st
Review for Chapters 1, 6, 8, & 21
Definitions: (The meanings of these words and phrases should be very familiar to you!)
System State of the System
Processes Surroundings
Macroscopic System
Open System
Isolated System
Closed System
Adiabatic
Reversible Process
Irreversible Process
Isothermal
Isochoric
Macroscopic System
Thermodynamic State
Equilibrium State
Ideal Gas Law
Kinetic Theory of Gases
Isobaric
Isenthalpic
Microscopic System
State Functions or State Variables
SI Units
Functions
Dependent Variable
Equation of State
Daltons Law of Partial Pressures
Supercritical Fluid
Coefficients
Independent Variable
Extensive Variable
Intensive Variable
Simple System
Boyle Temperature
Triple Point
Normal Melting Temp. (Point)
Normal Boiling Point
Critical Point
Isotherms
Reduced Variables
Internal Energy
Partial Derivative
Gibbs Phase Rule
Phase Transitions
Coexistence Curve
Normal Freezing Point
Normal Sublimation Temp.
Critical Constants
Gibbs Phase Rule
Kinetic Energy
Potential Energy
Thermodynamic Energy
Boltzmann Distribution
Probability Distribution
Identity for Change of Variables
Average Values (like average speed) Assumptions for Kinetic Theory
Calculations: (you should be familiar with the following in terms of calculations)
Van der Waals Eqn. of State Relation between molar volume and density
Multiplication Rule for Derivatives
Virial Equation of State
Compression Factor Z
Ideal Gas Law and Daltons Law
Deriving Eqns. from the Ideal Gas Law
Virial Coefficients
Compressibility of Liquid vs Gas
Critical Points from Eqn. of State
Pressure Volume Curves Pressure Temperature Curves
Calculation of Degrees of Freedom from Phase Rule
Pressure as a Function of Molar Volume at Various Constant Temp. Values

Review for Chapters 2 and 3
Definitions: (The meanings of these words and phrases should be very familiar to you!)
System Work
External Pressure
Reversible
Exactness
Heat
Standard State
Inexact Differential
Criteria for Exactness
Specific Heat
Heat Capacity
Calorimeter Constant
Law of Conservation of Energy
Isothermal
Isobaric
Reversible
Isolated System
1st Law of Thermodynamics
Closed System
Isochoric
Adiabatic
State Variable
Path Dependent
Expression for dH for Ideal Gas
Expression for dU for Ideal Gas
Molar Internal Energy
Exothermic
Hess' Law
Kirchoff's Law
STP (Standard Temperature and Pressure)
Isenthalpic
Internal Energy
Enthalpy
Endothermic
Enthalpy of Formation
H= U + PV dU = dq + dw
Cpm = Cvm + R
Bomb Calorimeter
Joule Thompson Coefficient
Cp = Cv + nR
Enthalpy of Combustion
Average Bond Energies
Permuter
Inverter Isothermal compressibility
Expressing U in Terms of Heat and Work
Use and/or derivation of Expressions for Work, Heat, Enthalpy etc of adiabatic, reversible, isothermal etc. process
Calculations with a Constant External Pressure
Calculations with a External Pressure in Equil with Internal Pressure
Use of 1st Law to Test for Allowedness of a Process
Dependence of Enthalpy on Temperature
Use of Heat Capacity to Determine the Enthalpy Change
Calculation of dH and dU from Heats of Formation, Heats of Combustion, i.e. from Hess' Law
Expression for dH and dU for ideal gas
Converting between H and U
Exactness
Relation of Exactness and State Function
Manipulation of partial derivatives with the Permuter, Inverter, Exactness, etc. giving new
Thermodynamic Expressions
The use of the Joule Thompson coefficient to determine if a substance can be used as a coolant

nd
Definitions: (The meanings of these words and phrases should be very familiar to you!)
1st and 2nd Laws of Thermodynamics Carnot Heat Engine & Assumptions dS = dq rev
/ T
State Function dS > dq

S irr
/ T rev cycl
=
 surr
State Variable cycle
dq re
/T = 0 Entropy is a State Function dS > 0 (irr adiabatic)
S = k ln W
Efficiency
Steady State
Absolute zero
Coefficient of Performance
Absolute Entropy
Troutons Rule
Heat Pump
Standard State for Entropy
G = H - TS
Claussius Inequality
Spontaneity
3rd Law of Thermodynamics
Refrigeration
Spontaneous
A = U-TS
Gibbs Energy of Reaction
Criteria, based on System Variables, for
H = U + PV
How does the entropy change as a function of temperature for a constant pressure or constant volume process (ideal gas) What is the expression for the entropy change of a constant pressure process which goes through several changes of state? In any irreversible process the entropy of the universe increases. Because entropy is a state variable, the entropy change of the system can be calculated using a reversible process having the same initial and final states as the irreversible process. In any reversible process the entropy of the universe remains constant and therefore this constitutes the minimal or maximal work case.
The Carnot efficiency tells about the maximum efficiency realizable for a process which converts heat to work. Know what the coefficient of performance is and how it related to give the work necessary in refrigeration or for a heat pump. The expression for the
Second Law of Thermodynamics  S tot
> 0 for an irreversible process is strictly applicable to the system and the surroundings. HOWEVER, new thermodynamic state variables, named the Helmholtz Free Energy and Gibbs Free Energy, were defined which determine the spontaneity of a process based on the system properties and the mechanical variables. Thus (dA)
T,V
< 0 and (dG)
T,p
< 0
(closed; PV).
Which property tells about the max non PV work? Which tells about the maximum amount of work the system can do?
Calculations: (These types of problems should be workable)
Use of the 2nd Law to test if a process is realizable (spontaneous).
General procedure for calculating entropy change in a closed system.
Calculation of entropy changes for processes that include phase transitions
Calculation of the Carnot efficiency.
Calculation of the Standard Entropy Change for a chemical reaction.

Calculation of the Entropy change for many types of processes.
Specification of the sign of the entropy change for a process.
Calculation of the Standard Gibbs Energy of Reaction
Derivation of Equations used to calculate the entropy changes for various processes.
Definitions: (The meanings of these words and phrases should be very familiar to you!)
1 st
and 2 nd
Laws of Thermodynamics Chain Rule and Product Rule
Fundamental Eqn. of Thermodynamics Gibbs Free energy
Maxwell Relation
Fugacity
Fugacity coefficient
Gibbs-Helmholtz eqn.
Defining Equation for the fugacity
Chemical Potential
Exact Differential
Standard State molar Gibbs Free Energy dU = TdS - PdV
Pressure Dependence of G
Natural Thermodynamic Variables for U,H,G, & A
Calculations: (you should be familiar with the following in terms of calculations)
Relation of Exactness and State Function to Deriving Maxwell’s Relations
Use of the Combination of 1 st and 2 nd Laws to find other Thermodynamic Relationships
Manipulation of Thermodynamic relations using partial derivative identities
Calculations to find

G at a different temperature
Calculations to find

G at a different pressure
Coming up with a value for the fugacity given the pressure.
Sketch of the graphs of the molar Gibbs Free energy as a function of the Pressure and
Fugacity
Use of the definition of the Chemical Potentials
Expression for the molar Gibbs free energy of a gas
Expression for the Fundamental Eqn of Chemical Thermodynamics or an open system.
General:
Know how the expression for the Fundamental Equation of Thermodynamics. It is essentially a combination of the 1st and 2nd Laws. Remember it was derived based on a reversible process, but is it applicable to any process within the restrictions of its derivation? What are the best thermodynamic variables for U? Be able to show how more Thermodynamic information comes from the definition of the exact differential and the exactness criterion. What are Maxwell’s relations?
How are they derived from the equation for dU? What about dG? What are the natural thermodynamic variables for dG? Be able to write out dG for these variables?
These variables are derived from the basic definition of G=H-TS. What is the Maxwell relation derived from dG?

G sys
tells us about the spontaneity of the process, and whether the system is at equilibrium. How does it change with P and T? What is the
Gibbs Helmholtz relation? Be able to use it. Does the Gibbs free energy change very

much with a pressure change on a solid or liquid? Why?
What is the fundamental equation of chemical thermodynamics? The chemical potential of a pure substance is (

G/
 n) p,T
and for a perfect gas

=
 o
+ RT ln(p/p o
) how does this change for a real gas. In general

=
 o
+ RT ln a where a is the activity. For ideal gas a = p/p o . For real gas a = f/p o . What is the fugacity coefficient? What is this eqn., f=
 p. Not only does the chemical potential show how G changes when the composition changes but also H and A when the appropriate thermodynamic variables are held constant. What needs to be held constant? Does the name chemical potential fit?
Definitions: vapor pressure phase boundaries pressure condensed phase critical point
Clapeyron Eqn. phase diagram triple point
Clausius Clapeyron partial vapor
Solid normal boiling point fluid p = p
* exp(V m
(l)

P/(RT)) (d

/dP)
T
= V m
Liquid
(d

/dT)
P
= -S m
Gas normal melting pt. standard freezing pt. supercritical
Know how the chemical potential or the Gibbs Free energy behaves before a phase change, during the phase change and afterward. How does the chemical potential vary with pressure?
How does the chemical potential vary with temperature. Why does the chemical potential of a pure substance decrease as the temp is raised. Be able to explain how the boiling pt of a pure substance increases with increasing pressure using a

vs T diagram and the fact that (d

/dP)
T
= V m
and V m
is much higher for a liquid than a solid. Why can an ice skater skate on H
2
O ice but not “dry” ice.
You should know how to find the location of phase boundaries since at equilibrium, their chemical potentials will be equal. This can be used to derive the Clapeyron Eqn. The general Clapeyron equation tells us the slope of the phase boundary at any point, and can be made more specific depending on the type of phase boundary.
You should be able to derive the Clapeyron equation that describes the specific case of a liquid-vapor, or solid-vapor boundary. What are the assumptions?
Calculations
P, T phase diagram p = p
* exp(V m
(l)

P/(RT))
(d

/dP)
T
= V m
and (d

/dT)
P
= -S m
Clapeyron Eqn
Claussius Clapeyron Eqn
The exam will cover mainly Chp.5,6,7. The following review guide should give you an

idea of what to expect on the exam. The format of the exam will follow that of the previous exams, some multiple choice or true false followed by a problem section.
Terminology
Chp. 5
Partial Molar quantity
Gibbs Duhem Eqn.
Thermodynamics
Std. Chemical Potential
Partial Molar volume
Fundamental Eqn. of Chemical
Raoults Law
Ideal Solution
Thermodynamics of Mixing
Henry’s Law
Ideal Dilute Solution positive deviation from Raoults Law Negative deviation form Raoults Law
Excess Thermodynamic Functions

- interaction parameter boiling pt. elevation
Regular Solution
Colligative properties freezing pt. depression
Boiling pt. constant
Osmosis activity in terms of molality activities of ions
Debye Huckel Limiting Law
Ionic strength
Activity solute activity chemical potential of solvent
Freezing pt. constante
Osmotic Pressure
Activity coefficient solvent activity chemical potential of solute
Margules Eqns. mean activity coefficient extended Debye-Huckel law
Chp 6
Gibbs Phase Rule phase variance, degrees of freedom ` partial pressure of components isopleth temperature-composition diagram fractional distillation theoretical plates high boiling azeotrope immiscible liquids upper critical solution temp. eutectic composition incongruent melting liquid crystals zone levelling
Chp 7. extent of reaction exergonic rxn. component phase diagram vapor pressure diagram lever rule tie line simple distillation steam distillation azeotrope low boiling azeotrope partially miscible liquids lower critical solution temp. congruent melting peritectic line zone refining reaction Gibbs energy endergonic rxn.

std. Gibbs free energy of reaction reaction quotient thermodynamic equilibrium constant dissociation
K x
, K
P
, K

, K b
or K m van’t Hoff Eqn.
Ellingham diagram equilibrium equilibrium constant degree of dissociation, fractional
Le Chatelier’s Principle
Extraction of metals from oxides
BACKGROUND INFO
The chemical potential of a pure substance is (dG/dn) p,T
and for a perfect gas

=
 o
+ RT ln(p/p o
) how does this change for a real gas. In general

=
 o
+ RT ln a where a is the activity. For ideal gas a = p/p o . For real gas a = f/p o . What is the fugacity coefficient f=
 p. Not only does the chemical potential show how G changes when the composition changes but also H and A when the appropriate thermodynamic variables are held constant. What needs to be held constant? Does the name chemical potential fit?
Know how the chemical potential or the Gibbs Free energy behaves before a phase change, during the phase change and afterward.
Chapter 5
In previous chapters applications of the thermodynamics were applied to the case of pure substances. In Chapter 5, applications were extended to the thermodynamics of mixtures.
You should know that the total volume of a mixture is not necessarily the sum of the volumes of the pure substances since the volume occupied by the molecules is dependent upon the interactions of the molecules surrounding them. The partial molar volume describes this dependence and is the slope of the graph of the total volume as the amount of a component is changed. Once the partial molar volume is known at any composition, the total volume can be obtained. How?
Partial molar quantities can be extended to other state functions. One example is the partial molar Gibbs function, the chemical potential. Remember the Gibbs Duhem equation relates the changes in the chemical potentials of the components.
Systems tend toward a lower Gibbs Free energy. This allows us to apply thermodynamics to the discussion of spontaneous changes of composition. One specific application is the mixing of two gases. Look at

G of mixing to decide if the mixing will occur. You should be able to determine

S,

H,

V of mixing given

G of mixing.
What are the Excess functions?
When deriving the

G of mixing for binary solutions, we must decide if the solution is ideal, or ideal dilute, etc. You should know what the partial vapor pressure and total pressure vs liquid mole fraction or vapor mole fraction look like for Raoult’s
Law or Henry’s Law obeying mixtures. Why is there deviation from Raoults Law? Why do deviations occur both positively and negatively?
All colligative properties, freezing point depression, boiling point elevation, etc

stem from the reduction of the chemical potential of the liquid solvent as a result of the addition of a nonvolatile component being added. To derive the colligative constants, we simply set the chemical potentials of the solvent plus solute equal to the pure gas or solid chemical potential. Be familiar with freezing point depression, boiling point elevation, and osmotic pressure calculations. You should be familiar with the ebullioscopic and cryoscopic constants and the definitions of molality, mole fraction, and molarity. What is the Vant Hoff’s constant, i?
How can the activity of a solvent or solute be determined? How is the standard state for the activity different depending on whether we are talking about a solute or a solvent ideal; solution or ideal dilute solution. What is the relationship between the activity coefficients and the activity? The activity coefficients go to one in the limit of high solvent concentration or very low solute concentration. Again this is specific to the standard state applicable in a particular situation. What are the Margules factors. How do they relate to the Gibbs Free energy of mixing and the Excess Gibbs energy. How are they related to the activities and activity coefficients? Remember that for an ionic solution, the activity coefficients used are the mean activity coefficients. Replacing the activities with the molalities works only for dilute solutions. The mean activity coefficient can be derived from the Debye Huckel limiting law or the extended law with works at slightly higher ionic strengths, but remains poor near 1molal concentraiton.
Chapter 6
Background for this is the ability to use the Gibss phase rule and identify the various phase regions of a one component P vs T phase diagram.
You should be able to interpret pressure-composition, and temp. vs composition diagrams. You should be able to determine the composition of the vapor phase given information about the liquid phase composioin and the pure component vapor pressures for an ideal two component mixture. What does a plot of y
A
vs x a
look like? How about a P vs y
A
diagram or a P vs x
A
diagram from and ideal solution. In a two component phase diagram, we plot pressure as a function of z
A
(both y
A
and x
A
) or temp. as a function of z
A
(both y
A
and x
A
). Remember that pressure is being set for a temp. vs composition diagram and temp. is fixed for a pressure vs composition diagram. How do you know which component’s mole fraction is going from 0 to 1 from left to right on the x axis? Be able to describe the various regions on the phase diagrams, the number of phases, the number of degrees of freedom, and the various compositions present if an isopleth is followed. The lever rule can be used to determine the relative amounts
(moles) of the two phases present at a given overall concentration and temperature or pressure. Know how to show a simple distillation or fractional distillation on a temperature vs composition diagram. What is an azeotrope. How does it look on a temp vs composition diagram, and how does the way it looks correspond to a positive or negative deviation from Raoults law? Why are there both low boiling and high boiling azeotropes?
Be able to interpret liquid-liquid diagrams (temp-composition) for partially miscible liquids, liquid vapor temp vs composition diagrams for partially miscible liquids, and liquid-solid phase diagrams (temp-composition). Again, as before, you should know what phases are present in a particular region of the diagram as well as the # of degrees of freedom. Know how to label the axes and points on the axes. Know the

terms lower critical temperature, upper critical temperature, and eutectic point. Be able to identify incongruent melting and the presence of a product compound. Although there are three constituents there are only two components, why? What is the eutectic halt. Be able to construct a two component phase diagram from information about boiling point, melting points, eutectic points and general lines. Know how to use the triangular phase diagrams for three component systems, What is the difference in the diagram for miscible and partially miscible systems. Know what the constant two component molar ratio looks like on the diagram.
CHAPTER 7
Spontaneous change at constant temperature and pressure is in the direction of lower values of the Gibbs function. Why is a plot of G vs extent,
 a curve with a minimum rather than just a straight line. Know the definition of the extent of a reaction. How is it related to the coefficients in the balanced chemical eqn? What is the definition of the rxn
Gibbs function? Is it a partial molar function? What is the difference between the reaction Gibbs function and the standard reaction Gibbs free energy change, G o rxn
?
Remember for the reaction A
<-
- > B, the change per mole when 1 mol of A in its std state is converted to 1 mol of B in its std state is G o rxn
and these values come from the G o f
.
What do the terms endergonic and exergonic mean? Know the definition of the reaction
Gibbs function in terms of the reaction quotient Q. What is the value of
 r
G at equilibrium? How is the equilibrium constant related to G o rxn
? What is the relationship between the reaction quotient and the equilibrium constant?
You should be able to express the reaction quotient in most general terms as a function of the activities of the reactants and products, and more specifically in terms of pressures, fugacities, activity coefficients times molality, etc. The form of the expression depends on whether the system is an ideal gas or solution, or real gas or solution. What is the relationship between the various equilibrium constants such as K = K x
K p
or K = K

K m
Does the equilibrium constant change with pressure? Does the equilibrium constant change with temperature? Know the relations. What is the Van’t Hoff’s equation?. How does Le Chatelier’s principle fit with the effects of pressure and temperature on equilibrium?
You should be familiar with the Ellingham diagram and its use for determining whether or not the metal oxides can be converted to metal by adding carbon to them or at what temperature that this could occur.
Electrochemistry will not be covered on this exam.
Exam Specifics
Problems will be similar to homework problems. In Chp. 5 we focused on how thermodynamic parameters were a function of the composition. Calculations related to partial molar quantities or use of Henry’s Law or Raults law to determine the solubility or vapor pressure of a component or total pressure as a function of composition or a plot of pressure vs mole fraction for ideal and ideal dilute solutions. As you have seen many of

the problems from Chp 6 include relating x
A
and y
A,
interpretation of regions on various types of temperature vs. composition, or pressure vs. composition diagrams for a two component system exhibiting varying miscibility. You should be familiar with the determination of distillation parameters such as # of theoretical plates, eutectic compositions, incogruent melting, upper critical points, etc. Also, there may be a question on ternary diagrams, three component phase diagrams
Expect a problem or two calculating the equilibrium constant or the equilibrium composition for a particular reaction, and the temperature and pressure dependence of the equilibrium constant or the mole fraction part of the equilibrium constant
Look at quizzes for potential problems.

You should recognize the difference between a galvanic or voltaic cell and an electrolytic cell. What are the anode and the cathode. Be able to identify the type of half reactions that take place at the anode and cathode. You should be familiar with cell notation. How is DG = -nFE ? What about the equilibrium constant?
The Nernst Eqn. gives the cell potential at various concentrations.
the Nernst Eqn. How can one use the definition of the Standard Hydrogen Electrode to determine the cell potentials of other half reactions? These half reactions are in an electrolyte.
The thermodynamic properties of electrolyte solutions are discussed in terms of the chemical potentials and activities in much the same way as non-electrolyte solutions, but ions interact strongly with one another through their electric charges and deviations from ideality are important even at very low concentrations. Many ion reactions involve the transfer of electrons, and can be studied by allowing them to take place in an electrochemical cell. In this way we can tabulate thermodynamic information about solutions much like was done using Hess’s law and heats of formation or enthalpies of combustion in chapter 2.