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)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  2. Describe the state of chemical equilibrium and the meaning of the equilibrium constant. (13.28)
  3. 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 )
  1. 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)
  1. 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)
  1. 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)
  1. Given the concentration of a strong acid or strong base, calculate the pH. (14.10-11, 14.62, 14.64)
  1. 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)
  1. 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)
  1. Identify the strongest acid in a group of binary or oxoacids. (14.26, 14.38, 14.96, 14.98, 14.100)
  1. 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. (15.3-5, 15.52, 15.54, 15.56, 15.58, 15.26, 15.98, 15.102)
  1. 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)
  1. 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)
  1. 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 DG = DH - T DS to calculate the free energy change of a reaction, determine if a reaction is spontaneous based on signs of DH, DS, and DG and to estimate the temperature at which a reaction becomes spontaneous. (16.7-11, 16.72, 16.78)
  1. Calculate the equilibrium constant from the standard free energy change or calculate the standard free energy change from the equilibrium constant using DG° = -RTlnK. (16.86. 16.88, 16.90, 16.92)
  1. 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)
  1. Calculate the free energy change of a reaction under nonstandard conditions using DG = DG° + RTlnQ. (16.82, 16.84)
  1. Calculate standard molar entropy changes for reactions from the standard molar entropies of reactants and products. (16.5, 16.48, 16.52)
  1. Sketch a graph of free energy versus reaction composition (reactants → products) given DG° 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)
  1. 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)
  1. 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)
  1. 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)
  1. Calculate the standard free energy change from the cell potential and vice versa. (17.52, 17.54)
  1. 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)
  1. 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)
  1. 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)
  1. 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)
  1. Describe the types of radioactivity and their properties. (22.24, 22.26, 22.28)