A Problem-Solving Approach to Aquatic Chemistry
Description
Enables civil and environmental engineers to understand the theory and application of aquatic equilibrium chemistry
The second edition of A Problem-Solving Approach to Aquatic Chemistry provides a detailed introduction to aquatic equilibrium chemistry, calculation methods for systems at equilibrium, applications of aquatic chemistry, and chemical kinetics. The text directly addresses two required ABET program outcomes in environmental engineering: "... chemistry (including stoichiometry, equilibrium, and kinetics)" and "material and energy balances, fate and transport of substances in and between air, water, and soil phases."
The book is very student-centered, with each chapter beginning with an introduction and ending with a summary that reviews the chapter's main points. To aid in reader comprehension, important terms are defined in context and key ideas are summarized. Many thought-provoking discussion questions, worked examples, and end of chapter problems are also included. Each part of the text begins with a case study, a portion of which is addressed in each subsequent chapter, illustrating the principles of that chapter. In addition, each chapter has an Historical Note exploring connections with the people and cultures connected to topics in the text.
A Problem-Solving Approach to Aquatic Chemistry includes:
* Fundamental concepts, such as concentration units, thermodynamic basis of equilibrium, and manipulating equilibria
* Solutions of chemical equilibrium problems, including setting up the problems and algebraic, graphical, and computer solution techniques
* Acid-base equilibria, including the concepts of acids and bases, titrations, and alkalinity and acidity
* Complexation, including metals, ligands, equilibrium calculations with complexes, and applications of complexation chemistry
* Oxidation-reduction equilibria, including equilibrium calculations, graphical approaches, and applications
* Gas-liquid and solid-liquid equilibrium, with expanded coverage of the effects of global climate change
* Other topics, including chemical kinetics of aquatic systems, surface chemistry, and integrative case studies
For advanced/senior undergraduates and first-year graduate students in environmental engineering courses, A Problem-Solving Approach to Aquatic Chemistry serves as an invaluable learning resource on the topic, with a variety of helpful learning elements included throughout to ensure information retention and the ability to apply covered concepts in practical settings.
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James N. Jensen is Professor in the Department of Civil, Structural and Environmental Engineering at the State University of New York at Buffalo.
Content
Preface xix
Part I Fundamental Concepts
1 Getting Started with the Fundamental Concepts 3
1.1 Introduction 3
1.2 Why Calculate Chemical Species Concentrations at Equilibrium? 3
1.3 Primary Variables: Importance of pH and pe 6
1.4 Properties of Water 7
1.5 Part I Roadmap 9
1.6 Chapter Summary 9
1.7 Part I Case Study: Can Methylmercury be Formed Chemically in Water? 10
Chapter Key Ideas 11
Chapter Glossary 11
Historical Note: S.P.L. Sørensen and the p in pH 11
Chapter References 12
2 Concentration Units 13
2.1 Introduction 13
2.2 Units Analysis 13
2.3 Molar Concentration Units 14
2.4 Mass Concentration Units 19
2.5 Dimensionless Concentration Units 24
2.6 Equivalents 25
2.7 Review of Units Interconversion 26
2.8 Common Concentration Units in the Gas Phase 27
2.9 Common Concentration Units in the Solid Phase 28
2.10 Activity 28
2.11 Chapter Summary 30
2.12 Part I Case Study: Can Methylmercury Be Formed Chemically in Water? 30
Chapter Key Ideas 31
Chapter Glossary 31
Historical Note: Amadea Avogadro and Avogadro's Number 32
Problems 33
Chapter References 34
3 Thermodynamic Basis of Equilibrium 35
3.1 Introduction 35
3.2 Thermodynamic Properties 36
3.3 Why Do We Need Thermodynamics to Calculate Species Concentrations? 39
3.4 Thermodynamic Laws 42
3.5 Gibbs Free Energy 45
3.6 Properties of Thermodynamic Functions 48
3.7 Changes in Thermodynamic Properties During Chemical Reactions 50
3.8 Relating Gibbs Free Energy to Species Concentrations 55
3.9 Chemical Equilibrium and the Equilibrium Constant 60
3.10 Chapter Summary 62
3.11 Part I Case Study: Can Methylmercury Be Formed Chemically in Water? 63
Chapter Key Ideas 63
Chapter Glossary 64
Historical Note: Josiah Willard Gibbs 66
Problems 67
Chapter References 68
4 Manipulating Equilibrium Expressions 69
4.1 Introduction 69
4.2 Chemical and Mathematical Forms of Equilibria 69
4.3 Units of Equilibrium Constants 73
4.4 Reversing Equilibria 75
4.5 Effects of Stoichiometry 76
4.6 Adding Equilibria 78
4.7 Creating Equilibria 81
4.8 Chapter Summary 87
4.9 Part I Case Study: Can Methylmercury Be Formed Chemically in Water? 87
Chapter Key Ideas 88
Chapter Glossary 89
Historical Note: Henri- Louis Le Châtelier and Le Châtelier's Principle 89
Problems 90
Chapter References 91
Part II Solving Chemical Equilibrium Problems
5 Getting Started withSolving Equilibrium Problems 95
5.1 Introduction 95
5.2 A Framework for Solving Chemical Equilibrium Problems 95
5.3 Introduction to Defining the Chemical System 97
5.4 Introduction to Enumerating Chemical Species 98
5.5 Introduction to Defining the Constraints on Species Concentrations 98
5.6 Part II Roadmap 100
5.7 Chapter Summary 100
5.8 Part II Case Study: Have You Had Your Zinc Today? 101
Chapter Key Ideas 101
Chapter Glossary 101
Historical Note: "Active Mass" and Familial Relations 102
Chapter References 103
6 Setting Up Chemical Equilibrium Calculations 105
6.1 Introduction 105
6.2 Defining the Chemical System 105
6.3 Enumerating Chemical Species 106
6.4 Defining the Constraints on Species Concentrations 112
6.5 Review of Procedures for Setting up Equilibrium Systems 120
6.6 Concise Mathematical Form for Equilibrium Systems 121
6.7 Chapter Summary 122
6.8 Part II Case Study: Have You Had Your Zinc Today? 123
Chapter Key Ideas 126
Chapter Glossary 126
Historical Note: Salts of the Ocean 127
Problems 129
Chapter References 130
7 Algebraic Solutions to Chemical Equilibrium Problems 131
7.1 Introduction 131
7.2 Background on Algebraic Solutions 131
7.3 Method of Substitution 133
7.4 Method of Approximation 139
7.5 Chapter Summary 148
7.6 Part II Case Study: Have You Had Your Zinc Today? 148
Chapter Key Ideas 152
Historical Note: What's in a Name? 152
Problems 153
8 Graphical Solutions to Chemical Equilibrium Problems 155
8.1 Introduction 155
8.2 Log Concentration and pC- pH Diagrams 156
8.3 Using pC- pH Diagrams with More Complex Systems 162
8.4 Special Shortcuts for Monoprotic Acids 167
8.5 When Graphical Methods Fail: The Proton Condition 171
8.6 Chapter Summary 177
8.7 Part II Case Study: Have You Had Your Zinc Today? 178
Chapter Key Ideas 179
Chapter Glossary 180
Historical Note: Who Was First? 180
Problems 181
Chapter Reference 182
9 Computer Solutions to Chemical Equilibrium Problems 183
9.1 Introduction 183
9.2 Chapter Problem 183
9.3 Spreadsheet Solutions 184
9.4 Equilibrium Calculation Software 188
9.5 Nanoql SE 190
9.6 The Tableau Method and Other Equilibrium Calculation Apps 192
9.7 Visual MINTEQ 201
9.8 Chapter Summary 202
9.9 Part II Case Study: Have You Had Your Zinc Today? 202
Chapter Key Ideas 203
Chapter Glossary 203
Historical Note: ALGOL to VBA 203
Problems 204
Chapter References 205
Part III Acid-Base Equilibria in Homogenous Aqueous Systems
10 Getting Started with Acid-Base Equilibrium in Homogenous Aqueous Systems 209
10.1 Introduction 209
10.2 Homogeneous Systems 209
10.3 Types of Reactions in Homogeneous Systems 211
10.4 The Wonderful World of Acids and Bases 212
10.5 Part III Roadmap 215
10.6 Chapter Summary 215
10.7 Part III Case Study: Acid Rain 215
Chapter Key Ideas 216
Chapter Glossary 216
Historical Note: "An Evil of the Highest Magnitude" 217
Chapter References 218
11 Acids and Bases 219
11.1 Introduction 219
11.2 Definitions of Acids and Bases 219
11.3 Acid and Base Strength 223
11.4 Polyprotic Acids 228
11.5 Alpha Values (Distribution Functions) 236
11.6 Chapter Summary 239
11.7 Part II Case Study: Acid Rain 239
Chapter Key Ideas 241
Chapter Glossary 242
Historical Note: Why Is a Base a Base? 242
Problems 243
Addendum: A Surprising Exact Solution 245
Chapter References 248
12 Acid-Base Titrations 249
12.1 Introduction 249
12.2 Principles of Acid-Base Titrations 250
12.3 Equivalence Points 255
12.4 Titration of Polyprotic Acids 265
12.5 Buffers 269
12.6 Interpretation of Acid-Base Titration Curves with Complex Mixtures 277
12.7 Chapter Summary 279
12.8 Part III Case Study: Acid Rain 280
Chapter Key Ideas 282
Chapter Glossary 283
Historical Note: Mohr about Titrations 284
Problems 285
Chapter References 286
13 Alkalinity and Acidity 287
13.1 Introduction 287
13.2 Alkalinity and the Acid Neutralizing Capacity 287
13.3 Alkalinity and the Charge Balance 290
13.4 Characteristics of Alkalinity and Acidity 292
13.5 Using the Definitions of Alkalinity to Solve Problems 302
13.6 Effects of Other Weak Acids and Bases on Alkalinity 308
13.7 Chapter Summary 310
13.8 Part III Case Study: Acid Rain 310
Chapter Key Ideas 311
Chapter Glossary 312
Historical Note: Can You Pass the Litmus Test? 313
Problems 314
Chapter References 316
Part IV Other Equilibria in Homogenous Aqueous Systems
14 Getting Started with Other Equilibria in Homogeneous Aqueous Systems 319
14.1 Introduction 319
14.2 Electron- Sharing Reactions 319
14.3 Electron Transfer 321
14.4 Part IV Roadmap 323
14.5 Chapter Summary 323
14.6 Part IV Case Study: Which Form of Copper Plating Should You Use? 323
Chapter Key Ideas 324
Historical Note: Hauptvalenz and Nebenvalenz 324
Chapter References 325
15 Complexation 327
15.1 Introduction 327
15.2 Metals 327
15.3 Ligands 330
15.4 Equilibrium Calculations with Complexes 335
15.5 Systems with Several Metals and Ligands 345
15.6 Applications of Complexation Chemistry 357
15.7 Chapter Summary 361
15.8 Part IV Case Study: Which Form of Copper Plating Should You Use? 362
Chapter Key Ideas 364
Chapter Glossary 365
Historical Note: British Anti- Lewisite - A WMD- Inspired Ligand 366
Problems 368
Chapter References 369
16 Oxidation and Reduction 371
16.1 Introduction 371
16.2 A Few Definitions 371
16.3 Balancing Redox Reactions 374
16.4 Which Redox Reactions Occur? 383
16.5 Redox Thermodynamics and Oxidant and Reductant Strength 386
16.6 Manipulating Half Reactions 393
16.7 Algebraic Equilibrium Calculations in Systems Undergoing Electron Transfer 396
16.8 Graphical Representations of Systems Undergoing Electron Transfer 399
16.9 Applying Redox Equilibrium Calculations to the Real World 413
16.10 Chapter Summary 414
16.11 Part IV Case Study: Which Form of Copper Plating Should You Use? 415
Chapter Key Ideas 417
Chapter Glossary 418
Historical Note: Walther Hermann Nernst 419
Problems 420
Chapter References 422
Part V Heterogeneous Systems
17 Getting Started with Heterogeneous Systems 425
17.1 Introduction 425
17.2 Equilibrium Exchange Between Gas and Aqueous Phases 426
17.3 Equilibrium Exchange Between Solid and Aqueous Phases 427
17.4 Part V Roadmap 428
17.5 Chapter Summary 428
17.6 Part V Case Study: The Killer Lakes 428
Chapter Key Ideas 429
Historical Note: "A Spirit Case and a Gasogene" 429
Chapter References 430
18 Gas-Liquid Equilibria 431
18.1 Introduction 431
18.2 Raoult's Law and Henry's Law 431
18.3 Equilibrium Calculations Involving Gas-Liquid Equilibria 438
18.4 Dissolved Carbon Dioxide 449
18.5 Chapter Summary 456
18.6 Part V Case Study: The Killer Lakes 456
Chapter Key Ideas 457
Chapter Glossary 458
Historical Note: A Brief History of Carbon Dioxide 459
Problems 460
Chapter References 462
19 Solid-Liquid Equilibria 463
19.1 Introduction 463
19.2 Saturation and the Activity of Pure Solids 463
19.3 Equilibrium Calculations with Solid-Liquid Equilibria 466
19.4 Factors Affecting Metal Solubility 474
19.5 Solubility of Calcium Carbonate 480
19.6 Models for the Acid-Base Chemistry of Natural Waters 484
19.7 Chapter Summary 491
19.8 Part V Case Study: The Killer Lakes 491
Chapter Key Ideas 492
Chapter Glossary 493
Historical Note: Black Smokers and White Smokers 493
Problems 494
Addendum: Information Requirements 497
Chapter References 498
Part VI Beyond Dilute Solutions at Equilibrium
20 Getting Started with Beyond Dilute Solutions at Equilibrium 501
20.1 Introduction 501
20.2 Extensions to Nonideal and Nonstandard Conditions 502
20.3 The Strange World of Surfaces 503
20.4 Nonequilibrium Conditions 504
20.5 Integrated Case Studies 504
20.6 Part VI Roadmap 505
20.7 Chapter Summary 505
Chapter Key Ideas 506
Chapter Glossary 506
Historical Note: "Harcourt, Come to Me!" 506
Chapter References 507
21 Thermodynamics Revisited: The Effects of Ionic Strength, Temperature, and Pressure 509
21.1 Introduction 509
21.2 Effects of Ionic Strength 510
21.3 Effects of Temperature on Equilibrium Constants 522
21.4 Effects of Pressure on Equilibrium Constants 528
21.5 Chapter Summary 529
Chapter Key Ideas 530
Chapter Glossary 531
Historical Note: Jacobus Henricus van't Hoff 531
Problems 532
Chapter References 534
22 Aquatic Chemistry ofSurfaces 535
22.1 Introduction 535
22.2 Nomenclature 535
22.3 Isotherms and Ion Exchange 538
22.4 Introduction to Surface Complexation Modeling 543
22.5 Surface Complexation Modeling 546
22.6 Chapter Summary 552
Chapter Key Ideas 553
Chapter Glossary 553
Historical Note: From "Cat's Cradle" to the "Swiss Model" to Surface Complexation Modeling 554
Problems 555
Addendum: The Freundlich Isotherm and Adsorption Equilibria 556
Chapter References 557
23 Chemical Kinetics of Aquatic Systems 559
23.1 Introduction 559
23.2 The Need for Chemical Kinetics 560
23.3 Reaction Rates 561
23.4 Common Rate Expressions 569
23.5 More Complex Kinetic Forms 577
23.6 Effects of Temperature and Ionic Strength on Kinetics 582
23.7 Chapter Summary 587
Chapter Key Ideas 587
Chapter Glossary 588
Historical Note: Arrhenius, Chick, and Foote 589
Problems 590
Chapter References 592
24 Putting It All Together: Integrated Case Studies in Aquatic Chemistry 593
24.1 Introduction 593
24.2 Integrated Case Study 1: Metal Finishing 594
24.3 Integrated Case Study 2: Oxidation of Fe(+II) by Oxygen 598
24.4 Integrated Case Study 3: Inorganic Mercury Chemistry in Natural Waters 603
24.5 Integrated Case Study 4: Phosphate Buffers 607
24.6 Integrated Case Study 5: Global Climate Change 610
24.7 Chapter Summary 613
Historical Note: Stumm and Morgan 614
Chapter References 614
Appendix A: Background Information 617
A.1 Introduction 617
A.2 Chemical Principles 617
A.3 Mathematical Principles 619
A.4 Spreadsheet Skills 620
Chapter Key Ideas 623
Chapter Glossary 623
Useful Physical Constants and Conversions 623
Appendix B: Equilibrium Revisited 625
B.1 Introduction 625
B.2 Equilibrium and Steady State 625
B.3 Energy Minimization and Algebraic Solutions 628
Chapter Key Ideas 631
Chapter Glossary 631
Appendix C: Summary of Procedures 633
C.1 Oxidation States and Balancing Reactions 633
C.2 Setting Up Chemical Equilibrium Systems (Section 6.5) 634
C.3 Algebraic Solution Techniques 635
C.4 Graphical Solutions 635
C.5 Computer Solutions: Tableau Method (Section 9.6.6) 637
C.6 Acid-Base Titrations 638
C.7 Complexation (Section 15.4.4) 638
C.8 Ionic Strength Effects (Section 21.2.7) 639
C.9 Surface Complexation Modeling Method (Section 22.5.4) 639
C.10 Chemical Kinetics (Section 23.3.4) 639
Appendix D: Selected Equilibrium Constants 641
Chapter References 651
Appendix E: Animations and Example Spreadsheet Files 653
E.1 Introduction to Animations 653
E.2 Variation of the Equilibrium pH of a Monoprotic Acid Solution with the Total Acid Concentration and K a 653
E.3 How to Draw pC- pH Diagrams for Monoprotic Acids 654
E.4 Equilibrium pH During the Titration of a Monoprotic Acid with a Strong Base 656
E.5 Spreadsheet Examples 657
Appendix F: Nanoql SE 661
F.1 Introduction 661
F.2 Entering Your System 661
F.3 How to Solve Systems and Vary System Parameters 663
F.4 Nanoql SE Examples 666
Chapter Reference 668
Index 669
Biographical Index 677
CHAPTER 1
Getting Started with the Fundamental Concepts
1.1 INTRODUCTION
The first part of this text reviews fundamental concepts that must be mastered prior to learning how to calculate and interpret species concentrations in aquatic systems. In this chapter, the motivation for studying chemical species and a few general principles concerning aquatic systems are presented.
In Section 1.2, the motivation for why engineers and scientists are interested in individual chemical species concentrations at equilibrium will be discussed. Important water quality parameters, called primary variables, are introduced in Section 1.3. It is impossible to study water chemistry without a little knowledge of the structure of water. A few of the unique properties of water will be explored in Section 1.4, especially as they relate to the chemical reactions that occur in water. In Section 1.5, a road map for Part I of the text is presented and discussed. Finally, the Part I case study is presented at the conclusion of this chapter. Before beginning Part I of the text, you are urged to review the chemistry background material in Appendix A (Section A.2).
1.2 WHY CALCULATE CHEMICAL SPECIES CONCENTRATIONS AT EQUILIBRIUM?
1.2.1 Overview
The bulk of this book is dedicated to the calculation of species concentrations at equilibrium. The focus here is on chemical species that undergo chemical reactions; in other words, reactive species. More specifically, the emphasis here is on chemical species which react with water. Reactions with water are called hydrolysis reactions (from the Greek hydor water + lyein to loosen). When substances react with water, numerous other compounds can be formed. Indeed, the richness of aquatic chemistry stems from the large number of substances that react not only with water but also with the products of myriad other hydrolysis reactions.
hydrolysis reactions :
reactions with water
Key idea: Hydrolysis reactions produce a wealth of dissolved chemical species
This richness is illustrated in Figure 1.1. Inputs of chemical species (from aqueous discharge, runoff, atmospheric deposition, and dissolution from sediments) react with water to form hydrolysis products. The hydrolysis products and input chemicals react further to increase the complexity of aquatic systems.
FIGURE 1.1 Complexity of Aquatic Systems
(rain cloud image: OpenClipart-Vectors/Pixabay)
So why so much interest in calculating the equilibrium concentrations of chemical species? This question is really two questions. First, why calculate the concentrations of individual chemical species? Second, why calculate species concentrations at equilibrium?
1.2.2 Importance of Individual Chemical Species
Throughout this text, you will see that knowing the concentrations of individual chemical species is critically important in analyzing many environmental problems. At first glance, this statement may not make sense. After all, many environmental regulations are based on total concentrations of classes of compounds rather than on the concentrations of individual species. Should you be more concerned about the total amount of mercury or phenol or ammonia than about individual species stemming from the hydrolysis of mercury, phenol, or ammonia?
Key idea: The ability to calculate the concentrations of individual chemical species is critically important in analyzing many environmental problems
In fact, you will find that individual species frequently are more important. Three general examples will illustrate this point. First, adverse impacts on human health and ecosystem viability may be due to only one or several of a large number of related hydrolysis products. A prime example is the transition metals (such as mercury, copper, zinc, cadmium, iron, and lead), in which toxicity varies dramatically among the hydrolysis products. Another example is cyanide. Hydrogen cyanide (HCN) is much more toxic to humans than cyanide ion (CN-).
Second, the success of engineered treatment systems may depend on knowledge of the concentrations of key individual species. Since hydrolysis products vary in their physical, chemical, and biochemical properties, the design and operation of treatment processes depend on quantitative models for the concentrations of individual chemical species. For example, the addition of gaseous chlorine to wastewater for disinfection results in the formation of many chemical species (including HOCl, OCl-, NH2Cl, and NHCl2), each of which differs in its ability to inactivate (i.e., kill) microorganisms.
Third, individual species vary greatly in how readily they cross cell membranes or cell walls and are assimilated by aquatic biota. Thus, understanding the cycling of trace nutrients in the aquatic environment (and humankind's impact on nutrient cycling) requires knowledge of the concentrations of individual chemical species.
As an example of the importance of the concentrations of individual chemical species, consider the soup created when copper sulfate crystals, CuSO4(s),1 are added to a reservoir for algae control. The CuSO4(s) dissolves in water to form a copper-containing ion (called the aquo cupric ion) and sulfate. The structure of the aquo cupric ion is usually abbreviated as Cu2+. The Cu2+ ions thus formed react very quickly with water to form a number of hydrolysis products, including CuOH+, Cu(OH)2(aq), Cu(OH)3-, Cu(OH)42-, and Cu2(OH)22+. Under certain chemical conditions, copper may precipitate as CuO(s). As you spread the copper sulfate from the back of a boat, carbon dioxide in the atmosphere is equilibrating with the reservoir water to form its own hydrolysis products. The hydrolysis products of carbon dioxide are H2CO3, HCO3-, and CO32-. The aquo cupric ion will react to some extent with the hydrolysis products of carbon dioxide to form CuCO3(aq), Cu(CO3)22-, and perhaps even solids containing copper and carbonate (CO32-). By adding one copper compound to a natural water body, you may be faced with accounting for as many as 10 copper-containing species even in a relatively simple chemical model.
Of course, the real world is even more complex. The reservoir water contains many more species that can react with copper than just hydroxide (OH-) and carbonate. A realistic model for copper in the reservoir would have to include the reactions of Cu2+ with (among other chemical species) chloride, amino acids, ammonia, particulates, and microorganisms. In reality, the act of throwing copper sulfate crystals into the reservoir will produce dozens of chemical species containing copper.
Key idea: Doses depend on both the required concentration of the target individual species and the chemistry of the water
Why should you care that copper sulfate forms many copper-containing species in a lake? Remember that copper is added to kill algae. It is well-established that copper toxicity to algae is due almost entirely to one chemical species: Cu2+ (Jackson and Morgan 1978). Thus, to determine the copper sulfate dose, you must be able to calculate the concentration of Cu2+ after a certain amount of copper sulfate is added to the reservoir. Since the Cu2+ concentration usually is exceedingly small, this is akin to counting needles of Cu2+ in this haystack of copper-containing species. In practice, you would back-calculate the copper sulfate dose required to achieve a required level of Cu2+. Even if two reservoirs had the same amounts of algae, different water chemistries in the reservoirs may lead to very different copper sulfate doses to achieve the same Cu2+ concentration. The chemistry of the water determines how the required concentration of one species (Cu2+) is translated back into a copper sulfate dose.
The process of relating a dose to a required concentration of an individual chemical species is illustrated in Figure 1.2. The arrows in Figure 1.2 indicate chemical reactions that must be included in a mathematical model to allow for the determination of the copper sulfate dose. In this text, you will learn the tools to make quantitative decisions to solve similar problems in the aquatic environment.
FIGURE 1.2 Qualitative Relationship Between the Dose Required and End Species Concentrations Desired
1.2.3 Importance of Equilibrium
An entire chapter of this book (Chapter 3) is devoted to developing the thermodynamic basis of equilibrium. For the present, you can think of the equilibrium state as the condition in which the concentrations of all chemical species do not change with time. To impose equilibrium on a chemical system, the interesting and important time-dependent nature of chemical concentrations are excluded. The study of the rates of chemical reaction is called chemical kinetics and is covered in Chapter 23. Why constrain the discussion mainly to the equilibrium state here, with shorter coverage of chemical kinetics?
chemical kinetics: the study of chemical reaction rates
There are two reasons for focusing on equilibrium. First, many of the chemical reactions you will examine in this text are fast. For example, the reaction of H+ and OH- to form water occurs on the time scale of 10-5 s at...
