Learning Goals for Test 1
Ch. 12
1. Use a table of concentrations and times to calculate an average rate of a reaction
over a given time period. (12.2, 12.32)
2. Use a graph of concentration versus time to calculate an instantaneous rate at a
given time. (12.34)
3. Express the relative rates of appearance of products and disappearance of
reactants from the coefficients in the balanced chemical equation. (12.1, 12.36,
12.38)
4. Explain how reaction rates depend on reactant concentrations for zero, first and
second order reactions. Explain the concept of half-life and how half-lives differ
for zero, first and second order reactions. (12.3, 12.20, 12.22, 12.24, 12.40,
12.42, 12.44)
5. Given initial concentrations and initial rates, determine the order of a reaction
with respect to each reactant, the overall order of the reaction, the rate constant,
and the initial rate for any other initial conditions. (This is the initial rate method.)
(12.4-6, 12.26, 12.46, 12.48)
6. Determine the order of a reaction and the rate constant from plots of ln
concentration versus time, 1/concentration versus time, and concentration versus
time. From the appropriate graph, calculate the rate constant. (12.8, 12.11, 12.60)
7. Use integrated rate laws for 0th, 1st and 2nd order reactions to find one variable
given values of the other variables. (12.7, 12.11, 12.50, 12.56, 12.64)
8. Use the expressions for half-life of 1st and 2nd order reactions to find the half-life
from the rate constant or vice-versa. Use the half-life to determine the amount of
reactant remaining at some point in time. Estimate the half-life of 0th , 1st and 2nd
order reactions using a graph of concentration versus time. Distinguish reaction
orders by examining these graphs. (12.9, 12.10, 12.52, 12.54, 12.58, 12.62)
9. Given a reaction mechanism, identify the reaction intermediates and catalysts,
determine the molecularity of each elementary reaction, and write a rate law for
each elementary reaction. Deduce the overall rate law predicted from the
mechanism. If given the experimentally determined rate law, determine if the
reaction mechanism is consistent with the experimental rate law. (12.12-15,
12.25, 12.28, 12.66, 12.68, 12.70, 12.72, 12.76, 12.78)
10. Explain how activation energy, collision frequency and orientation affect reaction
rates. Explain potential energy profiles and suggest structures for transition
states. (12.16, 12.28, 12.80, 12.88)
11. Graph ln k versus 1/T (Arrhenius plot) and determine the activation energy from
the slope. Use the Arrhenius equation to solve for any variable given the other
variables. (12.17, 12.82, 12.84, 12.86)
12. Sketch a potential energy profile and show the activation energy of the steps in
the reaction and how they are affected by the addition of a catalyst. Use the
diagram to explain which steps are faster. (12.18, 12.90, 12.94, 12.96)
13. Describe how concentration of reactants, temperature, activation energy, and
catalysts affect reaction rates. (12.44, 12.92)
Ch. 13
1. Write the equilibrium constant expression given any balanced equation. (13.1,
13.7, 13.60)
2. Calculate the equilibrium constant, Kc, given the equilibrium concentrations of
reactants and products or Kc for the reverse reaction. (13.2-4, 13.40, 13.44, 13.46,
13.48, 13.50, 13.52, 13.54, 13.70)
3. Calculate the equilibrium constant, Kp, given the equilibrium partial pressures of
reactants and products. (13.5, 13.42)
4. Calculate Kc from Kp and vice-versa. (13.6-7, 13.56, 13.58)
5. Given the value of Kc, determine if the reaction is product-favored, reactantfavored, or neither. (13.8, 13.62, 13.64, 13.66)
6. Describe the state of chemical equilibrium and the meaning of the equilibrium
constant. (13.28)
7. Given concentrations of reactants and products, calculate the reaction quotient to
determine if the reaction is at equilibrium. If the system is not at equilibrium,
determine the direction the reaction must proceed to reach equilibrium. (13.9-10,
13.20, 13.32, 13.68)