CHEM 345.001 - Principles of Physical Chemistry 1 - Fall 2003
11/14/03
COURSE CONTENT GUIDE: Keyed to Textbook Chapters
Physical Chemistry, David W. Ball, Cole/Brooks (2003), ISBN 0-534-26658-4

Chapter 1: Gases and the Zeroth Law of Thermodynamics (1.1-1.6)
  1. At-home review: SI units and conversions between them; pressure unit conversions (atm, Pa, torr); kelvin and centigrade scale temperature conversions; ideal gas law; Dalton's Law of Partial Pressures.
  2. Basic curve fitting, numerical integration and numerical differentiation: Worksheet #1
  3. Ideal gas law including written in terms of the molar mass and density of the gas.
  4. Relationships between Avogadro's number (NA), Boltzmann's constant (kB) and the ideal gas constant (R)
  5. P-V diagram of an ideal gas (non-condensing).
  6. Expressing the composition (mole fractions) of an ideal gas mixture in terms of partial pressures: xi = Pi/PT .
  7. The use of partial derivatives to obtain slopes of PV isotherms, PT isochors, TV isobars, etc for the ideal gas.
  8. P-V diagram of a (condensing) real gas and be able to identify: liquid region, gas region, two-phase region, critical point.
  9. The compression factor (Z) gas and how this factor can be used to describe the deviation of a real gas from ideal gas behavior.
  10. Equations of state: ideal gas, van der Waals, virial (expanded in P and in 1/V).
  11. The significance of the two parameters of the van der Waals equation of state.
  12. Solving the van der Waals equation for P, T or V
Chapter 2: The First Law of Thermodynamics (2.1-2.7;2.8-2.11)
  1. At-home review: Hess' Law and its application to chemical reactions
  2. Expansion work: given by dW = - PextdV; reversible expansion work given by dW = - PdV ; calculating isothermal expansion work.
  3. Numerical calculation of work: Worksheet #2.
  4. Reversible, isothermal expansion of an ideal gas.
  5. Operational definitions of constant pressure heat capacity (CP = dqP/dT) and constant volume heat capacity (CV = dqV/dT).
  6. Explanation of why CP = CV for liquids and solids but CP >> CV for gases.
  7. Empirical descriptions of the variation of the heat capacity of a substance with temperature.
  8. Calculation of the amount of heat required to change the temperature of a substance, given its heat capacity.
  9. The First Law stated in differential and in integral form:
                   dU = dq + dw
                   DU = q + w
  10. Definition of enthalpy H = U + PV and be able to use it to calculate DH.
  11. Enthalpy changes, internal energy changes and heat flow: dH = dqP and that dU = dqV.
  12. Formal definitions of constant pressure heat capacity and constant volume heat capacity: CP=(dH/dT)P   and   CV=(dU/dT)V
  13. Calculating enthalpy changes associated with heating a substance at constant pressure.
  14. Calculation of standard enthalpies of reaction (DHo) at 298K from tables of standard enthalpies of formation (DHfo) at 298 K.
  15. Analytical method for estimating standard enthalpies of reaction (DHo) at any temperature using temperature independent heat capacities
  16. Numerical method for determining DHo at any temperature using temperature dependent heat capacity data: Worksheet #3.
Chapter 3: The Second and Third Laws of Thermodynamics (3.1;3.4-3.5;3.7-3.8)
  1. Classical definition of entropy by: dS = dqrev/T.
  2. Calculating entropy changes of substances during isobaric heating from their heat capacity functions.
  3. Enthalpy changes accompanying phase transitions.
  4. Calculation of the standard entropy of reaction (DSo ) at 298K from tables of absolute entropies (So) at 298 K.
  5. Analytical method for estimating the standard entropy of reaction (DSo ) at any temperature from tables of absolute entropies (So) at 298 K and temperature independent heat capacities.
  6. Numerical method for calculating the standard entropy of reaction (DSo ) at any temperature using temperature dependent heat capacities: Worksheet #4.
  7. Entropy changes associated with the isothermal expansion of an ideal gas.
  8. Obtaining entropies of mixing of ideal gases.
Chapter 4: Free Energy and Chemical Potential (4.1-4.8)
  1. General criterion for reaction spontaneity.
  2. Definition the Gibbs free energy (G) and its relationship to non-expansion work.
  3. Second Law in terms of the Gibbs energy change applied to a constant T,P process.
  4. Using the Summary of Equations resulting from the Laws of Thermodynamics
  5. Change of the Gibbs energy of liquids, solids and gases with pressure.
  6. Using the temperature variation of the Gibbs free energy of reaction (DG) to estimate DS of reaction and DH of reaction.
  7. Calculation of the standard Gibbs free energy of reaction (DGo ) at 298K from tables of standard Gibbs free energies of oformation (DGfo) at 298 K.
  8. Analytical estimation of the standard Gibbs free energy of reaction (DGo ) at any temperature using temperature independent enthalpy of reactions (DH).
  9. Numerical determination of the standard Gibbs free energy of reaction (DGo ) at any temperature using temperature dependent heat capacities: Worksheet #4.
  10. Partial molar quantities: partial molar volume and partial molar Gibbs free energy ("chemical potential")
Chapter 5: Introduction to Chemical Equilibrium (5.1-5.3; 5.5)
  1. Equilibrium: thermal (DT=0), mechanical (DP=0) and chemical (Dm=0 and DG=0).
    2. .
  2. The general equlibrium constant for real systems: gases, solutes, solvents, pure liquids, pure solid
  3. Relationship between DG and the reaction quotient
  4. Relationship between DGo and the equlibrium constant
  5. Estimating changes in the equilibrium constant with temperature assuming a constant enthalpy of reaction.
Chapter 7: Equilibria in Multiple-Component Systems (7.1-7.7)
  1. Raoult's Law and its application to relating the composition of an ideal binary (non-electrolyte) solution to the composition of the vapor in equilibrium with the solution.
  2. Henry's Law and its application to relating the composition of a dilute binary (non-electrolyte) solution to the composition of the vapor in equilibrium with the solution.
  3. Real Solutions: deviations from Raoult's Law and from Henry's Law.
  4. Calculating activity coefficients from vapor pressure data.
  5. Writing the chemical potential for components of a non-electrolyte solution
  6. Estimating solubilities of solids in liquid solvents
  7. Colligative Properties: Boiling point elevation, freezing point depression, osmotic presure
Chapter 8: Electrochemistry and Ionic Solutions (8.3-8.7)
  1. Standard electrodes, half-reactions and standard electrode potentials
  2. Standard cell potentials, overall cell reactions, DGo and Keq for the cell reaction.
  3. Nernst equation: variation of the cell potential with composition.
  4. Non-standard cell potentials, overall cell reactions and DG for the cell reaction.
  5. Electrolyte dissociation in solution.
  6. Activities of ions in solution: empirical variation with ionic strength and the predictions of various theories.
Chapter 10: Introduction to Quantum Mechanics (10.1-10.12)
  1. Heisenberg Uncertainty Principle: position/momentum, time/energy.
  2. Basic assumptions of quantum mechanics.
  3. Normalizing the wave function.
  4. Use of the wave function and operators to obtain system properties.
  5. General use of the Schrödinger equation and its features.
  6. Results for a free particle in a one-dimensional box: wave functions, energies, transitions, zero point energy.
  7. Results for a free particle in a three-dimensional box:wave functions, energies, transitions, degeneracy.
  8. Quantum mechanical tunneling: finite potential energy barriers.
Chapter 11: Model Systems (11.1; 11.3-11.5;11.7)
  1. Diatomic vibrations as a harmonic oscillator: wave functions, energies, transitions, zero point energy, tunneling.
  2. Rotation as a free rigid rotor: wave functions, energies, transitions, zero point energy.
Chapter 14: Rotational and Vibrational Spectroscopy (14.4-14.6;14.8-14.12
  1. General features of spectroscopy: Bohr condition, line width, line intensity, line energy (frequency), selection rules, lifetimes of excited states, line width.
  2. Boltzmann population distributions.
  3. Requirements in order to observe an infrared, Raman, microwave absorption bands.
  4. Electron transitions in a linear conjugated hydrocarbon treated as free particles in a box
  5. The infrared spectrum of a diatomic molecule treated as a non-rotating harmonic oscillator.
  6. The microwave spectrum of a diatomic molecule treated as a rigid free rotor.
  7. Calculating the number of normal modes of vibration in a polyatomic species: Worksheet #5
Chapter 19: The Kinetic Theory of Gases (19.3-19.6)
  1. The Maxwell speed distribution.
  2. Most probable speed, mean speed, mean square speed, root-mean-square speed.
  3. Collision rates between particles.
  4. Collision rates with the wall.
  5. Effusion and diffusion.
Chapter 20: Kinetics (20.1-20.7;20.10-20.11)
  1. Definition of the reaction rate in terms of reactant or product appearance rates.
  2. The experimental rate law.
  3. The 1st and 2nd order reaction in one component: rate expressions; integrated rate laws for product and reactant; half-lives.
  4. Numerical estimations of reactant/product concentrations directly from the rate expressions: Worksheet #6.
  5. Elementary reactions: rate expression derived from its stoichiometry.
  6. Reaction mechanisms: collections of elementary reactions.
  7. Reversible unimolecular reaction: rate expression; integrated rate laws for product and reactant.
  8. Two consecutive irreversible unimolecular reactions: rate expression; integrated rate laws for product, reactant and intermediate.
  9. Approximations for predicting experimental rate laws: rate determining step; fast equilibrium followed by a rate determining step; steady state assumption.
  10. Transition state theory.