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Physical and Chemical Equilibrium for Chemical Engineers 2nd edition [Kõva köide]

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(University of Utah, Salt Lake City)
  • Formaat: Hardback, 384 pages, kõrgus x laius x paksus: 287x222x27 mm, kaal: 1098 g
  • Ilmumisaeg: 10-Apr-2012
  • Kirjastus: John Wiley & Sons Inc
  • ISBN-10: 0470927100
  • ISBN-13: 9780470927106
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Physical and Chemical Equilibrium for Chemical Engineers 2nd editionSuurem pilt
  • Formaat: Hardback, 384 pages, kõrgus x laius x paksus: 287x222x27 mm, kaal: 1098 g
  • Ilmumisaeg: 10-Apr-2012
  • Kirjastus: John Wiley & Sons Inc
  • ISBN-10: 0470927100
  • ISBN-13: 9780470927106
Teised raamatud teemal:
Püsilink: https://www.kriso.ee/db/9780470927106.html
Märksõnad:
"New to this edition is an appendix covering the Bridgman table which includes a basic set of thermodynamic equations Has a new chapter on the thermodynamics of biochemical reactions Updates naming and notation used in the first edition; where the first edition used the traditional names for the Gibbs Free Energy and for Partial Molal Properties, this edition uses the more popular Gibbs Energy and Partial Molar Properties Teaches Physical and chemical equilibrium which deals with calculating the thermodynamic properties for mixtures. Includes many problems in the text to help the reader understand the material covered and includes a solutions manual to these problems"--



This book concentrates on the topic of physical and chemical equilibrium. Using the simplest mathematics along with numerous numerical examples it accurately and rigorously covers physical and chemical equilibrium in depth and detail. It continues to cover the topics found in the first edition however numerous updates have been made including: Changes in naming and notation (the first edition used the traditional names for the Gibbs Free Energy and for Partial Molal Properties, this edition uses the more popular Gibbs Energy and Partial Molar Properties,) changes in symbols (the first edition used the Lewis-Randal fugacity rule and the popular symbol for the same quantity, this edition only uses the popular notation,) and new problems have been added to the text. Finally the second edition includes an appendix about the Bridgman table and its use.
Preface xiii
About the Author xv
Nomenclature xvii
1 Introduction to Equilibrium
1(12)
1.1 Why Study Equilibrium?
1(3)
1.2 Stability and Equilibrium
4(1)
1.3 Time Scales and the Approach to Equilibrium
5(1)
1.4 Looking Ahead, Gibbs Energy
5(1)
1.5 Units, Conversion Factors, and Notation
6(2)
1.6 Reality and Equations
8(1)
1.7 Phases and Phase Diagrams
8(2)
1.8 The Plan of this Book
10(1)
1.9 Summary
10(3)
References
11(2)
2 Basic Thermodynamics
13(22)
2.1 Conservation and Accounting
13(1)
2.2 Conservation of Mass
14(1)
2.3 Conservation of Energy; the First Law of Thermodynamics
15(2)
2.4 The Second Law of Thermodynamics
17(2)
2.4.1 Reversibility
17(1)
2.4.2 Entropy
18(1)
2.5 Convenience Properties
19(1)
2.6 Using the First and Second Laws
19(2)
2.7 Datums and Reference States
21(1)
2.8 Measurable and Immeasurable Properties
22(1)
2.9 Work and Heat
22(1)
2.10 The Property Equation
23(1)
2.11 Equations of State (EOS)
24(2)
2.11.1 EOSs Based on Theory
25(1)
2.11.2 EOSs Based on Pure Data Fitting
25(1)
2.12 Corresponding States
26(2)
2.13 Departure Functions
28(1)
2.14 The Properties of Mixtures
28(1)
2.15 The Combined First and Second Law Statement; Reversible Work
29(2)
2.16 Summary
31(4)
References
33(2)
3 The Simplest Phase Equilibrium Examples and Some Simple Estimating Rules
35(14)
3.1 Some General Statements About Equilibrium
35(2)
3.2 The Simplest Example of Phase Equilibrium
37(1)
3.2.1 A Digression, the Distinction between Vapor and Gas
37(1)
3.2.2 Back to the Simplest Equilibrium
37(1)
3.3 The Next Level of Complexity in Phase Equilibrium
37(2)
3.4 Some Simple Estimating Rules: Raoult's and Henry's "Laws"
39(4)
3.5 The General Two-Phase Equilibrium Calculation
43(1)
3.6 Some Simple Applications of Raoult's and Henry's Laws
43(3)
3.7 The Uses and Limits of Raoult's and Henry's Laws
46(1)
3.8 Summary
46(3)
References
48(1)
4 Minimization of Gibbs Energy
49(12)
4.1 The Fundamental Thermodynamic Criterion of Phase and Chemical Equilibrium
49(2)
4.2 The Criterion of Equilibrium Applied to Two Nonreacting Equilibrium Phases
51(2)
4.3 The Criterion of Equilibrium Applied to Chemical Reactions
53(1)
4.4 Simple Gibbs Energy Diagrams
54(4)
4.4.1 Comparison with Enthalpy and Entropy
55(1)
4.4.2 Gibbs Energy Diagrams for Pressure-Driven Phase Changes
55(2)
4.4.3 Gibbs Energy Diagrams for Chemical Reactions
57(1)
4.5 Le Chatelier's Principle
58(1)
4.6 Summary
58(3)
References
60(1)
5 Vapor Pressure, the Clapeyron Equation, and Single Pure Chemical Species Phase Equilibrium
61(12)
5.1 Measurement of Vapor Pressure
61(1)
5.2 Reporting Vapor-Pressure Data
61(1)
5.2.1 Normal Boiling Point (NBP)
61(1)
5.3 The Clapeyron Equation
62(1)
5.4 The Clausius-Clapeyron Equation
63(1)
5.5 The Accentric Factor
64(2)
5.6 The Antoine Equation and Other Data-Fitting Equations
66(1)
5.6.1 Choosing a Vapor-Pressure Equation
67(1)
5.7 Applying the Clapeyron Equation to Other Kinds of Equilibrium
67(1)
5.8 Extrapolating Vapor-Pressure Curves
68(1)
5.9 Vapor Pressure of Solids
69(1)
5.10 Vapor Pressures of Mixtures
69(1)
5.11 Summary
69(4)
References
72(1)
6 Partial Molar Properties
73(16)
6.1 Partial Molar Properties
73(1)
6.2 The Partial Molar Equation
74(1)
6.3 Tangent Slopes
74(3)
6.4 Tangent Intercepts
77(1)
6.5 The Two Equations for Partial Molar Properties
78(1)
6.6 Using the Idea of Tangent Intercepts
79(1)
6.7 Partial Mass Properties
80(1)
6.8 Heats of Mixing and Partial Molar Enthalpies
80(2)
6.8.1 Differential Heat of Mixing
80(1)
6.8.2 Integral Heat of Mixing
81(1)
6.9 The Gibbs-Duhem Equation and the Counterintuitive Behavior of the Chemical Potential
82(2)
6.10 Summary
84(5)
References
87(2)
7 Fugacity, Ideal Solutions, Activity, Activity Coefficient
89(18)
7.1 Why Fugacity?
89(1)
7.2 Fugacity Defined
89(1)
7.3 The Use of the Fugacity
90(1)
7.4 Pure Substance Fugacities
90(5)
7.4.1 The Fugacity of Pure Gases
91(3)
7.4.2 The Fugacity of Pure Liquids and Solids
94(1)
7.5 Fugacities of Species in Mixtures
95(1)
7.6 Mixtures of Ideal Gases
95(1)
7.7 Why Ideal Solutions?
95(1)
7.8 Ideal Solutions Defined
96(2)
7.8.1 The Consequences of the Ideal Solution Definition
96(2)
7.9 Why Activity and Activity Coefficients?
98(1)
7.10 Activity and Activity Coefficients Defined
98(2)
7.11 Fugacity Coefficient for Pure Gases and Gas Mixtures
100(1)
7.12 Estimating Fugacities of Individual Species in Gas Mixtures
100(4)
7.12.1 Fugacities from Gas PvT Data
100(2)
7.12.2 Fugacities from an EOS for Gas Mixtures
102(1)
7.12.3 The Lewis and Randall (L-R) Fugacity Rule
102(1)
7.12.4 Other Mixing Rules
103(1)
7.13 Liquid Fugacities from Vapor-Liquid Equilibrium
104(1)
7.14 Summary
104(3)
References
105(2)
8 Vapor-Liquid Equilibrium (VLE) at Low Pressures
107(38)
8.1 Measurement of VLE
107(3)
8.2 Presenting Experimental VLE Data
110(1)
8.3 The Mathematical Treatment of Low-Pressure VLE Data
110(2)
8.3.1 Raoult's Law Again
111(1)
8.4 The Four Most Common Types of Low-Pressure VLE
112(10)
8.4.1 Ideal Solution Behavior (Type I)
114(1)
8.4.2 Positive Deviations from Ideal Solution Behavior (Type II)
114(1)
8.4.3 Negative Deviations from Ideal Solution Behavior (Type III)
115(2)
8.4.4 Azeotropes
117(1)
8.4.5 Two-Liquid Phase or Heteroazeotropes (Type IV)
118(2)
8.4.6 Zero Solubility and Steam Distillation
120(1)
8.4.7 Distillation of the Four Types of Behavior
121(1)
8.5 Gas-Liquid Equilibrium, Henry's Law Again
122(1)
8.6 The Effect of Modest Pressures on VLE
122(2)
8.6.1 Liquids
123(1)
8.6.2 Gases, the L-R Rule
123(1)
8.7 Standard States Again
124(1)
8.8 Low-Pressure VLE Calculations
125(7)
8.8.1 Bubble-Point Calculations
127(1)
8.8.1.1 Temperature-Specified Bubble Point
127(1)
8.8.1.2 Pressure-Specified Bubble Point
128(1)
8.8.2 Dew-Point Calculations
129(1)
8.8.2.1 Temperature-Specified Dew Point
129(1)
8.8.2.2 Pressure-Specified Dew Point
129(1)
8.8.3 Isothermal Flashes (T- and P-Specified Flashes)
130(1)
8.8.4 Adiabatic Flashes
131(1)
8.9 Traditional K-Factor Methods
132(1)
8.10 More Uses for Raoult's Law
132(4)
8.10.1 Nonvolatile Solutes, Boiling-Point Elevation
132(3)
8.10.2 Freezing-Point Depression
135(1)
8.10.3 Colligative Properties of Solutions
136(1)
8.11 Summary
136(9)
References
143(2)
9 Correlating and Predicting Nonideal VLE
145(24)
9.1 The Most Common Observations of Liquid-Phase Activity Coefficients
145(2)
9.1.1 Why Nonideal Behavior?
145(1)
9.1.2 The Shapes of In, γ-x Curves
146(1)
9.2 Limits on Activity Coefficient Correlations, the Gibbs-Duhem Equation
147(1)
9.3 Excess Gibbs Energy and Activity Coefficient Equations
148(2)
9.4 Activity Coefficients at Infinite Dilution
150(1)
9.5 Effects of Pressure and Temperature on Liquid-Phase Activity Coefficients
151(2)
9.5.1 Effect of Pressure Changes on Liquid-Phase Activity Coefficients
151(1)
9.5.2 Effect of Temperature Changes on Liquid-Phase Activity Coefficients
152(1)
9.6 Ternary and Multispecies VLE
153(2)
9.6.1 Liquid-Phase Activity Coefficients for Ternary Mixtures
154(1)
9.7 Vapor-Phase Nonideality
155(3)
9.8 VLE from EOS
158(1)
9.9 Solubility Parameter
158(2)
9.10 The Solubility of Gases in Liquids, Henry's Law Again
160(3)
9.11 Summary
163(6)
References
167(2)
10 Vapor-Liquid Equilibrium (VLE) at High Pressures
169(12)
10.1 Critical Phenomena of Pure Species
169(1)
10.2 Critical Phenomena of Mixtures
170(4)
10.3 Estimating High-Pressure VLE
174(4)
10.3.1 Empirical K-Value Correlations
175(1)
10.3.2 Estimation Methods for Each Phase Separately Not Based on Raoult's Law
175(1)
10.3.3 Estimation Methods Based on Cubic EOSs
176(2)
10.4 Computer Solutions
178(1)
10.5 Summary
178(3)
References
179(2)
11 Liquid-Liquid, Liquid-Solid, and Gas-Solid Equilibrium
181(36)
11.1 Liquid-Liquid Equilibrium (LLE)
181(1)
11.2 The Experimental Determination of LLE
181(6)
11.2.1 Reporting and Presenting LLE Data
182(1)
11.2.2 Practically Insoluble Liquid Pairs at 25°C
183(1)
11.2.3 Partially Soluble Liquid Pairs at 25°C
183(1)
11.2.4 Miscible Liquid Pairs at 25°C
183(1)
11.2.5 Ternary LLE at 25°C
184(2)
11.2.6 LLE at Temperatures Other Than 25°C
186(1)
11.3 The Elementary Theory of LLE
187(3)
11.4 The Effect of Pressure on LLE
190(1)
11.5 Effect of Temperature on LLE
191(3)
11.6 Distribution Coefficients
194(1)
11.7 Liquid-Solid Equilibrium (LSE)
195(5)
11.7.1 One-Species LSE
195(1)
11.7.2 The Experimental Determination of LSE
195(1)
11.7.3 Presenting LSE Data
195(2)
11.7.4 Eutectics
197(2)
11.7.5 Gas Hydrates (Clathrates)
199(1)
11.8 The Elementary Thermodynamics of LSE
200(2)
11.9 Gas-Solid Equilibrium (GSE) at Low Pressures
202(1)
11.10 GSE at High Pressures
203(1)
11.11 Gas-Solid Adsorption, Vapor-Solid Adsorption
204(7)
11.11.1 Langmuir's Adsorption Theory
205(2)
11.11.2 Vapor-solid Adsorption, BET Theory
207(1)
11.11.3 Adsorption from Mixtures
208(1)
11.11.4 Heat of Adsorption
209(1)
11.11.5 Hysteresis
210(1)
11.12 Summary
211(6)
References
215(2)
12 Chemical Equilibrium
217(26)
12.1 Introduction to Chemical Reactions and Chemical Equilibrium
217(1)
12.2 Formal Description of Chemical Reactions
217(1)
12.3 Minimizing Gibbs Energy
218(1)
12.4 Reaction Rates, Energy Barriers, Catalysis, and Equilibrium
219(1)
12.5 The Basic Thermodynamics of Chemical Reactions and Its Convenient Formulations
220(3)
12.5.1 The Law of Mass Action and Equilibrium Constants
222(1)
12.6 Calculating Equilibrium Constants from Gibbs Energy Tables and then Using Equilibrium Constants to Calculate Equilibrium Concentrations
223(4)
12.6.1 Change of Reactant Concentration, Reaction Coordinate
224(3)
12.6.2 Reversible and Irreversible Reactions
227(1)
12.7 More on Standard States
227(2)
12.8 The Effect of Temperature on Chemical Reaction Equilibrium
229(5)
12.9 The Effect of Pressure on Chemical Reaction Equilibrium
234(4)
12.9.1 Ideal Solution of Ideal Gases
235(1)
12.9.2 Nonideal Solution, Nonideal Gases
236(1)
12.9.3 Liquids and Solids
237(1)
12.10 The Effect of Nonideal Solution Behavior
238(1)
12.10.1 Liquid-Phase Nonideality
238(1)
12.11 Other Forms of K
238(1)
12.12 Summary
239(4)
References
242(1)
13 Equilibrium in Complex Chemical Reactions
243(22)
13.1 Reactions Involving Ions
243(1)
13.2 Multiple Reactions
244(5)
13.2.1 Sequential Reactions
244(1)
13.2.2 Simultaneous Reactions
245(1)
13.2.3 The Charge Balance Calculation Method and Buffers
246(3)
13.3 Reactions with More Than One Phase
249(3)
13.3.1 Solubility Product
249(1)
13.3.2 Gas-Liquid Reactions
249(3)
13.4 Electrochemical Reactions
252(3)
13.5 Chemical and Physical Equilibrium in Two Phases
255(2)
13.5.1 Dimerization (Association)
255(2)
13.6 Summary
257(8)
References
262(3)
14 Equilibrium with Gravity or Centrifugal Force, Osmotic Equilibrium Equilibrium with Surface Tension
265(14)
14.1 Equilibrium with Other Forms of Energy
265(1)
14.2 Equilibrium in the Presence of Gravity
266(3)
14.2.1 Centrifuges
268(1)
14.3 Semipermeable Membranes
269(2)
14.3.1 Osmotic Pressure
270(1)
14.4 Small is Interesting! Equilibrium with Surface Tension
271(4)
14.4.1 Bubbles, Drops and Nucleation
271(4)
14.4.2 Capillary Condensation
275(1)
14.5 Summary
275(4)
References
278(1)
15 The Phase Rule
279(14)
15.1 How Many Phases Can Coexist in a Given Equilibrium Situation?
279(1)
15.2 What Does the Phase Rule Tell Us? What Does It Not Tell Us?
280(1)
15.3 What is a Phase?
280(1)
15.4 The Phase Rule is Simply Counting Variables
281(1)
15.5 More On Components
282(3)
15.5.1 A Formal Way to Find the Number of Independent Equations
285(1)
15.6 The Phase Rule for One- and Two-Component Systems
285(3)
15.7 Harder Phase Rule Problems
288(1)
15.8 Summary
288(5)
References
291(2)
16 Equilibrium in Biochemical Reactions
293(10)
16.1 An Example, the Production of Ethanol from Sugar
293(1)
16.2 Organic and Biochemical Reactions
293(1)
16.3 Two More Sweet Examples
294(1)
16.4 Thermochemical Data for Biochemical Reactions
295(1)
16.5 Thermodynamic Equilibrium in Large Scale Biochemistry
296(1)
16.6 Translating between Biochemical and Chemical Engineering Equilibrium Expressions
296(2)
16.6.1 Chemical and Biochemical Equations
297(1)
16.6.2 Equilibrium Constants
297(1)
16.6.3 pH and Buffers
298(1)
16.6.4 Ionic Strength
298(1)
16.7 Equilibrium in Biochemical Separations
298(1)
16.8 Summary
299(4)
References
300(3)
Appendix A Useful Tables and Charts
303(16)
A.1 Useful Property Data for Corresponding States Estimates
303(2)
A.2 Vapor-Pressure Equation Constants
305(1)
A.3 Henry's Law Constants
306(1)
A.4 Compressibility Factor Chart (z Chart)
307(1)
A.5 Fugacity Coefficient Charts
307(1)
A.6 Azeotropes
308(4)
A.7 Van Laar Equation Constants
312(1)
A.8 Enthalpies and Gibbs Energies of Formation from the Elements in the Standard States, at T = 298.15 K = 25°C and P = 1.00 bar
313(4)
A.9 Heat Capacities of Gases in the Ideal Gas State
317(2)
Appendix B Equilibrium with other Restraints Other Approaches to Equilibrium
319(4)
Appendix C The Mathematics of Fugacity, Ideal Solutions Activity and Activity Coefficients
323(6)
C.1 The Fugacity of Pure Substances
323(1)
C.2 Fugacities of Components of Mixtures
324(2)
C.3 The Consequences of the Ideal Solution Definition
326(1)
C.4 The Mathematics of Activity Coefficients
326(3)
Appendix D Equations of State for Liquids and Solids Well Below their Critical Temperatures
329(6)
D.1 The Taylor Series EOS and Its Short Form
329(1)
D.2 Effect of Temperature on Density
330(1)
D.3 Effect of Pressure on Density
331(1)
D.4 Summary
332(3)
References
333(2)
Appendix E Gibbs Energy of Formation Values
335(4)
E.1 Values "From the Elements"
335(1)
E.2 Changes in Enthalpy, Entropy, and Gibbs Energy
335(1)
E.2.1 Enthalpy Changes
335(1)
E.2.2 Entropy Changes
336(1)
E.3 Ions
337(1)
E.4 Presenting these Data
337(2)
References
337(2)
Appendix F Calculation of Fugacities from Pressure-Explicit EOSs
339(8)
F.1 Pressure-Explicit and Volume-Explicit EOSs
339(1)
F.2 f/P of Pure Species Based on Pressure-Explicit EOSs
339(1)
F.3 Cubic Equations of State
340(2)
F.4 fi/Pyi for Individual Species in Mixtures, Based on Pressure-Explicit EOSs
342(1)
F.5 Mixing Rules for Cubic EOSs
343(1)
F.6 VLE Calculations with a Cubic EOS
344(1)
F.7 Summary
345(2)
References
346(1)
Appendix G Thermodynamic Property Derivatives and the Bridgman Table
347(4)
References
350(1)
Appendix H Answers to Selected Problems
351(2)
Index 353
NOEL de NEVERS, PhD, followed five years of working for Chevron with thirty-seven years as a Professor in the Chemical Engineering Department of the University of Utah. His textbooks (and research papers) are in fluid mechanics, thermodynamics, and air pollution control engineering. He regularly consults as an expert on explosions, fires, and toxic exposures. In addition to technical work, he has three "de Nevers's Laws" in a Murphy's Laws compilation and won the title "Poet Laureate of Jell-O Salad" in a Salt Lake City competition, with three limericks and a quatrain. He has climbed the Grand Teton, Mt. Rainier, Mt. Whitney, Kala Pattar, and Mt. Kilimanjaro, and is the official discoverer of Private Arch in Arches National Park.