explanation for the fact that NH3, like H2O, has an unexpectedly high specific heat. Hint: See Figure 3.12 for the Lewis structure and bond geometry in NH3. 39. a. What type of bond holds together the two hydrogen atoms in the hydrogen molecule, H2? b. Explain why the term hydrogen bonding does not apply to the bond within H2. 40. Consider ethanol, an alcohol with the chemical formula of C2H5OH. a. Draw the Lewis structure for ethanol. b. A cube of solid ethanol sinks rather than floats in liquid ethanol. Explain this behavior.
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41. The unusually high specific heat of water helps keep our body temperature within a normal range despite age, activity, and environmental factors. Consider some of the ways the body produces and loses heat. How would these differ if water had a low specific heat? 42. Health goals for contaminants in drinking water are expressed as MCLG, or maximum contaminant level goals. Legal limits are given as MCL, or maximum contaminant levels. How are MCLG and MCL related for a given contaminant? 43. Some areas have a higher than normal amount of THMs (trihalomethanes) in the drinking water. Suppose that you are considering moving to such an area. Write a letter to the local water district asking relevant questions about the drinking water. 44. Infants are highly susceptible to elevated nitrate levels because bacteria in their digestive tract convert nitrate ion into nitrite ion, a much more toxic substance. a. Give chemical formulas for both the nitrate ion and nitrite ion. b. Nitrite ion can interfere with the ability of blood to carry oxygen. Explain the role of oxygen in respiration. Hint: Review Sections 1.1 and 3.5 for more about respiration. c. Boiling nitrate-containing water will not remove nitrate ion. Explain. 45. Water quality in a chemistry building on campus was continuously monitored because testing indicated water from drinking fountains in the building had dissolved lead levels above those established by the Safe Drinking Water Act. a. What is the likely major source of the lead in the drinking water? b. Do the research activities carried out in this chemistry building account for the elevated lead levels found in the drinking water? Explain. 46. Explain why desalination techniques, despite proven technological effectiveness, are not used more widely to produce potable drinking water. Exploring Extensions 47. In 2001, four countries in South America reached a historic agreement to share the immense Guarani Aquifer. This is significant because underground sources are routinely not considered by international law. a. Where is the Guarani Aquifer, and which four countries joined in this international “underground concordat”? b. What concerns did each country have about this aquifer that led to the agreement? 48. Liquid CO2 has been used successfully for many years to decaffeinate coffee. Explain how and why this works. 49. How can you purify your water when you are hiking? Name two or three possibilities. Compare these
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methods in terms of cost and effectiveness. Are any of these methods similar to those used to purify municipal water supplies? Explain. 50. Hydrogen bonds vary in strength from about 4 to 40 kJ/mol. Given that the hydrogen bonds between water molecules are at the high end of this range, how does the strength of a hydrogen bond between water molecules compare with the strength of a H–O covalent bond within a water molecule? Do your values bear out the assertion made in Section 5.2 that hydrogen bonds are about one tenth as strong as covalent bonds? Hint: Consult Table 4.4 for covalent bond energies. 51. Levels of naturally occurring mercury in surface water are usually less than 0.5 mg/L. The average intake of mercury from food is 2–20 mg daily, but may be much higher in regions where fish is a staple of the diet. a. Name three human activities that add Hg21 (“inorganic mercury”) to water. b. What is “organic mercury”? This chemical form of mercury tends to accumulate in the fatty tissues of fish. Explain why. 52. We all have the amino acid glycine in our bodies. Here is the structural formula.
H
H
O
N
C
C
H
H
O
H
a. Is glycine a polar or nonpolar molecule? Explain. b. Can glycine exhibit hydrogen bonding? Explain. c. Is glycine soluble in water? Explain. 53. Hard water may contain Mg21 and Ca21 ions. The process of water softening involves removing these ions. a. How hard is the water in your local area? One way to answer this question is to determine the number of water-softening companies in your area. Use the resources of the web, as well as ads in your local newspapers and yellow pages, to find out if your area is targeted for marketing water-softening devices. b. If you chose to treat your hard water, what are the options? 54. Suppose that you are in charge of regulating an industry in your area that manufactures agricultural pesticides. How will you decide if this plant is obeying necessary environmental controls? Which criteria affect the success of this plant? 55. Before the U.S. EPA banned their manufacture in 1979, PCBs were regarded as useful chemicals. What properties made them desirable? Besides being persistent in the environment, they bioaccumulate in the fatty tissues of animals. Use the electronegativity concept to show why PCB molecules are nonpolar and thus fat-soluble.
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Neutralizing the Threat of Acid Rain
“I tell my son, go to see the corals now because soon it will be too late.” James Orr, Laboratory of Sciences of the Climate and Environment, France. Source: New Scientist, August 5, 2006. Ocean Acidification: the Other CO2 Problem.
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Coral reefs are massive structures found in shallow ocean waters. Often called “the rainforests of the ocean,” reefs provide habitat for countless species of marine creatures. Attached to a reef or in the water nearby you may find sponges, mollusks, clams, crab, shrimp, sea urchins, sea worms, jellyfish, and many species of fish. Of benefit to humans, reefs protect fragile coastlines from powerful ocean waves. In some countries, they single-handedly support the tourism industry. Coral reefs are alive. With lifetimes of hundreds of thousands of years, reefs can grow thousands of kilometers long, millimeter by millimeter. The tiny animals that form the reef grow and develop using nutrients and other essential chemical compounds dissolved in the ocean water. Why does James Orr want his son to go see the corals? Simply put, today’s reefs may not exist tomorrow. Scientists estimate that at least a quarter of the reefs are now lost; very few are pristine. As energy use worldwide has climbed, changes in both the temperature and chemical composition of our atmosphere have occurred. In turn, changes have occurred in our oceans. At best, these changes slow the growth of coral reefs; at worst they damage the ecosystems established within the reefs. Here we are witness to another example of the tragedy of the commons. As you learned in earlier chapters, we burn fuels to harness their energy. Emissions from combustion include carbon dioxide, nitrogen oxides, and sulfur dioxide. These gases, especially the oxides of nitrogen and sulfur, are soluble in water, including the salt water of our oceans. They dissolve to produce acids; as a result, the water becomes more acidic. For example, over time oceans absorb 25–40% of the carbon dioxide that is emitted as a result of human activities. As more carbon dioxide is emitted, more carbon dioxide dissolves in the oceans. The resulting changes in the seawater have significant effects on the ocean ecosystems. For example, the increase in acidity leads to a reduction in the amount of carbonate ion available to build and maintain coral reefs. Not only do the emissions from combustion dissolve in the oceans, but they also dissolve in water anywhere on the planet, including the rain, snow, and mist of our atmosphere. For example, when SO2 and NOx dissolve in rainwater, they fall back to the Earth in the form of acid rain. Just as added acidity damages ocean ecosystems, it damages the ecosystems of rivers and lakes as well. In Chapter 5 you learned about the special properties of water, what water is used for, how it tends to get dirty, and how we clean it. In this chapter, you will learn how certain compounds dissolve in water to produce acidic and basic solutions. In the right context, acids and bases are extremely useful compounds. We depend on them in many agricultural and manufacturing processes. Acids also impart flavors to the foods we eat. However, in the wrong place at the wrong time, acids can have devastating effects. Ocean acidification and acid rain are two examples. This chapter tells the story of both with an eye to helping you understand what is happening and why. But this story would not be complete without a discussion of bases and of pH, so we touch on these topics as well. Since people tend to be more familiar with acids than with bases, we begin with a discussion of acids and their properties. In what contexts have you already encountered acids? Before you read about acids in the next section, take a moment to do this activity.
Your Turn 6.1
Over time, healthy reefs repair themselves. So in part, the question becomes how healthy the reefs were to begin with.
The tragedy of the commons was discussed in Chapter 1.
Ocean acidification is one of several factors that harm coral reefs. Physical damage occurs with storms and warming ocean temperature; chemical damage through waste products dumped in the ocean.
Recall from Chapter 5 that water is a polar compound. Nonpolar compounds such as CO2 dissolve in water only to a small extent. In contrast, SO2 is a polar compound and very soluble.
Recall from Chapter 1 that NOx is a shorthand notation for NO and NO2.
Acids You Have Encountered
a. List the names for any three compounds that are acids. b. In what context do you know of these acids? For example, is it from reading ingredient labels on foods? Did you run across these acids in some sport or activity? Did you read about these acids in a news article?
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6.1
Figure 6.1 Citrus fruit contains both citric acid and ascorbic acid.
The plant dye litmus changes from blue to pink in acid. The term litmus test has also come to refer to something that quickly reveals a politician’s point of view.
An antacid tablet dissolved in water generates a “fizz” that results from the carbonate in the tablet reacting with citric acid, also present in the tablet.
The notation (aq) is short for aqueous. Revisit equations 5.4 and 5.5 that show the formation of ions as a solute dissolves in water.
| What Is an Acid?
We can approach acids either by listing their observable properties or by describing their behavior at the molecular level. Either way, the information is useful to our discussion, so we employ both approaches. Historically, chemists identified acids by properties such as their sour taste. Although tasting is not a smart way to identify chemicals, you undoubtedly know the sour taste of acetic acid in vinegar. The sour taste of lemons comes from acids as well (Figure 6.1). Acids also show a characteristic color change with indicators such as litmus. Another way to identify an acid is by its chemical properties. For example, under certain conditions acids can react with marble, eggshell, or the shells of marine creatures, causing them to dissolve. These materials all contain the carbonate ion (CO322) either as calcium carbonate or magnesium carbonate. An acid reacts with a carbonate to produce carbon dioxide. This gas is the “burp” when carbonate-containing stomach antacid tablets react with acids in your stomach. This chemical reaction also explains the dissolution of the skeletons of carbonate-based sea creatures such as coral in acidified oceans, as we will see in a later section. At the molecular level, an acid is a compound that releases hydrogen ions, H1, in aqueous solution. Remember that a hydrogen atom is electrically neutral and consists of one electron and one proton. If the electron is lost, the atom becomes a positively charged ion (H1). Because only a proton remains, sometimes H1 is referred to as a proton. For example, consider hydrogen chloride (HCl), a compound that is a gas at room temperature. Hydrogen chloride is composed of HCl molecules. These dissolve readily in water to produce a solution that we name hydrochloric acid. As the polar HCl molecules dissolve, they become surrounded by polar water molecules. Once dissolved, these molecules break apart into two ions: H1(aq) and Cl2(aq). This equation represents the two steps of the reaction.
HCl(g)
H2O
H(aq) Cl(aq)
HCl(aq) 1
You will see the term proton again in Chapter 8 (proton exchange membranes) and Chapter 10 (the protonated form of drug molecules).
We also could say that HCl dissociates into H and Cl . No HCl molecules remain in solution because they dissociate completely. Hydrochloric acid is a strong acid; that is, an acid that dissociates completely in aqueous solution. There is a slight complication with the definition of acids as substances that release H1 ions (protons) in aqueous solutions. By themselves, H1 ions are much too reactive to exist as such. Rather, they attach to something else, such as water molecules. When dissolved in water, each HCl molecule donates a proton (H1) to an H2O molecule, forming H3O1, a hydronium ion. Here is a representation of the overall reaction. HCl(aq) H2O(l)
Here is the Lewis structure for the hydronium ion:
H H O H
[6.1]
2
H3O(aq) Cl(aq)
[6.2]
The solution represented on the product side in both equations 6.1 and 6.2 is called hydrochloric acid. It has the characteristic properties of an acid because of the presence of H3O1 ions. Chemists often simply write H1 when referring to acids (for example, in equation 6.1), but understand this to mean H3O1 (hydronium ion) in aqueous solutions.
It obeys the octet rule.
Your Turn 6.2
Acidic Solutions
For each of these strong acids dissolved in water, write a chemical equation that shows the release of a hydrogen ion, H1. Hint: Remember to include the charges on the ions. The net charge on both sides of the equation should be the same. a. HI(aq ), hydroiodic acid Answer c. H2SO4(aq)
b. HNO3(aq ), nitric acid
H(aq) HSO4(aq)
c. H2SO4(aq ), sulfuric acid
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Consider This 6.3
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Are All Acids Harmful?
Although the word acid may conjure up all sorts of pictures in your mind, every day you eat or drink various acids. Check the labels of foods or beverages and make a list of the acids you find. Speculate on the purpose of each acid.
Hydrogen chloride is but one of several gases that dissolves in water to produce an acidic solution. Sulfur dioxide and nitrogen dioxide are two others. These two gases are emitted during the combustion of certain fuels (particularly coal) to produce heat and electricity. As we mentioned in the introduction to this chapter, SO2 and NO2 both dissolve in rain and mist. When they do so, they form acids that in turn fall back to the Earth’s surface in rain and snow. The increased acidity of the Earth’s water due to anthropogenic emissions is the central focus of this chapter. But before delving into the acidity in rain caused by nitrogen oxides and sulfur dioxide, let’s focus on carbon dioxide. With an atmospheric concentration of about 390 ppm in 2010, and rising, carbon dioxide is at a far higher concentration than either sulfur dioxide or nitrogen dioxide. Just as solids vary in their solubility in water, so do gases. Compared with more polar compounds such as SO2 and NO2, carbon dioxide is far less soluble in water. Even so, it dissolves to produce a weakly acidic solution. At this stage, you as a Skeptical Chemist should be raising an important question. Given that an acid is defined as a substance that releases hydrogen ions in water, how can carbon dioxide act as an acid? There are no hydrogen atoms in carbon dioxide! The explanation is that when CO2 dissolves in water, it produces carbonic acid, H2CO3(aq). Here is a way to represent the process. CO2(g)
H2O
CO2(aq) H2O(l)
CO2(aq) H2CO3(aq)
Refer back to Chapter 5 to review solubility. In general, “like dissolves like.”
[6.3a] [6.3b]
The carbonic acid dissociates to produce the H1 ion and the hydrogen carbonate ion. H(aq) HCO3(aq)
H2CO3(aq)
[6.3c] 1
This reaction occurs only to a limited extent, producing only tiny amounts of H and HCO32. Accordingly, we say that carbonic acid is a weak acid; that is, an acid that dissociates only to a small extent in aqueous solution. Although carbon dioxide is only slightly soluble in water and then only a tiny amount of the dissolved carbonic acid dissociates to produce H1, these reactions are happening on a large scale across the planet. The carbon dioxide can dissolve in water in the upper atmosphere (making acids that may fall as acidic rain) or in the planet’s oceans, lakes, and streams. We will return to this topic after we introduce bases.
6.2 |
What Is a Base?
No discussion of acids would be complete without discussing their chemical counterparts— bases. For our purposes, a base is a compound that releases hydroxide ions, OH2, in aqueous solution. Aqueous solutions of bases have their own characteristic properties attributable to the presence of OH2(aq). Unlike acids, bases generally taste bitter and do not lend an appealing flavor to foods. Aqueous solutions of bases have a slippery, soapy feel. Common examples of bases include household ammonia (an aqueous solution of NH3) and NaOH, sometimes called lye. The cautions on oven cleaners (Figure 6.2) warn that lye can cause severe damage to eyes, skin, and clothing. Many common bases are compounds containing the hydroxide ion. For example, sodium hydroxide (NaOH), a water-soluble ionic compound, dissolves in water to produce sodium ions (Na1) and hydroxide ions (OH2). NaOH(s)
H2O
Na(aq) OH(aq)
[6.4]
Dilute basic solutions have a soapy feel because bases can react with the oils of your skin to produce a tiny bit of soap.
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Chapter 6
Figure 6.2 Oven cleaning products may contain NaOH, commonly called lye.
Although sodium hydroxide is very soluble in water, most compounds containing the hydroxide ion are not. Table 5.9 summarized the solubility trends for compounds containing specific types of anions. Bases that dissociate completely in water, such as NaOH, are called strong bases.
Your Turn 6.4
Basic Solutions
These solids dissolve in water to release hydroxide ions. For each, write a balanced chemical equation. a. KOH(s), potassium hydroxide b. LiOH(s), lithium hydroxide c. Ca(OH)2(s), calcium hydroxide
In some industrial applications, ammonia (rather than HCFCs) is used as a refrigerant gas. Great care needs to be taken to prevent the exposure of workers to ammonia, as the gas can dissolve in moist lung tissue and injure or kill.
The ammonium ion, NH41, is analogous to the hydronium ion, H3O1, in that each was formed by the addition of a proton (H1) to a neutral compound.
Some bases, however, do not contain the hydroxide ion, OH–, but rather react with water to form it. One example is ammonia, a gas with a distinctive sharp odor. Unlike carbon dioxide, ammonia is very soluble in water. It rapidly dissolves in water to form an aqueous solution. NH3(g)
H2O
NH3(aq)
[6.5a]
On a supermarket shelf, you may see a 5% (by mass) aqueous solution of ammonia with the name of “household ammonia.” This cleaning agent has an unpleasant odor; if it gets on your skin, you should wash it off with plenty of water. The chemical behavior of aqueous ammonia is difficult to simplify, but we will do our best to represent it for you with a chemical equation. Think in terms of an ammonia molecule reacting with a water molecule. In essence, the water molecule transfers an H1 ion to the aqueous NH3 molecule to form an aqueous ammonium ion, NH41(aq), and a hydroxide ion. However, this reaction only occurs to a small extent; that is, only a tiny amount of OH2(aq) is produced. NH3(aq) H2O(l)
only to a small extent
NH4(aq) OH(aq)
[6.5b]
The source of the hydroxide ion in household ammonia now should be apparent. When ammonia dissolves in water, it releases small amounts of the hydroxide ion and the ammonium ion. Aqueous ammonia is an example of a weak base, a base that dissociates only to a small extent in aqueous solution. In order to more clearly indicate that aqueous ammonia is a base, some people use the representation NH4OH(aq). If you add up the atoms (and their charges), you
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will see that NH4OH(aq) is equivalent to the left-hand side of equation 6.5b. It is unlikely, however, that this species exists as such in an aqueous solution of ammonia.
6.3 |
Neutralization: Bases Are Antacids
Acids and bases react with each other—often very rapidly. Not only does this happen in laboratory test tubes, but also in your home and in almost every ecological niche of our planet. For example, if you put lemon juice on fish, an acid–base reaction occurs. The acids found in lemons neutralize the ammonia-like compounds that produce the “fishy smell.” Similarly, if the ammonia fertilizer on a corn field comes in contact with the acidic emissions of a power plant nearby, an acid–base reaction occurs. Let us first examine the acid–base reaction of solutions of hydrochloric acid and sodium hydroxide. When the two are mixed, the products are sodium chloride and water. HCl(aq) NaOH(aq)
NaCl(aq) H2O(l)
[6.6]
This is an example of a neutralization reaction, a chemical reaction in which the hydrogen ions from an acid combine with the hydroxide ions from a base to form water molecules. The formation of water can be represented like this. H(aq) OH(aq)
H2O(l)
Recall from Section 5.8 that NaCl is an ionic compound that dissolves in water to produce Na1(aq) and Cl2(aq).
[6.7]
What about the sodium and chloride ions? Recall from equations 6.1 and 6.4 that the HCl(g) and NaOH(s), when dissolved in water, completely dissociate into ions. We can rewrite equation 6.6 to show this. H(aq) Cl(aq) Na(aq) OH(aq)
Na(aq) Cl(aq) H2O(l) [6.8]
Neither Na1(aq) nor Cl2(aq) take part in the neutralization reaction; they remain unchanged. Canceling these ions from both sides again gives us equation 6.7, which summarizes the chemical changes taking place in an acid–base neutralization reaction.
Your Turn 6.5
Neutralization Reactions
For each acid–base pair, write a neutralization reaction. Then rewrite the equation in ionic form and eliminate ions common to both sides. What is the relevance of the final simplified step in each case? a. HNO3(aq) and KOH(aq) b. H2SO4(aq) and NH4OH(aq) c. HBr(aq) and Ba(OH)2(aq) Answer c. 2 HBr(aq) Ba(OH)2(aq)
BaBr2(aq) 2 H2O(l) 2
2 H (aq) 2 Br (aq) Ba 2 H(aq) 2 OH(aq)
(aq) 2 OH(aq)
Ba2(aq) 2 Br(aq) 2 H2O(l)
2 H2O(l)
Divide by 2 to simplify this last equation. H(aq) OH(aq)
H2O(l)
In each case, the final step summarizes the reaction; that is, it shows that the hydrogen ion from the acid and the hydroxide ion from the base react with each other to form water.
Neutral solutions are neither acidic nor basic; that is, they have equal concentrations of H1 and OH2 ions. Pure water is a neutral solution. Some salt solutions also are neutral, such as the one formed by dissolving solid NaCl in water. In contrast, acidic solutions contain a higher concentration of H1 than OH2 ions, and basic solutions contain a higher concentration of OH2 than H1 ions.
There is no such thing as “pure” water. Remember from Chapter 5 that water always has some kind of impurities in it.
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Chapter 6
It may seem strange that acidic and basic solutions contain both hydroxide ions and hydrogen ions. But when water is involved, it is not possible to have H1 without OH2 (or vice versa). A simple, useful, and very important relationship exists between the concentration of hydrogen ion and hydroxide ion in any aqueous solution. The product [H1][OH2] is dependent on temperature. The value 1 3 10214 is valid at 25 °C.
Refer to Section 5.5 for the definition of molarity.
[H1][OH2] 5 1 3 10214
The square brackets indicate that the ion concentrations are expressed in molarity, and [H1] is read as “the hydrogen ion concentration.” When [H1] and [OH2] are multiplied together, the product is a constant with a value of 1 3 10214 as shown in mathematical equation 6.9. This shows that the concentrations of H1 and OH2 depend on each other. When [H1] increases, [OH2] decreases, and when [H1] decreases, [OH2] increases. Both ions are always present in aqueous solutions. Knowing the concentration of H1, we can use equation 6.9 to calculate the concentration of OH2 (or vice versa). For example, if a rain sample has a H1 concentration of 1 3 1025 M, we can calculate the OH2 concentration by substituting in 1 3 1025 M for [H1]. (1 3 1025) 3 [OH2] 5 1 [OH2] 5 1 1 2 [OH ] 5 1
By definition, the product of the two concentrations is unitless.
[H1] > [OH2] Neutral solution [H1] 5 [OH2] Basic solution 1
2
[H ] < [OH ]
3 3 3 3
10214 10214 1025 1029
Since the hydroxide ion concentration (1 3 1029 M) is smaller than the hydrogen ion concentration (1 3 1025 M), the solution is acidic. In pure water or in a neutral solution, the molarities of the hydrogen and hydroxide ions both equal 1 3 1027 M. Applying mathematical expression 6.9, we can see that [H1][OH2] 5 (1 3 1027)(1 3 1027) 5 1 3 10214.
Your Turn 6.6
Acidic solution
[6.9]
Acidic and Basic Solutions
Classify these solutions as acidic, neutral, or basic. Then, for parts a and c, calculate [OH2]. For b, calculate [H1]. a. [H1] 5 1 3 1024 M
b. [OH2] 5 1 3 1026 M
c. [H1] 5 1 3 10210 M
Answer a. The solution is acidic because [H1] . [OH2]. [H1][OH2] 5 1 3 10214. Solving, [OH2] 5 1 3 10210 M.
Your Turn 6.7
Ions in Acidic and Basic Solutions
Classify each solution as acidic, basic, or neutral. All of the compounds are strong acids or strong bases. Then list all of the ions present in order of decreasing relative amounts in each solution. a. KOH(aq)
b. HNO2(aq)
c. H2SO3(aq)
d. Ca(OH)2(aq)
Answer d. When calcium hydroxide dissociates, two hydroxide ions are released for each calcium ion. The basic solution contains much more OH2 than H1. OH2(aq) . Ca21(aq) . H1(aq)
How can we know if the acidity of seawater and of rain is cause for concern? To make a judgment, we need a convenient way of reporting how acidic or basic a solution is. The pH scale is just such a tool because it relates the acidity of a solution to its H1 concentration.
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6.4 |
Introducing pH
The term pH already may be familiar to you. For example, test kits for soils and for the water in aquariums and swimming pools report the acidity in terms of pH. Deodorants and shampoos claim to be pH-balanced (Figure 6.3). And, of course, articles about acid rain make reference to pH. The notation pH is always written with a small p and a capital H and stands for “power of hydrogen.” In the simplest terms, pH is a number, usually between 0 and 14, that indicates the acidity (or basicity) of a solution. As the midpoint on the scale, pH 7 separates acidic from basic solutions. Solutions with a pH less than 7 are acidic, and those with a pH greater than 7 are basic (alkaline). Solutions of pH 7 (such as pure water) have equal concentrations of H1 and OH– and are said to be neutral. The pH values of common substances are displayed in Figure 6.4. You may be surprised that you eat and drink so many acids. Acids naturally occur in foods and contribute distinctive tastes. For example, the tangy taste of McIntosh apples comes from malic acid. Yogurt gets its sour taste from lactic acid, and cola soft drinks contain several acids, including phosphoric acid. Tomatoes are well known for their acidity, but with a pH of about 4.5, they are in fact less acidic than many other fruits.
Consider This 6.8
Is the water on the planet acidic, basic, or neutral? Water would be expected to have a pH of 7.0, but Figure 6.4 shows that the pH of water depends on where it is found. “Normal” rain is slightly acidic, with a pH value between 5 and 6. Even though the acid formed by dissolved carbon dioxide is a weak acid, enough H1 is produced to lower the pH of rain. Seawater is slightly basic, with a pH of approximately 8.2. As you might have guessed, pH values are related to the hydrogen ion concentration. For solutions in which [H1] is 10 raised to some power, the pH value is this power (the exponent) with its sign changed. For example, if [H1] 5 1 3 1023 M, then the pH is 3. Similarly, for [H1] 5 1 3 1029 M, the pH is 9. Appendix 3 describes the relation between pH and [H1] in more detail. Equation 6.9 shows that the hydrogen ion concentration multiplied by the hydroxide ion concentration is a constant, 1 3 10214. When the concentration of H1 is high (and Stomach Lemon Coca- Tomato Pure Sea- Milk of Household Oven cleaner acid juice Cola juice Milk water Blood water magnesia ammonia (lye)
1
2
3
4
Acid rain and fog
For highly acidic or basic solutions, the pH may lie outside the 0-to-14 range.
Universal indicator paper is a quick way to estimate the pH of a solution. For more accurate results, pH meters are used.
Acidity of Foods
a. Rank tomato juice, lemon juice, milk, cola, and pure water in order of increasing acidity. Check your order against Figure 6.4. b. Pick any other five foods and make a similar ranking. Look up the actual pH values. A helpful link is provided at the textbook’s website.
pH
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5
6
7
8
9
10
11
12
13
14
Normal rain
Figure 6.4 Common substances and their pH values.
Figures Alive! Visit the textbook’s website to learn more about acids, bases, and the pH scale.
Figure 6.3 This shampoo claims to be “pHbalanced,” that is, adjusted to be closer to neutral. Soaps tend to be basic, which can be irritating to the skin.
Carbonated water is more acidic than rainwater, because more CO2 gas is forced under pressure to dissolve. The pH is about 4.7.
The mathematical relationship is pH 5 2log[H1] More information can be found in Appendix 3.
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Chapter 6 [H⫹]
10
10
10
10
10
10
10
10
10
pH
1
2
3
4
5
6
7
8
9
⫺1
⫺2
⫺3
⫺4
⫺5
⫺6
⫺7
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10
10
Acidic
⫺11
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⫺12
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11
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⫺13
10
⫺14
10
13
14
Basic Neutral
Figure 6.5 The relationship between pH and the concentration of H1 in moles per liter (M). As pH increases, [H1] decreases.
the pH is low), the concentration of OH2 is low. Likewise, as pH values rise above 7.0, the concentration of hydrogen ions decreases and the concentration of hydroxide ions increases. As the pH value decreases, the acidity increases. For example, a sample of water with a pH of 5.0 is 10 times less acidic than one with a pH of 4.0. This is because a pH of 4 means that the [H1] is 0.0001 M. By contrast, a solution with a pH of 5 is more dilute, with a [H1] 5 0.00001 M. This second solution is less acidic with only 1/10 the hydrogen ion concentration of a solution of pH 4. Figure 6.5 shows the relationship between pH and the hydrogen ion concentration.
Your Turn 6.9
Small Changes, Big Effects
Compare the pairs of samples below. For each, which one is more acidic? Include the relative difference in hydrogen ion concentration between the two pH values. a. Rain sample, pH 5 5, and lake water sample, pH 5 4. b. Ocean water sample, pH 5 8.3, and tap water sample, pH 5 5.3. c. Tomato juice sample, pH 5 4.5, and milk sample, pH 5 6.5. Answer c. Although the pH values differ only by 2, the tomato juice sample is 100 times more acidic and (for an equal volume) would have 100 times more H1 than the milk sample.
Consider This 6.10
On the Record
A legislator from the Midwest is on record with an impassioned speech in which he argued that the environmental policy of the state should be to bring the pH of rain all the way down to zero. Assume that you are an aide to this legislator. Draft a tactful memo to your boss to save him from additional public embarrassment.
6.5 |
Ocean pH can vary by 60.3 pH units, depending on latitude and region.
Ocean Acidification
How can seawater be basic when rain is naturally acidic? Indeed this is the case, as shown in Figure 6.4. Ocean water contains small amounts of three chemical species that all play roles in maintaining the ocean pH at approximately 8.2. These three species—the carbonate ion, the bicarbonate ion, and carbonic acid—interact with each other as well. All arise from dissolved carbon dioxide (equations 6.3a, b, and c). These species also help maintain your blood at a pH of about 7.4. O
Only one resonance form is shown for the bicarbonate ion and for the carbonate ion. For more about resonance, see Section 2.3.
O
C
O
O O
carbonate ion CO32 (aq)
C HO
O
bicarbonate ion HCO3 (aq)
C HO
OH
carbonic acid H2CO3 (aq)
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Many organisms, such as mollusks, sea urchins, and coral, have connections to this bit of ocean chemistry because they build their shells out of calcium carbonate, CaCO3. Changing the amount of one chemical species in the ocean (such as carbonic acid) can affect the concentration of the others, in turn affecting marine life. Human industrial activity has rapidly increased the amount of carbon dioxide released into the atmosphere over the past 200 years. As a result, more carbon dioxide is dissolving into the oceans and forming carbonic acid. In turn, the pH of seawater has dropped by roughly 0.1 pH unit since the early 1800s. This may sound like a small number. Remember, though, that each full unit of pH represents a 10-fold difference in the concentration of H1 ions. A decrease of 0.1 pH unit corresponds to a 26% increase in the amount of H1 in seawater. The lowering of the ocean pH due to increased atmospheric carbon dioxide is called ocean acidification. How can such a seemingly small change in pH pose a danger to marine organisms? Part of the answer lies in the chemical interactions between CO322(aq), HCO32(aq), and H2CO3(aq). The H1 produced by the dissociation of carbonic acid reacts with carbonate ion in seawater to form the bicarbonate ion. HCO3(aq)
H(aq) CO32(aq)
[6.10]
The net effect is to reduce the concentration of carbonate ion in seawater. The calcium carbonate in the shells of sea creatures begins to dissolve in order to maintain the concentration of carbonate ions in seawater. CaCO3(s)
H2O
Ca2(aq) CO32(aq)
[6.11]
The interaction of carbonic acid, bicarbonate ion, and carbonate ion are summarized in Figure 6.6. As carbon dioxide dissolves in ocean water, it forms carbonic acid. This in turn dissociates to produce “extra” acidity in the chemical form of H1. The H1 reacts with the carbonate ion, depleting it and producing more bicarbonate ion. Calcium carbonate then dissolves to replace the carbonate that was depleted. Ocean scientists predict that within the next 40 years, the carbonate ion concentration will reach a low enough level that the shells of sea creatures near the ocean surface will begin to dissolve. In fact, one study has shown that the Great Barrier Reef off the coast of Australia is already growing at slower and slower rates. However, other factors could be to blame, including ocean warming. One can examine growth rings in a slice of coral, much as one can view tree rings (Figure 6.7). To date, only a small number of researchers have focused on the effects of thinning shells on sea creatures. However, negative effects on whole ecosystems have been
CO2
CO2
H2O
H2CO3
H CO32
CaCO3 (coral)
Figure 6.6 Chemistry of CO2 in the ocean.
H
HCO3
Ca2 CO32
HCO3
251
Recall (Section 3.2) that in recent decades, the concentration of CO2 in the atmosphere has risen a few parts per million each year. The increase has been steady and relentless.
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Figure 6.7 A thin slice of coral. Special lighting reveals annual growth rings. A recent study has shown that some corals have seen a dramatic decrease in their growth rate over the past 20 years.
projected. For example, weaker (or missing) coral reefs could fail to protect coastlines from harsh ocean waves. Coral reefs also provide fish species with their habitat, and damage to the reefs would translate into losses of marine life. Finally, a weakening of the reefs would make them more susceptible to further damage from storms and predators. Can the ocean heal itself? Although we don’t know the answer for sure, nonetheless we can speculate from what we know of past events. When changes in ocean pH have occurred over a very long period of time, the ocean has been able to compensate. This happens because large collections of sediment at the bottom of the ocean contain massive amounts of calcium carbonate, mostly from the shells of long-deceased marine creatures. Over long time periods, these sediments dissolve to replenish the carbonate lost to reaction with excess H1. But today’s changes in the pH of the oceans have happened rapidly on the geologic time scale. In just 200 years, the pH of the ocean has dropped to a level not seen in the past 400 million years. Because the acidification is occurring over a relatively short time and in water close to the surface, the sediment reserve has not had time to dissolve and counteract the effects of the added acidity. Even if the amount of carbon dioxide in the atmosphere were to immediately level off, the oceans would take thousands of years to return to the pH measured in preindustrial times. Coral reefs would take even longer to regenerate, and any species lost to extinction, of course, would not return.
Consider This 6.11
International Response to Ocean Acidification
In October, 2008, a group of scientists met in Monaco to raise awareness about ocean acidification. They issued the Monaco Declaration, calling on the countries of the world to reverse carbon dioxide emissions trends by 2020. Have more recent gatherings of scientists and negotiators created a worldwide policy to address ocean acidification? Do research of your own and summarize your findings.
6.6
|
The Challenges of Measuring the pH of Rain
Ours is a wet planet, as we saw in Chapter 5. Carbon dioxide not only has the opportunity to dissolve in the ocean but also can dissolve in rain and fresh water everywhere. The result is the same—when CO2 dissolves, the pH drops slightly due to the formation
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Figure 6.8 A pH meter with a digital display.
of carbonic acid. In the case of rain, the resulting pH is between 5 and 6. But at times the pH of rain can be even lower than this. We refer to this type of precipitation as acid rain, that is, rain with a pH below 5 that is more acidic than “normal” rain. What are the levels of acidity in rain across the mainland United States, Alaska, Hawaii, and Puerto Rico? To measure acidity levels anywhere in the world, we need an analytical tool, the pH meter. Many types of pH meters are available, depending both on the conditions under which you wish to use them and how much you are willing to pay. The pH meter that you are most likely to encounter has a special probe capped with a membrane that is sensitive to H1. When the probe is immersed in a sample, the difference in H1 ion concentrations between the solution and the probe creates a voltage across the membrane. The meter measures this voltage and converts it to a pH value (Figure 6.8). Although it is straightforward to measure the pH of a rain sample, certain procedures are necessary to ensure accurate results. For example, the electrode of the pH meter needs to be carefully calibrated. Another challenge is to collect and measure rain samples without contaminating them. The collection containers must be scrupulously clean and free of minerals from the water in which they were washed. When a container is placed on site, it must be set high enough to prevent splash contamination either from the ground or surrounding objects. Even if elevated, contamination may still occur from the pollen of nearby plants, insects, bird droppings, leaves, soil dust, or even the ash of a fire. One way to minimize contamination is to fit a rain collection bucket with a lid and a moisture sensor that opens this lid when it begins to rain. This is how samples are collected at the approximately 250 sites of the National Atmospheric Deposition Program/National Trends Network (NADP/NTN). Figure 6.9a shows the sensor and two buckets at a NADP/NTN monitoring station in Illinois that has been in operation for over 30 years. One bucket is for dry deposition (open when it is not raining) and the other is covered. A sensor opens this bucket (closing the other) when it rains. Deciding where to locate the collection sites also is a challenge. Due to budgetary constraints, the test sites cannot go in as many places as researchers might like. The relative advantages of widely dispersing the sites versus putting several near each other in a special ecosystem (such as a national park) need to be weighed. Currently there are more collection sites in the eastern United States, as historically the acidity levels have been higher there due in large part to coal-fired electric power plants. Rain samples have been collected routinely in the United States and Canada since about 1970. Since 1978, NADP/NTN has collected over 250,000 samples, analyzing them for pH and for these ions: SO422, NO32, Cl2, NH41, Ca21, Mg21, K1, and Na1. Figure 6.9b shows the five active NADP/NTN sites in the state of Illinois. To discover how many sites are in your state, complete the following activity.
Section 8.5 describes how hydrogen fuel cells also create a voltage across a membrane.
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IL18 IL19 IL78 IL11
ILLINOIS IL46
IL99
IL47 IL35 IL63
(a)
(b)
Figure 6.9 (a) The Bondville Monitoring Station in central Illinois (IL11) has been in operation since 1979. The black moisture sensor connected to the left of the table controls which bucket is open. It is not raining, so the right bucket for wet deposition is closed. (b) The five active NTN precipitation monitoring sites in Illinois, including IL11 in Bondville. The sites marked with triangles are inactive. Source: National Atmospheric Deposition Program 2009. NADP Program Office, Illinois State Water Survey. http://nadp.sws.uiuc.edu/sites/sitemap.asp?state=il
Consider This 6.12
The Rain in Maine . . . Oregon or Florida
Thanks to the NADP/NTN, almost every state plus Puerto Rico and the Virgin Islands has one or more precipitation monitoring sites. a. In Figure 6.9a, name the precautions you see being taken to preserve the integrity of the rain samples. b. How many monitoring sites are in your state? A map with links is provided at the textbook’s website. c. Do you think the number and placement of collection sites in your state fairly represent the acidic deposition? d. On the web, select a collection site in your state (or in a neighboring one) that provides a photograph. Compare the picture with Figure 6.9a. What additional ways of minimizing contamination (for example, a fence or signage) can you spot, if any? Answer a. The collection buckets are located up off the ground, one is fitted with a lid that opens when it rains, and the area around the site is mowed. Also, the location is far from people and roads.
Each week, researchers at the Central Analytical Laboratory in Champaign, Illinois, receive hundreds of rain samples. The photographs in Figure 6.10 indicate the magnitude of the operation. At the left are sample collection buckets waiting to be cleaned prior to being shipped back to the collection sites. The center photo shows a set of rain samples in the queue to be analyzed, each assigned an alphanumeric label. A small portion of each sample is saved after analysis and stored under refrigeration. The photo on the right shows Chris Lehmann, director of the laboratory, standing inside a cold room in which samples are archived. Each year, researchers at the Central Analytical Laboratory use the analytical data to construct maps such as the one shown in Figure 6.11. From these maps, we observe that all rain is slightly acidic. Remember that rain contains a small amount of dissolved carbon dioxide, and that the carbon dioxide dissolved in rainwater produces a weakly acidic solution.
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Figure 6.10 Photographs from the Central Analytical Laboratory (CAL), Champaign, Illinois. Left: Sample collection buckets waiting to be cleaned. Center: Rain samples in the queue to be analyzed. Right: Chris Lehmann, Director of the CAL.
Recall that rain naturally has a pH between 5 and 6. Figure 6.11 shows that the pH of rain samples is well below normal in the eastern third of the United States, especially in the Ohio River valley. Carbon dioxide is not the only source of H1 in rain. Chemical analysis of rainwater confirms the presence of other substances that result in the formation of hydrogen ion: sulfur dioxide (SO2), sulfur trioxide (SO3), nitrogen monoxide (NO), and nitrogen dioxide (NO2). These compounds are affectionately known as “Nox and Sox.” Chemically, we write SOx and NOx, where x 5 2 and 3 for SOx, and x 5 1 and 2 for NOx. Let’s examine SOx and NOx in turn. First, here is a way to represent the process by which sulfur trioxide dissolves in water to form sulfuric acid. SO3(g) H2O(l)
H2SO4(aq)
[6.12]
Equation 6.12 is analogous to the reaction of CO2 with water.
sulfuric acid
• 5.3 5.3 5.5 • • • 5.3 • 5.3 • 5.5
5.5 •
• 5.2
• 5.6
• 5.4 • 5.7
• 5.5 5.8 • • 5.4
5.4 5.6 • • • 5.6
5.5 •
• 5.7
• 5.6
5.4 •
5.5 5.5• 5.9 • • • 5.8 • • 5.4 5.7 • 6.2 • 5.6 5.3 • 5.4 6.1 • 5.1 •• 5.2 • 5.6 • 5.3 5.4 • • • 5.2 5.4 5.4 5.4 • 5.5 • • • • 6.1 5.6 5.3 5.1 5.9 5.4 • • • 5.2 5.2 5.4 • • • • • 5.6 • 5.5 • 5.7 5.3 • 5.3 5.1 • 5.3 •
• 5.3
5.4 • • 5.4
Sites not pictured: Alaska 5.3 Puerto Rico 5.2 Virgin Islands 5.3
5.4 •
• 4.9 4.9 • • • 4.8 5.9 • 5.0 • 5.0 • • 4.9 • 5.1 4.7 4.7 4.7 4.8 • 4.9 • 5.9 • 5.2 • • • • •5.1 • 4.7 • • 4.8 5.4• • • • 4.7 4.6 5.1 4.8 • 4.6 4.7 • 5.4 • • 5.7 4.6 • 5.3 • 5.0 6.2 • 4.8 • 4.6 • 5.2 • • • 4.6 4.6 • • • 4.6 4.7 4.8 5.6 5.2 5.1 • • 4.8 • 4.5 • • 4.7 • • 4.6 5.1 4.5 4.5 •• • • 4.6 4.8 4.6 4.7 • • 4.5 4.6 5.6 4.9 • 5.5 • • • • 4.5 • • 4.8 • • • 4.8 4.6 • 4.4 4.8 • 4.6 • 4.8 4.8 5.2 • • 4.6 • 4.8 4.6 • • • 4.8 4.7 4.6 • • • 4.7 • 4.8 • 4.7 • • 4.7 4.7 5.5 4.7 • • 4.7 • 5.2 • 4.7 4.7 5.3 4.8 4.7 4.8 • • • • 4.8 4.7 4.9 4.9 • 4.9 • 4.7 • 4.6 • • 4.8 • • 5.6 Lab pH 4.8 • 4.9 • • • • 4.9 • 5.0 5.1 4.8 • • 5.1 • 5.2 5.3 5.3 4.6 4.8 • 5.1 • • 4.9 5.2–5.3 4.8 4.9 • 5.0 • • 4.8 • • 5.1–5.2 • 4.9 5.2 4.9 • 5.0–5.1 5.0 4.8 • • • • • 5.0 4.8 4.9–5.0 5.0 • 5.0 4.9 • 4.8–4.9 • 4.9 • • • • 5.0 4.8 4.9 4.7–4.8 4.9 5.0 • 4.9 • • 4.8 • 4.6–4.7 5.2 4.5–4.6 • • 4.9 • 5.3 5.4 •
5.2 •
5.1 •
Figure 6.11 The pH of rain samples. Measurements made at the Central Analytical Laboratory, 2008. Values at stations in Alaska, Puerto Rico, and the Virgin Islands are given at the lower left. Hawaii data not available. Source: National Atmospheric Deposition Program Illinois State Water Survey.
4.4–4.5 4.3–4.4 4.3
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In water, sulfuric acid is a source of the hydrogen ion. H2SO4(aq)
H(aq) HSO4(aq)
[6.13a]
hydrogen sulfate ion
The hydrogen sulfate ion also can dissociate to yield another hydrogen ion. Equations 6.13b and 6.13c actually are more complex than shown here.
HSO4(aq)
H(aq) SO42(aq)
[6.13b]
sulfate ion
Adding equations 6.13a and 6.13b shows that sulfuric acid dissociates to yield two hydrogen ions and a sulfate ion (SO422). H2SO4(aq)
2 H(aq) SO42(aq)
[6.13c]
The sulfate ion can be detected in rainwater, providing a clue to the source of the additional acidity in the rain.
Your Turn 6.13
Sulfurous Acid
Sulfur dioxide dissolves in water to form sulfurous acid, H2SO3. Write equations for the formation of 2 H1(aq) from sulfurous acid, analogous to chemical equations 6.13a, 6.13b, and 6.13c for sulfuric acid. Visit Figures Alive! at the textbook’s website for a set of interactive activities relating to acids and bases.
Now let’s turn to NOx. Nitrogen oxides also dissolve in water to form acids, but the chemical reactions are more complex because the element O2 also is a reactant. For example, NO2 reacts in moist air to form nitric acid. This reaction is a simplification of the atmospheric chemistry that takes place. 4 NO2(g) 2 H2O(l) O2(g)
4 HNO3(aq)
[6.14]
nitric acid
In water, nitric acid dissociates to release H1. HNO3(aq)
Aerosols consist of tiny particles that remain suspended in our atmosphere (Section 1.11) and have a cooling effect on the Earth (Section 3.9).
H(aq) NO3(aq)
[6.15]
nitrate ion
The nitrate ion that is produced in this reaction can be detected in rainwater. Rain is only one of several ways that acids can be delivered to Earth’s surface waters. Snow and fog obviously are others. The term acid deposition refers to both wet and dry forms of delivery of acids from the upper atmosphere to the surface of the Earth. Examples of wet deposition include rain, snow, and fog. Mountaintops are particularly susceptible to wet deposition resulting from direct contact with clouds containing microscopic water droplets. Because these droplets contain acids that are more concentrated than those found in larger raindrops, they are often more acidic and damaging than acid rain. If SOx and NOx are indeed responsible for the increased acidity of rain falling in the eastern part of the United States, these regions should show elevated levels of sulfate ion and nitrate ion in rainwater, from SOx and NOx, respectively. Indeed, this is the case. Figure 6.12 shows the nitrate and sulfate ion content for wet deposition. Acid deposition also includes the “dry” forms of acids that deposit on land and water. For example, during dry weather, tiny solid particles (aerosols) of the acidic compounds ammonium nitrate, NH4NO3, and ammonium sulfate, (NH4)2SO4, are deposited. Dry deposition can be just as significant as the wet deposition of the acids in rain, snow, and fog. These aerosols also contribute to haze, as we will see in Section 6.12. Armed with the knowledge that oxides of sulfur and nitrogen contribute to acid rain, we now need to look more closely at how these oxides come to be released into the atmosphere. In the next section, we examine the chemistry of SO2 and its connection to the burning of coal. In the two sections after, we turn our attention to nitrogen chemistry, including that of NO and NO2.
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