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6.13: Additional Resources

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    The following experiments involve the experimental determination of equilibrium constants, the characterization of buffers, and, in some cases, demonstrations of the importance of activity effects.

    • “The Effect of Ionic Strength on an Equilibrium Constant (A Class Study)” in Chemical Principles in Practice, J. A. Bell, Ed., Addison-Wesley: Reading, MA, 1967.

    • “Equilibrium Constants for Calcium Iodate Solubility and Iodic Acid Dissociation” in Chemical Principles in Practice, J. A. Bell, Ed., Addison-Wesley: Reading, MA, 1967.

    • “The Solubility of Silver Acetate” in Chemical Principles in Practice, J. A. Bell, Ed., Addison-Wesley: Reading, MA, 1967.

    • Cobb, C. L.; Love, G. A. “Iron(III) Thiocyanate Revisited: A Physical Chemistry Equilibrium Lab Incorporating Ionic Strength Effects,” J. Chem. Educ. 1998, 75, 90–92.

    • Green, D. B.; Rechtsteiner, G.; Honodel, A. “Determination of the Thermodynamic Solubility Product, Ksp, of PbI2 Assuming Nonideal Behavior,” J. Chem. Educ. 1996, 73, 789–792.

    • Russo, S. O.; Hanania, I. H. “Buffer Capacity,” J. Chem. Educ. 1987, 64, 817–819.

    • Stolzberg, R. J. “Discovering a Change in Equilibrium Constant with Change in Ionic Strength,” J. Chem. Educ. 1999, 76, 640–641.

    • Wiley, J. D. “The Effect of Ionic Strength on the Solubility of an Electrolyte,” J. Chem. Educ. 200481, 1644–1646.

    A nice discussion of Berthollet’s discovery of the reversibility of reactions is found in

    • Roots-Bernstein, R. S. Discovering, Harvard University Press: Cambridge, MA, 1989.

    The following texts provide additional coverage of equilibrium chemistry.

    • Butler, J. N. Ionic Equilibria: A Mathematical Approach; Addison-Wesley: Reading, MA, 1964.

    • Butler, J. N. Solubility and pH Calculations; Addison-Wesley: Reading, MA, 1973.

    • Fernando, Q.; Ryan, M. D. Calculations in Analytical Chemistry, Harcourt Brace Jovanovich: New York, 1982.

    • Freiser, H.; Fernando, Q. Ionic Equilibria in Analytical Chemistry, Wiley: New York, 1963.

    • Freiser, H. Concepts and Calculations in Analytical Chemistry, CRC Press: Boca Raton, 1992.

    • Gordus, A. A. Schaum’s Outline of Analytical Chemistry; McGraw-Hill: New York, 1985.

    • Ramette, R. W. Chemical Equilibrium and Analysis, Addison-Wesley: Reading, MA, 1981.

    The following papers discuss a variety of general aspects of equilibrium chemistry.

    • Cepría, G.; Salvatella, L. “General Procedure for the Easy Calculation of pH in an Introductory Course of General or Analytical Chemistry,” J. Chem. Educ. 2014, 91, 524–530.

    • Gordus, A. A. “Chemical Equilibrium I. The Thermodynamic Equilibrium Concept,” J. Chem. Educ. 1991, 68, 138–140.

    • Gordus, A. A. “Chemical Equilibrium II. Deriving an Exact Equilibrium Equation,” J. Chem. Educ.  1991, 68, 215–217.

    • Gordus, A. A. “Chemical Equilibrium III. A Few Math Tricks,” J. Chem. Educ. 1991, 68, 291–293.

    • Gordus, A. A. “Chemical Equilibrium IV. Weak Acids and Bases,” J. Chem. Educ. 1991, 68, 397–399.

    • Gordus, A. A. “Chemical Equilibrium VI. Buffer Solutions,” J. Chem. Educ. 1991, 68, 656–658.

    • Gordus, A. A. “Chemical Equilibrium VII. Precipitates, “J. Chem. Educ. 1991, 68, 927–930.

    • Reijenga, J.; Van Hoof, A.; van Loon, A.; Teunissen, B. “Development of Methods for the Determination of pKa Values,” Analytical Chemistry Insights, 2013, 8, 53–71.

    • Thomson, B. M.; Kessick, M. A. “On the Preparation of Buffer Solutions,” J. Chem. Educ. 1981, 58, 743–746.

    • Weltin, E. “Are the Equilibrium Concentrations for a Chemical Reaction Always Uniquely Determined by the Initial Concentrations?” J. Chem. Educ. 1990, 67, 548.

    • Weltin, E. “Are the Equilibrium Compositions Uniquely Determined by the Initial Compositions? Properties of the Gibbs Free Energy Function,” J. Chem. Educ. 1995, 72, 508–511.

    Collected here are a papers that discuss a variety of approaches to solving equilibrium problems.

    • Ault, A. “Do pH in Your Head,” J. Chem. Educ. 1999, 76, 936–938.

    • Chaston, S. “Calculating Complex Equilibrium Concentrations by a Next Guess Factor Method,” J. Chem. Educ. 1993, 70, 622–624.

    • Donato, H. “Graphing Calculator Strategies for Solving Chemical Equilibrium Problems,” J. Chem. Educ. 1999, 76, 632–634.

    • Glaser, R. E. Delarosa, M. A.; Salau, A. O.; Chicone, C. “Dynamical Approach to Multiequilibria Problems for Mixtures of Acids and Their Conjugate Bases,” J. Chem. Educ. 2014, 91, 1009–1016.

    • Olivieri, A. C. “Solution of Acid-Base Equilibria by Successive Approximations,” J. Chem. Educ. 1990, 67, 229–231.

    • Weltin, E. “A Numerical Method to Calculate Equilibrium Concentrations for Single-Equation Systems,” J. Chem. Educ. 1991, 68, 486–487.

    • Weltin, E. “Calculating Equilibrium Concentrations,” J. Chem. Educ. 1992, 69, 393–396.

    • Weltin, E. “Calculating Equilibrium Concentrations for Stepwise Binding of Ligands and Polyprotic Acid-Base Systems,” J. Chem. Educ. 1993, 70, 568–571.

    • Weltin, E. “Equilibrium Calculations are Easier Than You Think - But You do Have to Think!” J. Chem. Educ. 1993, 70, 571–573.

    • Weltin, E. “Calculating Equilibrium Concentrations by Iteration: Recycle Your Approximations,” J. Chem. Educ. 1995, 72, 36–38.

    Additional historical background on the development of the Henderson-Hasselbalch equation is provided by the following papers.

    • de Levie, R. “The Henderson Approximation and the Mass Action Law of Guldberg and Waage,” Chem. Educator 2002, 7, 132–135.

    • de Levie, R. “The Henderson-Hasselbalch Equation: Its History and Limitations,” J. Chem. Educ. 2003, 80, 146.

    A simulation is a useful tool for helping students gain an intuitive understanding of a topic. Gathered here are some simulations for teaching equilibrium chemistry.

    • Edmonson, L. J.; Lewis, D. L. “Equilibrium Principles: A Game for Students,” J. Chem. Educ. 1999, 76, 502.
    • Huddle, P. A.; White, M. W.; Rogers, F. “Simulations for Teaching Chemical Equilibrium,” J. Chem. Educ. 2000, 77, 920–926.

    The following papers provide additional resources on ionic strength, activity, and the effect of ionic strength and activity on equilibrium reactions and pH.

    • Clark, R. W.; Bonicamp, J. M. “The Ksp-Solubility Conundrum,” J. Chem. Educ. 1998, 75, 1182– 1185.
    • de Levie, R. “On Teaching Ionic Activity Effects: What, When, and Where?” J. Chem. Educ. 2005, 82, 878–884.
    • McCarty, C. G.; Vitz, E. “pH Paradoxes: Demonstrating That It Is Not True That pH = –log[H+],”J. Chem. Educ. 2006, 83, 752–757.
    • Ramshaw, J. D. “Fugacity and Activity in a Nutshell,” J. Chem. Educ. 1995, 72, 601–603.
    • Sastre de Vicente, M. E. “The Concept of Ionic Strength Eighty Years After Its Introduction,” J. Chem. Educ. 2004, 81, 750–753.
    • Solomon, T. “The Definition and Unit of Ionic Strength,” J. Chem. Educ. 2001, 78, 1691–1692.

    For a contrarian’s view of equilibrium chemistry, please see the following papers.

    • Hawkes, S. J. “Buffer Calculations Deceive and Obscure,” Chem. Educator, 1996, 1, 1–8.
    • Hawkes, S. J. “What Should We Teach Beginners About Solubility and Solubility Products?” J. Chem. Educ. 1998, 75, 1179–1181.
    • Hawkes, S. J. “Complexation Calculations are Worse Than Useless,” J. Chem. Educ. 1999, 76, 1099–1100.
    • Hawkes, S. J. “Easy Deviation of pH ≈ (pKa1 + pKa2)/2 Using Autoprotolysis of HA: Doubtful Value of the Supposedly More Rigorous Equation,” J. Chem. Educ. 2000, 77, 1183–1184. See, also, an exchange of letters between J. J. Roberts and S. J. Hawkes, J. Chem. Educ. 2002, 79, 161–162.

    This page titled 6.13: Additional Resources is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by David Harvey.

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