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Electrochemistry

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    CHEM 2300, In-class Exercises, Class 25

    Name: _________________________

    1. Write the relevant half reactions and calculate Eo and K for the following reaction. Is the reaction spontaneous as written?

    \[\ce{I2(s) + 5 Br2(aq) + 6 H2O ↔ 2IO3- + 10Br- + 12H+}\nonumber\]

    \[\mathrm{E^o_{I_2} = 1.210}\nonumber\]

    \[\mathrm{E^o_{Br_2} = 1.098}\nonumber\]

     

     

     

     

     

     

     

     

     

     

     

    1. Write the line notation for a cell made up of the following half reactions if the indicator electrodes are platinum and a monoprotic acid is in solution in the cathode cell. If a balanced reaction is not given, assume that the larger Eo value is the cathode.

    \[\ce{Fe^3+ + e- ↔ Fe^2+} \hspace{50px} \mathrm{E^0 = 0.771}\nonumber\]

    \[\ce{Cr2O7^2- + 14H+ +6e- ↔ 2Cr^3+ + 7H2O} \hspace{30px} \mathrm{E^0 = 1.36}\nonumber\]

     

     

     

     

     

     
    1. Given the following reaction, assign oxidation states to each element and write the 2 half reactions (both as reductions).

    \[\ce{2 AgCl(s) + Pb(s) + 2F- (aq) ↔ 2 Ag(s) + 2Cl- (aq) + PbF2(s)}\nonumber\]

     

     

     

     

     

     

    1. Write the line notation for this cell.

     

     

     

     

     

     

    1. If the concentrations of NaF and KCl are both 0.10 M, calculate the cell potential for the cell made up of the above half reactions at 25 oC:

      Lead_HalfReactions,Potentials.pngSilver_HalfReactions,Potentials.png

     

     

     

     

     

     

     

     

     

     

     

     

     

     

    CHEM 2300, In-class Exercises, Class 26

    Name: _________________________

    1. The apparatus in the figure below was used to monitor the titration of 50.0 mL of 0.100 M AgNO3 with 0.200 M NaBr. Calculate the cell voltage at each volume of NaBr, below. 

      Fig14-7_Apparatus.png

      \[\mathrm{K_{sp}\: for\: AgBr(s) = 5.0 \times 10^{-13}}\nonumber\]

      \[\mathrm{E_{S.C.E.}=0.241\: V}\nonumber\]

      1. 1.00 mL
      2. 25.1

     

     

     

     

     

     

     

    1. Write the half reaction(s) for the following reactions, as reductions:
      1. Pt (s) | Cr3+ (0.10 M), Cr2+ (0.050 M)









         
      2. Cd(s) | Cd2+ (1 M) || H+(aq, 1 M), Mn2+ (l M), MnO4- (l M) | Pt











         
      3. Cr2O72-(aq), Cr3+(aq), HA(aq) |Pt(s)

     

     

     

     

     

     

     

     

     
    1. A solution contains 0.100 M Ce3+, 1.00x10-4 M Ce4+, 1.00x10-4 M Mn2+, 0.100 M MnO4-, and 1.00M HClO4.
      1. Write a balanced net reaction that can occur between species in this solution, assuming the Ce half cell is the cathode.
      2. If Ecell is measured at -0.019 V, calculate the equilibrium constant for the balanced net reaction.

    \[\mathrm{K=10^{n*E^o/0.05916}}\nonumber\]

     

     

     

     

     

     

     

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