Chemical Engineering Thermodynamics II

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  • Chemical Engineering Thermodynamics II

    Chemical Engineering

    Thermodynamics II

    (CHE 303 Course Notes)

    T.K. Nguyen

    Chemical and Materials Engineering

    Cal Poly Pomona

    (Winter 2009) Contents

    Chapter 1: Introduction

    1.1 Basic Definitions 1-1

    1.2 Property 1-2

    1.3 Units 1-3

    1.4 Pressure 1-4

    1.5 Temperature 1-6

    1.6 Energy Balance 1-7

    Example 1.6-1 : Gas in a piston-cylinder system 1-8

    Example 1.6-2 : Heat Transfer through a tube 1-10

    Chapter 2: Thermodynamic Property Relationships

    2.1 Type of Thermodynamic Properties 2-1

    Example 2.1-1 : Electrolysis of water 2-3

    Example 2.1-2 : Voltage of a hydrogen fuel cell 2-4

    2.2 Fundamental Property Relations 2-5

    Example 2.2-1 : Finding the saturation pressure 2-8

    2.3 Equations of State 2-11

    2.3-1 The Virial Equation of State 2-11

    Example 2.3-1 : Estimate the tank pressure 2-12

    2.3-2 The Van de Walls Equation of State 2-13

    Example 2.3-2 : Expansion work with Van de Walls EOS 2-15

    2.3-3 Soave-Redlick-Kwong (SRK) Equation 2-17

    Example 2.3-3 : Gas Pressure with SRK equation 2-17

    Example 2.3-4 : Volume calculation with SRK equation 2-18

    2.4 Properties Evaluations 2-21

    Example 2.4-1 : Exhaust temperature of a turbine 2-21

    Example 2.4-2 : Change in temperature with respect to pressure 2-24

    Example 2.4-3 : Estimation of thermodynamic property 2-26

    Example 2.4-4 : Heat required to heat a gas 2-27

    Chapter 3: Phase Equilibria

    3.1 Phase and Pure Substance 3-1

    3.2 Phase Behavior 3-4

    Example 3.2-1 : Specific volume from data 3-7

    3.3 Introduction to Phase Equilibrium 3-11

    3.4 Pure Species Phase Equilibrium 3-12

    3.4-1 Gibbs Free Energy as a Criterion for Chemical Equilibrium 3-12

    3.4-2 The Chemical Potential 3-13

    3.4-3 Vapor Liquid Phase Equilibrium 3-16

    Example 3.4-1 : Horsepower of a compressor 3-18

    Example 3.4-2 : Two-level refrigeration system 3-19

    3.4-4 The Clapeyron Equation 3-23

    Example 3.4-3 : Triple point estimation 3-27

    3.5 Refrigeration 3-30

    i Example 3.5-1 : A vapor compression refrigeration cycle 3-31

    Example 3.5-2 : A refrigeration cycle using refrigerant 134a 3-32

    Example 3.5-3 : A refrigeration cycle using refrigerant 134a 3-34

    Example 3.5-4 : A refrigeration cycle with nonideal compressor 3-35

    3.6 Partial Molar Properties 3-36

    Example 3.6-1 : Partial molar volume 3-37

    Example 3.6-2 : Enthalpy of mixing 3-38

    Example 3.6-3 : Enthalpy-concentration chart 3-39

    Chapter 4: Principle of Phase Equilibrium II

    4.1 The Phase Rule 4-1

    Example 4.1-1 : The degree of freedom 4-1

    Example 4.1-2 : Degree of freedom for vapor-liquid equilibrium 4-2

    4.2 The Fugacity 4-3

    Example 4.2-1 : Fugacity of liquid water 4-5

    Example 4.2-2 : Fugacity using Van der Waals EOS 4-6

    Example 4.2-3 : Fugacity of steam 4-10

    Example 4.2-4 : Change in Gibbs energy 4-11

    4.3 Fugacity of Species i in a Gas Mixture 4-15

    Example 4.3-1 : Fugacity using Peng-Robinson EOS 4-17

    4.4 Fugacity in the Liquid Phase 4-19:20

    Chapter 5: Applied Phase Equilibrium

    5.1 Vapor-Liquid Equilibrium for Ideal Systems 5-1

    Example 5.1-1 : Bubble point temperature calculation 5-3

    Example 5.1-2 : Bubble point pressure calculation 5-5

    Example 5.1-3 : Dew point temperature calculation 5-6

    Example 5.1-4 : Dew point pressure calculation 5-8

    Example 5.1-5 : Txy diagram for benzene-toluene mixture 5-9

    Example 5.1-6 : Dew point pressure using K-values 5-11

    5.2 Isothermal Flash Calculation Using K-values 5-15

    Example 5.2-1 : Isothermal flash with V/F specified 5-18

    Example 5.2-2 : Isothermal flash with T specified 5-20

    5.3 Vapor-Liquid Equilibrium with Non-ideal Liquid 5-23

    Example 5.3-1 : Bubble point pressure with Van Laar model 5-25

    Example 5.3-2 : Bubble point temperature with Wilson model 5-26

    Example 5.3-3 : Dew point temperature with Wilson model 5-28

    Example 5.3-4 : Txy diagram with Wilson model 5-31

    5.4 Fitting Activity Coefficient Models with VLE Data 5-35

    Example 5.4-1 : Evaluation of Margules and Van Laar parameters 5-37

    Example 5.4-2 : Evaluation of Wilson parameters 5-43

    5.5 Azeotropes 5-45

    Example 5.5-1 : Evaluation of Wilson parameters 5-50

    5.6 Estimation of Activity Coefficients 5-52

    Example 5.6-1 : Acetone mole fraction in a system with air 5-52

    Example 5.6-2 : Solubility of ethane in n-heptanol 5-54

    5.7 Phase Behavior in Partially Miscible Systems 5-55

    ii Example 5.7-1 : Liquid mixture in two separate phases 5-56

    Example 5.7-2 : Liquid mixture in a single phase 5-56

    Example 5.7-3 : LLE for a binary mixture 5-61

    5.8 Vapor-Liquid-Liquid Equilibrium: VLLE 5-63

    Example 5.8-1 : Composition and pressure for a VLLE system 5-64

    Example 5.8-2 : VLLE calculations 5-66

    Example 5.8-3 : VLLE estimation 5-68

    Example 5.8-4 : VLLE for water and hydrocarbon 5-71

    5.9 The Thermodynamics of Osmosis 5-73

    Example 5.9-1 : Estimation of the PVC molecular weight 5-76

    Example 5.9-2 : Dissociation of NaCl molecules 5-78

    5.10 Distribution of a Solute Between Liquid Phases 5-79

    5.10a Solubility of a Solid in a Liquid Phase 5-79

    Example 5.10-1: Solubility of a Drug 5-81

    5.10a Distribution of a Solute Between Liquid Phases 5-84

    Example 5.10-2: Drug Extraction from the Aqueous Phase 5-87

    Example 5.10-3: Purification of an Antibiotic 5-88

    5.10c Single-Stage Equilibrium Extraction 5-89

    Example 5.10-4: Drug Extraction from the Aqueous Stream 5-90:91

    Chapter 6: Chemical Equilibrium

    6.1 Introduction 6-1

    Example 6.1-1 : Extent of reaction and fractional conversion 6-2

    Example 6.1-2 : Time to reach 90% conversion in a batch reactor 6-5

    6.2 Chemical Reaction and Gibbs Energy 6-6

    6.3 The Condition of Equilibrium for a Chemical Reaction 6-9

    6.4 Calculation of Equilibrium Constant from Data 6-11

    Example 6.4-1 : Equilibrium constant for methanol reaction 6-11

    Example 6.4-2 : Equilibrium constant for nitrogen oxide 6-12

    6.5 Variation of Equilibrium Constant with Temperature 6-14

    Example 6.5-1 : Methanol reaction at 60oC 6-17

    Example 6.5-2 : Pyrolysis of methanol 6-18

    6.6 Homogeneous Gas Phase Reaction 6-20

    Example 6.6-1 : Decomposition of Ethane 6-21

    Example 6.6-2 : Partial pressure of monatomic hydrogen 6-23

    6.7 Heterogeneous Reaction 6-30

    Example 6.7-1 : Activity of water 6-31

    Example 6.7-2 : Equilibrium conversion for isomerization 6-32

    Example 6.7-3 : Dissociation of CaCO (s) 6-33


    6.8 Thermodynamics of Pack Cementation 6-34

    6.9 Equilibrium in Electrochemical Systems 6-37

    Example 6.9-1 : Copper etching 6-40

    6.10 Complex Chemical Equilibrium 6-41

    Example 6.10-1 : Number of independent reactions 6-42

    Example 6.10-2 : Number of independent reactions 6-43

    iii Appendix

    Appendix A: Solving Algebraic Equations

    A.1 The Newton-Raphson Method A-1

    Example A.1-1 : Newton-Raphson method for a root in [1, 2] A-2

    A.2 Newton’s Method for Systems of Nonlinear

    Algebraic Equations A-3

    Example A.2-1 : Newton method for 3 nonlinear equations A-4

    Solving set of nonlinea r equations with Excel A-6

    Appendix B: Curve Fitting

    B.1 Nonlinear Curve Fitting B-1

    Example B.1-1 : Fit the function T(t; ε, h) = ε(1 − e−ht) to the data B-4

    Appendix C: Process Simulator (Provision)

    Appendix D: Previous Exams

    Quiz 1 D-1

    Quiz 2 D-3

    Quiz 3 D-7

    Quiz 4 D-9

    Quiz 5 D-12

    Answer to Quizzes

    iv Chapter 1


    1.1 Basic Definitions

    Thermodynamics is the science that seeks to predict the amount of energy needed to

    bring about a change of state of a system from one equilibrium state to another. While

    thermodynamics tells us nothing about the mechanisms of energy transfer, rates of change,

    and time associated with a system changing from one equilibrium state to another, it is still

    the lynch-pin that allow us to answer these questions.

    • Definition of 'heat': Heat is energy in transit solely as a result of a temperature


    • Definition of 'work': Work is energy exchange between system and surroundings due

    to any phenomenon except a temperature difference.

    • Definition of 'temperature': Temperature is a measure of the mean kinetic energy of

    molecules. Absolute zero (0oK) is a state of complete motionless of molecules.

    • 'Rate': 'Rate' implies an element of speed, how fast an event happens, and time.

    • 'System': In thermodynamics, the universe can be divided into two parts. One part is

    the system, the other part is the rest of the universe called the surroundings. System

    can be classified as (1) isolated system where no mass or energy is transferred across

    the system boundaries, (2) closed system (system) where only energy is transferred

    across the system boundaries, or (3) open system (control volume) where mass and

    energy can be transferred across the system boundaries. A system is any designated

    region of a continuum of fixed mass. The boundaries of a system may be deformable

    but they always enclose the same mass.





    Figure 1.1 Schematic diagram of the "universe", showing a system and the


    1-1 • 'Control volume': A 'control volume' is also any designated region of a continuum

    except that it may permit matter to cross its boundaries. If the boundaries of a control

    volume are such that matter may not enter or leave the control volume, the control

    volume is identical to a system. In these respects, a 'system' is a subset of a 'control


    • 'Equilibrium': 'Equilibrium' means that there are no spatial differences in the variables

    that describe the condition of the system, also called the 'state' of a system, such as its

    pressure, temperature, volume, and mass (P, T, V, m), and that any changes which

    occur do so infinitesimally slowly.

    The laws of thermodynamics are applicable only to equilibrium states which means that the

    state does not really change significantly with time, differences in variables between the state

    of a system and its surroundings are of infinitesimal magnitude and that within the system

    itself there are no spatial variations of the variables that determine its state. Using

    thermodynamics, we can predict the amount of energy needed to change a system from an

    equilibrium state to another. For example it will take about 75 kJ to change 1 kg of air at

    25oC and 1 atm to 100oC and 1 atm. It will take much more energy, about 2257 kJ, to change

    1 kg of water at 100oC and 1 atm to water vapor (steam) at the same temperature and


    State 1 75 kJ Air, 1 atm State 2

    Air, 1 atm required 100oC

    25oC 1 kg

    1 kg


    1 atm Steam

    100oC 1 atm

    1 kg 2257 kJ 100oC

    State 1 State 2

    required 1 kg

    Figure 1.1 Energy required changing air or water from state 1 to state 2.

    1.2 Property

    A property is a macroscopic characteristic of a system such as pressure, temperature, volume,

    and mass. At a given state each property has a definite value independent of how the system

    arrived at that state. The properties of air in state 1 shown in Figure 1.1 are: pressure at 1 atm,

    temperature at 25oC, and mass of 1 kg.

    A property can be classified as extensive or intensive. An extensive property depends on the

    size of the system while an intensive property is independent on the size of the system.

    Consider systems (1) and (2) shown in Figure 1.2 both at 100oC and 1 atm containing 2 and 5

    kg of steam, respectively.

    1-2 System (2)

    System (1)



    1 atm

    1 atm

    5 kg

    2 kg

    Figure 1.2 Example of intensive and extensive properties.

    Temperature, pressure, and specific volume of both systems are intensive properties. Total

    mass and total volume of each system are extensive property. At 100oC and 1 atm, the

    specific volume v of each system is 1.674 m3/kg. The mass of system (1) is m = 2 kg and


    that of system (2) is m = 5 kg. The total volume of system (1) is V = m v = (2 kg)(1.674

    2 1 1

    m3/kg) = 3.348 m3. The total volume of system (2) is V = m v = (5 kg)(1.674 m3/kg) = 8.37

    2 2

    m3. An intensive property might be obtained from an extensive propery by dividing the

    extensive property by the mass of the system.

    1.3 Units

    The SI units (Système International d'Unitès, translated Internal System of Units) are used in

    this text. It happens that seven primary quantities are needed to completely describe all

    natural phenomena1. The decision as to which quantities are primary is arbitrary. The units of

    the primary quantities and their symbols are listed in Table 1.3-1 and are defined arbitrarily

    as follows:

    Meter: the length of the trajectory traveled by light in a vacuum per 1/299,792,458 s,

    Kilogram: the mass of the platinum cylinder deposited at the International Office for

    Weights and Measures, Sèvres, France,

    Second: 9,192,631,770 times the period of radiation in energy level transitions in the fine

    spectral structure of 133Cs,

    Kelvin: 1/273.16 of the triple point temperature of water with naturally occurring amounts of

    H and O isotopes,

    Amperes: the current which, on passing through two parallel infinite conducting wires of

    negligible cross section, separated by 1 m and in vacuum, induces a force (per unit length) of

    2×10-7 N/m,

    Mole: the amount of a matter containing the number of particles equal to the number of

    atoms in 0.012 kg of the pure isotope 12C,

    Candela: the amount of perpendicular light (luminosity) of 1/60×10-6 m2 of the surface of an

    absolute black body at the melting temperature of platinum and a pressure of 101,325 Pa.

    1-3 Table 1.3-1 The seven primary quantities and their units in SI

    Primary quantity Unit

    Length Meter (m)

    Mass Kilogram (kg)

    Time Second (s)

    Temperature Kelvin (K)

    Electric current Ampere (A)

    Amount of matter Mole (mol)

    Amount of light Candela (cd)

    Several of the derived quantities with units are listed in Table 1.3-2. A derived unit is a

    quantity expressed in terms of a product and/or quotient of two or more primary units.

    Table 1.3-2 The derived quantities and their units in SI

    Derived quantity Unit

    C , specific heat capacity J/kg·K


    E, energy J = N·m, joule

    F, force N = kg·m/s2, newton

    k, thermal conductivity W/m·K

    p, pressure Pa = N/m2, pascal

    q, heat transfer rate W = J/s = kg·m2/s3, watt

    q", heat flux W/m2 = J/s·m2

    q′′′, heat generation rate per unit volume W/m3


    μ, viscosity


    ρ, density

    1.4 Pressure

    Any force acting on a surface consists of a component perpendicular to the surface and a

    component parallel to the surface. These two components are called normal force and shear

    force as shown in Figure 1.4-1. Pressure is defined as a normal force per unit area on which

    the force acts. The SI pressure unit, N/m2, is called a pascal (Pa). Pressure at any point is a

    fluid is the same in any direction.

    Normal force

    Normal stress = Normal force/A F


    Shear force


    Area (A) s

    Figure 1.4-1 Normal and parallel components of a force on a surface.

    1-4 Consider a hole in the wall of a tank or a pipe as shown in Figure 1.4-2. The fluid pressure p

    may be defined as the ratio F/A, where F is the minimum force that would have to exerted on

    a frictionless plug in the hole to keep the fluid from emerging1





    F(N) P(N/m2)


    Fluid flowing through a pipe

    Figure 1.4-2 Fluid pressure in a tank and a pipe.

    The pressure at a given position measured relative to absolute zero pressure or absolute

    vacuum is called the absolute pressure. Most pressure-measuring devices are calibrated to

    read zero in the atmosphere as shown in Figure 1.4-3. These pressure gages indicate the

    difference between the absolute pressure and the local atmospheric pressure. Pressures below

    atmospheric pressure are called vacuum pressures and are measured by vacuum gages that

    indicate the difference between the atmospheric pressure and the absolute pressure. Absolute,

    gage, and vacuum pressures are all positive quantities and are related to each other by

    P = P − P

    gage abs atm

    P = P − P

    vac atm abs


    6 7 8 gage P

    5 9 abs

    4 10

    P P

    atm 3 atm

    P 2 kPa

    vac 1





    P atm


    P = 0 P = 0

    abs abs

    Figure 1.4-3 Absolute, gage, and vacuum pressures.

    Two common pressure units are the bar and standard atmosphere:

    1 bar = 105 Pa = 0.1 Mpa = 100 kPa

    1 atm = 101,325 Pa = 101.325 kPa = 1.01325 bar = 14.696 psi

    1 R. M. Felder and R. W. Rousseau, Elementary Principles of Chemical Processes, Wiley, 2000, p.54.


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