CVEN 3454/5404 Water Chemistry

Problem Set 7  Redox Reactions

1.  Use data from Tables 8.7 and 9.3 and Appendix A of Benjamin (2002) to write the full reactions described by the following situations.  Using half-cell reactions, calculate the equilibrium constants (K), standard pe values (pe0), and standard redox potentials (EH0) at 25°C of the full reactions.  Then, for the conditions given, calculate the pe and EH values for the non-standard state conditions.  State whether or not the reactions will proceed spontaneously under the conditions given.

  1. The oxidation of hydrogen sulfide (H2S) to sulfate by oxygen ([H2S(aq)] = 10-4 M, [SO42-] = 10-5 M, PO2 = 0.18 bar, pH 5.0).
  2. The reduction of nitrate to nitrogen gas (N2(g)) by the oxidation of Mn2+ to MnO2(s) ([NO3-] = 10-3 M, PN2 = 0.79 bar, [Mn2+] = 10-4 M, pH = 8.0).
  3. The reduction of chromate to Cr(OH)3(s) by the oxidation of zinc metal to Zn2+ ([CrO42-] = 10-6 M, [Zn2+] = 10-6 M, pH = 6.0).

2.  Determine the concentrations of the following species:

  1. the concentration of nitrate in a water at pH 8 containing [NH4+] = 10-6 M, [H2S] = 10-6 M, and [SO42-] = 10-3 M.
  2. the concentration of ferric iron in a water at pH 7 containing [Fe2+] = 10-6 M in equilibrium with the atmosphere containing PO2 = 0.21 bar.

3.  Problem 3, Chapter 9.  Treating a mercury solution with iron filings.

 

Last updated on November 11, 2007 at 08:56 PM by Joe Ryan