Introduction to Molecular Thermodynamics / Edition 1

by Robert M. Hanson, Susan Green

Overview

Starting with just a few basic principles of probabilityand the distribution of energy, this book takes students (and faculty!) on anadventure into the inner workings of the molecular world like no other.
“I wish I had learned thermodynamics this way!” That’s what the authors hear all the time from instructors using Introduction to Molecular Thermodynamics. Starting with just a few basic principles of probability and the distribution of energy, the book takes students (and faculty!) on an adventure into the inner workings of the molecular world like no other. Made to fit into a standard second-semester of a traditional first-year chemistry course, or as a supplement for more advanced learners, the book takes the reader from probability to Gibbs energy and beyond, following a logical step-by-step progression of ideas, each just a slight expansion of the previous. Filled with examples ranging from casinos to lasers, from the “high energy bonds” of ATP to endangered coral reefs, Introduction to Molecular Thermodynamics hits the mark for students and faculty alike who have an interest in understanding the world around them in molecular terms.

Product Details

ISBN-13:	9781891389498
Publisher:	MIT Press
Publication date:	07/21/2008
Edition description:	2008
Pages:	318
Product dimensions:	8.00(w) x 10.50(h) x (d)

About the Author

Robert M. Hanson (Author) 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.Susan Green (Author) Susan Green has had the privilege of being both a student and a professor at St. Olaf College in Northfield, Minnesota where she was first introduced to the idea of teaching thermodynamics to first-year students. She trained as a physical chemist at the University of Minnesota studying the vibrational and electronic structure of small metal oxides as well as trying her hand at analytical chemistry. When she in not chasing after her two children with the help of her husband Hans, she can be found with a book.

Table of Contents

Preface To the InstructorTo the Student: How to Study Thermodynamics Acknowledgments CHAPTER 1 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 Distributions1.10 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 Chatelier’s Principle Is Based on Probability 1.15 Determining Equilibrium Amounts and Constants Based onProbability 1.16 Summary CHAPTER 2 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 Energy2.8 The Most Probable Distribution Is the Boltzmann Distribution2.9 The Effect of Temperature vi Contents2.10 The Effect of Energy Level Separation 2.11 Why Is the Boltzmann Distribution the Most Probable? 2.12 Determining the Population of the Lowest Level2.13 Estimating the Fraction of Particles That Will React 2.14 Estimating How Many Levels Are Populated 2.15 Summary CHAPTER 3 Energy Levels in Real Chemical Systems 3.1 Historical Perspective 3.2 The Modern Viewpoint3.3 Planck, Einstein, and de Broglie 3.4 The “Wave” Can Be Thought of in Terms of Probability 3.5 Electronic Energy 3.6 Vibrational Energy3.7 Rotational Energy3.8 Translational Energy 3.9 Putting It All Together3.10 Chemical Reactions 3.11 Chemical Equilibrium and Energy Levels3.12 Color, Fluorescence, and Phosphorescence 3.13 Lasers and Stimulated Emission 3.14 Summary CHAPTER 4 Internal Energy (U) and the First Law4.1 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 + w4.6 Macroscopic Heat and Heat Capacity: q = CT 4.7 Macroscopic Work: w = −P V 4.8 In Chemical Reactions, Work Can Be Ignored 4.9 Calorimeters Allow the Direct Determination of U4.10 Don’t Forget the Surroundings! 4.11 Engines: Converting Heat into Work 4.12 Biological and Other Forms of Work4.13 Summary CHAPTER 5 Bonding and Internal Energy 5.1 The Chemical Bond5.2 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 Energy5.6 Estimating rU for Chemical Reactions Using Bond Dissociation Energies 5.7 Using Bond Dissociation Energies to Understand Chemical Reactions5.8 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 CHAPTER 6 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 SummaryCHAPTER 7 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 Entropy 7.10 Determining the Standard Change in Entropy for a Chemical Reaction 7.11 Another Way to Look at S 7.12 Summary CHAPTER 8 The Effect of Pressure and Concentrationon 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 andSx = 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 Quotient8.11 The Evaporation of Liquid Water 8.12 A Microscopic Picture of Pressure Effects on Entropy 8.13 Summary CHAPTER 9 Enthalpy (H) and the Surroundings 9.1 Heat Is Not a State Function 9.2 The Definition of Enthalpy: H = U + P V 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 CHAPTER 10 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 CHAPTER 11 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 Equilibrium11.6 Le Chatelier’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 is R ln K/Q 11.11 Beyond Ideality: “Activity” 11.12 Summary CHAPTER 12 Applications of Gibbs Energy: Phase Changes12.1 Review 12.2 Evaporation and Boiling 12.3 Sublimation and Vapor Deposition 12.4 Triple Points 12.5 Critical Points and Phase Diagrams12.6 Solubility: 0 = rH◦ − T (rS◦ − R ln[X]sat) 12.7 Impure Liquids: S = S◦ − R ln x 12.8 Summary CHAPTER 13 Applications of Gibbs Energy: Electrochemistry 13.1 Introduction 13.2 Review: Gibbs Energy and Entropy 13.3 Including Internal Energy and Electrical Work in the Big Picture 13.4 Electrical Work Is Limited by the Gibbs Energy13.5 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 = −nF Ecell 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 = −nF E◦ cell,T + RT ln K 13.12 Actual Cell Voltages and the Nernst Equation: −nF Ecell = −nF E◦ cell,T + RT ln Q 13.13 Detailed Examples13.14 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

Editorial Reviews

"Hanson and Green offer a very valuable work on molecular thermodynamics. Highly recommended." —Choice

"I am excited to see this material introduced in a first-year course. Statistics, as the driving force behind chemical equilibria and thermodynamics, is a profound concept that most students only get a taste of in physical chemistry. This book provides an excellent way to introduce these ideas at an early stage." —J. Matthew Hutchison, Swarthmore College

"Throughout, Introduction to Molecular Thermodynamics is a friendly and appealing book. There are not many textbooks that are a pleasure to read, and this is one of them. I would encourage its consideration for course adoption, even if you have to make up a new course." —J. Chem. Ed.

From the Publisher

Throughout, Introduction to Molecular Thermodynamics is a friendly and appealing book. There are not many textbooks that are a pleasure to read, and this is one of them. I would encourage its consideration for course adoption, even if you have to make up a new course.

J. Chem. Ed.

I am excited to see this material introduced in a first-year course. Statistics, as the driving force behind chemical equilibria and thermodynamics, is a profound concept that most students only get a taste of in physical chemistry. This book provides an excellent way to introduce these ideas at an early stage.

J. Matthew Hutchison

Hanson and Green offer a very valuable work on molecular thermodynamics. Highly recommended.

Choice