At the end of this course you should be able to do the following:Chapter 2

Define the terms open system, closed system, Isolated system, diathermic boundary, and adiabatic boundary.

Decide if a system is open, closed or isolated. Decide if a boundary is diathermic or adiabatic. List the three properties that define an ideal gas. Use the ideal gas equation to solve a variety of problems (e.g. find final volume or the

molar mass of the gas molecules etc). State Dalton’s Law for mixtures of ideal gases. Use Dalton’s Law to calculate the partial pressures of the individual ideal gases in a

mixture. Sketch Z (compressibility factor) vs P for one mole of an ideal gas at 25.00oC. State under what conditions of pressure real gases behave most ideally. Explain which will have the greater pressure, a real gas or an ideal gas (assuming n, V

and T are the same). Explain the effect that attractive forces will have on the pressure of a gas. Explain which van der Waals’ factor (a or b) seems to be responsible for the fact that Z >

1 at high pressures, and also which factor seems to be responsible for the fact that Z < 1 at moderately high pressures.

Which van der Waals’ factor seems to be responsible for the second situation above (a or b)? Please explain your answer

When given a list of gas molecules, predict which molecule will have the largest a van der Waals value, and which will have the smallest. Predict which will have the largest b van der Waals value, and which will have the smallest.

Use the van der Waals equation to determine the pressures of real gases. List the three things that define the kinetic molecular theory (KMT) model. Plot the approximate Maxwell speed distribution function for (i) various molecules at the

same temp, and (ii) one molecule at different temps. State that any ideal gas at a given T will always have the same Etrans (K.E.). Calculate root mean square speed, average speed and most probable speed for a pure

gaseous sample of molecules or atoms. State the three properties that collision frequency depends upon. Calculate the collision frequency, Z1 or Z11 for a sample of gaseous molecules. Calculate the mean free path for a sample of gaseous molecule.

Chapter 3 After a physical process is described, decide whether any work is done, if not state why

not, and if so, state if work was done on or by the system (and also state whether the work is positive or negative).

Calculate the work done when a gas expands, under a variety of conditions. Explain the difference between a reversible and an irreversible expansion. Using words, state the first law of thermodynamics. State the two common ways that the energy of a system can be altered. Calculate w, q, U and H for a variety of processes. Explain why Cp is greater than Cv. Calculate Cp for a substance. Use Hess’s Law to calculate Ho values from other Ho values. Calculate rHo for a reaction from the relevant rHf

o values. Explain why rHo may be different for a given chemical reaction at different

temperatures. Calculate rHo at T2 if you know (or can easily calculate) rHo at T1.

Chapter 4 State one version of the second law of thermodynamics (in words). Describe a simple example of a heat engine. State that not all heat can be turned into work. Predict whether rS of a reaction with be positive or negative by just looking at the

chemical equation for the reaction. Calculate mixS for a mixing process (for the mixing of ideal gases). Calculate fusS and vapS for phase transitions. Calculate S for a substance that is heated from one temp to another. Calculate S for an ideal gas that undergoes a reversible expansion. State the 3rd Law of thermodynamics Explain why all entropies are positive. Calculate rSo from tables of So values. Explain briefly why some So values are greater than others. For a reaction, calculate rS, Ssurr, and Stot and decide if the reaction will be

spontaneous or not. State that Gibbs Energy is for constant pressure processes and Helmoltz Energy is for

constant volume processes. Calculate rGo for a reaction, either from fGo or from rHo and rSo, and decide whether

or not the reaction will be spontaneous. Explain briefly the significance of rGo in terms of “free energy” and in terms of

“electrical work”. State (in words) what happens to the Gibbs Energy of a substance if the T is increased, or

if the P is increased.

Calculate rG for a reaction at a certain pressure if you know rGo (at 1 bar), assuming all gases involved in the reaction are behaving ideally.

Chapter 5 Calculate G for a mixture (if you know moles and chemical potentials). Calculate for a gas in a mixture of ideal gases. Calculate mixG and mixH for a mixing process (for the mixing of ideal gases). State the two characteristics of an ideal solution. For an ideal solution, calculate the difference between for liquid #1 in the mixture, and

for liquid #1 in its pure state. Use Raoult’s Law for an ideal solution to calculate the vapor pressures of the components

in the mixture, and to calculate the mole fractions of the components in the vapor phase over the mixture, if you know the composition of the mixture.

State the one characteristic of an ideal-dilute solution. Use Henry’s Law to calculate the molality of the solute in an ideal-dilute solution. Calculate the activity of a substance in a real solution if you know the concentration and

the activity co-efficient.

Chapter 6 Calculate K from rGo and vice versa. Give an equation that relates K for a reaction to the concentrations (or partial pressures)

of reactants and products. Convert partial pressure to fugacity for real gases if you know the fugacity coefficient. State that Kp and Kc do not depend on pressure. Calculate degree of dissociation for a reaction at 1bar if you know T and rGo. Calculate the partial pressures at equilibrium of all components in the reaction discussed

in the bullet above. Calculate the following, using the van’t Hoff equation: rHo if you know K at two temps,

or K at a temp if you know rHo and K at another temp. Predict what happens to the equilibrium of a reaction (and the K for a reaction) when the

T is changed, or when the P is changed (either by (i) reducing the volume or (ii) by adding an inert gas at const P or const V).

Chapter 9 Express the rate of reaction in terms of the rate of change of substance with time (e.g. rate

= - ½ d[A]/dt etc) Express the rate of reaction as a rate law, with the orders unknown and therefore given as

m, n etc (e.g. rate = k[A]m[B]n). State that the rate law (the k, m and n values) can only be determined by experiment. Define the “reaction half-life” in words.

Use the integrated rate laws and the half-life rate laws from table 9.2 (given in the exam) to calculate the following types of things: how long it will take to reduce the amount of reactant by a certain amount/percentage, how much reactant will be present at a certain time , what is the half-life of a compound if a certain percentage decomposes in a certain number of minutes etc.

Explain briefly how the integration method is used to find rate laws. Decide on the overall order of a reaction after looking at plots of e.g. [A]t vs t, 1/[A]t vs t

etc. Calculate the orders (m, n etc) for a reaction if you are given a table where each row is a

run of the same chemical reaction (though each run starts with different concentrations of reactants). The first column will be initial rate, the second column will be initial concentration of reactant 1, and the third column will be initial concentration of reactant 2. Since the exam is only 50 mins, you will be given examples that will not need a plot.

Define a “mechanism” in one sentence. Define an “elementary step” in one sentence. Decide whether an elementary step is “unimolecular”, “bimolecular” or “termolecular”. Explain briefly the difference between a “transition state” and an “intermediate” in a

chemical reaction. Decide which of two proposed mechanisms is most reasonable for a chemical reaction

after looking at the experimentally determined rate law for the reaction. Write the rate law (rate = k[A]2 etc) for an elementary step (whether it is a reversible

elementary step or not). State that for a reversible elementary step, k1/k-1 = K Sketch approximate curves for [A] vs t, [B] vs t, [C] vs t when you have the consecutive

reaction: A → B → C (where B is a reactive intermediate). Apply the steady state approximation to the above reaction and obtain equations for [A] =

? and [B] = ? Use the Arrhenius equation to calculate activation energies when you know rate

constants, and rate constants when you know activation energies etc. Look at a potential energy contour plot, and decide where the low and high P.E. areas are,

where the reactants and products are, where the transition state is, and what the probable path of the reaction will be.

Use Collision Theory (and appropriate equations) to calculate ZAA, ZAB, zAB, rate of reaction (assuming no steric hindrance issues), k for a reaction (assuming no steric hindrance issues).

Use Transition State Theory (and appropriate equations) to calculate k for a reaction (from K# or from Ho# and So#), or to calculate Go# from k.

Chapter 11 State one of the postulates of quantum mechanics. Operate on with H, and then obtain an equation for E.

Explain in a few words the meaning of Ψ? Explain in a few words the meaning of Ψ2 (or ΨΨ*)? Explain briefly in words how the Hamiltonion Operator is derived Explain what is meant by the phrase “wavefunctions are orthoganol” Mathematically prove that wavefunctions are orthoganol. Mathematically prove that wavefunctions are normalized. Use use equn 9 in worksheet #9 (from class 26) to find expectation (average) values. E.g.

<x> (particle-on-line), or <r> (hydrogen atomic orbitals). For the particle-on-a-line (or particle-in-a-1D-box) model specifically:

o state how nx and energy are related.o state how the length of line and energy are relatedo state how the number of nodes and are energy related.o sketch Ψnx for any nx, or sketch Ψ2nx for any nx

o Predict the average and most probable locations of the particle on the line for a given nx

o Calculate the probability of finding the particle between two values for length. For the hydrogen atomic orbitals specifically:

o State that Y = R, Y = .o State that all of the electronic energies (the energies of the electrons in the

orbitals) are negative.o State that energy depends on n, but not l or m (ie 2s has same E as 2p orbital in

H).o Find the values of r where a radial node exists for a given orbital.o Find the value of for a given r for a given orbital.o Sketch s orbitals and p-orbitals.o Calculate the probability of finding the electron between two values of r for a

given orbital.o Calculate the most probable distance of the electron from the nuclues (rmp) for a

given orbital.o Calculate the energy of an electron in an orbital.o Calculate the ionization energy for the hydrogen atom (if the atom is in the

ground state).o Decide from which orbitals it will be harder to remove an electron from, and

explain your answer briefly.o Explain which series (Lyman or Balmer) includes the higher energy photons.o Calculate the energy of the photon that is released or absorbed if an electron

undergoes a transition (e.g. if an electron moves from a 1s orbital to a 2s orbital). State that the fourth quantum number (ms) can have the values +1/2 and -1/2. State that no two electrons in an atom can have the same set of four quantum numbers. Give ground state electron configurations for many-electron atoms and ions (following

the general rules for filling and removing electrons from atomic orbitals). Decide whether many-electron atoms and ions are diamagnetic or paramagnetic.

Explain, using the concepts of shielding and effective nuclear charge, why an e- in a given orbital is higher/lower energy than another e- in a different given orbital (e.g. explain why an e- in a 2s orbital is lower in energy than an e- in a 2p orbital in C).

Explain, using the concepts of shielding and effective nuclear charge, why an e- in an H atom with a specific “n” value (an e- in a specific shell) always has the same energy regardless of the l and ml values (regardless of whether it is in an s or a p orbital etc).

State that for large atoms, the e- in the 4s orbital are often very similar in energy to the e- in the 3d orbitals.

Predict the lowest electron configuration if you have a theoretical atom where the 4s orbital is exactly the same energy as the 3d orbitals, and after filling the 3p orbitals there are 2 more electrons that need to go into these 3d/4s orbitals (and explain your prediction).

State that the only true way to find the ground state electron configuration for atoms/ions with more than 18 e- is with experiments.

Predict the number of kinetic energy terms, the number of positive potential energy terms and the number of negative potential energy terms in the Hamiltonian operator for a variety of atoms other than H.

State that the Schrodinger Equation cannot be solved for any atom with more than 1 e-. State that approximate solutions to the Schrodinger equation can be found using

computational methods such as variational theory and the self-consistent field method. Define First Ionization Energy in one sentence, and define Second Ionization Energy in

one sentence. Predict trends in Ionization Energies (and explain your predictions).

o E.g. if given several elements, arrange them in order of increasing first ionization energy, and also predict which elements have particularly large first ionization energies.

o E.g. If you are given a single element, predict which ionization energy will be the highest (1st , 2nd, 3rd Ionization Energy etc.), and predict which ionization energies might be significantly greater than which others.

Calculate the ionization energy (in J/mol) for an electron in an element, if you know the frequency of the light used to eject the electron and you know the K.E. of the ejected electron.

Chapter 12 Use the theory of Linear Combination of Atomic Orbitals (LCAO) to predict how many

bonding molecular orbitals and how many anti-bonding molecular orbitals will form from a given set of atomic orbitals are combined in a molecule.

For diatomic molecules, use LCAO to predict approximate shapes of bonding and anti-bonding Molecular orbitals that might form when two atomic orbitals combine in a molecule.

When given an appropriate MO energy level diagram for a diatomic molecule: fill the diagram with the appropriate number of electrons, give the electron configuration of the molecule, give the bond order of the molecule, and predict whether the molecule will be diamagnetic or paramagnetic.

If given several diatomic molecules (and their appropriate MO energy level diagrams), rank the molecules in order of increasing stability.

For a given delocalized system, predict the number of bonding molecular orbitals and the number of anti-bonding molecular orbitals formed from the atomic orbitals.

Predict the number of nodes in the lowest energy molecular orbital and the number of nodes in the highest energy molecular orbital that can be formed from a set of p-orbitals combining to form a -system (from sketching the way the atomic p-orbitals came together to form these molecular orbitals).

Explain why some atom pairs are unable to form bonds (and so cannot form double bonds or be involved in e- delocalization).