Introduction to Molecular Thermodynamics (Paperback)

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 .-