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Starting with just a few basic principles of probability and the distribution of energy, Introduction to Molecular Thermodynamics takes students on an adventure into the inner workings of the molecular world like no other, from probability to Gibbs energy and beyond, following a logical step-by-step progression of ideas.
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Starting with just a few basic principles of probability and the distribution of energy, Introduction to Molecular Thermodynamics takes students on an adventure into the inner workings of the molecular world like no other, from probability to Gibbs energy and beyond, following a logical step-by-step progression of ideas.
Produktdetails
- Produktdetails
- Verlag: Macmillan Education / University Science Books
- 2008
- Seitenzahl: 318
- Erscheinungstermin: 21. Juli 2008
- Englisch
- Abmessung: 269mm x 205mm x 22mm
- Gewicht: 788g
- ISBN-13: 9781891389498
- ISBN-10: 1891389491
- Artikelnr.: 24437137
- Verlag: Macmillan Education / University Science Books
- 2008
- Seitenzahl: 318
- Erscheinungstermin: 21. Juli 2008
- Englisch
- Abmessung: 269mm x 205mm x 22mm
- Gewicht: 788g
- ISBN-13: 9781891389498
- ISBN-10: 1891389491
- Artikelnr.: 24437137
Robert Hanson is a Professor of Chemistry at St. Olaf College, in Northfield, Minnesota, where he has been teaching since 1986. Trained as an organic chemist with Gilbert Stork at Columbia University, he shares a patent with 2001 Nobel Prize winner K.Barry Sharpless for the asymmetric epoxidation of allylic alcohols. His interest in thermodynamics goes back to early training at the California Institute of Technology, from which he got a B.S. degree in 1979. He spends his occasional moments of free time playing the violin in a community orchestra, piloting gliders, and designing new Sudoku strategies.
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 .-
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 .-