Module 2

• Oxidation/reduction titrations


  Talks on the preparation of standard solutions of oxidants and reductants, and their applications in analytical chemistry
• Auxiliary reagents


  Convert an analyte to a single state
• Steps that precede titration

  1. Dissolving the sample

  2. Separating interferences

  3. Convert the analyte to a mixture of different oxidation states
• The analyte in an oxidation/reduction titration must be in a single oxidation state at the outset or start of titration
• Auxiliary oxidizing and reducing agents

  • Achieve a single oxidation state sample
  • React quantitatively with the analyte
  • Any excess must be readily removable as they usually interfere with titration with standard solution
• Auxiliary Reducing Reagents


  • Sticks or coils of the metal can be immersed directly in the analyte solution and the solid removed manually after reduction is judged complete
  • Finely divided metal held in a vertical glass tube through which the solution is drawn under a mild vacuum (Jones reductor)

  
• Jones reductor


  • 2 cm diameter
  • 40 to 50 cm column of amalgamated zinc
  • Amalgamation accomplished by allowing zinc granules to stand briefly in mercury (II) chloride solution
• Amalgamated zinc

  Has the important virtue of inhibiting the reduction of hydrogen ions
• Acidic solutions can be passed through the Jones reductor without significant hydrogen formation
• Principal applications of the Jones reductor and reductions that can be accomplished with Walden reductor

  • Listed in Table 2-1
• Walden reductor

  • Granular metallic silver is the reductant
  • Silver is not a good reducing agent unless in the presence of silver salt of low solubility formed from chloride or other ions
  • Prereductions with Walden reductor are generally carried out from HCl solutions of the analyte
• Sodium Bismuthate, NaBiO3


  • A powerful oxidizing agent
  • Capable of converting Mn(II) to MnO4-
• Ammonium Peroxydisulfate/Ammonium Persulfate, (NH4)2S2O8


  • Also a powerful oxidizing agent
  • In acid solution and catalyzed by traces of silver ion, it converts Cr3+ to Cr2O72-, Ce3+ to Ce4+, and Mn2+ to MnO4-
• Excess persulfate reagent decomposition

  Brief period of boiling: 2S2O82- + 2H2O → 4SO42- + O2(g) + 4H+
• Sodium Peroxide and Hydrogen Peroxide


  Convenient oxidizing agents as the solid salt or as a dilute solution of the acid
• Excess hydrogen peroxide removal

  Boiling: 2H2O2 → 2H2O + O2(g)
• Standard solutions of most reductants tend to react with atmospheric oxygen
• Reductants are seldom used for direct titration of oxidizing analytes; indirect methods are used instead
• Iron (II) solutions


  • Easily prepared from Iron(II) ammonium sulfate (Mohr's salt) or Iron(II) ethylenediamine sulfate (Oesper's salt)
  • Air-oxidation of iron (II) is inhibited in the presence of acids, with the most stable preparations being about 0.5M in H2SO4
• Determining numerous oxidants


  Treatment of the analyte solution with a measured excess of standard iron (II) followed by immediate titration of the excess with a standard solution of K2Cr2O7 or Ce4+
• Oxidants determined by the above process


  • Organic peroxides, hydroxylamine, Cr6+, Ce4+, Mo6+, nitrate, chlorate, perchlorate and numerous other oxidants
• Sodium Thiosulfate


  A moderately strong reducing agent widely used to determine oxidizing agents by an indirect method in which iodine is an intermediate
• Scheme used to determine oxidizing agents with sodium thiosulfate


  1. An unmeasured excess of KI is added to a slightly acidic analyte solution

  2. Reduction of the analyte produces a stoichiometrically equivalent amount of iodine

  3. The liberated iodine is titrated with std Na2S2O3

  
• Determination of sodium hypochlorite in bleaches


  • OCl- + 2I- + 2H+ → Cl- + I2 + H2O

  I2 + 2S2O32- → 2I- + S4O62-

  
• Detecting end points in iodine/thiosulfate titrations

  1. Disappearance of the iodine color (brown color)

  2. With starch indicator → deep blue color

  
• Starch undergoes decomposition in solution with high I2 concentration, so addition of the indicator must be deferred/delayed until most of the I2 has been reduced
• Aqueous starch suspensions


  Decompose within a few days due to bacterial action

  Inhibit rate of decomposition by preparing and storing indicator under sterile conditions, and by adding HgI2 or CHCl3 as a bacteriostat

  Simplest alternative is to freshly prepare the indicator (on the day to be used)
• Stability of sodium thiosulfate solutions


  Tendency to decompose: S2O32- + H+ → HSO3- + S(s)

  Decomposition rate-influencing variables: pH, microorganisms' presence, solution concentration, presence of Cu(II) ion, and sunlight exposure
• Standardizing thiosulfate solutions


  Primary standard: potassium iodate, KIO3

  Weighed amount of primary-std-grade KIO3 reagent are dissolved in water containing excess of KI. When acidified with strong acid: IO3- + 5I- + 6H+ → 3I2 + 2H2O

  The liberated iodine is then titrated with the thiosulfate solution.
• The stoichiometry of the reaction is 1 mol IO3- = 3 mol I2 = 6 mol S2O32-
• Example 2-1: Standardizing a sodium thiosulfate solution

  • 0.1210 g KIO3 (214.00 g/mol) dissolved in water, excess KI added, acidified with HCl. The liberated I2 required 41.64 mL of the thiosulfate solution to decolorize. Calculate the molarity of the Na2S2O3.

  
• Other primary standards for thiosulfate

  • K2Cr2O7, KBrO3, KH(IO3)2, K3[Fe(CN)6], Cu
• Applications of sodium thiosulfate solutions


  • Listed in Table 2-2
• Properties of the 5 most widely used volumetric oxidizing agents


  • Summarized in Table 2-3
• Potassium Permanganate


  MnO4- + H+ + 5e- → Mn2+ + 4H2O, E⁰=1.51V (in ≥ 0.1 M strong acid)

  Color of permanganate intense enough to serve as indicator

  Modest cost

  
• Cerium(IV)


  Ce4+ + e- → Ce3+, E⁰=1.44 V (in 1 M H2SO4), E⁰=1.70 V (in 1 M HClO4), E⁰=1.61 V (in 1 M HNO3)

  Tendency to form precipitate of basic salts in < 0.1 M acidity solution
• Comparison of Potassium Permanganate and Cerium(IV)


  • Summarized in Table 2-4
• Detecting end points in permanganate titrations


  Indicators: KMnO4 (intense purple color), Diphenylamine sulfonic acid, 1,10-phenanthroline complex of Fe(II)

  Permanganate end point is not permanent because excess permanganate ions react slowly with manganese (II) ions present at end point

  
• Detecting end points in cerium(IV) titrations


  Indicator: Fe(II) complex of 1,10-phenanthroline
• E⁰


  Standard reduction potential