CHAPTER II. POTENTIOMETRY AND REDOX TITRATIONS I ...

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point, [Fe 2+ ] = [Fe 3+ ] and E + = E° for the Fe 3+ /Fe 2+ couple.The point at which V = V e /2 is analogous to the point at which pH = pK a when V =V e /2 in an acid-base titration.b) Region 2: At the Equivalent PointExactly enough Ce 4+ has been added to react with all the Fe 2+ . Virtually all ceriumis in the form Ce 3+ and virtually all iron is in the form Fe 3+ . Tiny amounts of Ce 4+and Fe 2+ are present at equilibrium.i.e., [Ce 3+ ] = [Fe 3+ ] and [Ce 4+ ] = [Fe 2+ ].At any time, Reactions 1 and 2 are both in equilibrium at the Pt electrode.Hence, both reactions contribute to the cell potential at the equivalent point,E + = 0.767 − 0.0592⋅log([Fe 2+ ]/[Fe 3+ ]) (1)E + = 1.70 − 0.0592⋅log([Ce 3+ ]/[Ce 4+ ]) (2)Adding Equations 1 and 2,2E + = 2.46 7 − 0.0592⋅log([Fe 2+ ][Ce 3+ ]/[Fe 3+ ][Ce 4+ ])Because [Ce 3+ ] = [Fe 3+ ] and [Ce 4+ ] = [Fe 2+ ] at the equivalent point,2E + = 2.46 7 ⇒ E + = 1.23 VThe cell voltage potential isE = E + − E(calomel) = 1.23 − 0.241 V = 0.99 VIn this particular titration, the equivalence-point potential is independent ofconcentrations and volumes of reactants.c) Region 3: After the Equivalent PointNow both [Ce 3+ ] and [Ce 4+ ] are known, it is convenient to calculate cell potentialby using Reaction 2 instead of Reaction 1.E = {1.70 − 0.0592⋅log([Ce 3+ ]/[Ce 4+ ])} − 0.241= 1.46 − 0.0592⋅log([Ce 3+ ]/[Ce 4+ ])}At the special point when V = 2V e , [Ce 3+ ] = [Ce 4+ ] and E + = E° for the Ce 4+ /Ce 3+couple.After the equivalence point, the cell potential levels off near 1.46 V.6
e.g., Titration of Tl + by IO − 3 in 1.00 M HCl.Titration reaction: IO − 3 + 2Tl + + 2Cl − + 6H + → ICl − 2 + 2Tl 3+ + 3H 2 OIndicator half-reaction: IO − 3 + 2Cl − + 6H + + 4e = ICl − 2 + 3H 2 O E° = 1.24 VIndicator half-reaction: Tl 3+ + 2e = Tl +E° = 0.77 VFIGURE 2-17. “Harris” Fig. 16-3 (p. 353).When the stoichiometry of the reaction is not 1:1, the curve is not symmetric. Still,negligible error is introduced if center of steepest portion is taken as end-point.Less change in voltage near equivalence point as compared to last titration curve!i.e., clearest results are achieved with strongest oxidizing and reducing agents.ii/. Redox IndicatorsA redox indicator changes color when it goes from its oxidized to its reduced state.TABLE 2-18. “Harris” Table 16-2 (p. 355).To predict the potential range over which the indicator color will change, write aNernst equation for the indicator.In(oxidized) + ne = In(reduced)E = E° − (0.0592/n)⋅log{[In(reduced)]/[In(oxidized)]}The color of In(reduced) will be observed when[In(reduced)]/[In(oxidized)] ≥ 10/1and the color of In(oxidized) will be observed when[In(reduced)]/[In(oxidized)] ≤ 1/10i.e., The color change will occur over the rangeE = (E° ± 0.0592/n) Ve.g., ferroin, E° = 1.147 VFIGURE 2-19. “Harris” (p. 354).The color change is expected to occur in the approximate range 1.088 V to 1.206 Vwith respect to standard hydrogen electrode.7
Page 1 and 2: CHAPTER II. POTENTIOMETRY AND REDOX
Page 3 and 4: ⇒ E = (0.05916/n) ⋅log(a 1 /a 2
Page 5: vi/. Gas-Sensing Probese.g., FIGURE
Page 9 and 10: Both SO − 4 + Ag 2+ are powerful
Page 11 and 12: MnO − 4 + 4H + + 3e = MnO 2 (s) +
Page 13 and 14: 4I − + O 2 (g) + 4H + → 2I 2 +

CHAPTER II. POTENTIOMETRY AND REDOX TITRATIONS I ...

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