Introduction to Molecular Thermodynamics

Preface.- To the Instructor.- To the Student: How to Study Thermodynamics.- Acknowledgments.- PART I: PROBABILITY, DISTRIBUTIONS, AND EQUILIBRIUM 1.1 Chemical Change.- 1.2 Chemical Equilibrium.- 1.3 Probability Is '(Ways of getting x) / (Ways total)'.- 1.4 AND Probability Multiplies.- 1.5 OR Probability Adds.- 1.6 AND and OR Probability Can Be Combined.- 1.7 The Probability of 'Not X' Is One Minus the Probability of 'X'.- 1.8 Probability Can Be Interpreted Two Ways.- 1.9 Distributions.- 1.10 For Large Populations, We Approximate.- 1.11 Relative Probability and Fluctuations.- 1.12 Equilibrium and the Most Probable Distribution.- 1.13 Equilibrium Constants Describe the Most Probable Distribution.- 1.14 Le Ch^atelier's Principle Is Based on Probability.- 1.15 Determining Equilibrium Amounts and Constants Based on Probability.- 1.16 Summary.- PART II: THE DISTRIBUTION OF ENERGY 2.1 Real Chemical Reactions.- 2.2 Temperature and Heat Energy.- 2.3 The Quantized Nature of Energy.- 2.4 Distributions of Energy Quanta in Small Systems.- 2.5 Calculating W Using Combinations.- 2.6 Why Equations 2.1 and 2.2 Work.- 2.7 Determining the Probability of a Particular Distribution of Energy.- 2.8 The Most Probable Distribution Is the Boltzmann Distribution.- 2.9 The Effect of Temperature.- 2.10 The Effect of Energy Level Separation.- 2.11 Why Is the Boltzmann Distribution the Most Probable?.- 2.12 Determining the Population of the Lowest Level.- 2.13 Estimating the Fraction of Particles That Will React.- 2.14 Estimating How Many Levels Are Populated.- 2.15 Summary.- PART III: ENERGY LEVELS IN REAL CHEMICAL SYSTEMS 3.1 Historical Perspective.- 3.2 The Modern Viewpoint.- 3.3 Planck, Einstein, and de Broglie.- 3.4 The 'Wave' Can Be Thought of in Terms of Probability.- 3.5 Electronic Energy.- 3.6 Vibrational Energy.- 3.7 Rotational Energy.- 3.8 Translational Energy.- 3.9 Putting It All Together.- 3.10 Chemical Reactions.- 3.11 Chemical Equilibrium and Energy Levels.- 3.12 Color, Fluorescence, and Phosphorescence.- 3.13 Lasers and Stimulated Emission.- 3.14 Summary.- PART IV: INTERNAL ENERGY (U) AND THE FIRST LAW 4.1 The Internal Energy (U).- 4.2 Internal Energy (U) Is a State Function.- 4.3 Microscopic Heat (q) and Work (w).- 4.4 'Heating' vs. 'Adding Heat'.- 4.5 The First Law of Thermodynamics: U = q + w.- 4.6 Macroscopic Heat and Heat Capacity: q = CT.- 4.7 Macroscopic Work: w =?PV.- 4.8 In Chemical Reactions, Work Can Be Ignored.- 4.9 Calorimeters Allow the Direct Determination of U.- 4.10 Don't Forget the Surroundings!.- 4.11 Engines: Converting Heat into Work.- 4.12 Biological and Other Forms of Work.- 4.13 Summary.- PART V: BONDING AND INTERNAL ENERGY 5.1 The Chemical Bond.- 5.2 Hess's Law.- 5.3 The Reference Point for Changes in Internal Energy Is 'Isolated Atoms'.- 5.4 Two Corollaries of Hess's Law.- 5.5 Mean Bond Dissociation Energies and Internal Energy.- 5.6 Estimating rU for Chemical Reactions Using Bond Dissociation Energies.- 5.7 Using Bond Dissociation Energies to Understand Chemical Reactions.- 5.8 The 'High-Energy Phosphate Bond' and Other Anomalies.- 5.9 Computational Chemistry and the Modern View of Bonding.- 5.10 Beyond Covalent Bonding.- 5.11 Summary.- PART VI: THE EFFECT OF TEMPERATURE ON EQUILIBRIUM 6.1 Chemical Reactions as Single Systems: Isomerizations.- 6.2 The Temperature Effect on Isomerizations.- 6.3 K vs. T for Evenly Spaced Systems.- 6.4 Experimental Data Can Reveal Energy Level Information.- 6.5 Application to Real Chemical Reactions.- 6.6 The Solid/Liquid Problem.- 6.7 Summary.- PART VII: ENTROPY (S) AND THE SECOND LAW 7.1 Energy Does Not Rule.- 7.2 The Definition of Entropy: S = k ln W.- 7.3 Changes in Entropy: S = k ln(W2/W1).- 7.4 The Second Law of Thermodynamics: Suniverse > 0.- 7.5 Heat and Entropy Changes in the Surroundings: Ssur = qsur/T .- 7.6 Measuring Entropy Changes.- 7.7 Standard Molar Entropy: S? .- 7.8 Entropy Comparisons Are Informative.- 7.9 The Effect of Ground State Electronic Degeneracy on Molar .- 7.10 Determining the Standard Change in Entropy for a Chemical Reaction.- 7.11 Another Way to Look at S.- 7.12 Summary.- PART VIII: THE EFFECT OF PRESSURE AND CONCENTRATION ON ENTROPY 8.1 Introduction.- 8.2 Impossible? or Just Improbable?.- 8.3 Ideal Gases and Ideal Solutions.- 8.4 The Volume Effect on Entropy: S = nR ln(V2/V1).- 8.5 The Entropy of Mixing Is Just the Entropy of Expansion.- 8.6 The Pressure Effect for Ideal Gases: S =?nR ln(P2/P1).- 8.7 Concentration Effect for Solutions: S =?nR ln([X]2/[X]1).- 8.8 Adjustment to the Standard State: Sx = S?x ? R ln Px and Sx = S?x ? R ln[X].- 8.9 The Reaction Quotient: rS = rS? ? R ln Q.- 8.10 Solids and Liquids Do Not Appear in the Reaction Quotient.- 8.11 The Evaporation of Liquid Water.- 8.12 A Microscopic Picture of Pressure Effects on Entropy.- 8.13 Summary.- PART IX: ENTHALPY (H) AND THE SURROUNDINGS 9.1 Heat Is Not a State Function.- 9.2 The Definition of Enthalpy: H = U + PV.- 9.3 Standard Enthalpies of Formation, fH?.- 9.4 Using Hess's Law and fH? to Get rH? for a Reaction.- 9.5 Enthalpy vs. Internal Energy.- 9.6 High Temperature Breaks Bonds.- 9.7 Summary.- PART X: GIBBS ENERGY (G) 10.1 The Second Law Again, with a Twist.- 10.2 The Definition of Gibbs Energy: G = H ? T S.- 10.3 Plotting G vs. T (G-T Graphs).- 10.4 Comparing Two or More Substances Using G-T Graphs.- 10.5 Equilibrium Is Where rG = 0.- 10.6 The 'Low Enthalpy/High Entropy Rule'.- 10.7 A Quantitative Look at Melting Points: 0 = fusH ? TmpfusS.- 10.8 The Gibbs Energy of a Gas Depends upon Its Pressure.- 10.9 Vapor Pressure, Barometric Pressure, and Boiling.- 10.10 Summary.- PART XI: THE EQUILIBRIUM CONSTANT (K ) 11.1 Introduction.- 11.2 The Equilibrium Constant.- 11.3 Determining the Values of rH? and rS? Experimentally.- 11.4 The Effect of Temperature on Keq.- 11.5 A Qualitative Picture of the Approach to Equilibrium.- 11.6 Le Ch^atelier's Principle Revisited.- 11.7 Determining Equilibrium Pressures and Concentrations.- 11.8 Equilibration at Constant Pressure (optional).- 11.9 Standard Reaction Gibbs Energies, rG?T.- 11.10 The Potential for Change in Entropy of the Universe isR ln K/Q.- 11.11 Beyond Ideality: 'Activity'.- 11.12 Summary.- PART XII: APPLICATIONS OF GIBBS ENERGY: PHASE CHANGES 12.1 Review.- 12.2 Evaporation and Boiling.- 12.3 Sublimation and Vapor Deposition.- 12.4 Triple Points.- 12.5 Critical Points and Phase Diagrams.- 12.6 Solubility: 0 = rH? ? T (rS? ? R ln[X]sat).- 12.7 Impure Liquids: S = S? ? R ln x.- 12.8 Summary.- PART XIII: APPLICATIONS OF GIBBS ENERGY: ELECTROCHEMISTRY 13.1 Introduction.- 13.2 Review: Gibbs Energy and Entropy.- 13.3 Including Internal Energy and ElectricalWork in the Big Picture.- 13.4 Electrical Work Is Limited by the Gibbs Energy.- 13.5 The Gibbs Energy Change Can Be Positive.- 13.6 The Electrical Connection: ?G = Qelec × Ecell = I × t × Ecell.- 13.7 The Chemical Connection: Qrxn = n × F.- 13.8 Gibbs Energy and Cell Potential: rG=?nFEcell.- 13.9 Standard State for Cell Potential: E?cell,T.- 13.10 Using Standard Reduction Potentials to Predict Reactivity.- 13.11 Equilibrium Constants from Cell Potentials: 0=?nFE?cell,T + RT ln K.- 13.12 Actual Cell Voltages and the Nernst Equation: ?nFEcell =?nFE?cell,T + RT ln Q.- 13.13 Detailed Examples.- 13.14 Summary.- APPENDIX A Symbols and Constants.- APPENDIX B Mathematical Tricks.- APPENDIX C Table of Standard Reduction Potentials.- APPENDIX D Table of Standard Thermodynamic Data (25°C and 1 bar).- APPENDIX E Thermodynamic Data for the Evaporation of Liquid Water.- Answers to Selected Exercises.- Index.