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Summary:
Publisher Summary 1
For undergraduate science and engineering students, Metiu (chemistry and physics, U. of California, Santa Barbara) provides a text on thermodynamics and physical chemistry that uses software to perform the complicated mathematics. The CD-ROM that accompanies the book contains this software in the Mathematica and Mathcad Workbooks, used for solving the problems and illustrating calculations. Calculations are also explained in the text. Topics discussed are temperature, pressure, molar volume, and equilibrium; the equation of a state; thermodynamic transformations; work; heat; path-dependent and path-independent quantities; the first and second laws; free energies; calculating the change of entropy; enthalpy and energy change; thermochemistry; chemical potential; mixtures; chemical equilibrium; phase transitions; electrolyte solutions; and galvanic cells. Metiu is also author of the three other volumes in the four-volume series, including Physical Chemistry: Statistical Mechanics, Physical Chemistry: Kinetics, and Physical Chemistry: Quantum Mechanics. Annotation 漏2006 Book News, Inc., Portland, OR (booknews.com)
目录
Table Of Contents:
Preface xxv
How to use the workbooks, exercises, and problems xxxi
Temperature, pressure, molar volume, and equilibrium 1(20)
Introduction 1(1)
System and environment 1(1)
Temperature and thermal equilibrium 2(5)
Thermometer 2(1)
Temperature scales 3(4)
Pressure and mechanical equilibrium 7(4)
Calculating the pressure in a cylinder 8(1)
The units of pressure 8(1)
The units Torr, atm, bar, and psi 9(1)
Conversion between units 10(1)
Volume, density, and molar volume 11(1)
Intensive and extensive quantities 12(1)
Equilibrium 13(1)
Equilibrium and environment 13(1)
Supplement 1.1 An excerpt from Fahrenheit's article 14(2)
Supplement 1.2 Origin of the pressure units atmosphere and Torr 16(3)
Problems 19(2)
The equation of state 21(12)
Introduction 21(3)
The state of a gas or a liquid at equilibrium 21(1)
The equation of state 22(1)
Solids are different 22(2)
The ideal gas equation 24(4)
Units 24(2)
When and why the ideal gas law is valid 26(2)
The van der Waals equation of state 28(1)
Accurate equations of state 29(1)
Summary and perspective 30(1)
Problems 31(2)
How to use the equation of state 33(26)
Calculate pressure when you know molar volume and temperature 33(3)
Why we calculate pressure 33(1)
How to calculate pressure 34(2)
Calculate molar volume when you know pressure and temperature 36(4)
Why we calculate the molar volume 36(2)
An example of molar volume calculation 38(2)
Calculate temperature when you know molar volume and pressure 40(1)
When such calculations are needed 40(1)
Summary of Chapters 1--3 40(1)
Problems 41(4)
Supplement 3.1 How to get your own equation of state 45(11)
Least squares fitting 50(1)
Why we use the square of the error 50(1)
Minimizing the global error 51(2)
Determining the parameters in the van der Waals equation 53(3)
Problems 56(3)
Thermodynamic transformations 59(10)
Definition and examples of thermodynamic transformations 59(2)
Non-equilibrium transformations 61(1)
Initial and final states 61(1)
The path of the transformation 62(1)
Equilibrium transformations 62(2)
Why we study equilibrium transformations 64(1)
Supplement 4.1. More about equilibrium transformations and their paths 65(4)
Two equilibrium transformations with the same initial and final state but different paths 65(4)
Work 69(22)
Introduction 69(1)
The definition of work 70(3)
The sign convention 71(1)
Units 71(1)
Work is an extensive quantity 72(1)
No change in volume, no work 72(1)
Work is performed against an opposing force 73(1)
There are many ways of exchanging work 73(1)
How to calculate the work in a finite transformation 73(3)
The work performed in a finite transformation 73(3)
Work performed in an isothermal transformation 76(5)
What an isothermal transformation is 76(1)
The work performed in an isothermal expansion 77(1)
A numerical calculation of isothermal work 78(3)
The work performed in an isobaric transformation 81(2)
What an isobaric transformation is 81(1)
A numerical calculation of the work performed in an isobaric transformation 81(2)
The work performed in a transformation depends on the path 83(4)
Work is not a function of state 86(1)
Problems 87(4)
Heat 91(52)
What is heat? 91(4)
The caloric theory 91(2)
What is transferred from a hot body to a cold one 93(1)
How to measure the amount of heat 94(1)
Heat and work are equivalent 94(1)
The amount of heat has a sign 95(1)
Thermal coefficients: definitions 95(4)
Heat exchanged at constant pressure 96(1)
Heat exchanged at constant volume 97(1)
The heat exchanged when pressure changes and temperature is constant 97(1)
The heat exchanged in a general transformation 98(1)
Information regarding the thermal coefficients 99(7)
Heat capacity at constant pressure 99(2)
The temperature dependence of Cp 101(2)
The pressure dependence of Cp 103(1)
Heat capacity at constant volume 104(2)
Calculations of the heat exchanged in simple transformations 106(8)
The heat transferred when a system is heated at constant volume 108(1)
Heat exchange when two bodies of different temperatures are brought into contact 109(2)
The number of moles 111(3)
Problems 114(1)
Supplement 6.1. Heat is a form of motion: an experiment in boring cannon 115(7)
Supplement 6.2. Joseph Black, heat capacity 122(5)
Supplement 6.3. A more extensive look at heat theory and calculations 127(13)
Information regarding lp and lv 129(1)
lp and lv for an ideal gas 130(1)
The heat exchanged in an infinitesimal transformation 131(1)
The heat exchanged in a finite transformation 131(1)
The heat exchanged in an isothermal compression: general equation 132(1)
The heat exchanged in an isothermal compression of an ideal gas 133(1)
The heat exchanged in an isothermal compression of a van der Waals gas 134(3)
The case of a real gas: implementation 137(1)
A numerical evaluation of qT 138(2)
Problems 140(3)
Reversible and irreversible transformations 143(6)
Introduction 143(1)
Definition 144(1)
Heat transfer 144(1)
Diffusion 145(1)
An equilibrium transformation is reversible 146(3)
Path-dependent and path-independent quantities 149(16)
Path-independent line integrals 149(12)
Line integrals in thermodynamics 149(4)
Most line integrals depend on the path (the line) 153(1)
Path-independent integrals: an example 154(2)
Exact differentials 156(1)
Path-independent line integrals: theorems 157(4)
Applications to thermodynamics 161(4)
Work and heat depend on path 161(1)
A proof that ∫δW depends on path 162(1)
Exact differentials and functions of state 163(2)
First and second laws of thermodynamics 165(14)
The formulation of the laws 165(1)
Introduction 165(1)
The First Law 165(1)
The Second Law 166(1)
The Third Law 166(1)
Common features of energy and entropy 167(3)
∫dU and ∫dS are path-independent 167(1)
U and S are functions of state 168(1)
Maxwell relations 168(1)
Adding a constant to entropy or energy causes no measurable change 169(1)
U and S are extensive properties 170(1)
A few comments about the First Law 170(5)
Energy conservation 170(4)
The First Law today 174(1)
A few comments about the Second Law 175(4)
The direction of a transformation 175(2)
Impossible processes 177(1)
Why this is useful 177(2)
Helmholtz and Gibbs free energies 179(8)
Introduction 179(1)
Why do we need other functions besides entropy? 179(3)
A convenient form for the First and Second Laws 179(1)
Second Law for a transformation keeping U and V constant 180(2)
Helmholtz free energy A 182(2)
Second Law in terms of changes in V and T 182(1)
Transformations at constant T and V 183(1)
Using A: a hint 184(1)
Gibbs free energy G 184(3)
The definition of Gibbs free energy 184(1)
Some properties of Gibbs free energy 185(1)
The change of G in a transformation in which T and p are held constant 185(1)
Using Gibbs free energy: a hint 186(1)
How to calculate the change of entropy in an equilibrium transformation 187(32)
Introduction 187(1)
Which variable to use 187(1)
The variables T and p: theory 188(5)
The change ds of entropy when pressure changes by dp and temperature by dT 189(1)
Notation for partial derivatives 189(2)
Maxwell's method 191(1)
The derivative (∂s/∂p)T 191(1)
The derivative (∂s/∂T)p 192(1)
Combine these results to get an expression for ds 193(1)
The change of entropy in a transformation in which T and p change by a finite amount 193(1)
The variables T and v: theory 193(3)
The calculation of (∂s/∂v)T 194(1)
The change in entropy when both T and v change 195(1)
The change in entropy in a finite transformation 195(1)
Calculations of entropy change in various transformations 196(9)
Entropy change in an isobaric transformation 196(2)
The change of entropy in an isothermal transformation 198(4)
The change of entropy in a general transformation 202(2)
The change of entropy: numerical results 204(1)
How to use the tables of data to calculate entropy changes 204(1)
Problems 205(3)
Supplement 11.1. Using Maxwell's method to derive useful equations 208(3)
Using Maxwell equations 208(3)
Supplement 11.2. Obtaining new equations by changing variables 211(4)
An equation for (∂v/∂T)p 212(1)
A more general method for calculating (∂v/∂T)p 213(1)
An equation for Cp
Cv 214(1)
Supplement 11.3. Adiabatic transformations 215(4)
The transformation path in an adiabatic transformation 216(1)
The equation for the final temperature 217(1)
An example of adiabatic compression 218(1)
Enthalpy and energy change during a thermodynamic transformation 219(28)
Introduction 219(4)
Heat and enthalpy 221(1)
The connection between the enthalpy change and heat 222(1)
Heat and energy 222(1)
How to calculate the enthalpy change in a transformation: theory 223(4)
The change of enthalpy in an infinitesimal transformation 223(2)
The change of enthalpy in a finite transformation 225(1)
Choosing a path 225(2)
How to calculate the enthalpy change in a transformation: an example 227(11)
The change of enthalpy ΔhA when T changes and p is constant (path A) 228(1)
The change of enthalpy ΔhB when pressure changes and temperature is held constant 229(1)
The first difficulty 229(2)
The second difficulty 231(2)
The final result for ΔhB 233(1)
An example of a calculation of ΔhB 234(1)
The order of magnitude of ΔhB 234(1)
The use of tables to calculate enthalpy changes with temperature 235(3)
Supplement 12.1. Energy changes in a thermodynamic transformation 238(2)
The change of energy in an infinitesimal transformation in which T and v are changed 238(1)
The change of energy in a finite transformation in which T and v are changed 238(1)
Another way of calculating energy changes 239(1)
Supplement 12.2. Isenthalpic transformations 240(7)
This transformation takes place without a change of enthalpy 241(1)
The change of temperature caused by an isenthalpic transformation 242(1)
Isenthalpic transformations are used for cooling 243(3)
Ideal gas 246(1)
Thermochemistry 247(38)
Introduction 247(1)
Definition of the heat of reaction 248(3)
Two reactions used as examples 248(1)
The definition of the heat of reaction 249(1)
The initial state 249(1)
The final state 249(1)
Sign convention 250(1)
The presentation of data: standard state 250(1)
The connection between the heat of reaction and the enthalpy of the participants 251(4)
Enthalpy change in a reaction 252(3)
The dependence of the heat of reaction on temperature and pressure 255(3)
Enthalpy change in a transformation from σi to σf 255(1)
The heat of reaction at Tf and pf 256(1)
The change of the heat of reaction with temperature 257(1)
The change of heat of reaction with pressure 257(1)
An example: the heat of reaction for ammonia synthesis 258(6)
The heat of reaction ΔH 259(1)
The temperature dependence of the heat of reaction 260(2)
The heat of reaction at 623.15 K and 394.8 atm 262(2)
Calculating the heat of a reaction from heats of formation or combustion 264(10)
Heats of formation: definition 266(1)
The connection between the heat of reaction and the heats of formation of the compounds 266(3)
Where to get heats of formation 269(1)
The use of heats of combustion to calculate heats of reaction 270(4)
Supplement 13.1. Calculating the heat of one reaction from the heats of other reactions 274(7)
Problems 281(4)
The change of chemical potential during an equilibrium transformation 285(16)
Introduction 285(1)
The change of chemical potential calculated with Eq. 14.4 286(9)
How to evaluate Δμ1 ≡ μ(Tf, pi) - μ(Ti, pi) 287(1)
An example of evaluation of Δμ1 ≡ μ(Tf, pi) - μ(Ti, pi) 288(1)
How to evaluate Δμ2 ≡ μ(Tf, pf) - μ(Tf, pi) 289(4)
Calculating Δμ = μ(130 K, 600 atm) - μ(298.15 K. 1 atm) 293(2)
Additional material about chemical potential 295(3)
The pressure dependence of the chemical potential of an ideal gas 295(1)
Fugacity 295(1)
The dependence of chemical potential on temperature 296(1)
Calculating μ from μ = h - Ts 297(1)
Problems 298(3)
The chemical potential of a compound in a mixture 301(10)
General remarks 301(2)
Infinitesimal transformations 302(1)
The chemical potential of a compound in a mixture 303(8)
The change of Gibbs free energy when I change temperature, pressure, and composition 303(1)
Change of variables 304(3)
The chemical potential of ideal mixtures 307(1)
The partial pressure of a gas in an ideal mixture 307(1)
The chemical potential of a gas in an ideal mixture 308(1)
A few words about Josiah Willard Gibbs 309(2)
Mixtures: partial molar quantities and activities 311(24)
Partial molar quantities 311(16)
The correct formula 314(2)
Partial molar volume is an intensive quantity 316(1)
Other partial molar quantities 317(1)
Partial molar enthalpy and the heat of mixing 318(1)
Chemical potential as a partial molar quantity: the Gibbs--Duhem equation 319(2)
Equations similar to the Gibbs--Duhem equation for other partial molar quantities 321(1)
How to determine partial molar quantities from measurements 322(3)
Relations among partial molar quantities 325(2)
The composition dependence of chemical potential: ideal solutions 327(3)
The definition of an ideal solution 327(1)
The change of volume when we mix compounds to form an ideal solution 327(1)
The enthalpy of an ideal mixture 328(1)
The heat of mixing to form an ideal solution 329(1)
The entropy of mixing to form an ideal solution 329(1)
Chemical potential of real solutions: the activity and the reference potential 330(5)
Chemical equilibrium 335(24)
Introduction 335(4)
Reactants and products 336(1)
Stable and metastable chemical equilibrium 337(2)
The extent of reaction and the composition of a reacting mixture 339(7)
The extent of reaction 339(2)
Mass conservation 341(1)
Molar fractions 341(1)
Examples of the use of the extent of reaction 342(3)
Some properties of the extent of reaction 345(1)
The equilibrium conditions and the direction of a reaction 346(6)
The equilibrium conditions 346(1)
The direction of a chemical reaction 347(1)
Chemical affinity of a reaction 348(1)
Geometric interpretation 349(2)
How to use affinity to answer practical chemistry questions 351(1)
The equilibrium condition in terms of chemical potentials and equilibrium constant 352(7)
The reaction affinity in terms of chemical potential 352(1)
The affinity of a reaction for ideal mixtures 353(3)
The equilibrium conditions for real mixtures 356(3)
Chemical equilibrium: the connection between the equilibrium constant and composition 359(24)
Introduction 359(2)
How to calculate the equilibrium constant from a measurement of the equilibrium concentration of one species 361(2)
How to calculate the equilibrium composition when you know the equilibrium constant and the initial number of moles 363(2)
The dependence of the extent of reaction on the initial number of moles 365(1)
The role of the initial number of moles 365(6)
Supplement 18.1. Another example of composition calculations 371(2)
Supplement 18.2. A second example and error analysis 373(6)
Supplement 18.3. The direction of a reaction 379(4)
A metastable state 381(2)
Chemical equilibrium: how to calculate K from ΔG0 =
RT In K 383(26)
Calculate ΔG0(T,p) from ΔG of formation of the compounds 384(5)
An analogy with ΔH 384(1)
How to calculate ΔG0 385(4)
Calculate ΔG0 from ΔG0(T, p) = ΔH0(T, p) - T ΔS0(T, p): the method 389(4)
Calculation of ΔH0(T,p) 391(1)
Calculation of ΔS0 391(2)
Calculate ΔG0 from ΔG0 = ΔH0
TΔS0: the change of K with temperature 393(5)
Calculation of ΔH0 394(2)
Calculation of ΔS0(T, p) 396(1)
Calculation of ΔG0(T, p) 397(1)
Calculate ΔG0 and K: the change of equilibrium constant with pressure 398(6)
Calculation of IH(H2O; 1000 K,p) 400(2)
Numerical results for the molar volumes and IH 402(1)
Calculation of IS(i; 1000 K,p) 402(2)
Calculate ΔG0(T, p) from ΔG0(T, p) = Σi v(i) μ0 (i; T, p) 404(2)
Problems 406(3)
Chemical equilibrium: dependence of equilibrium constant on temperature and pressure 409(22)
The change of the equilibrium constant K with temperature and pressure: the equations 410(1)
The change of equilibrium constant with temperature 410(1)
The change of equilibrium constant with pressure 411(1)
Le Chatelier's Principle 411(3)
The formulation of the principle 411(1)
The derivation of the principle 412(2)
Calculations of the change of the equilibrium constant K with temperature 414(5)
The equilibrium composition at different temperatures and initial compositions 419(4)
The change of the equilibrium constant K with pressure 423(5)
The Use of (∂ In K/∂p)T, n = - ΔV0/RT (Eq. 20.5) 423(2)
Calculation of the molar volumes 425(1)
Calculation of Iv(i; T = 1000 K,p) 426(1)
Calculation of the equilibrium constant at various pressures 427(1)
How to calculate the change of equilibrium constant when both the temperature and pressure are changed 428(2)
Summary 430(1)
Chemical equilibrium of coupled reactions 431(20)
Introduction 431(1)
Mass balance and equilibrium conditions for coupled reactions 432(3)
Mass balance 432(1)
Equilibrium conditions 432(2)
The equilibrium constants 434(1)
The calculation of ΔG01, ΔG02, K1, and K2 435(1)
Calculation of equilibrium composition for coupled reactions 435(6)
The molar fractions 436(1)
The number of moles 436(3)
The equilibrium constants 439(1)
Calculation of the equilibrium extents of reaction 440(1)
Equilibrium composition 440(1)
Some interesting complications for heterogeneous reactions 440(1)
Generalization 441(3)
Stoichiometric coefficients v(i, α) 442(1)
Number of moles and molar fractions 442(1)
The equilibrium constants 443(1)
Another application 444(7)
The results of the calculations 444(3)
Using Le Chatelier's principle 447(4)
Phase transitions in one-component systems: the phenomena 451(12)
Introduction 451(1)
The phenomena taking place during a phase transition 452(7)
This process is reversible 455(1)
The coexistence curve 455(2)
The complete phase diagram 457(2)
How we use phase diagrams 459(1)
The vapor pressure 459(4)
Phase transitions in one-component systems: the equilibrium conditions 463(18)
The equilibrium condition for coexisting phases 463(5)
The equilibrium (coexistence) condition 464(1)
How to use this equation 465(1)
Phase stability 466(2)
The analogy to a chemical reaction 468(1)
The Clapevron equation 468(4)
The entropy of transformation is connected to the heat of transformation 471(1)
Supplement 23.1. The evaporation of droplets and bubbles, and the mysteries of nucleation 472(9)
There is an interface between the phases 472(1)
The stability of a liquid droplet 473(1)
The phase transition is a change of radius 474(1)
If μ(g) < μ(l), the droplet is unstable 475(1)
The strange ``stability'' when μ(l;T.p) < μ(g: T,p) 476(1)
A metastable equilibrium 477(1)
Nucleation 478(1)
Heterogeneous nucleation 479(2)
Phase transitions in one-component systems: how to use the equilibrium conditions 481(20)
Introduction 481(2)
Transitions involving two condensed phases 483(4)
Melting 483(1)
An example of a solid--solid phase transformation 484(3)
Transitions involving one condensed phase and one vapor phase 487(14)
A simplified theory 488(7)
Using the Clapeyron equation to calculate the vapor pressure of NH3 495(2)
Another kind of phase diagram 497(4)
Phase equilibria in binary systems: the phenomena 501(18)
Introduction 501(3)
The independent variables 502(2)
How to make a constant-pressure phase diagram for the liquid--vapor equilibrium in a two-component system 504(12)
Making a constant-pressure phase-diagram 504(5)
The full liquid-vapor phase diagram at constant pressure 509(2)
How much material is there in each coexisting phase 511(3)
Systems with an azeotrope 514(2)
How to make a constant-temperature liquid-vapor phase diagram 516(3)
Equilibrium conditions for binary systems with two phases: application to vapor-liquid equilibrium 519(28)
Introduction 519(1)
The equilibrium conditions for two coexisting phases in a binary system 520(4)
Derivation of the equilibrium condition 520(2)
The number of independent variables 522(1)
The phase rule 523(1)
Application to liquid-vapor equilibrium: ideal mixtures 524(11)
From equilibrium conditions to phase diagrams 524(1)
Phase diagrams for ideal binary mixtures 525(1)
Replace μ0i(g) and μ0i(l) with measurable quantities 525(1)
Raoult's Law 526(1)
The vapor pressure p0i (T) of the pure compound 527(2)
Calculation of the dew line and bubble line for a mixture of carbon disulphide and benzene 529(1)
The dew line 529(1)
The bubble line 529(1)
Numerical calculations and comparison 530(5)
Supplement 26.1. The liquid-vapor equilibrium when the liquid is a real solution and the vapor is an ideal mixture of ideal gases 535(12)
Activity coefficients 536(1)
The dependence of activity coefficient on molar fraction 537(1)
How to calculate the phase diagram at constant temperature if you know the activity coefficients 538(3)
The calculation of the activity coefficients from data on liquid-vapor equilibrium 541(1)
What is wrong with the Margules equation? 541(6)
Electrolyte solutions 547(24)
Introduction 547(3)
Electrolyte solutions 547(1)
Long-range interactions 548(2)
Mass balance and independent variables 550(2)
Notation, mass balance, and charge conservation 550(1)
Charge neutrality 551(1)
Equilibrium conditions 552(7)
The change in Gibbs free energy 552(1)
A new set of variables 553(1)
The dissociation equilibrium 554(1)
Equilibria involving two phases 555(2)
A discussion of various chemical potentials 557(1)
Why μ(+) and μ(-) are not relevant individually 558(1)
Activity and activity coefficient 559(7)
Molality 559(2)
The definition of activity in the molality scale 561(1)
The chemical potential of the electrolyte 562(4)
The Debye--Huckel theory of electrolyte solutions 566(5)
Formulae and use 566(5)
Galvanic cells: phenomena 571(26)
Introduction 571(1)
Galvanic cells 572(14)
How to make a Daniell cell 572(1)
Daniell cell: the cell reactions 573(1)
Daniell cell: the half reaction at the Cu electrode 574(1)
Daniell cell: the half reaction at the Zn electrode 575(1)
Contact potentials 576(1)
At equilibrium, the voltage inside each conducting phase is constant (in space) 577(1)
Daniell cell in an open circuit 578(2)
Daniell cell in a closed circuit 580(1)
The battery turns chemical energy into electric work 580(2)
Some practice with cell symbols 582(3)
Electrolysis 585(1)
Supplement 28.1. Fuel cells 586(9)
A modern fuel cell 590(2)
The electrolyte 592(1)
The cathode 592(2)
The macro factors 594(1)
Supplement 28.2. A brief history of electrolytic galvanic cells 595(1)
Supplement 28.3. One- and two-euro coins 596(1)
Galvanic cells: equilibrium conditions 597(74)
Introduction 597(1)
The equilibrium conditions for a charged species in two conducting phases in contact 598(5)
A review of the thermodynamics of mixtures 600(1)
Electrolyte solutions in the absence of an electric field 601(1)
Electrolyte solutions in the presence of an electrostatic potential 601(2)
The change of the Gibbs free energy 603(1)
The equilibrium conditions for charged particles in two adjoining phases 603(2)
The contact potential of two metals 603(2)
The equilibrium conditions for a galvanic cell 605(7)
When a galvanic cell is in equilibrium 605(1)
The equilibrium condition in a cell: example 605(3)
The interpretation of the equilibrium condition 608(1)
The equilibrium composition in the cell 609(3)
The determination of the standard electromotive force E0 612(6)
The experimental determination of E0 612(2)
Determination of E0 for the Cell Pt(I), H2(1 atm) | HCl(m) | Hg2Cl2(s) 614(1)
Determine the standard emf of the Cell Pt(I), H2(1 atm) | HCl(m) | HCl(m) | AgCl(s), Ag | Pt(H) 615(3)
Going beyond the Debye--Huckel theory 618(1)
Half-cell electromotive force 618(8)
The rules for defining and using standard half-cell emfs 619(2)
Why half-cell emfs are useful 621(2)
The propensity of an oxidation-reduction reaction 623(1)
Why the rule giving the cell from the half-cell emfs works 624(2)
The use of electromotive-force measurements to determine the activity coefficients of an electrolyte solution 626(4)
Using emf measurements to determine activity coefficients 626(1)
The activity coefficient γ (±; HCl) from the electromotive force of the cell Pt(I), H2(1 atm) | HCl(m) | AgCl(s) | Ag | Pt(II) 627(3)
The use of the activity coefficients of electrolytes to calculate how the emf of a cell changes with the molality of the electrolyte solution 630(2)
The connection between the equilibrium composition of a reaction performed in a cell and that of the same reaction performed in a beaker 632(13)
A review of the results needed 633(1)
The gas-liquid equilibrium also matters 634(1)
Developing the equilibrium conditions 635(3)
Using measurements on the open-circuit cell to calculate the gas composition in the short-circuited cell 638(1)
The composition and the partial pressures 639(6)
Appendices
A1. Conversion factors for pressure units 645(1)
A2. Ethane data 645(5)
A3. Gas constant R in units of energy/mol K 650(1)
A4. van der Waals constants 650(1)
A5. Value of the constants appearing in the Beattie-Bridgeman equation 651(1)
A6. The value of the constants appearing in the Benedict--Webb--Rubin equation 652(3)
A7. Units for work, energy, and heat 655(1)
A8. The dependence of heat capacity on temperature 655(3)
A9. Thermodynamic properties at several temperatures and 1 bar 658(13)
Further Reading 671(4)
Index 675
Preface xxv
How to use the workbooks, exercises, and problems xxxi
Temperature, pressure, molar volume, and equilibrium 1(20)
Introduction 1(1)
System and environment 1(1)
Temperature and thermal equilibrium 2(5)
Thermometer 2(1)
Temperature scales 3(4)
Pressure and mechanical equilibrium 7(4)
Calculating the pressure in a cylinder 8(1)
The units of pressure 8(1)
The units Torr, atm, bar, and psi 9(1)
Conversion between units 10(1)
Volume, density, and molar volume 11(1)
Intensive and extensive quantities 12(1)
Equilibrium 13(1)
Equilibrium and environment 13(1)
Supplement 1.1 An excerpt from Fahrenheit's article 14(2)
Supplement 1.2 Origin of the pressure units atmosphere and Torr 16(3)
Problems 19(2)
The equation of state 21(12)
Introduction 21(3)
The state of a gas or a liquid at equilibrium 21(1)
The equation of state 22(1)
Solids are different 22(2)
The ideal gas equation 24(4)
Units 24(2)
When and why the ideal gas law is valid 26(2)
The van der Waals equation of state 28(1)
Accurate equations of state 29(1)
Summary and perspective 30(1)
Problems 31(2)
How to use the equation of state 33(26)
Calculate pressure when you know molar volume and temperature 33(3)
Why we calculate pressure 33(1)
How to calculate pressure 34(2)
Calculate molar volume when you know pressure and temperature 36(4)
Why we calculate the molar volume 36(2)
An example of molar volume calculation 38(2)
Calculate temperature when you know molar volume and pressure 40(1)
When such calculations are needed 40(1)
Summary of Chapters 1--3 40(1)
Problems 41(4)
Supplement 3.1 How to get your own equation of state 45(11)
Least squares fitting 50(1)
Why we use the square of the error 50(1)
Minimizing the global error 51(2)
Determining the parameters in the van der Waals equation 53(3)
Problems 56(3)
Thermodynamic transformations 59(10)
Definition and examples of thermodynamic transformations 59(2)
Non-equilibrium transformations 61(1)
Initial and final states 61(1)
The path of the transformation 62(1)
Equilibrium transformations 62(2)
Why we study equilibrium transformations 64(1)
Supplement 4.1. More about equilibrium transformations and their paths 65(4)
Two equilibrium transformations with the same initial and final state but different paths 65(4)
Work 69(22)
Introduction 69(1)
The definition of work 70(3)
The sign convention 71(1)
Units 71(1)
Work is an extensive quantity 72(1)
No change in volume, no work 72(1)
Work is performed against an opposing force 73(1)
There are many ways of exchanging work 73(1)
How to calculate the work in a finite transformation 73(3)
The work performed in a finite transformation 73(3)
Work performed in an isothermal transformation 76(5)
What an isothermal transformation is 76(1)
The work performed in an isothermal expansion 77(1)
A numerical calculation of isothermal work 78(3)
The work performed in an isobaric transformation 81(2)
What an isobaric transformation is 81(1)
A numerical calculation of the work performed in an isobaric transformation 81(2)
The work performed in a transformation depends on the path 83(4)
Work is not a function of state 86(1)
Problems 87(4)
Heat 91(52)
What is heat? 91(4)
The caloric theory 91(2)
What is transferred from a hot body to a cold one 93(1)
How to measure the amount of heat 94(1)
Heat and work are equivalent 94(1)
The amount of heat has a sign 95(1)
Thermal coefficients: definitions 95(4)
Heat exchanged at constant pressure 96(1)
Heat exchanged at constant volume 97(1)
The heat exchanged when pressure changes and temperature is constant 97(1)
The heat exchanged in a general transformation 98(1)
Information regarding the thermal coefficients 99(7)
Heat capacity at constant pressure 99(2)
The temperature dependence of Cp 101(2)
The pressure dependence of Cp 103(1)
Heat capacity at constant volume 104(2)
Calculations of the heat exchanged in simple transformations 106(8)
The heat transferred when a system is heated at constant volume 108(1)
Heat exchange when two bodies of different temperatures are brought into contact 109(2)
The number of moles 111(3)
Problems 114(1)
Supplement 6.1. Heat is a form of motion: an experiment in boring cannon 115(7)
Supplement 6.2. Joseph Black, heat capacity 122(5)
Supplement 6.3. A more extensive look at heat theory and calculations 127(13)
Information regarding lp and lv 129(1)
lp and lv for an ideal gas 130(1)
The heat exchanged in an infinitesimal transformation 131(1)
The heat exchanged in a finite transformation 131(1)
The heat exchanged in an isothermal compression: general equation 132(1)
The heat exchanged in an isothermal compression of an ideal gas 133(1)
The heat exchanged in an isothermal compression of a van der Waals gas 134(3)
The case of a real gas: implementation 137(1)
A numerical evaluation of qT 138(2)
Problems 140(3)
Reversible and irreversible transformations 143(6)
Introduction 143(1)
Definition 144(1)
Heat transfer 144(1)
Diffusion 145(1)
An equilibrium transformation is reversible 146(3)
Path-dependent and path-independent quantities 149(16)
Path-independent line integrals 149(12)
Line integrals in thermodynamics 149(4)
Most line integrals depend on the path (the line) 153(1)
Path-independent integrals: an example 154(2)
Exact differentials 156(1)
Path-independent line integrals: theorems 157(4)
Applications to thermodynamics 161(4)
Work and heat depend on path 161(1)
A proof that ∫δW depends on path 162(1)
Exact differentials and functions of state 163(2)
First and second laws of thermodynamics 165(14)
The formulation of the laws 165(1)
Introduction 165(1)
The First Law 165(1)
The Second Law 166(1)
The Third Law 166(1)
Common features of energy and entropy 167(3)
∫dU and ∫dS are path-independent 167(1)
U and S are functions of state 168(1)
Maxwell relations 168(1)
Adding a constant to entropy or energy causes no measurable change 169(1)
U and S are extensive properties 170(1)
A few comments about the First Law 170(5)
Energy conservation 170(4)
The First Law today 174(1)
A few comments about the Second Law 175(4)
The direction of a transformation 175(2)
Impossible processes 177(1)
Why this is useful 177(2)
Helmholtz and Gibbs free energies 179(8)
Introduction 179(1)
Why do we need other functions besides entropy? 179(3)
A convenient form for the First and Second Laws 179(1)
Second Law for a transformation keeping U and V constant 180(2)
Helmholtz free energy A 182(2)
Second Law in terms of changes in V and T 182(1)
Transformations at constant T and V 183(1)
Using A: a hint 184(1)
Gibbs free energy G 184(3)
The definition of Gibbs free energy 184(1)
Some properties of Gibbs free energy 185(1)
The change of G in a transformation in which T and p are held constant 185(1)
Using Gibbs free energy: a hint 186(1)
How to calculate the change of entropy in an equilibrium transformation 187(32)
Introduction 187(1)
Which variable to use 187(1)
The variables T and p: theory 188(5)
The change ds of entropy when pressure changes by dp and temperature by dT 189(1)
Notation for partial derivatives 189(2)
Maxwell's method 191(1)
The derivative (∂s/∂p)T 191(1)
The derivative (∂s/∂T)p 192(1)
Combine these results to get an expression for ds 193(1)
The change of entropy in a transformation in which T and p change by a finite amount 193(1)
The variables T and v: theory 193(3)
The calculation of (∂s/∂v)T 194(1)
The change in entropy when both T and v change 195(1)
The change in entropy in a finite transformation 195(1)
Calculations of entropy change in various transformations 196(9)
Entropy change in an isobaric transformation 196(2)
The change of entropy in an isothermal transformation 198(4)
The change of entropy in a general transformation 202(2)
The change of entropy: numerical results 204(1)
How to use the tables of data to calculate entropy changes 204(1)
Problems 205(3)
Supplement 11.1. Using Maxwell's method to derive useful equations 208(3)
Using Maxwell equations 208(3)
Supplement 11.2. Obtaining new equations by changing variables 211(4)
An equation for (∂v/∂T)p 212(1)
A more general method for calculating (∂v/∂T)p 213(1)
An equation for Cp
Cv 214(1)
Supplement 11.3. Adiabatic transformations 215(4)
The transformation path in an adiabatic transformation 216(1)
The equation for the final temperature 217(1)
An example of adiabatic compression 218(1)
Enthalpy and energy change during a thermodynamic transformation 219(28)
Introduction 219(4)
Heat and enthalpy 221(1)
The connection between the enthalpy change and heat 222(1)
Heat and energy 222(1)
How to calculate the enthalpy change in a transformation: theory 223(4)
The change of enthalpy in an infinitesimal transformation 223(2)
The change of enthalpy in a finite transformation 225(1)
Choosing a path 225(2)
How to calculate the enthalpy change in a transformation: an example 227(11)
The change of enthalpy ΔhA when T changes and p is constant (path A) 228(1)
The change of enthalpy ΔhB when pressure changes and temperature is held constant 229(1)
The first difficulty 229(2)
The second difficulty 231(2)
The final result for ΔhB 233(1)
An example of a calculation of ΔhB 234(1)
The order of magnitude of ΔhB 234(1)
The use of tables to calculate enthalpy changes with temperature 235(3)
Supplement 12.1. Energy changes in a thermodynamic transformation 238(2)
The change of energy in an infinitesimal transformation in which T and v are changed 238(1)
The change of energy in a finite transformation in which T and v are changed 238(1)
Another way of calculating energy changes 239(1)
Supplement 12.2. Isenthalpic transformations 240(7)
This transformation takes place without a change of enthalpy 241(1)
The change of temperature caused by an isenthalpic transformation 242(1)
Isenthalpic transformations are used for cooling 243(3)
Ideal gas 246(1)
Thermochemistry 247(38)
Introduction 247(1)
Definition of the heat of reaction 248(3)
Two reactions used as examples 248(1)
The definition of the heat of reaction 249(1)
The initial state 249(1)
The final state 249(1)
Sign convention 250(1)
The presentation of data: standard state 250(1)
The connection between the heat of reaction and the enthalpy of the participants 251(4)
Enthalpy change in a reaction 252(3)
The dependence of the heat of reaction on temperature and pressure 255(3)
Enthalpy change in a transformation from σi to σf 255(1)
The heat of reaction at Tf and pf 256(1)
The change of the heat of reaction with temperature 257(1)
The change of heat of reaction with pressure 257(1)
An example: the heat of reaction for ammonia synthesis 258(6)
The heat of reaction ΔH 259(1)
The temperature dependence of the heat of reaction 260(2)
The heat of reaction at 623.15 K and 394.8 atm 262(2)
Calculating the heat of a reaction from heats of formation or combustion 264(10)
Heats of formation: definition 266(1)
The connection between the heat of reaction and the heats of formation of the compounds 266(3)
Where to get heats of formation 269(1)
The use of heats of combustion to calculate heats of reaction 270(4)
Supplement 13.1. Calculating the heat of one reaction from the heats of other reactions 274(7)
Problems 281(4)
The change of chemical potential during an equilibrium transformation 285(16)
Introduction 285(1)
The change of chemical potential calculated with Eq. 14.4 286(9)
How to evaluate Δμ1 ≡ μ(Tf, pi) - μ(Ti, pi) 287(1)
An example of evaluation of Δμ1 ≡ μ(Tf, pi) - μ(Ti, pi) 288(1)
How to evaluate Δμ2 ≡ μ(Tf, pf) - μ(Tf, pi) 289(4)
Calculating Δμ = μ(130 K, 600 atm) - μ(298.15 K. 1 atm) 293(2)
Additional material about chemical potential 295(3)
The pressure dependence of the chemical potential of an ideal gas 295(1)
Fugacity 295(1)
The dependence of chemical potential on temperature 296(1)
Calculating μ from μ = h - Ts 297(1)
Problems 298(3)
The chemical potential of a compound in a mixture 301(10)
General remarks 301(2)
Infinitesimal transformations 302(1)
The chemical potential of a compound in a mixture 303(8)
The change of Gibbs free energy when I change temperature, pressure, and composition 303(1)
Change of variables 304(3)
The chemical potential of ideal mixtures 307(1)
The partial pressure of a gas in an ideal mixture 307(1)
The chemical potential of a gas in an ideal mixture 308(1)
A few words about Josiah Willard Gibbs 309(2)
Mixtures: partial molar quantities and activities 311(24)
Partial molar quantities 311(16)
The correct formula 314(2)
Partial molar volume is an intensive quantity 316(1)
Other partial molar quantities 317(1)
Partial molar enthalpy and the heat of mixing 318(1)
Chemical potential as a partial molar quantity: the Gibbs--Duhem equation 319(2)
Equations similar to the Gibbs--Duhem equation for other partial molar quantities 321(1)
How to determine partial molar quantities from measurements 322(3)
Relations among partial molar quantities 325(2)
The composition dependence of chemical potential: ideal solutions 327(3)
The definition of an ideal solution 327(1)
The change of volume when we mix compounds to form an ideal solution 327(1)
The enthalpy of an ideal mixture 328(1)
The heat of mixing to form an ideal solution 329(1)
The entropy of mixing to form an ideal solution 329(1)
Chemical potential of real solutions: the activity and the reference potential 330(5)
Chemical equilibrium 335(24)
Introduction 335(4)
Reactants and products 336(1)
Stable and metastable chemical equilibrium 337(2)
The extent of reaction and the composition of a reacting mixture 339(7)
The extent of reaction 339(2)
Mass conservation 341(1)
Molar fractions 341(1)
Examples of the use of the extent of reaction 342(3)
Some properties of the extent of reaction 345(1)
The equilibrium conditions and the direction of a reaction 346(6)
The equilibrium conditions 346(1)
The direction of a chemical reaction 347(1)
Chemical affinity of a reaction 348(1)
Geometric interpretation 349(2)
How to use affinity to answer practical chemistry questions 351(1)
The equilibrium condition in terms of chemical potentials and equilibrium constant 352(7)
The reaction affinity in terms of chemical potential 352(1)
The affinity of a reaction for ideal mixtures 353(3)
The equilibrium conditions for real mixtures 356(3)
Chemical equilibrium: the connection between the equilibrium constant and composition 359(24)
Introduction 359(2)
How to calculate the equilibrium constant from a measurement of the equilibrium concentration of one species 361(2)
How to calculate the equilibrium composition when you know the equilibrium constant and the initial number of moles 363(2)
The dependence of the extent of reaction on the initial number of moles 365(1)
The role of the initial number of moles 365(6)
Supplement 18.1. Another example of composition calculations 371(2)
Supplement 18.2. A second example and error analysis 373(6)
Supplement 18.3. The direction of a reaction 379(4)
A metastable state 381(2)
Chemical equilibrium: how to calculate K from ΔG0 =
RT In K 383(26)
Calculate ΔG0(T,p) from ΔG of formation of the compounds 384(5)
An analogy with ΔH 384(1)
How to calculate ΔG0 385(4)
Calculate ΔG0 from ΔG0(T, p) = ΔH0(T, p) - T ΔS0(T, p): the method 389(4)
Calculation of ΔH0(T,p) 391(1)
Calculation of ΔS0 391(2)
Calculate ΔG0 from ΔG0 = ΔH0
TΔS0: the change of K with temperature 393(5)
Calculation of ΔH0 394(2)
Calculation of ΔS0(T, p) 396(1)
Calculation of ΔG0(T, p) 397(1)
Calculate ΔG0 and K: the change of equilibrium constant with pressure 398(6)
Calculation of IH(H2O; 1000 K,p) 400(2)
Numerical results for the molar volumes and IH 402(1)
Calculation of IS(i; 1000 K,p) 402(2)
Calculate ΔG0(T, p) from ΔG0(T, p) = Σi v(i) μ0 (i; T, p) 404(2)
Problems 406(3)
Chemical equilibrium: dependence of equilibrium constant on temperature and pressure 409(22)
The change of the equilibrium constant K with temperature and pressure: the equations 410(1)
The change of equilibrium constant with temperature 410(1)
The change of equilibrium constant with pressure 411(1)
Le Chatelier's Principle 411(3)
The formulation of the principle 411(1)
The derivation of the principle 412(2)
Calculations of the change of the equilibrium constant K with temperature 414(5)
The equilibrium composition at different temperatures and initial compositions 419(4)
The change of the equilibrium constant K with pressure 423(5)
The Use of (∂ In K/∂p)T, n = - ΔV0/RT (Eq. 20.5) 423(2)
Calculation of the molar volumes 425(1)
Calculation of Iv(i; T = 1000 K,p) 426(1)
Calculation of the equilibrium constant at various pressures 427(1)
How to calculate the change of equilibrium constant when both the temperature and pressure are changed 428(2)
Summary 430(1)
Chemical equilibrium of coupled reactions 431(20)
Introduction 431(1)
Mass balance and equilibrium conditions for coupled reactions 432(3)
Mass balance 432(1)
Equilibrium conditions 432(2)
The equilibrium constants 434(1)
The calculation of ΔG01, ΔG02, K1, and K2 435(1)
Calculation of equilibrium composition for coupled reactions 435(6)
The molar fractions 436(1)
The number of moles 436(3)
The equilibrium constants 439(1)
Calculation of the equilibrium extents of reaction 440(1)
Equilibrium composition 440(1)
Some interesting complications for heterogeneous reactions 440(1)
Generalization 441(3)
Stoichiometric coefficients v(i, α) 442(1)
Number of moles and molar fractions 442(1)
The equilibrium constants 443(1)
Another application 444(7)
The results of the calculations 444(3)
Using Le Chatelier's principle 447(4)
Phase transitions in one-component systems: the phenomena 451(12)
Introduction 451(1)
The phenomena taking place during a phase transition 452(7)
This process is reversible 455(1)
The coexistence curve 455(2)
The complete phase diagram 457(2)
How we use phase diagrams 459(1)
The vapor pressure 459(4)
Phase transitions in one-component systems: the equilibrium conditions 463(18)
The equilibrium condition for coexisting phases 463(5)
The equilibrium (coexistence) condition 464(1)
How to use this equation 465(1)
Phase stability 466(2)
The analogy to a chemical reaction 468(1)
The Clapevron equation 468(4)
The entropy of transformation is connected to the heat of transformation 471(1)
Supplement 23.1. The evaporation of droplets and bubbles, and the mysteries of nucleation 472(9)
There is an interface between the phases 472(1)
The stability of a liquid droplet 473(1)
The phase transition is a change of radius 474(1)
If μ(g) < μ(l), the droplet is unstable 475(1)
The strange ``stability'' when μ(l;T.p) < μ(g: T,p) 476(1)
A metastable equilibrium 477(1)
Nucleation 478(1)
Heterogeneous nucleation 479(2)
Phase transitions in one-component systems: how to use the equilibrium conditions 481(20)
Introduction 481(2)
Transitions involving two condensed phases 483(4)
Melting 483(1)
An example of a solid--solid phase transformation 484(3)
Transitions involving one condensed phase and one vapor phase 487(14)
A simplified theory 488(7)
Using the Clapeyron equation to calculate the vapor pressure of NH3 495(2)
Another kind of phase diagram 497(4)
Phase equilibria in binary systems: the phenomena 501(18)
Introduction 501(3)
The independent variables 502(2)
How to make a constant-pressure phase diagram for the liquid--vapor equilibrium in a two-component system 504(12)
Making a constant-pressure phase-diagram 504(5)
The full liquid-vapor phase diagram at constant pressure 509(2)
How much material is there in each coexisting phase 511(3)
Systems with an azeotrope 514(2)
How to make a constant-temperature liquid-vapor phase diagram 516(3)
Equilibrium conditions for binary systems with two phases: application to vapor-liquid equilibrium 519(28)
Introduction 519(1)
The equilibrium conditions for two coexisting phases in a binary system 520(4)
Derivation of the equilibrium condition 520(2)
The number of independent variables 522(1)
The phase rule 523(1)
Application to liquid-vapor equilibrium: ideal mixtures 524(11)
From equilibrium conditions to phase diagrams 524(1)
Phase diagrams for ideal binary mixtures 525(1)
Replace μ0i(g) and μ0i(l) with measurable quantities 525(1)
Raoult's Law 526(1)
The vapor pressure p0i (T) of the pure compound 527(2)
Calculation of the dew line and bubble line for a mixture of carbon disulphide and benzene 529(1)
The dew line 529(1)
The bubble line 529(1)
Numerical calculations and comparison 530(5)
Supplement 26.1. The liquid-vapor equilibrium when the liquid is a real solution and the vapor is an ideal mixture of ideal gases 535(12)
Activity coefficients 536(1)
The dependence of activity coefficient on molar fraction 537(1)
How to calculate the phase diagram at constant temperature if you know the activity coefficients 538(3)
The calculation of the activity coefficients from data on liquid-vapor equilibrium 541(1)
What is wrong with the Margules equation? 541(6)
Electrolyte solutions 547(24)
Introduction 547(3)
Electrolyte solutions 547(1)
Long-range interactions 548(2)
Mass balance and independent variables 550(2)
Notation, mass balance, and charge conservation 550(1)
Charge neutrality 551(1)
Equilibrium conditions 552(7)
The change in Gibbs free energy 552(1)
A new set of variables 553(1)
The dissociation equilibrium 554(1)
Equilibria involving two phases 555(2)
A discussion of various chemical potentials 557(1)
Why μ(+) and μ(-) are not relevant individually 558(1)
Activity and activity coefficient 559(7)
Molality 559(2)
The definition of activity in the molality scale 561(1)
The chemical potential of the electrolyte 562(4)
The Debye--Huckel theory of electrolyte solutions 566(5)
Formulae and use 566(5)
Galvanic cells: phenomena 571(26)
Introduction 571(1)
Galvanic cells 572(14)
How to make a Daniell cell 572(1)
Daniell cell: the cell reactions 573(1)
Daniell cell: the half reaction at the Cu electrode 574(1)
Daniell cell: the half reaction at the Zn electrode 575(1)
Contact potentials 576(1)
At equilibrium, the voltage inside each conducting phase is constant (in space) 577(1)
Daniell cell in an open circuit 578(2)
Daniell cell in a closed circuit 580(1)
The battery turns chemical energy into electric work 580(2)
Some practice with cell symbols 582(3)
Electrolysis 585(1)
Supplement 28.1. Fuel cells 586(9)
A modern fuel cell 590(2)
The electrolyte 592(1)
The cathode 592(2)
The macro factors 594(1)
Supplement 28.2. A brief history of electrolytic galvanic cells 595(1)
Supplement 28.3. One- and two-euro coins 596(1)
Galvanic cells: equilibrium conditions 597(74)
Introduction 597(1)
The equilibrium conditions for a charged species in two conducting phases in contact 598(5)
A review of the thermodynamics of mixtures 600(1)
Electrolyte solutions in the absence of an electric field 601(1)
Electrolyte solutions in the presence of an electrostatic potential 601(2)
The change of the Gibbs free energy 603(1)
The equilibrium conditions for charged particles in two adjoining phases 603(2)
The contact potential of two metals 603(2)
The equilibrium conditions for a galvanic cell 605(7)
When a galvanic cell is in equilibrium 605(1)
The equilibrium condition in a cell: example 605(3)
The interpretation of the equilibrium condition 608(1)
The equilibrium composition in the cell 609(3)
The determination of the standard electromotive force E0 612(6)
The experimental determination of E0 612(2)
Determination of E0 for the Cell Pt(I), H2(1 atm) | HCl(m) | Hg2Cl2(s) 614(1)
Determine the standard emf of the Cell Pt(I), H2(1 atm) | HCl(m) | HCl(m) | AgCl(s), Ag | Pt(H) 615(3)
Going beyond the Debye--Huckel theory 618(1)
Half-cell electromotive force 618(8)
The rules for defining and using standard half-cell emfs 619(2)
Why half-cell emfs are useful 621(2)
The propensity of an oxidation-reduction reaction 623(1)
Why the rule giving the cell from the half-cell emfs works 624(2)
The use of electromotive-force measurements to determine the activity coefficients of an electrolyte solution 626(4)
Using emf measurements to determine activity coefficients 626(1)
The activity coefficient γ (±; HCl) from the electromotive force of the cell Pt(I), H2(1 atm) | HCl(m) | AgCl(s) | Ag | Pt(II) 627(3)
The use of the activity coefficients of electrolytes to calculate how the emf of a cell changes with the molality of the electrolyte solution 630(2)
The connection between the equilibrium composition of a reaction performed in a cell and that of the same reaction performed in a beaker 632(13)
A review of the results needed 633(1)
The gas-liquid equilibrium also matters 634(1)
Developing the equilibrium conditions 635(3)
Using measurements on the open-circuit cell to calculate the gas composition in the short-circuited cell 638(1)
The composition and the partial pressures 639(6)
Appendices
A1. Conversion factors for pressure units 645(1)
A2. Ethane data 645(5)
A3. Gas constant R in units of energy/mol K 650(1)
A4. van der Waals constants 650(1)
A5. Value of the constants appearing in the Beattie-Bridgeman equation 651(1)
A6. The value of the constants appearing in the Benedict--Webb--Rubin equation 652(3)
A7. Units for work, energy, and heat 655(1)
A8. The dependence of heat capacity on temperature 655(3)
A9. Thermodynamic properties at several temperatures and 1 bar 658(13)
Further Reading 671(4)
Index 675
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