Chemistry the central science 13e by theodore l brown 1

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Chemistry T h e C e n t r a l S c i e n c e 13 TH Edition Chemistry T h e C e n t r a l S c i e n c e 13 TH Edition Theodore L Brown University of Illinois at Urbana-Champaign H Eugene LeMay, Jr University of Nevada, Reno Bruce E Bursten University of Tennessee, Knoxville Catherine J Murphy University of Illinois at Urbana-Champaign Patrick M Woodward The Ohio State University Matthew W Stoltzfus The Ohio State University Boston Columbus Indianapolis New York San Francisco Upper Saddle River Amsterdam Cape Town Dubai London Madrid Milan Munich Paris Montréal Toronto Delhi Mexico City São Paulo Sydney Hong Kong Seoul Singapore Taipei Tokyo Editor in Chief, Chemistry: Adam Jaworski Senior Acquisitions Editor: Terry Haugen Acquisitions Editor: Chris Hess, Ph.D Executive Marketing Manager: Jonathan Cottrell Associate Team Lead, Program Management, Chemistry and Geoscience: Jessica Moro Editorial Assistant: Lisa Tarabokjia/Caitlin Falco Marketing Assistant: Nicola Houston Director of Development: Jennifer Hart Development Editor, Text: Carol Pritchard-Martinez Team Lead, Project Management, Chemistry and Geosciences: Gina M Cheselka Project Manager: Beth Sweeten Full-Service Project Management/Composition: Greg Johnson, PreMediaGlobal Operations Specialist: Christy Hall Illustrator: Precision Graphics Art Director: Mark Ong Interior / Cover Designer: Tamara Newnam Image Lead: Maya Melenchuk Photo Researcher: Kerri Wilson, PreMediaGlobal Text Permissions Manager: Alison Bruckner Text Permission Researcher: Jacqueline Bates, GEX Publishing Services Senior Content Producer: Kristin Mayo Production Supervisor, Media: Shannon Kong Electrostatic Potential Maps: Richard Johnson, Chemistry Department, University of New Hampshire Cover Image Credit: “Metal-Organic Frameworks” by Omar M Yaghi, University of California, Berkeley Credits and acknowledgments borrowed from other sources and reproduced, with permission, in this textbook appear on the appropriate page within the text or on pp P-1–P-2 Copyright © 2015, 2012, 2009, 2006, 2003, 2000, 1997, 1994, 1991, 1988, 1985, 1981, 1977 Pearson Education, Inc All rights reserved Manufactured in the United States of America This publication is protected by Copyright, and permission should be obtained from the publisher prior to any prohibited reproduction, storage in a retrieval system, or transmission in any form or by any means, electronic, mechanical, photocopying, recording, or likewise To obtain permission(s) to use material from this work, please submit a written request to Pearson Education, Inc., Permissions Department, Lake Street, Department 1G, Upper Saddle River, NJ 07458 Many of the designations used by manufacturers and sellers to distinguish their products are claimed as trademarks Where those designations appear in this book, and the publisher was aware of a trademark claim, the designations have been printed in initial caps or all caps Library of Congress Cataloging-In Publication Data Brown, Theodore L (Theodore Lawrence), 1928- author  Chemistry the central science.—Thirteenth edition / Theodore L Brown, University of Illinois at Urbana-Chanmpaign, H Euguene LeMay, Jr., University of Nevada, Reno, Bruce E Bursten, University of Tennessee, Knoxville, Catherine J Murphy, University of Illinois at Urbana-Chanmpaign, Patrick M Woodward, The Ohio State University, Matthew W Stoltzfus, The Ohio State University    pages cm   Includes index   ISBN-13: 978-0-321-91041-7   ISBN-10: 0-321-91041-9   Chemistry Textbooks I Title   QD31.3.B765 2014  540—dc23 2013036724 10—CRK— 17 16 15 14 www.pearsonhighered.com Student Edition: 0-321-91041-9 / 978-0-321-91041-7 Instructor’s Resource Copy: 0-321-96239-7 / 978-0-321-96239-3 To our students, whose enthusiasm and curiosity have often inspired us, and whose questions and suggestions have sometimes taught us Brief Contents Preface  xx Introduction: Matter and Measurement  2 Atoms, Molecules, and Ions  40 Chemical Reactions and Reaction Stoichiometry  80 Reactions in Aqueous Solution  122 Thermochemistry  164 Electronic Structure of Atoms  212 Periodic Properties of the Elements  256 Basic Concepts of Chemical Bonding  298 Molecular Geometry and Bonding Theories  342 10 Gases  398 11 Liquids and Intermolecular Forces  442 12 Solids and Modern Materials  480 13 Properties of Solutions  530 14 Chemical Kinetics  574 15 Chemical Equilibrium  628 16 Acid–Base Equilibria  670 17 Additional Aspects of Aqueous Equilibria  724 18 Chemistry of the Environment  774 19 Chemical Thermodynamics  812 20 Electrochemistry  856 21 Nuclear Chemistry  908 22 Chemistry of the Nonmetals  952 23 Transition Metals and Coordination Chemistry  996 24 The Chemistry of Life: Organic and Biological Chemistry  1040 Appendices   A Mathematical Operations  1092 B Properties of Water  1099 C Thermodynamic Quantities for Selected Substances at 298.15 K (25 °C)  1100 D Aqueous Equilibrium Constants  1103 E Standard Reduction Potentials at 25 °C  1105 Answers to Selected Exercises  A-1 Answers to Give It Some Thought  A-31 Answers to Go Figure  A-38 Answers to Selected Practice Exercises  A-44 Glossary  G-1 Photo/Art Credits  P-1 Index  I-1 vi Contents Preface  xx 1 Introduction: Matter and Measurement  2 1.1 The Study of Chemistry  2 The Atomic and Molecular Perspective of Chemistry  4  Why Study Chemistry?  5 1.2 Classifications of Matter  6 States of Matter  7  Pure Substances  7  Elements  7 Compounds  8 Mixtures  10 1.3 Properties of Matter  11 Physical and Chemical Changes  12  Separation of Mixtures  13 1.4 Units of Measurement  14 SI Units  15  Length and Mass  17  Temperature  17 Derived SI Units  19 Volume  19 Density  19 1.5 Uncertainty in Measurement  22 Precision and Accuracy  22  Significant Figures  22  Significant Figures in Calculations  22 1.6 Dimensional Analysis  27 Using Two or More Conversion Factors  28  Conversions Involving Volume  29 Chapter Summary and Key Terms  32 Learning Outcomes  32 Key Equations  32 Exercises  32 Additional Exercises  37 Chemistry Put to Work  Chemistry and the Chemical Industry  6 A Closer Look  The Scientific Method  14 Chemistry Put to Work  Chemistry in the News  20 Strategies in Chemistry  Estimating Answers  28 Strategies in Chemistry  The Importance of Practice  31 Strategies in Chemistry  The Features of This Book  32 2 Atoms, Molecules, and Ions  40 2.1 The Atomic Theory of Matter  42 2.2 The Discovery of Atomic Structure  43 Cathode Rays and Electrons  43  Radioactivity  45  The Nuclear Model of the Atom  46 2.3 The Modern View of Atomic Structure  47 Atomic Numbers, Mass Numbers, and Isotopes  49 2.4 Atomic Weights  50 The Atomic Mass Scale  50  Atomic Weight  51 2.5 The Periodic Table  52 2.6  Molecules and Molecular Compounds  56 Molecules and Chemical Formulas  56  Molecular and Empirical Formulas  56 Picturing Molecules  57 2.7 Ions and Ionic Compounds  58 Predicting Ionic Charges  59  Ionic Compounds  60 2.8 Naming Inorganic Compounds  62 Names and Formulas of Ionic Compounds  62 Names and Formulas of Acids  67  Names and Formulas of Binary Molecular Compounds  68 2.9 Some Simple Organic Compounds  69 Alkanes  69  Some Derivatives of Alkanes  70 Chapter Summary and Key Terms  72 Learning Outcomes  72 Key Equations  73 Exercises  73 Additional Exercises  78 A Closer Look  Basic Forces  49 A Closer Look  The Mass Spectrometer  52 A Closer Look  What Are Coins Made Of?  54 Chemistry and Life  Elements Required by Living Organisms  61 Strategies in Chemistry  How to Take a Test  71 vii viii Contents Electrolytes and Nonelectrolytes  124  How Compounds Dissolve in Water  125  Strong and Weak Electrolytes  126 3 Chemical Reactions and Reaction Stoichiometry  80 4.2 Precipitation Reactions  128 Solubility Guidelines for Ionic Compounds  129  Exchange (Metathesis) Reactions  130  Ionic Equations and Spectator Ions  131 Reactions  132 Acids  132  Bases  133  Strong and Weak Acids and Bases  133  Identifying Strong and Weak Electrolytes  135  Neutralization Reactions and Salts  135  Neutralization Reactions with Gas Formation  138 3.1 Chemical Equations  82 Balancing Equations  82  Indicating the States of Reactants and Products  85 3.2 Simple Patterns of Chemical Reactivity  86 Combination and Decomposition Reactions  86  Combustion Reactions  89 4.4 Oxidation–Reduction Reactions  138 Oxidation and Reduction  138  Oxidation Numbers  140  Oxidation of Metals by Acids and Salts  142  The Activity Series  143 3.3 Formula Weights  89 Formula and Molecular Weights  90  Percentage Composition from Chemical Formulas  91 4.5 Concentrations of Solutions  146 Molarity  146  Expressing the Concentration of an Electrolyte  147  Interconverting Molarity, Moles, and Volume  148 Dilution  149 3.4 Avogadro’s Number and the Mole  91 Molar Mass  93  Interconverting Masses and Moles  95  Interconverting Masses and Numbers of Particles  96 3.5 Empirical Formulas from Analyses  98 Molecular Formulas from Empirical Formulas  100  Combustion Analysis  101 3.6  Quantitative Information from Balanced Equations  103 3.7 Limiting Reactants  106 Theoretical and Percent Yields  109 Chapter Summary and Key Terms  111 Learning Outcomes  111 Key Equations  112 Exercises  112 Additional Exercises  118 Integrative Exercises  120 Design an Experiment  120 4.3  Acids, Bases, and Neutralization 4.6  Solution Stoichiometry and Chemical Analysis  151 Titrations  152 Chapter Summary and Key Terms  155 Learning Outcomes  156 Key Equations  156 Exercises  156 Additional Exercises  161 Integrative Exercises  161 Design an Experiment  163 Chemistry Put to Work Antacids  139 Strategies in Chemistry  Analyzing Chemical Reactions  146 Strategies in Chemistry  Problem Solving  92 Chemistry and Life  Glucose Monitoring  95 Strategies in Chemistry  Design an Experiment  110 5 Thermochemistry   164 5.1 Energy  166 4 Reactions in Aqueous Solution  122 4.1  General Properties of Aqueous Solutions  124 Kinetic Energy and Potential Energy  166  Units of Energy  168  System and Surroundings  169  Transferring Energy: Work and Heat  169 5.2 The First Law of Thermodynamics  170 Internal Energy  171  Relating ∆E to Heat and Work  172  Endothermic and Exothermic Processes  173  State Functions  174 Contents Orbitals and Quantum Numbers  228 5.3 Enthalpy  175 Pressure–Volume Work  175  Enthalpy Change  177 5.4 Enthalpies of Reaction  179 5.5 Calorimetry  181 Heat Capacity and Specific Heat  181  Constant-Pressure Calorimetry  183  Bomb Calorimetry (Constant-Volume Calorimetry)  185 6.6 Representations of Orbitals  230 The s Orbitals  230  The p Orbitals  233  The d and f Orbitals  233 6.7 Many-Electron Atoms  234 Orbitals and Their Energies  234  Electron Spin and the Pauli Exclusion Principle  235 6.8 Electron Configurations  237 Hund’s Rule  237  Condensed Electron Configurations  239 Transition Metals  240  The Lanthanides and Actinides  240 5.6 Hess’s Law  187 5.7 Enthalpies of Formation  189 Using Enthalpies of Formation to Calculate Enthalpies of Reaction  192 5.8 Foods and Fuels  194 Foods  194 Fuels  197 Other Energy Sources  198 Chapter Summary and Key Terms  200 Learning Outcomes  201 Key Equations  202 Exercises  202 Additional Exercises  209 Integrative Exercises  210 Design an Experiment  211 A Closer Look  Energy, Enthalpy, and P–V Work  178 Strategies in Chemistry  Using Enthalpy as a Guide  181 Chemistry and Life  The Regulation of Body Temperature  186 ix 6.9  Electron Configurations and the Periodic Table  241 Anomalous Electron Configurations  245 Chapter Summary and Key Terms  246 Learning Outcomes  247 Key Equations  247 Exercises  248 Additional Exercises  252 Integrative Exercises  255  Design an Experiment  255 A Closer Look  Measurement and the Uncertainty Principle  225 A Closer Look  Thought Experiments and Schrödinger’s Cat  227 A Closer Look  Probability Density and Radial Probability Functions  232 Chemistry and Life  Nuclear Spin and Magnetic Resonance Imaging  236 Chemistry Put to Work  The Scientific and Political Challenges of Biofuels  198 6 Electronic Structure of Atoms  212 6.1 The Wave Nature of Light  214 6.2 Quantized Energy and Photons  216 Hot Objects and the Quantization of Energy  216  The Photoelectric Effect and Photons  217 6.3 Line Spectra and the Bohr Model  219 Line Spectra  219  Bohr’s Model  220  The Energy States of the Hydrogen Atom  221  Limitations of the Bohr Model  223 6.4 The Wave Behavior of Matter  223 The Uncertainty Principle  225 6.5  Quantum Mechanics and Atomic Orbitals  226 7 Periodic Properties of the Elements  256 7.1 Development of the Periodic Table  258 7.2 Effective Nuclear Charge  259 7.3 Sizes of Atoms and Ions  262 Periodic Trends in Atomic Radii  264  Periodic Trends in Ionic Radii  265 7.4 Ionization Energy  268 Variations in Successive Ionization Energies  268  Periodic Trends in First Ionization Energies  268  Electron Configurations of Ions  271 7.5 Electron Affinity  272 7.6 Metals, Nonmetals, and Metalloids  273 Metals  274 Nonmetals  276 Metalloids  277 section 15.6 Applications of Equilibrium Constants 647 the reaction quotient in terms of molar concentrations is Qc = 3D4d3E4e 3A4a3B4b [15.23] (A related quantity Qp can be written for any reaction that involves gases by using partial pressures instead of concentrations.) Although we use what looks like the equilibrium-constant expression to calculate the reaction quotient, the concentrations we use may or may not be the equilibrium concentrations For example, when we substituted the starting concentrations into the equilibrium-constant expression of Equation 15.22, we obtained Qc = 0.500 whereas Kc = 0.105 The equilibrium constant has only one value at each temperature The reaction quotient, however, varies as the reaction proceeds Of what use is Q? One practical thing we can with Q is tell whether our reaction really is at equilibrium, which is an especially valuable option when a reaction is very slow We can take samples of our reaction mixture as the reaction proceeds, separate the components, and measure their concentrations Then we insert these numbers into Equation 15.23 for our reaction To determine whether we are at equilibrium, or in which direction the reaction proceeds to achieve equilibrium, we compare the values of Qc and Kc or Qp and Kp Three possible situations arise: • Q K: The concentration of products is too small and that of reactants too large The reaction achieves equilibrium by forming more products; it proceeds from left to right • Q = K: The reaction quotient equals the equilibrium constant only if the system is at equilibrium • Q K: The concentration of products is too large and that of reactants too small The reaction achieves equilibrium by forming more reactants; it proceeds from right to left These relationships are summarized in ▶ Figure 15.8 At equilibrium Q< K Q K Reaction proceeds to form more products Q Q=K K Equilibrium Q Q >K K Reaction proceeds to form more reactants ▲ Figure 15.8  Predicting the direction of a reaction by comparing Q and K at a given temperature Sample Exercise 15.9  Predicting the Direction of Approach to Equilibrium At 448 °C, the equilibrium constant Kc for the reaction H21g2 + I21g2 ∆ HI1g2 is 50.5 Predict in which direction the reaction proceeds to reach equilibrium if we start with 2.0 * 10-2 mol of HI, 1.0 * 10-2 mol of H2, and 3.0 * 10-2 mol of I2 in a 2.00-L container Solution Analyze We are given a volume and initial molar amounts of the species in a reaction and asked to determine in which direction the reaction must proceed to achieve equilibrium Plan We can determine the starting concentration of each species in the reaction mixture We can then substitute the starting concentrations into the equilibrium-constant expression to calculate the reaction quotient, Qc Comparing the magnitudes of the equilibrium constant, which is given, and the reaction quotient will tell us in which direction the reaction will proceed Solve The initial concentrations are 3HI4 = 2.0 * 10-2 mol>2.00 L = 1.0 * 10-2 M 3H24 = 1.0 * 10-2 mol>2.00 L = 5.0 * 10-3 M 3I24 = 3.0 * 10-2 mol>2.00 L = 1.5 * 10-2 M The reaction quotient is therefore Qc = 3HI42 3H243I24 = 11.0 * 10-222 15.0 * 10-3211.5 * 10-22 = 1.3 Because Qc Kc, the concentration of HI must increase and the concentrations of H2 and I2 must decrease to reach equilibrium; the reaction as written proceeds left to right to attain equilibrium Practice Exercise Which of the following statements accurately describes what would happen to the direction of the reaction described in the sample exercise above, if the size of the container were different from 2.00 L? (a) The reaction would proceed in the opposite direction (from right to left) if the container volume were reduced sufficiently (b) The reaction would proceed in the opposite direction if the container volume were expanded sufficiently (c) The direction of this reaction does not depend on the volume of the container Practice Exercise At 1000 K, the value of Kp for the reaction SO31g2 ∆ SO21g2 + O21g2 is 0.338 Calculate the value for Qp, and predict the direction in which the reaction proceeds toward equilibrium if the initial partial pressures are PSO3 = 0.16 atm; PSO2 = 0.41 atm; PO2 = 2.5 atm 648 chapter 15 Chemical Equilibrium Calculating Equilibrium Concentrations Chemists frequently need to calculate the amounts of reactants and products present at equilibrium in a reaction for which they know the equilibrium constant The approach in solving problems of this type is similar to the one we used for evaluating equilibrium constants: We tabulate initial concentrations or partial pressures, changes in those concentrations or pressures, and final equilibrium concentrations or partial pressures Usually, we end up using the equilibrium-constant expression to derive an equation that must be solved for an unknown quantity, as demonstrated in Sample Exercise 15.10 Sample Exercise 15.10  Calculating Equilibrium Concentrations For the Haber process, N21g2 + H21g2 ∆ NH31g2, Kp = 1.45 * 10-5, at 500 °C In an equilibrium mixture of the three gases at 500 °C, the partial pressure of H2 is 0.928 atm and that of N2 is 0.432 atm What is the partial pressure of NH3 in this equilibrium mixture? Solution Analyze We are given an equilibrium constant, Kp, and the equilibrium partial pressures of two of the three substances in the equation 1N2 and H22, and we are asked to calculate the equilibrium partial pressure for the third substance 1NH32 Plan We can set Kp equal to the equilibrium-constant expression and substitute in the partial pressures that we know Then we can solve for the only unknown in the equation Solve We tabulate the equilibrium pressures: N21g2 + H21g2 ∆ NH31g2 Equilibrium pressure 1atm2 0.432 0.928 x Because we not know the equilibrium pressure of NH3, we represent it with x At equilibrium, the pressures must satisfy the equilibriumconstant expression: Kp = 1PNH322 PN21PH22 = x2 = 1.45 * 10-5 10.432210.92823 We now rearrange the equation to solve for x: x2 = 11.45 * 10-5210.432210.92823 = 5.01 * 10-6 x = 25.01 * 10-6 = 2.24 * 10-3 atm = PNH3 Check We can always check our answer by using it to recalculate the value of the equilibrium constant: 12.24 * 10-322 Kp = = 1.45 * 10-5 10.432210.92823 Practice Exercise At 500 K, the reaction NO1g2 + Cl21g2 ∆ NOCl1g2 has Kp = 51 In an equilibrium mixture at 500 K, the partial pressure of NO is 0.125 atm and Cl2 is 0.165 atm What is the partial pressure of NOCl in the equilibrium mixture? (a) 0.13 atm, (b) 0.36 atm, (c) 1.0 atm, (d) 5.1 * 10-5 atm, (e) 0.125 atm Practice Exercise At 500 K, the reaction PCl51g2 ∆ PCl31g2 + Cl21g2 has Kp = 0.497 In an equilibrium mixture at 500 K, the partial pressure of PCl5 is 0.860 atm and that of PCl3 is 0.350 atm What is the partial pressure of Cl2 in the equilibrium mixture? In many situations, we know the value of the equilibrium constant and the initial amounts of all species We must then solve for the equilibrium amounts Solving this type of problem usually entails treating the change in concentration as a variable The stoichiometry of the reaction gives us the relationship between the changes in the amounts of all the reactants and products, as illustrated in Sample Exercise 15.11 The calculations frequently involve the quadratic formula, as you will see in this exercise Sample Exercise 15.11  Calculating Equilibrium Concentrations from Initial Concentrations A 1.000-L flask is filled with 1.000 mol of H21g2 and 2.000 mol of I21g2 at 448 °C The value of the equilibrium constant Kc for the reaction H21g2 + I21g2 ∆ HI1g2 at 448 °C is 50.5 What are the equilibrium concentrations of H2, I2, and HI in moles per liter? Solution Analyze We are given the volume of a container, an equilibrium constant, and starting amounts of reactants in the container and are asked to calculate the equilibrium concentrations of all species Plan In this case, we are not given any of the equilibrium concentra- tions We must develop some relationships that relate the initial concentrations to those at equilibrium The procedure is similar in many regards to that outlined in Sample Exercise 15.8, where we calculated an equilibrium constant using initial concentrations 649 section 15.6 Applications of Equilibrium Constants Solve (1)  We note the initial concentrations of H2 and I2: (2)  We construct a table in which we tabulate the initial concentrations: 3H24 = 1.000 M and 3I24 = 2.000 M Initial concentration (M) H21g2 I21g2 ∆ H21g2 + I21g2 ∆ + 1.000 2.000 HI1g2 Change in concentration (M) Equilibrium concentration (M) (3)  We use the stoichiometry of the reaction to determine the changes in concentration that occur as the reaction proceeds to equilibrium The H2 and I2 concentrations will decrease as equilibrium is established and that of HI will increase Let’s represent the change in concentration of H2 by x The balanced chemical equation tells us the relationship between the changes in the concentrations of the three gases For each x mol of H2 that reacts, x mol of I2 are consumed and 2x mol of HI are produced: (4)  We use initial concentrations and changes in concentrations, as dictated by stoichiometry, to express the equilibrium concentrations With all our entries, our table now looks like this: (5)  We substitute the equilibrium concentrations into the equilibrium-constant expression and solve for x: If you have an equation-solving calculator, you can solve this equation directly for x If not, expand this expression to obtain a quadratic equation in x: Initial concentration (M) Change in concentration (M) 1.000 2.000 -x -x + 2x Equilibrium concentration (M) H21g2 + Initial concentration (M) 1.000 -x Change in concentration (M) 1.000 - x Equilibrium concentration (M) Kc = 3HI42 3H243I24 = 12x22 11.000 - x212.000 - x2 I21g2 ∆ -x + 2x 2.000 - x 2x = 50.5 4x2 = 50.51x2 - 3.000x + 2.0002 When we substitute x = 2.323 into the expressions for the equilibrium concentrations, we find negative concentrations of H2 and I2 Because a negative concentration is not chemically meaningful, we reject this solution We then use x = 0.935 to find the equilibrium concentrations: Check We can check our solution by putting these numbers into the equilibrium-constant expression to assure that we correctly calculate the equilibrium constant: x = - 1-151.52 { 21- 151.522 - 4146.521101.02 2146.52 3H24 = 1.000 - x = 0.065 M 3I24 = 2.000 - x = 1.065 M 3HI4 = 2x = 1.87 M Kc = 3HI42 3H243I24 = 11.8722 10.065211.0652 = 51 Comment Whenever you use a quadratic equation to solve an equilibrium problem, one of the solutions to the equation will give you a value that leads to negative concentrations and thus is not chemically meaningful Reject this solution to the quadratic equation HI1g2 2.000 46.5x2 - 151.5x + 101.0 = Solving the quadratic equation (Appendix A.3) leads to two solutions for x: HI1g2 = 2.323 or 0.935 650 chapter 15 Chemical Equilibrium Practice Exercise For the equilibrium Br21g2 + Cl21g2 ∆ BrCl1g2, the equilibrium constant Kp is 7.0 at 400 K If a cylinder is charged with BrCl(g) at an initial pressure of 1.00 atm and the system is allowed to come to equilibrium what is the final (equilibrium) pressure of BrCl? (a) 0.57 atm, (b) 0.22 atm, (c) 0.45 atm, (d) 0.15 atm, (e) 0.31 atm Practice Exercise For the equilibrium PCl51g2 ∆ PCl31g2 + Cl21g2, the equilibrium constant Kp is 0.497 at 500 K A gas cylinder at 500 K is charged with PCl51g2 at an initial pressure of 1.66 atm What are the equilibrium pressures of PCl5, PCl3, and Cl2 at this temperature? 15.7 | Le Châtelier’s Principle Many of the products we use in everyday life are obtained from the chemical industry Chemists and chemical engineers in industry spend a great deal of time and effort to maximize the yield of valuable products and minimize waste For example, when Haber developed his process for making ammonia from N2 and H2, he examined how reaction conditions might be varied to increase yield Using the values of the equilibrium constant at various temperatures, he calculated the equilibrium amounts of NH3 formed under a variety of conditions Some of Haber’s results are shown in ▼ Figure 15.9 Notice that the percent of NH3 present at equilibrium decreases with increasing temperature and increases with increasing pressure We can understand these effects in terms of a principle first put forward by Henri-Louis Le Châtelier* (1850–1936), a French industrial chemist: If a system at equilibrium is disturbed by a change in temperature, pressure, or a component concentration, the ­system will shift its equilibrium position so as to counteract the effect of the disturbance Go Figure At what combination of pressure and temperature should you run the reaction to maximize NH3 yield? Percent NH3 produced 7% 47 8% 9% 0% 35 9% 9% 2% 42 9% 54 9% 16 32 20 8% 37 48 8% 8% 550 60 6% re 500 450 400 Te m 200 To tal 300 pr ess 400 ur e( atm 500 ) ) 4% 12 26 (°C 27 38 % 9% atu 8.8 18 Percent of NH3 decreases with increasing temperature pe r Percent of NH3 increases with increasing pressure ▲ Figure 15.9  Effect of temperature and pressure on NH3 yield in the Haber process Each mixture was produced by starting with a 3 : 1 molar mixture of H2 and N2 *Pronounced “le-SHOT-lee-ay.” section 15.7  Le Châtelier’s Principle Le Châtelier’s Principle If a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, the system will shift its equilibrium position so as to counter the effect of the disturbance Concentration: adding or removing a reactant or product If a substance is added to a system at equilibrium, the system reacts to consume some of the substance If a substance is removed from a system, the system reacts to produce more of substance Initial equilibrium + Substance added Equilibrium reestablished + + Substances react Pressure: changing the pressure by changing the volume At constant temperature, reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas Pressure Initial volume System shifts to direction of fewer moles of gas Temperature: If the temperature of a system at equilibrium is increased, the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic reaction The equilibrium shifts in the direction that consumes the “excess reactant,” namely heat Endothermic Exothermic Increasing T Reaction shifts right Increasing T Reaction shifts left Decreasing T Reaction shifts left Decreasing T Reaction shifts right In this section, we use Le Châtelier’s principle to make qualitative predictions about how a system at equilibrium responds to various changes in external conditions We consider three ways in which a chemical equilibrium can be disturbed: (1) adding or removing a reactant or product, (2) changing the pressure by changing the volume, and (3) changing the temperature Change in Reactant or Product Concentration A system at dynamic equilibrium is in a state of balance When the concentrations of species in the reaction are altered, the equilibrium shifts until a new state of balance is attained What does shift mean? It means that reactant and product concentrations change over time to accommodate the new situation Shift does not mean that the equilibrium constant itself is altered; the equilibrium constant remains the same Le Châtelier’s principle states that the shift is in the direction that minimizes or reduces the effect of the change Therefore, if a chemical system is already at equilibrium and the concentration of any substance in the mixture is increased (either reactant or product), the system reacts to consume some of that substance Conversely, if the concentration of a substance is decreased, the system reacts to produce some of that substance There is no change in the equilibrium constant when we change the concentrations of reactants or products As an example, consider our familiar equilibrium mixture of N2, H2, and NH3: N21g2 + H21g2 ∆ NH31g2 651 652 chapter 15 Chemical Equilibrium Go Figure Why does the nitrogen concentration decrease after hydrogen is added? N2(g) + H2(g) Initial equilibrium NH3(g) Equilibrium reestablished H2 added Partial pressure H2 NH3 N2 Time ▲ Figure 15.10  Effect of adding H2 to an equilibrium mixture of N2, H2, and NH3 Adding H2 causes the reaction as written to shift to the right, consuming some N2 to produce more NH3 Adding H2 causes the system to shift so as to reduce the increased concentration of H2 (▲ Figure 15.10) This change can occur only if the reaction consumes H2 and simultaneously consumes N2 to form more NH3 Adding N2 to the equilibrium mixture likewise causes the reaction to shift toward forming more NH3 Removing NH3 also causes a shift toward producing more NH3, whereas adding NH3 to the system at equilibrium causes the reaction to shift in the direction that reduces the increased NH3 concentration: Some of the added ammonia decomposes to form N2 and H2 All of these “shifts” are entirely consistent with predictions that we would make by comparing the reaction quotient Q with the equilibrium constant K In the Haber reaction, therefore, removing NH3 from an equilibrium mixture of N2, H2, and NH3 causes the reaction to shift right to form more NH3 If the NH3 can be removed continuously as it is produced, the yield can be increased dramatically In the industrial production of ammonia, the NH3 is continuously removed by selectively liquefying it (▶ Figure 15.11) (The boiling point of NH3, -33 °C, is much higher than those of N2, -196 °C, and H2, -253 °C.) The liquid NH3 is removed, and the N2 and H2 are recycled to form more NH3 As a result of the product being continuously removed, the reaction is driven essentially to completion Give It Some Thought Does the equilibrium NO1g2 + O21g2 ∆ NO21g2 shift to the right (more products) or left (more reactants) if (a) O2 is added to the system? (b) NO is removed? Effects of Volume and Pressure Changes If a system containing one or more gases is at equilibrium and its volume is decreased, thereby increasing its total pressure, Le Châtelier’s principle indicates that the system responds by shifting its equilibrium position to reduce the pressure A system can reduce its pressure by reducing the total number of gas molecules (fewer molecules of gas section 15.7  Le Châtelier’s Principle Hot gases N2 H2 Incoming N2 and H2 gases Heat exchanger Heat exchanger Unreacted N2 and H2 recycled As gas mixture cools, NH3(g) liquifies Heated gases pass over catalyst, NH3 forms N2 and H2 gases heated to approximately 500 °C Liquid NH3 outlet ▲ Figure 15.11  Diagram of the industrial production of ammonia Incoming N21g2 and H21g2 are heated to approximately 500 °C and passed over a catalyst When the resultant N2, H2, and NH3 mixture is cooled, the NH3 liquefies and is removed from the mixture, shifting the reaction to produce more NH3 exert a lower pressure) Thus, at constant temperature, reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas Increasing the volume causes a shift in the direction that produces more gas molecules (▼ Figure 15.12) A(g) B(g) Decrease volume, increase pressure New equilibrium favors products to reduce total moles of gas Increase volume, decrease pressure Initial volume New equilibrium favors reactants to increase total moles of gas ▲ Figure 15.12  Pressure and Le Châtelier’s principle 653 654 chapter 15 Chemical Equilibrium Give It Some Thought What happens to the equilibrium SO21g2 + O21g2 ∆ SO31g2, if the volume of the system is increased? In the reaction N21g2 + H21g2 ∆ NH31g2, four molecules of reactant are consumed for every two molecules of product produced Consequently, an increase in pressure (caused by a decrease in volume) shifts the reaction in the direction that produces fewer gas molecules, which leads to the formation of more NH3, as indicated in Figure 15.9 In the reaction H21g2 + I21g2 ∆ HI1g2, the number of molecules of gaseous products (two) equals the number of molecules of gaseous reactants; therefore, changing the pressure does not influence the position of equilibrium Keep in mind that, as long as temperature remains constant, pressure–volume changes not change the value of K Rather, these changes alter the partial pressures of the gaseous substances In Sample Exercise 15.7, we calculated Kp = 2.79 * 10-5 for the Haber reaction, N21g2 + H21g2 ∆ NH31g2, in an equilibrium mixture at 472 °C containing 7.38 atm H2, 2.46 atm N2, and 0.166 atm NH3 Consider what happens when we suddenly reduce the volume of the system by one-half If there were no shift in equilibrium, this volume change would cause the partial pressures of all substances to double, giving PH2 = 14.76 atm, PN2 = 4.92 atm, and PNH3 = 0.332 atm The reaction quotient would then no longer equal the equilibrium constant: Qp = 1PNH322 PN21PH22 = 10.33222 14.922114.7623 = 6.97 * 10-6 ≠ Kp Because Qp Kp, the system would no longer be at equilibrium Equilibrium would be reestablished by increasing PNH3 and/or decreasing PN2 and PH2 until Qp = Kp = 2.79 * 10-5 Therefore, the equilibrium shifts to the right in the reaction as written, as Le Châtelier’s principle predicts It is possible to change the pressure of a system in which a chemical reaction is running without changing its volume For example, pressure increases if additional amounts of any reacting components are added to the system We have already seen how to deal with a change in concentration of a reactant or product However, the total pressure in the reaction vessel might also be increased by adding a gas that is not involved in the equilibrium For example, argon might be added to the ammonia equilibrium system The argon would not alter the partial pressures of any of the reacting components and therefore would not cause a shift in equilibrium Effect of Temperature Changes Changes in concentrations or partial pressures shift equilibria without changing the value of the equilibrium constant In contrast, almost every equilibrium constant changes as the temperature changes For example, consider the equilibrium established when cobalt(II) chloride 1CoCl22 is dissolved in hydrochloric acid, HCl(aq), in the endothermic reaction Co1H2O262+ 1aq2 + Cl-1aq2 ∆ CoCl42- 1aq2 + H2O1l2  ∆H 0 [15.24] Pale pink Deep blue Because Co1H2O262+ is pink and CoCl42- is blue, the position of this equilibrium is readily apparent from the color of the solution (▶ Figure 15.13) When the solution is heated it turns blue, indicating that the equilibrium has shifted to form more CoCl42- Cooling the solution leads to a pink solution, indicating that the equilibrium has shifted to produce more Co1H2O262+ We can monitor this reaction by spectroscopic methods, measuring the concentration of all species at the different  (Section 14.2) We can then calculate the equilibrium constant at temperatures each temperature How we explain why the equilibrium constants and therefore the position of equilibrium both depend on temperature? section 15.7  Le Châtelier’s Principle ∆H > 0, endothermic reaction Heat + Co(H2O)62+(aq) + Cl−(aq) Pink CoCl42−(aq) + H2O(l) Blue CoCl42+ Co(H2O)62– Cool Solution appears pink because lowering the temperature shifts the equilibrium to favor formation of the pink Co(H2O)62+ ion Heat Solution appears violet because appreciable amounts of both pink Co(H2O)62+ and blue CoCl42– are present ▲ Figure 15.13  Temperature and Le Châtelier’s principle In the molecular level views, only the CoCl42 - and Co1H2O262+ ions are shown for clarity We can deduce the rules for the relationship between K and temperature from Le Châtelier’s principle We this by treating heat as a chemical reagent In an endothermic (heat-absorbing) reaction, we consider heat a reactant, and in an exothermic (heat-releasing) reaction, we consider heat a product: Endothermic: Exothermic: Reactants + heat ∆ products Reactants ∆ products + heat When the temperature of a system at equilibrium is increased, the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic reaction The equilibrium shifts in the direction that consumes the excess reactant (or product), namely heat Give It Some Thought Use Le Châtelier’s principle to explain why the equilibrium vapor pressure of a liquid increases with increasing temperature In an endothermic reaction, such as Equation 15.24, heat is absorbed as reactants are converted to products Thus, increasing the temperature causes the equilibrium to shift to the right, in the direction of making more products, and K increases In an exothermic reaction, the opposite occurs: Heat is produced as reactants are converted to Solution appears blue because raising the temperature shifts the equilibrium to favor formation of the blue CoCl42– ion 655 656 chapter 15 Chemical Equilibrium products Thus, increasing the temperature in this case causes the equilibrium to shift to the left, in the direction of making more reactants, and K decreases Endothermic: Increasing T results in higher K value Exothermic: Increasing T results in lower K value Cooling a reaction has the opposite effect As we lower the temperature, the equilibrium shifts in the direction that produces heat Thus, cooling an endothermic reaction shifts the equilibrium to the left, decreasing K, as shown in Figure 15.13, and cooling an exothermic reaction shifts the equilibrium to the right, increasing K Sample Exercise 15.12  Using Le Châtelier’s Principle to Predict Shifts in Equilibrium Consider the equilibrium N2O41g2 ∆ NO21g2 ∆H ° = 58.0 kJ In which direction will the equilibrium shift when (a) N2O4 is added, (b) NO2 is removed, (c) the pressure is increased by addition of N21g2, (d) the volume is increased, (e) the temperature is decreased? Solution Analyze We are given a series of changes to be made to a system at equilibrium and are asked to predict what effect each change will have on the position of the equilibrium (e) The reaction is endothermic, so we can imagine heat as a reagent on the reactant side of the equation Decreasing the temperature will shift the equilibrium in the direction that produces heat, so the equilibrium shifts to the left, toward the formation of more N2O4 Note that only this last change also affects the value of the equilibrium constant, K Plan Le Châtelier’s principle can be used to determine the effects of Practice Exercise For the reaction Solve NH31g2 + O21g2 ∆ NO1g2 + H2O1g2 each of these changes (a) The system will adjust to decrease the concentration of the added N2O4, so the equilibrium shifts to the right, in the direction of product (b) The system will adjust to the removal of NO2 by shifting to the side that produces more NO2; thus, the equilibrium shifts to the right (c) Adding N2 will increase the total pressure of the system, but N2 is not involved in the reaction The partial pressures of NO2 and N2O4 are therefore unchanged, and there is no shift in the position of the equilibrium (d) If the volume is increased, the system will shift in the direction that occupies a larger volume (more gas molecules); thus, the equilibrium shifts to the right ∆H ° = - 904 kJ which of the following changes will shift the equilibrium to the right, toward the formation of more products? (a) Adding more water vapor, (b) Increasing the temperature, (c) Increasing the volume of the reaction vessel, (d) Removing O2(g), (e) Adding atm of Ne(g) to the reaction vessel Practice Exercise For the reaction PCl51g2 ∆ PCl31g2 + Cl21g2 ∆H° = 87.9 kJ in which direction will the equilibrium shift when (a) Cl21g2 is removed, (b) the temperature is decreased, (c) the volume of the reaction system is increased, (d) PCl31g2 is added? Sample Exercise 15.13  Predicting the Effect of Temperature on K (a) Using the standard heat of formation data in Appendix C, determine the standard enthalpy change for the reaction N21g2 + H21g2 ∆ NH31g2 (b) Determine how the equilibrium constant for this reaction should change with temperature Solution Analyze We are asked to determine the standard enthalpy change of a reaction and how the equilibrium constant for the reaction varies with temperature Plan (a) We can use standard enthalpies of formation to calculate ∆H° for the reaction (b) We can then use Le Châtelier’s principle to determine what effect temperature will have on the equilibrium constant Solve (a) Recall that the standard enthalpy change for a reaction is given by the sum of the standard molar enthalpies of formation of the products, each multiplied by its coefficient in the balanced chemical equation, minus the same quantities for the reactants  (Section 5.7) At 25 °C, ∆Hf° for NH31g2 is -46.19 kJ>mol The ∆Hf° values for H21g2 and N21g2 are zero by definition because the enthalpies of formation of the elements in their normal section 15.7  Le Châtelier’s Principle 657 with changes in temperature and that it is larger at lower temperatures states at 25 °C are defined as zero  (Section 5.7) Because mol of NH3 is formed, the total enthalpy change is Comment The fact that Kp for the formation of NH3 from N2 and H2 12 mol21-46.19 kJ>mol2 - = - 92.38 kJ (b) Because the reaction in the forward direction is exothermic, we can consider heat a product of the reaction An increase in temperature causes the reaction to shift in the direction of less NH3 and more N2 and H2 This effect is seen in the values for Kp presented in ▼ Table 15.2 Notice that Kp changes markedly decreases with increasing temperature is a matter of great practical importance To form NH3 at a reasonable rate requires higher temperatures At higher temperatures, however, the equilibrium constant is smaller, and so the percentage conversion to NH3 is smaller To compensate for this, higher pressures are needed because high pressure favors NH3 formation Table 15.2  Variation in K p with Temperature for n2 + 3h ∆ nh Temperature °C Kp 300 4.34 * 10-3 400 1.64 * 10-4 450 4.51 * 10 -5 500 1.45 * 10-5 550 5.38 * 10-6 600 2.25 * 10-6 Practice Exercise The standard enthalpy of formation of HCl(g) is - 92.3 kJ>mol Given only this information, in which direction would you expect the equilibrium for the reaction H21g2 + Cl21g2 ∆ HCl1g2 to shift as the temperature increases: (a) to the left, (b) to the right, (c) no shift in equilibrium? Practice Exercise Using the thermodynamic data in Appendix C, determine the enthalpy change for the reaction POCl31g2 ∆ PCl31g2 + O21g2 Use this result to determine how the equilibrium constant for the reaction should change with temperature The Effect of Catalysts What happens if we add a catalyst to a chemical system that is at equilibrium? As shown in ▼ Figure 15.14, a catalyst lowers the activation barrier between reactants and products The activation energies for both the forward and reverse reactions are lowered The catalyst thereby increases the rates of both forward and reverse reactions Since K is the ratio of the forward and reverse rate constants for a reaction, you can predict, correctly, that the presence of a catalyst, even though it changes the reaction rate, does not affect the numeric value of K (Figure 15.14) As a result, a catalyst increases the rate at which equilibrium is achieved but does not change the composition of the equilibrium mixture The rate at which a reaction approaches equilibrium is an important practical consideration As an example, let’s again consider the synthesis of ammonia from Go Figure What quantity dictates the speed of a reaction: (a) the energy difference between the initial state and the transition state or (b) the energy difference between the initial state and the final state? Catalyzed reaction proceeds more rapidly Transition states Catalyzed reaction has lower activation energy [B] Energy [B]eq A The two reactions reach the same equilibrium mixture, but the catalyzed reaction achieves equilibrium faster B Reaction pathway Time ▲ Figure 15.14  An energy profile for the reaction A ∆ B (left), and the change in concentration of B as a function of time (right), with and without a catalyst Green curves show the reaction with a catalyst; black curves show the reaction without a catalyst 658 chapter 15 Chemical Equilibrium N2 and H2 In designing his process, Haber had to deal with a rapid decrease in the equilibrium constant with increasing temperature (Table 15.2) At temperatures sufficiently high to give a satisfactory reaction rate, the amount of ammonia formed was too small The solution to this dilemma was to develop a catalyst that would produce a reasonably rapid approach to equilibrium at a sufficiently low temperature, so that the equilibrium constant remained reasonably large The development of a suitable catalyst thus became the focus of Haber’s research efforts After trying different substances to see which would be most effective, Carl Bosch settled on iron mixed with metal oxides, and variants of this catalyst formulation are still used today  (Section 15.2, “The Haber Process”) These catalysts make it possible to obtain a reasonably rapid approach to equilibrium at around 400 to 500 °C and 200 to 600 atm The high pressures are needed to obtain a satisfactory equilibrium amount of NH3 If a catalyst could be found that leads to sufficiently rapid reaction at temperatures lower than 400 °C, it would be possible to obtain the same extent of equilibrium conversion at pressures much lower than 200 to 600 atm This would result in great savings in both the cost of the high-pressure equipment and the energy consumed in the production of ammonia It is estimated that the Haber process consumes approximately 1% of the energy generated in the world each year Not surprisingly chemists and chemical engineers are actively searching for improved catalysts for the Haber process A breakthrough in this field would not only increase the supply of ammonia for fertilizers, it would also reduce the global consumption of fossil fuels in a significant way Give It Some Thought Can a catalyst be used to increase the amount of product produced for a reaction that reaches equilibrium quickly without a catalyst? Sample Integrative Exercise   Putting Concepts Together At temperatures near 800 °C, steam passed over hot coke (a form of carbon obtained from coal) reacts to form CO and H2: C1s2 + H2O1g2 ∆ CO1g2 + H21g2 The mixture of gases that results is an important industrial fuel called water gas (a) At 800 °C the equilibrium constant for this reaction is Kp = 14.1 What are the equilibrium partial pressures of H2O, CO, and H2 in the equilibrium mixture at this temperature if we start with solid carbon and 0.100 mol of H2O in a 1.00-L vessel? (b) What is the minimum amount of carbon required to achieve equilibrium under these conditions? (c) What is the total pressure in the vessel at equilibrium? (d) At 25 °C the value of Kp for this reaction is 1.7 * 10-21 Is the reaction exothermic or endothermic? (e) To produce the maximum amount of CO and H2 at equilibrium, should the pressure of the system be increased or decreased? Solution (a)  To determine the equilibrium partial pressures, we use the ideal-gas equation, first determining the starting partial pressure of water PH2O = nH2ORT V = 10.100 mol210.08206 L@atm>mol@K211073 K2 1.00 L = 8.81 atm We then construct a table of initial partial pressures and their changes as equilibrium is achieved: C1s2 + H2O1g2 8.81 H21g2 -x +x +x 8.81 - x x x Initial partial pressure (atm) Change in partial pressure (atm) Equilibrium partial pressure (atm) ∆ CO1g2 + There are no entries in the table under C(s) because the reactant, being a solid, does not appear in the equilibrium-constant expression Substituting the equilibrium partial pressures of the other species into the equilibrium-constant expression for the reaction gives Kp = PCOPH2 PH2O = 1x21x2 18.81 - x2 = 14.1 section 15.7  Le Châtelier’s Principle 659 Multiplying through by the denominator gives a quadratic equation in x:                 x2 = 114.1218.81 - x2 x2 + 14.1x - 124.22 = Solving this equation for x using the quadratic formula yields x = 6.14 atm Hence, the equilibrium partial pressures are PCO = x = 6.14 atm, PH2 = x = 6.14 atm, and PH2O = 18.81 - x2 = 2.67 atm (b)  Part (a) shows that x = 6.14 atm of H2O must react for the system to achieve equilibrium We can use the ideal-gas equation to convert this partial pressure into a mole amount n = 16.14 atm211.00 L2 PV = 0.0697 mol = RT 10.08206 L@atm>mol@K211073 K2 Thus, 0.0697 mol of H2O and the same amount of C must react to achieve equilibrium As a result, there must be at least 0.0697 mol of C (0.836 g C) present among the reactants at the start of the reaction (c)  The total pressure in the vessel at equilibrium is simply the sum of the equilibrium partial pressures: Ptotal = PH2O + PCO + PH2 = 2.67 atm + 6.14 atm + 6.14 atm = 14.95 atm (d)  In discussing Le Châtelier’s principle, we saw that endothermic reactions exhibit an increase in Kp with increasing temperature Because the equilibrium constant for this reaction increases as temperature increases, the reaction must be endothermic From the enthalpies of formation given in Appendix C, we can verify our prediction by calculating the enthalpy change for the reaction, ∆H° = ∆Hf°1CO1g22 + ∆Hf°1H21g22 - ∆Hf°1C(s, graphite2 - ∆Hf°1H2O1g22 = +131.3 kJ The positive sign for ∆H° indicates that the reaction is endothermic (e)  According to Le Châtelier’s principle, a decrease in the pressure causes a gaseous equilibrium to shift toward the side of the equation with the greater number of moles of gas In this case, there are mol of gas on the product side and only one on the reactant side Therefore, the pressure should be decreased to maximize the yield of the CO and H2 Chemistry Put to Work Controlling Nitric Oxide Emissions The formation of NO from N2 and O2, N21g2 + O21g2 ∆ NO1g2 Go Figure ∆H ° = 90.4 kJ [15.25] provides an interesting example of the practical importance of the fact that equilibrium constants and reaction rates change with temperature By applying Le Châtelier’s principle to this endothermic reaction and treating heat as a reactant, we deduce that an increase in temperature shifts the equilibrium in the direction of more NO The equilibrium constant Kp for formation of mol of NO from its elements at 300 K is only about * 10-15 (▶ Figure 15.15) At 2400 K, however, the equilibrium constant is about 0.05, which is 1013 times larger than the 300 K value Figure 15.15 helps explain why NO is a pollution problem In the cylinder of a modern high-compression automobile engine, the temperature during the fuel-burning part of the cycle is approximately 2400 K Also, there is a fairly large excess of air in the cylinder These conditions favor the formation of NO After combustion, however, the gases cool quickly As the temperature drops, the equilibrium in Equation 15.25 shifts to the left (because the reactant heat is being removed) However, the lower temperature also means that the reaction rate decreases, so the NO formed at 2400 K is essentially “trapped” in that form as the gas cools The gases exhausting from the cylinder are still quite hot, perhaps 1200 K At this temperature, as shown in Figure 15.15, the equilibrium constant for formation of NO is about * 10-4, much smaller than the Estimate the value of Kp at 1200 K, the exhaust gas temperature 1 N2(g) + O2(g) × 10−5 Kp Exhaust gas temperature × 10−10 NO(g) Cylinder temperature during combustion × 10−15 1000 2000 Temperature (K) ▲ Figure 15.15  Equilibrium and temperature The equilibrium constant increases with increasing temperature because the reaction is endothermic It is necessary to use a log scale for Kp because the values vary over such a large range 660 chapter 15 Chemical Equilibrium value at 2400 K However, the rate of conversion of NO to N2 and O2 is too slow to permit much loss of NO before the gases are cooled further As discussed in the “Chemistry Put to Work” box in Section 14.7, one of the goals of automotive catalytic converters is to achieve rapid conversion of NO to N2 and O2 at the temperature of the exhaust gas Some catalysts developed for this reaction are reasonably effective under the grueling conditions in automotive exhaust systems Nevertheless, scientists and engineers are continuously searching for new materials that provide even more effective catalysis of the decomposition of nitrogen oxides Chapter Summary and Key Terms The Concept of Equilibrium (Section 15.1)  A chemical reaction can achieve a state in which the forward and reverse processes are occurring at the same rate This condition is called chemical equilibrium, and it results in the formation of an equilibrium mixture of the reactants and products of the reaction The composition of an equilibrium mixture does not change with time if temperature is held constant The Equilibrium Constant (Section 15.2)  An equilib- rium that is used throughout this chapter is the reaction N21g2 + H21g2 ∆ NH31g2 This reaction is the basis of the Haber process for the production of ammonia The relationship between the concentrations of the reactants and products of a system at equilibrium is given by the law of mass action For an equilibrium equation of the form a A + b B ∆ d D + e E, the equilibriumconstant expression is written as Kc = 3D4d3E4e 3A4a3B4b where Kc is a dimensionless constant called the equilibrium constant When the equilibrium system of interest consists of gases, it is often convenient to express the concentrations of reactants and products in terms of gas pressures: Kp = 1PD2d1PE2e 1PA2a1PB2b Kc and Kp are related by the expression Kp = Kc1RT2∆n To this conversion properly, use R = 0.08206 L@atm>mol@K and temperature in kelvins Understanding and Working with Equilibrium Constants (Section 15.3)  The value of the equilibrium constant changes with temperature A large value of Kc indicates that the equilibrium mixture contains more products than reactants and therefore lies toward the product side of the equation A small value for the equilibrium constant means that the equilibrium mixture contains less products than reactants and therefore lies toward the reactant side The equilibrium-constant expression and the equilibrium constant of the reverse of a reaction are the reciprocals of those of the forward reaction If a reaction is the sum of two or more reactions, its equilibrium constant will be the product of the equilibrium constants for the individual reactions Heterogeneous Equilibria (Section 15.4)  Equilibria for which all substances are in the same phase are called homogeneous equilibria; in heterogeneous equilibria, two or more phases are present Because their activities are exactly the concentrations of pure solids and liquids are left out of the equilibrium-constant expression for a heterogeneous equilibrium Calculating Equilibrium Constants (Section 15.5)  If the concentrations of all species in an equilibrium are known, the equilibrium-constant expression can be used to calculate the equilibrium constant The changes in the concentrations of reactants and products on the way to achieving equilibrium are governed by the stoichiometry of the reaction Applications of Equilibrium Constants (Section 15.6)  The reaction quotient, Q, is found by substituting reactant and product concentrations or partial pressures at any point during a reaction into the equilibrium-constant expression If the system is at equilibrium, Q = K If Q ≠ K, however, the system is not at equilibrium When Q K, the reaction will move toward equilibrium by forming more products (the reaction proceeds from left to right); when Q K, the reaction will move toward equilibrium by forming more reactants (the reaction proceeds from right to left) Knowing the value of K makes it possible to calculate the equilibrium amounts of reactants and products, often by the solution of an equation in which the unknown is the change in a partial pressure or concentration Le Châtelier’s Principle (Section 15.7)  Le Châtelier’s prin- ciple states that if a system at equilibrium is disturbed, the equilibrium will shift to minimize the disturbing influence Therefore, if a reactant or product is added to a system at equilibrium, the equilibrium will shift to consume the added substance The effects of removing reactants or products and of changing the pressure or volume of a reaction can be similarly deduced For example, if the volume of the system is reduced, the equilibrium will shift in the direction that decreases the number of gas molecules While changes in concentration or pressure lead to shifts in the equilibrium concentrations they not change the value of the equilibrium constant, K Changes in temperature affect both the equilibrium concentrations and the equilibrium constant We can use the enthalpy change for a reaction to determine how an increase in temperature affects the equilibrium: For an endothermic reaction, an increase in temperature shifts the equilibrium to the right; for an exothermic reaction, a temperature increase shifts the equilibrium to the left Catalysts affect the speed at which equilibrium is reached but not affect the magnitude of K Learning Outcomes  After studying this chapter, you should be able to: • Explain what is meant by chemical equilibrium and how it relates to reaction rates (Section 15.1) • Write the equilibrium-constant expression for any reaction (Section 15.2) • Given the value of Kc convert to Kp and vice versa (Section 15.2) • Relate the magnitude of an equilibrium constant to the relative amounts of reactants and products present in an equilibrium mixture (Section 15.3) • Manipulate the equilibrium constant to reflect changes in the chemical equation (Section 15.3) Exercises 661 • Write the equilibrium-constant expression for a heterogeneous re- • Calculate equilibrium concentrations given the equilibrium con- • Calculate an equilibrium constant from concentration measure- • Calculate equilibrium concentrations, given the equilibrium con- • Predict the direction of a reaction given the equilibrium constant • Use Le Châtelier’s principle to predict how changing the concentra- action (Section 15.4) ments (Section 15.5) and the concentrations of reactants and products (Section 15.6) stant and all but one equilibrium concentration (Section 15.6) stant and the starting concentrations (Section 15.6) tions, volume, or temperature of a system at equilibrium affects the equilibrium position (Section 15.7) Key Equations 3D4d3E4e • Kc = • Kp = • Kp = Kc1RT2∆n [15.14] Relating the equilibrium constant based on pressures to the equilibrium constant based on concentration • Qc = [15.23] The reaction quotient The concentrations are for any time during a reaction If the concentrations are equilibrium concentrations, then Qc = Kc 3A4a3B4b 1PD2d1PE2e 1PA2a1PB2b 3D4d3E4e 3A4a3B4b [15.8] The equilibrium-constant expression for a general reaction of the type a A + b B ∆ d D + e E, the concentrations are equilibrium concentrations only [15.11] The equilibrium-constant expression in terms of equilibrium partial pressures Exercises Visualizing Concepts Potential energy 15.1 (a) Based on the following energy profile, predict whether kf kr or kf kr (b) Using Equation 15.5, predict whether the equilibrium constant for the process is greater than or less than [Section 15.1] Reactants Products Reaction progress 15.2 The following diagrams represent a hypothetical reaction A ¡ B, with A represented by red spheres and B represented by blue spheres The sequence from left to right represents the system as time passes Does the system reach equilibrium? If so, in which diagram is the system in equilibrium? [Sections 15.1 and 15.2] greater or smaller than if the volume is L and each atom/ molecule in the diagram represents mol? [Section 15.2] 15.4 The following diagram represents a reaction shown going to completion Each molecule in the diagram represents 0.1 mol and the volume of the box is 1.0 L (a) Letting A = red spheres and B = blue spheres, write a balanced equation for the reaction (b) Write the equilibrium-constant expression for the reaction (c) Calculate the value of Kc (d) Assuming that all of the molecules are in the gas phase, calculate ∆n, the change in the number of gas molecules that accompanies the reaction (e) Calculate the value of Kp [Section 15.2] 15.3 The following diagram represents an equilibrium mixture produced for a reaction of the type A + X ∆ AX Is K 15.5 Snapshots of two hypothetical reactions, A1g2 + B1g2 ∆ AB1g2 and X1g2 + Y1g2 ∆ XY1g2 at five different times ... Unit Cells and Crystal Lattices  483  Filling the Unit Cell  485 12 .3 Metallic Solids  486 The Structures of Metallic Solids  487  Close Packing  488 Alloys  4 91 12.4 Metallic Bonding  494 Electron-Sea... ISBN -13 : 978-0-3 21- 910 41- 7   ISBN -10 : 0-3 21- 910 41- 9   Chemistry Textbooks I Title   QD 31. 3.B765 2 014  540—dc23 2 013 036724 10 —CRK— 17 16 15 14 www.pearsonhighered.com Student Edition: 0-3 21- 910 41- 9... in Living Systems  10 11 A Closer Look  Charge-Transfer Color  10 28 24 The Chemistry of Life: Organic and Biological Chemistry 10 40 24 .1  General Characteristics of Organic Molecules  10 42 The

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  • Cover

  • Title Page

  • Copyright Page

  • CHEMICAL APPLICATIONS AND ESSAYS

  • Acknowledgments

  • List of Resources

  • About the Authors

  • Data-Driven Analytics: A New Direction in Chemical Education

  • Helping Students Think Like Scientists

  • Active and Visual

  • Adaptive

  • Contents

  • Preface

  • 1 Introduction: Matter and Measurement

    • 1.1 The Study of Chemistry

      • The Atomic and Molecular Perspective of Chemistry

      • Why Study Chemistry?

    • 1.2 Classifications of Matter

      • States of Matter

      • Pure Substances

      • Elements

      • Compounds

      • Mixtures

    • 1.3 Properties of Matter

      • Physical and Chemical Changes

      • Separation of Mixtures

    • 1.4 Units of Measurement

      • SI Units

      • Length and Mass

      • Temperature

      • Derived SI Units

      • Volume

      • Density

    • 1.5 Uncertainty in Measurement

      • Precision and Accuracy

      • Significant Figures

      • Significant Figures in Calculations

    • 1.6 Dimensional Analysis

      • Using Two or More Conversion Factors

      • Conversions Involving Volume

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Chemistry Put to Work: Chemistry and the Chemical Industry

    • A Closer Look: The Scientific Method

    • Chemistry Put to Work: Chemistry in the News

    • Strategies in Chemistry: Estimating Answers

    • Strategies in Chemistry: The Importance of Practice

    • Strategies in Chemistry: The Features of This Book

  • 2 Atoms, Molecules, and Ions

    • 2.1 The Atomic Theory of Matter

    • 2.2 The Discovery of Atomic Structure

      • Cathode Rays and Electrons

      • Radioactivity

      • The Nuclear Model of the Atom

    • 2.3 The Modern View of Atomic Structure

      • Atomic Numbers, Mass Numbers, and Isotopes

    • 2.4 Atomic Weights

      • The Atomic Mass Scale

      • Atomic Weight

    • 2.5 The Periodic Table

    • 2.6 Molecules and Molecular Compounds

      • Molecules and Chemical Formulas

      • Molecular and Empirical Formulas

      • Picturing Molecules

    • 2.7 Ions and Ionic Compounds

      • Predicting Ionic Charges

      • Ionic Compounds

    • 2.8 Naming Inorganic Compounds

      • Names and Formulas of Ionic Compounds

      • Names and Formulas of Acids

      • Names and Formulas of Binary Molecular Compounds

    • 2.9 Some Simple Organic Compounds

      • Alkanes

      • Some Derivatives of Alkanes

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • A Closer Look: Basic Forces

    • A Closer Look: The Mass Spectrometer

    • A Closer Look: What Are Coins Made Of?

    • Chemistry and Life: Elements Required by Living Organisms

    • Strategies in Chemistry How to Take a Test

  • 3 Chemical Reactions and Reaction Stoichiometry

    • 3.1 Chemical Equations

      • Balancing Equations

      • Indicating the States of Reactants and Products

    • 3.2 Simple Patterns of Chemical Reactivity

      • Combination and Decomposition Reactions

      • Combustion Reactions

    • 3.3 Formula Weights

      • Formula and Molecular Weights

      • Percentage Composition from Chemical Formulas

    • 3.4 Avogadro's Number and the Mole

      • Molar Mass

      • Interconverting Masses and Moles

      • Interconverting Masses and Numbers of Particles

    • 3.5 Empirical Formulas from Analyses

      • Molecular Formulas from Empirical Formulas

      • Combustion Analysis

    • 3.6 Quantitative Information from Balanced Equations

    • 3.7 Limiting Reactants

      • Theoretical and Percent Yields

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Strategies in Chemistry: Problem Solving

    • Chemistry and Life: Glucose Monitoring

    • Strategies in Chemistry: Design an Experiment

  • 4 Reactions in Aqueous Solution

    • 4.1 General Properties of Aqueous Solutions

      • Electrolytes and Nonelectrolytes

      • How Compounds Dissolve in Water

      • Strong and Weak Electrolytes

    • 4.2 Precipitation Reactions

      • Solubility Guidelines for Ionic Compounds

      • Exchange (Metathesis) Reactions

      • Ionic Equations and Spectator Ions

    • 4.3 Acids, Bases, and Neutralization Reactions

      • Acids

      • Bases

      • Strong and Weak Acids and Bases

      • Identifying Strong and Weak Electrolytes

      • Neutralization Reactions and Salts

      • Neutralization Reactions with Gas Formation

    • 4.4 Oxidation–Reduction Reactions

      • Oxidation and Reduction

      • Oxidation Numbers

      • Oxidation of Metals by Acids and Salts

      • The Activity Series

    • 4.5 Concentrations of Solutions

      • Molarity

      • Expressing the Concentration of an Electrolyte

      • Interconverting Molarity, Moles, and Volume

      • Dilution

    • 4.6 Solution Stoichiometry and Chemical Analysis

      • Titrations

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry Put to Work: Antacids

    • Strategies in Chemistry: Analyzing Chemical Reactions

  • 5 Thermochemistry

    • 5.1 Energy

      • Kinetic Energy and Potential Energy

      • Units of Energy

      • System and Surroundings

      • Transferring Energy: Work and Heat

    • 5.2 The First Law of Thermodynamics

      • Internal Energy

      • Relating &#916;E to Heat and Work

      • Endothermic and Exothermic Processes

      • State Functions

    • 5.3 Enthalpy

      • Pressure–Volume Work

      • Enthalpy Change

    • 5.4 Enthalpies of Reaction

    • 5.5 Calorimetry

      • Heat Capacity and Specific Heat

      • Constant-Pressure Calorimetry

      • Bomb Calorimetry (Constant-Volume Calorimetry)

    • 5.6 Hess's Law

    • 5.7 Enthalpies of Formation

      • Using Enthalpies of Formation to Calculate Enthalpies of Reaction

    • 5.8 Foods and Fuels

      • Foods

      • Fuels

      • Other Energy Sources

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Energy, Enthalpy, and P–V Work

    • Strategies in Chemistry: Using Enthalpy as a Guide

    • Chemistry and Life: The Regulation of Body Temperature

    • Chemistry Put to Work: The Scientific and Political Challenges of Biofuels

  • 6 Electronic Structure of Atoms

    • 6.1 The Wave Nature of Light

    • 6.2 Quantized Energy and Photons

      • Hot Objects and the Quantization of Energy

      • The Photoelectric Effect and Photons

    • 6.3 Line Spectra and the Bohr Model

      • Line Spectra

      • Bohr's Model

      • The Energy States of the Hydrogen Atom

      • Limitations of the Bohr Model

    • 6.4 The Wave Behavior of Matter

      • The Uncertainty Principle

    • 6.5 Quantum Mechanics and Atomic Orbitals

      • Orbitals and Quantum Numbers

    • 6.6 Representations of Orbitals

      • The s Orbitals

      • The p Orbitals

      • The d and f Orbitals

    • 6.7 Many-Electron Atoms

      • Orbitals and Their Energies

      • Electron Spin and the Pauli Exclusion Principle

    • 6.8 Electron Configurations

      • Hund's Rule

      • Condensed Electron Configurations

      • Transition Metals

      • The Lanthanides and Actinides

    • 6.9 Electron Configurations and the Periodic Table

      • Anomalous Electron Configurations

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Measurement and the Uncertainty Principle

    • A Closer Look: Thought Experiments and Schrödinger's Cat

    • A Closer Look: Probability Density and Radial Probability Functions

    • Chemistry and Life: Nuclear Spin and Magnetic Resonance Imaging

  • 7 Periodic Properties of the Elements

    • 7.1 Development of the Periodic Table

    • 7.2 Effective Nuclear Charge

    • 7.3 Sizes of Atoms and Ions

      • Periodic Trends in Atomic Radii

      • Periodic Trends in Ionic Radii

    • 7.4 Ionization Energy

      • Variations in Successive Ionization Energies

      • Periodic Trends in First Ionization Energies

      • Electron Configurations of Ions

    • 7.5 Electron Affinity

    • 7.6 Metals, Nonmetals, and Metalloids

      • Metals

      • Nonmetals

      • Metalloids

    • 7.7 Trends for Group 1A and Group 2A Metals

      • Group 1A: The Alkali Metals

      • Group 2A: The Alkaline Earth Metals

    • 7.8 Trends for Selected Nonmetals

      • Hydrogen

      • Group 6A: The Oxygen Group

      • Group 7A: The Halogens

      • Group 8A: The Noble Gases

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Effective Nuclear Charge

    • Chemistry Put to Work: Ionic Size and Lithium-Ion Batteries

    • Chemistry and Life: The Improbable Development of Lithium Drugs

  • 8 Basic Concepts of Chemical Bonding

    • 8.1 Lewis Symbols and the Octet Rule

      • The Octet Rule

    • 8.2 Ionic Bonding

      • Energetics of Ionic Bond Formation

      • Electron Configurations of Ions of the s- and p-Block Elements

      • Transition Metal Ions

    • 8.3 Covalent Bonding

      • Lewis Structures

      • Multiple Bonds

    • 8.4 Bond Polarity and Electronegativity

      • Electronegativity

      • Electronegativity and Bond Polarity

      • Dipole Moments

      • Differentiating Ionic and Covalent Bonding

    • 8.5 Drawing Lewis Structures

      • Formal Charge and Alternative Lewis Structures

    • 8.6 Resonance Structures

      • Resonance in Benzene

    • 8.7 Exceptions to the Octet Rule

      • Odd Number of Electrons

      • Less Than an Octet of Valence Electrons

      • More Than an Octet of Valence Electrons

    • 8.8 Strengths and Lengths of Covalent Bonds

      • Bond Enthalpies and the Enthalpies of Reactions

      • Bond Enthalpy and Bond Length

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Calculation of Lattice Energies: The Born–Haber Cycle

    • A Closer Look: Oxidation Numbers, Formal Charges, and Actual Partial Charges

    • Chemistry Put to Work: Explosives and Alfred Nobel

  • 9 Molecular Geometry and Bonding Theories

    • 9.1 Molecular Shapes

    • 9.2 The VSEPR Model

      • Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

      • Molecules with Expanded Valence Shells

      • Shapes of Larger Molecules

    • 9.3 Molecular Shape and Molecular Polarity

    • 9.4 Covalent Bonding and Orbital Overlap

    • 9.5 Hybrid Orbitals

      • sp Hybrid Orbitals

      • sp[sup(2)] and sp[sup(3)] Hybrid Orbitals

      • Hypervalent Molecules

      • Hybrid Orbital Summary

    • 9.6 Multiple Bonds

      • Resonance Structures, Delocalization, and &#960; Bonding

      • General Conclusions about &#963; and &#960; Bonding

    • 9.7 Molecular Orbitals

      • Molecular Orbitals of the Hydrogen Molecule

      • Bond Order

    • 9.8 Period 2 diatomic Molecules

      • Molecular Orbitals for Li[sub(2)] and Be[sub(2)]

      • Molecular Orbitals from 2[sub(p)] Atomic Orbitals

      • Electron Configurations for B[sub(2)] through Ne[sub(2)]

      • Electron Configurations and Molecular Properties

      • Heteronuclear Diatomic Molecules

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry and Life: The Chemistry of Vision

    • A Closer Look: Phases in Atomic and Molecular Orbitals

    • Chemistry Put to Work: Orbitals and Energy

  • 10 Gases

    • 10.1 Characteristics of Gases

    • 10.2 Pressure

      • Atmospheric Pressure and the Barometer

    • 10.3 The Gas Laws

      • The Pressure–Volume Relationship: Boyle's Law

      • The Temperature–Volume Relationship: Charles's Law

      • The Quantity–Volume Relationship: Avogadro's Law

    • 10.4 The Ideal-Gas Equation

      • Relating the Ideal-Gas Equation and the Gas Laws

    • 10.5 Further Applications of the Ideal-Gas Equation

      • Gas Densities and Molar Mass

      • Volumes of Gases in Chemical Reactions

    • 10.6 Gas Mixtures and Partial Pressures

      • Partial Pressures and Mole Fractions

    • 10.7 The Kinetic-Molecular Theory of Gases

      • Distributions of Molecular Speed

      • Application of Kinetic-Molecular Theory to the Gas Laws

    • 10.8 Molecular Effusion and Diffusion

      • Graham's Law of Effusion

      • Diffusion and Mean Free Path

    • 10.9 Real Gases: Deviations from Ideal Behavior

      • The van der Waals Equation

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Strategies in Chemistry: Calculations Involving Many Variables

    • A Closer Look: The Ideal-Gas Equation

    • Chemistry Put to Work: Gas Separations

  • 11 Liquids and Intermolecular Forces

    • 11.1 A Molecular Comparison of Gases, Liquids, and Solids

    • 11.2 Intermolecular Forces

      • Dispersion Forces

      • Dipole–Dipole Forces

      • Hydrogen Bonding

      • Ion–Dipole Forces

      • Comparing Intermolecular Forces

    • 11.3 Select Properties of Liquids

      • Viscosity

      • Surface Tension

      • Capillary Action

    • 11.4 Phase Changes

      • Energy Changes Accompanying Phase Changes

      • Heating Curves

      • Critical Temperature and Pressure

    • 11.5 Vapor Pressure

      • Volatility, Vapor Pressure, and Temperature

      • Vapor Pressure and Boiling Point

    • 11.6 Phase diagrams

      • The Phase Diagrams of H[sub(2)]O and CO[sub(2)]

    • 11.7 Liquid Crystals

      • Types of Liquid Crystals

      • Chapter Summary and Key Terms

      • Learning Outcomes

      • Exercises

      • Additional Exercises

      • Integrative Exercises

      • Design an Experiment

    • Chemistry Put to Work: Ionic Liquids

    • A Closer Look: The Clausius–Clapeyron Equation

  • 12 Solids and Modern Materials

    • 12.1 Classification of Solids

    • 12.2 Structures of Solids

      • Crystalline and Amorphous Solids

      • Unit Cells and Crystal Lattices

      • Filling the Unit Cell

    • 12.3 Metallic Solids

      • The Structures of Metallic Solids

      • Close Packing

      • Alloys

    • 12.4 Metallic Bonding

      • Electron-Sea Model

      • Molecular–Orbital Model

    • 12.5 Ionic Solids

      • Structures of Ionic Solids

    • 12.6 Molecular Solids

    • 12.7 Covalent-Network Solids

      • Semiconductors

      • Semiconductor Doping

    • 12.8 Polymers

      • Making Polymers

      • Structure and Physical Properties of Polymers

    • 12.9 Nanomaterials

      • Semiconductors on the Nanoscale

      • Metals on the Nanoscale

      • Carbons on the Nanoscale

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equation

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: X-ray Diffraction

    • Chemistry Put to Work: Alloys of Gold

    • Chemistry Put to Work: Solid-State Lighting

    • Chemistry Put to Work: Recycling Plastics

  • 13 Properties of Solutions

    • 13.1 The Solution Process

      • The Natural Tendency toward Mixing

      • The Effect of Intermolecular Forces on Solution Formation

      • Energetics of Solution Formation

      • Solution Formation and Chemical Reactions

    • 13.2 Saturated Solutions and Solubility

    • 13.3 Factors Affecting Solubility

      • Solute–Solvent Interactions

      • Pressure Effects

      • Temperature Effects

    • 13.4 Expressing Solution Concentration

      • Mass Percentage, ppm, and ppb

      • Mole Fraction, Molarity, and Molality

      • Converting Concentration Units

    • 13.5 Colligative Properties

      • Vapor-Pressure Lowering

      • Boiling-Point Elevation

      • Freezing-Point Depression

      • Osmosis

      • Determination of Molar Mass from Colligative Properties

    • 13.6 Colloids

      • Hydrophilic and Hydrophobic Colloids

      • Colloidal Motion in Liquids

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry and Life: Fat-Soluble and Water-Soluble Vitamins

    • Chemistry and Life: Blood Gases and Deep-Sea Diving

    • A Closer Look: Ideal Solutions with Two or More Volatile Components

    • A Closer Look: The Van't Hoff Factor

    • Chemistry and Life: Sickle-Cell Anemia

  • 14 Chemical Kinetics

    • 14.1 Factors that Affect Reaction Rates

    • 14.2 Reaction Rates

      • Change of Rate with Time

      • Instantaneous Rate

      • Reaction Rates and Stoichiometry

    • 14.3 Concentration and Rate Laws

      • Reaction Orders: The Exponents in the Rate Law

      • Magnitudes and Units of Rate Constants

      • Using Initial Rates to Determine Rate Laws

    • 14.4 The Change of Concentration with Time

      • First-Order Reactions

      • Second-Order Reactions

      • Zero-Order Reactions

      • Half-Life

    • 14.5 Temperature and Rate

      • The Collision Model

      • The Orientation Factor

      • Activation Energy

      • The Arrhenius Equation

      • Determining the Activation Energy

    • 14.6 Reaction Mechanisms

      • Elementary Reactions

      • Multistep Mechanisms

      • Rate Laws for Elementary Reactions

      • The Rate-Determining Step for a Multistep Mechanism

      • Mechanisms with a Slow Initial Step

      • Mechanisms with a Fast Initial Step

    • 14.7 Catalysis

      • Homogeneous Catalysis

      • Heterogeneous Catalysis

      • Enzymes

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Using Spectroscopic Methods to Measure Reaction Rates: Beer's Law

    • Chemistry Put to Work: Methyl Bromide in the Atmosphere

    • Chemistry Put to Work: Catalytic Converters

    • Chemistry and Life: Nitrogen Fixation and Nitrogenase

  • 15 Chemical Equilibrium

    • 15.1 The Concept of Equilibrium

    • 15.2 The Equilibrium Constant

      • Evaluating K[sub(c)]

      • Equilibrium Constants in Terms of Pressure, K[sub(p)]

      • Equilibrium Constants and Units

    • 15.3 Understanding and Working with Equilibrium Constants

      • The Magnitude of Equilibrium Constants

      • The Direction of the Chemical Equation and K

      • Relating Chemical Equation Stoichiometry and Equilibrium Constants

    • 15.4 Heterogeneous Equilibria

    • 15.5 Calculating Equilibrium Constants

    • 15.6 Applications of Equilibrium Constants

      • Predicting the Direction of Reaction

      • Calculating Equilibrium Concentrations

    • 15.7 Le Châtelier's Principle

      • Change in Reactant or Product Concentration

      • Effects of Volume and Pressure Changes

      • Effect of Temperature Changes

      • The Effect of Catalysts

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry Put to Work: The Haber Process

    • Chemistry Put to Work: Controlling Nitric Oxide Emissions

  • 16 Acid–Base Equilibria

    • 16.1 Acids and Bases: A Brief Review

    • 16.2 Br&#216;nsted–Lowry Acids and Bases

      • The H[sup(+)] Ion in Water

      • Proton-Transfer Reactions

      • Conjugate Acid–Base Pairs

      • Relative Strengths of Acids and Bases

    • 16.3 The Autoionization of Water

      • The Ion Product of Water

    • 16.4 The pH Scale

      • pOH and Other "p" Scales

      • Measuring pH

    • 16.5 Strong Acids and Bases

      • Strong Acids

      • Strong Bases

    • 16.6 Weak Acids

      • Calculating K[sub(a)] from pH

      • Percent Ionization

      • Using K[sub(a)] to Calculate pH

      • Polyprotic Acids

    • 16.7 Weak Bases

      • Types of Weak Bases

    • 16.8 Relationship between K[sub(a)] and K[sub(b)]

    • 16.9 Acid–Base Properties of Salt Solutions

      • An Anion's Ability to React with Water

      • A Cation's Ability to React with Water

      • Combined Effect of Cation and Anion in Solution

    • 16.10 Acid–Base Behavior and Chemical Structure

      • Factors That Affect Acid Strength

      • Binary Acids

      • Oxyacids

      • Carboxylic Acids

    • 16.11 Lewis Acids and Bases

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry Put to Work: Amines and Amine Hydrochlorides

    • Chemistry and Life: The Amphiprotic Behavior of Amino Acids

  • 17 Additional Aspects of Aqueous Equilibria

    • 17.1 The Common-Ion Effect

    • 17.2 Buffers

      • Composition and Action of Buffers

      • Calculating the pH of a Buffer

      • Buffer Capacity and pH Range

      • Addition of Strong Acids or Bases to Buffers

    • 17.3 Acid–Base Titrations

      • Strong Acid–Strong Base Titrations

      • Weak Acid–Strong Base Titrations

      • Titrating with an Acid–Base Indicator

      • Titrations of Polyprotic Acids

    • 17.4 Solubility Equilibria

      • The Solubility-Product Constant, K[sub(sp)]

      • Solubility and K[sub(sp)]

    • 17.5 Factors That Affect Solubility

      • Common-Ion Effect

      • Solubility and pH

      • Formation of Complex Ions

      • Amphoterism

    • 17.6 Precipitation and Separation of Ions

      • Selective Precipitation of Ions

    • 17.7 Qualitative Analysis for Metallic Elements

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry and Life: Blood as a Buffered Solution

    • A Closer Look: Limitations of Solubility Products

    • Chemistry and Life: Ocean Acidification

    • Chemistry and Life: Tooth Decay and Fluoridation

  • 18 Chemistry of the Environment

    • 18.1 Earth's Atmosphere

      • Composition of the Atmosphere

      • Photochemical Reactions in the Atmosphere

      • Ozone in the Stratosphere

    • 18.2 Human Activities and Earth's Atmosphere

      • The Ozone Layer and Its Depletion

      • Sulfur Compounds and Acid Rain

      • Nitrogen Oxides and Photochemical Smog

      • Greenhouse Gases: Water Vapor, Carbon Dioxide, and Climate

    • 18.3 Earth's Water

      • The Global Water Cycle

      • Salt Water: Earth's Oceans and Seas

      • Freshwater and Groundwater

    • 18.4 Human Activities and Water Quality

      • Dissolved Oxygen and Water Quality

      • Water Purification: Desalination

      • Water Purification: Municipal Treatment

    • 18.5 Green Chemistry

      • Supercritical Solvents

      • Greener Reagents and Processes

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Other Greenhouse Gases

    • A Closer Look: The Ogallala Aquifer—A Shrinking Resource

    • A Closer Look: Fracking and Water Quality

  • 19 Chemical Thermodynamics

    • 19.1 Spontaneous Processes

      • Seeking a Criterion for Spontaneity

      • Reversible and Irreversible Processes

    • 19.2 Entropy and the Second Law of Thermodynamics

      • The Relationship between Entropy and Heat

      • &#916;S for Phase Changes

      • The Second Law of Thermodynamics

    • 19.3 The Molecular Interpretation of Entropy and the Third Law of Thermodynamics

      • Expansion of a Gas at the Molecular Level

      • Boltzmann's Equation and Microstates

      • Molecular Motions and Energy

      • Making Qualitative Predictions about &#916;S

      • The Third Law of Thermodynamics

    • 19.4 Entropy Changes in Chemical Reactions

      • Entropy Changes in the Surroundings

    • 19.5 Gibbs Free Energy

      • Standard Free Energy of Formation

    • 19.6 Free Energy and Temperature

    • 19.7 Free Energy and the Equilibrium Constant

      • Free Energy under Nonstandard Conditions

      • Relationship between &#916;G&#176; and K

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: The Entropy Change When a Gas Expands Isothermally

    • Chemistry and Life: Entropy and Human Society

    • A Closer Look: What's "Free" about Free Energy?

    • Chemistry and Life: Driving Nonspontaneous Reactions: Coupling Reactions

  • 20 Electrochemistry

    • 20.1 Oxidation States and Oxidation–Reduction Reactions

    • 20.2 Balancing Redox Equations

      • Half-Reactions

      • Balancing Equations by the Method of Half-Reactions

      • Balancing Equations for Reactions Occurring in Basic Solution

    • 20.3 Voltaic Cells

    • 20.4 Cell Potentials Under Standard Conditions

      • Standard Reduction Potentials

      • Strengths of Oxidizing and Reducing Agents

    • 20.5 Free Energy and Redox Reactions

      • Emf, Free Energy, and the Equilibrium Constant

    • 20.6 Cell Potentials Under Nonstandard Conditions

      • The Nernst Equation

      • Concentration Cells

    • 20.7 Batteries and Fuel Cells

      • Lead–Acid Battery

      • Alkaline Battery

      • Nickel–Cadmium and Nickel–Metal Hydride Batteries

      • Lithium-Ion Batteries

      • Hydrogen Fuel Cells

    • 20.8 Corrosion

      • Corrosion of Iron (Rusting)

      • Preventing Corrosion of Iron

    • 20.9 Electrolysis

      • Quantitative Aspects of Electrolysis

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Electrical Work

    • Chemistry and Life: Heartbeats and Electrocardiography

    • Chemistry Put to Work: Batteries for Hybrid and Electric Vehicles

    • Chemistry Put to Work: Electrometallurgy of Aluminum

  • 21 Nuclear Chemistry

    • 21.1 Radioactivity and Nuclear Equations

      • Nuclear Equations

      • Types of Radioactive Decay

    • 21.2 Patterns of Nuclear Stability

      • Neutron-to-Proton Ratio

      • Radioactive Decay Chains

      • Further Observations

    • 21.3 Nuclear Transmutations

      • Accelerating Charged Particles

      • Reactions Involving Neutrons

      • Transuranium Elements

    • 21.4 Rates of Radioactive Decay

      • Radiometric Dating

      • Calculations Based on Half-Life

    • 21.5 Detection of Radioactivity

      • Radiotracers

    • 21.6 Energy Changes in Nuclear Reactions

      • Nuclear Binding Energies

    • 21.7 Nuclear Power: Fission

      • Nuclear Reactors

      • Nuclear Waste

    • 21.8 Nuclear Power: Fusion

    • 21.9 Radiation in the Environment and Living Systems

      • Radiation Doses

      • Radon

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Key Equations

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry and Life: Medical Applications of Radiotracers

    • A Closer Look: The Dawning of the Nuclear Age

    • A Closer Look: Nuclear Synthesis of the Elements

    • Chemistry and Life: Radiation Therapy

  • 22 Chemistry of the Nonmetals

    • 22.1 Periodic Trends and Chemical Reactions

      • Chemical Reactions

    • 22.2 Hydrogen

      • Isotopes of Hydrogen

      • Properties of Hydrogen

      • Production of Hydrogen

      • Uses of Hydrogen

      • Binary Hydrogen Compounds

    • 22.3 Group 8A: The Noble Gases

      • Noble-Gas Compounds

    • 22.4 Group 7A: The Halogens

      • Properties and Production of the Halogens

      • Uses of the Halogens

      • The Hydrogen Halides

      • Interhalogen Compounds

      • Oxyacids and Oxyanions

    • 22.5 Oxygen

      • Properties of Oxygen

      • Production of Oxygen

      • Uses of Oxygen

      • Ozone

      • Oxides

      • Peroxides and Superoxides

    • 22.6 The Other Group 6A Elements: S, Se, Te, and Po

      • General Characteristics of the Group 6A Elements

      • Occurrence and Production of S, Se, and Te

      • Properties and Uses of Sulfur, Selenium, and Tellurium

      • Sulfides

      • Oxides, Oxyacids, and Oxyanions of Sulfur

    • 22.7 Nitrogen

      • Properties of Nitrogen

      • Production and Uses of Nitrogen

      • Hydrogen Compounds of Nitrogen

      • Oxides and Oxyacids of Nitrogen

    • 22.8 The Other Group 5A Elements: P, As, Sb, and Bi

      • General Characteristics of the Group 5A Elements

      • Occurrence, Isolation, and Properties of Phosphorus

      • Phosphorus Halides

      • Oxy Compounds of Phosphorus

    • 22.9 Carbon

      • Elemental Forms of Carbon

      • Oxides of Carbon

      • Carbonic Acid and Carbonates

      • Carbides

    • 22.10 The Other Group 4A Elements: Si, Ge, Sn, and Pb

      • General Characteristics of the Group 4A Elements

      • Occurrence and Preparation of Silicon

      • Silicates

      • Glass

      • Silicones

    • 22.11 Boron

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: The Hydrogen Economy

    • Chemistry and Life: Nitroglycerin, Nitric Oxide, and Heart Disease

    • Chemistry and Life: Arsenic in Drinking Water

    • Chemistry Put to Work: Carbon Fibers and Composites

  • 23 Transition Metals and Coordination Chemistry

    • 23.1 The Transition Metals

      • Physical Properties

      • Electron Configurations and Oxidation States

      • Magnetism

    • 23.2 Transition-Metal Complexes

      • The Development of Coordination Chemistry: Werner's Theory

      • The Metal–Ligand Bond

      • Charges, Coordination Numbers, and Geometries

    • 23.3 Common Ligands in Coordination Chemistry

      • Metals and Chelates in Living Systems

    • 23.4 Nomenclature and Isomerism in Coordination Chemistry

      • Isomerism

      • Structural Isomerism

      • Stereoisomerism

    • 23.5 Color and Magnetism in Coordination Chemistry

      • Color

      • Magnetism of Coordination Compounds

    • 23.6 Crystal-Field Theory

      • Electron Configurations in Octahedral Complexes

      • Tetrahedral and Square-Planar Complexes

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • A Closer Look: Entropy and the Chelate Effect

    • Chemistry and Life: The Battle for Iron in Living Systems

    • A Closer Look: Charge-Transfer Color

  • 24 The Chemistry of Life: Organic and Biological Chemistry

    • 24.1 General Characteristics of Organic Molecules

      • The Structures of Organic Molecules

      • The Stabilities of Organic Substances

      • Solubility and Acid–Base Properties of Organic Substances

    • 24.2 Introduction to Hydrocarbons

      • Structures of Alkanes

      • Structural Isomers

      • Nomenclature of Alkanes

      • Cycloalkanes

      • Reactions of Alkanes

    • 24.3 Alkenes, Alkynes, and Aromatic Hydrocarbons

      • Alkenes

      • Alkynes

      • Addition Reactions of Alkenes and Alkynes

      • Aromatic Hydrocarbons

      • Stabilization of &#960; Electrons by delocalization

      • Substitution Reactions

    • 24.4 Organic Functional Groups

      • Alcohols

      • Ethers

      • Aldehydes and Ketones

      • Carboxylic Acids and Esters

      • Amines and Amides

    • 24.5 Chirality in Organic Chemistry

    • 24.6 Introduction to Biochemistry

    • 24.7 Proteins

      • Amino Acids

      • Polypeptides and Proteins

      • Protein Structure

    • 24.8 Carbohydrates

      • Disaccharides

      • Polysaccharides

    • 24.9 Lipids

      • Fats

      • Phospholipids

    • 24.10 Nucleic Acids

    • Chapter Summary and Key Terms

    • Learning Outcomes

    • Exercises

    • Additional Exercises

    • Integrative Exercises

    • Design an Experiment

    • Chemistry Put to Work: Gasoline

    • A Closer Look: Mechanism of Addition Reactions

    • Strategies in Chemistry: What Now?

  • Appendices

    • A: Mathematical Operations

    • B: Properties of Water

    • C: Thermodynamic Quantities for Selected Substances AT 298.15 K (25 °C)

    • D: Aqueous Equilibrium Constants

    • E: Standard Reduction Potentials at 25 °C

  • Answers to Selected Exercises

  • Answers to Give It Some Thought

  • Answers to Go Figure

  • Answers to Selected Practice Exercises

  • Glossary

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    • P

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  • Photo/Art Credits

  • Index

    • A

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