Chapter Learning Goals Assessed on Final Exam
Ch. 12
1. 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)
2. 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)
3. 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)
4. Determine the order of a reaction and the rate constant from plots of ln[A] versus
time, 1/[A] versus time, and [A] versus time. From the appropriate graph,
calculate the rate constant. (12.8, 12.11, 12.60)
5. 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)
6. 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)
7. 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)
8. Describe how concentration of reactants, temperature, activation energy, and
catalysts affect reaction rates. Sketch a potential energy profile to illustrate how
reaction rates depend on activation energy and the effect of a catalyst on the
activation energy. (12.44, 12.92, 12.18, 12.90, 12.94, 12.96)
9. Use the Arrhenius equation to solve for any variable given the other variables.
(12.17, 12.82, 12.84, 12.86)
Ch. 13
1. Given Kc or Kp and initial concentrations or partial pressure of reactants and
products, calculate the equilibrium concentrations or partial pressures of all
reactants and products. (13.11-13.15, 13.72, 13.74, 13.76, 13.78)
2. Determine the direction of a reaction when a stress is placed on a system at
equilibrium. Consider changes in concentrations of reactants and products, pressure,
volume, and temperature, and the addition of a catalyst. Describe how temperature
and the presence of a catalyst affect the equilibrium constant. (13.16-22, 13.34,
13.36, 13.38, 13.80, 13.82, 13.84, 13.86, 13.88)
3. Describe the state of chemical equilibrium and the meaning of the equilibrium
constant. (13.28)
4. Write the equilibrium constant expression given any balanced equation and solve
for K given the equilibrium concentrations or partial pressures. (13.1, 13.7, 13.60,
13.2-4, 13.40, 13.44, 13.46, 13.48, 13.50, 13.52, 13.54, 13.70 )
5. Explain how the ratio of the rate constants for the forward and reverse reactions is
related to the equilibrium constant and solve problems using this relationship.
(13.23, 13.90, 13.92, 13.94)
6. Solve for the reaction quotient and determine the direction of reaction to reach
equilibrium. (13.9-10, 13.20, 13.32, 13.68)
Ch. 14
1. Know which acids and bases are strong and which are weak. Define an acid and
base based on the Bronsted-Lowry definitions. Write chemical equations for
proton transfer reactions, identify the conjugate acid-base pairs. Write the formula
of the conjugate form of an acid or base. For weak acids and bases, relate the
strength to the Ka or Kb. (14.32, 14.48, 14.68, 14.1-3, 14.31, 14.42, 14.44, 14.46)
2. Calculate the [H3O+] from the [OH-] and vice versa. Calculate the pH and pOH
from the [H3O+] and the [OH-], respectively, and vice versa. Classify the solution
as acidic, basic or neutral. (14.6-9, 14.52, 14.54, 14.56, 14.58, 14.60)
3. Given the concentration of a strong acid or strong base, calculate the pH. (14.1011, 14.62, 14.64)
4. Given the pH of a solution of a weak acid, calculate the Ka. Given the Ka and
concentration of a weak acid, calculate the pH and percent dissociation. Write a
balanced net ionic equation for the dissociation of weak acids. Solve similar
problems for a weak base. (14.12-16, 14.19-20, 14.32, 14.66, 14.68, 14.70, 14.72,
14.74, 14.82, 14.84, 14.86)
5. Classify salt solutions as either acidic, basic, or neutral. Calculate the Ka of an
acid using the Kb of the conjugate base and vice-versa. Calculate the pH of the
solutions. Write net ionic equations for hydrolysis reactions. (14.21, 14.88,
14.22-25, 14.36, 14.90, 14.92, 14.94)
6. Identify the strongest acid in a group of binary or oxoacids. (14.26, 14.38, 14.96,
14.98, 14.100)
7. Identify substances that can act as Lewis acids and bases. (14.27-28, 14.102,
14.104, 14.106)
Ch. 15
1. Describe the common ion effect and predict changes in pH and solubility. (15.35, 15.52, 15.54, 15.56, 15.58, 15.26, 15.98, 15.102)
2. Identify solutions which can act as buffers. Given the initial concentrations of a
weak acid or base and its conjugate, calculate the pH of a buffer. Use the
Henderson-Hasselbalch equation to calculate the pH of a buffer or to calculate the
concentrations of an acid-base conjugate pair to make a buffer of a particular pH.
Calculate the pH of a buffer after the addition of acid or base. (15.6-12, 15.60,
15.66, 15.68, 15.70, 15.72)
3. Calculate pH values and explain the composition of the solution at any point
during a titration. From the relative strengths of the acid and base in a
neutralization reaction, determine whether the solution will be acidic, basic or
neutral at the equivalence point. Write balanced net ionic equations for
neutralization reactions Calculate the concentrations of all species present at any
point during a titration. (15.1-2, 15.44, 15.48, (15.13-16, 15.18-19, 15.37-39,
15.78, 15.80, 15.82, 15.84, 15.86, 15.88)
4. Write the solubility equilibrium and solubility product constant for a given ionic
compound. Given the Ksp of an ionic compound, calculate the solubility and
concentration of ions present. Given the solubility of an ionic compound,
calculate the Ksp. Given the Ksp, determine whether a precipitate will form on
mixing solutions of ionic compounds. (15.20, 15.40, 15.90, 15.21-24, 15.92,
15.94, 15.96, 15.29-30, 15.110, 15.112-114)
Ch. 16
1. State and explain the three laws of thermodynamics. Discuss how enthalpy and
entropy changes contribute to the driving force of chemical reactions. (16.20,
16.22, 16.24, 16.26, 16.28, 16.54, 16.62)
2. Use the equation G = H - T S to calculate the free energy change of a
reaction, determine if a reaction is spontaneous based on signs of H, S, and G
and to estimate the temperature at which a reaction becomes spontaneous. (16.711, 16.72, 16.78)
3. Calculate the equilibrium constant from the standard free energy change or
calculate the standard free energy change from the equilibrium constant using
G° = -RTlnK. (16.86. 16.88, 16.90, 16.92)
4. Calculate the standard free energy change of a reaction from the standard free
energies of formation of reactants and products. Write a balanced equation for a
formation reaction. (16.12, 16.70, 16.74, 16.76, 16.78, 16.80)
5. Calculate the free energy change of a reaction under nonstandard conditions using
G = G° + RTlnQ. (16.82, 16.84)
6. Calculate standard molar entropy changes for reactions from the standard molar
entropies of reactants and products. (16.5, 16.48, 16.52)
7. Sketch a graph of free energy versus reaction composition (reactants → products)
given G° and indicate the point of equilibrium. (16.27)
Ch. 11
1. Explain how temperature and pressure affect solubility and do calculations using
Henry's law. (11.11-12, 11.70, 11.72, 11.74)
2. Explain the colligative properties: vapor pressure lowering, boiling point
elevation, freezing point depression and osmotic pressure. Calculate these
properties for solutions.(11.13-26, 11.76, 11.78, 11.80, 11.82, 11.84, 11.86, 11.88,
11.90, 11.92, 11.94, 11.96, 11.98)
3. Define density, molarity, mole fraction, weight percent, ppm, ppb, and molality
and perform calculations using these quantities. Be able to convert between
concentration units. (11.3-10, 11.48, 11.50, 11.52, 11.54, 11.56, 11.58, 11.60,
11.62, 11.64, 11.66, 11.68)
Ch. 17
1. Sketch a galvanic cell and identify the anode and cathode reactions, the sign of
each electrode, and the direction of flow of electrons and ions. Write balanced
chemical equations for reactions occurring in a galvanic cell. Write and interpret
shorthand notations for galvanic cells. (17.1-4, 17.26, 17.28, 17.36, 17.38, 17.40,
17.42, 17.44, 17.46)
2. Use a table of standard reduction potentials to calculate the standard cell potential.
Use the table to rank substances in order of increasing oxidizing or reducing
strength and to determine if reactions are spontaneous. Use the table to predict
the half-reactions when an aqueous solution of a salt is electrolyzed. (17.6-9,
17.56, 17.58, 17.60, 17.64, 17.66, 17.96, 17.98)
3. Calculate the standard free energy change from the cell potential and vice versa.
(17.52, 17.54)
4. Use the Nernst equation (E = E° - (0.0592 V)lnQ/n) to calculate the cell potential
for a reaction under nonstandard conditions. (17.10-12, 17.32, 17.68, 17.70,
17.72, 17.74)
5. Calculate an equilibrium constant given the standard cell potential (and vice
versa) for a redox reaction using E° = (0.0592 V/n)log K. (17.13-14, 17.76,
17.78, 17.80, 17.82, 17.84, 17.86)
6. For an electrolytic cell, interconvert quantities of current, time, charge, moles of
electrons and moles and grams or liters of products. (17.22-23, 17.96, 17.98,
17.100, 17.102, 17.104)
Ch. 22
1. Write balanced equations for nuclear reactions and identify the types of radiation
and nuclides involved. (22.1-2, 22.30, 22.32, 22.34, 22.36, 22.38, 22.40, 22.74)
2. Use the integrated rate law for first order reactions to solve for the half-life, rate
constant, ratio of nuclei at time t to nuclei initially present, time required for the
decay of a certain percentage of reactivity, and age of an object based on rates of
radioactive decay. (22.3-7, 22.42, 22.44, 22.46, 22.48, 22.50, 22.52, 22.54, 22.56,
22.58)
3. Describe the types of radioactivity and their properties. (22.24, 22.26, 22.28)