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Homework 3: Electrochemistry

These homework problems are suggested and will not be turned in for review. However, answers will be available for them the following week by your class TAs. For more homework feel free to go to the Homework page.

Q19

For an electrochemical cell in which the spontaneous reaction is

3Cu2+ + 2Al → 2Al3+ + 3Cu

what will be the qualitative effect on the cell potential if we add ethylenediamine, a ligand that coordinates strongly with Cu2+ but not with Al3+?

Q20

A copper-zinc battery is set up under standard conditions with all species at unit activity. Initially, the voltage developed by this cell is 1.10 V. As the battery is used, the concentration of the cupric ion gradually decreases, and that of the zinc ion increases. According to Le Chatelier’s principle, should the voltage of the cell increase or decrease? What is the ratio, Q, of the concentrations of zinc and copper ions when the cell voltage is 1.00 V?

Q21

A galvanic cell consists of a rod of copper immersed in a 2.0M solution of CuSO4 and a rod of iron immersed in a 0.10M solution of FeSO4. Using the Nernst equation and the reduction potentials for

Fe2+ + 2e- → Fe
Cu2+ + 2e- → Cu

Calculate the voltage for the cell as described.

Q22

What voltage will be generated by a cell that consists of a rod of iron immersed in a 1.00M solution of FeSO4, and a rod of manganese immersed in a 0.10M solution of MnSO4?

Q23

Consider the cell: Zn | Zn2+(0.0010M) || Cu2+ (0.0010M) l Cu for which E° = +1.10 V.

Does the cell voltage, E, increase, decrease, or remain unchanged when each of the following changes is made?

  1. Excess 1.0M ammonia is added to the cathode compartment.
  2. Excess 1.0M ammonia is added to the anode compartment.
  3. Excess 1.0M ammonia is added to both compartments at the same time.
  4. H2S gas is bubbled into the Zn2+ solution.
  5. H2S gas is bubbled into the Zn2+ solution at the same time that excess 1.0M ammonia is added to the other solution.

Q24

Consider the galvanic cell Zn | Zn2+ || Cu2+ | Cu
Calculate the ratio of [Zn2+] to [Cu2+] when the voltage of the cell has dropped to 1.05 V, 1.00 V, and 0.90 V. Notice that a small drop in voltage parallels a large change in concentration. Therefore, a battery that registers 1.00 V is quite "run down."

Q25

25. Show that hydrogen peroxide (H2O2) is thermodynamically unstable and should disproportionate to water and oxygen.

Q26

A silver electrode is immersed in a 1.00M solution of AgNO3. This half-cell is connected to a hydrogen half-cell in which the hydrogen pressure is 1.00 atm and the H+ concentration is unknown. The voltage of the cell is 0.78 V. Calculate the pH of the solution.

Q27

Using half-reactions, show that Ag+ and I- spontaneously form AgI(s) when mixed directly at unit initial activity. Show that Ksp for Agl is 10-16.

Q28

A standard hydrogen half-cell is coupled to a standard silver half-cell. Sodium bromide is added to the silver half-cell, causing precipitation of AgBr, until a concentration of 1.00M Br- is reached. The voltage of the cell at this point is 0.072 V. Calculate the Ksp for silver bromide.

Extra Electrochemisty Problems

Electroplating of copper During the electroplating of copper from a CuSO4 solution, a current of 0.50 A passes through the cell for one hour. How many electrons pass through the cell? How many grams of copper will be deposited?

Fuel cell A current of 7.50 A is produced in a fuel cell when CO is oxidized to CO2 at the anode while oxygen is reduced at the cathode. How many grams of CO and of O2 does the cell consume per hour?

Electrolysis of salt solutions Predict the main product at each electrode when each of the following 1M solutions is subjected to electrolysis:

  1. NiBr2 with inert electrodes;
  2. NiSO4 with gold electrodes;
  3. Na2SO4 with copper electrodes.

Prediction of redox reactions Predict whether the following reactions would tend to occur to a significant extent in aqueous solution. Assume that the reactive concentrations of all species are 1M.

  1. Sn + Cd2+ → Cd + Sn2+
  2. 2I¡ + Sn4+ → I2(s) + Sn2+
  3. 2Fe2+ + Br2(l) → 2Fe3+ + 2Br-
  4. Cl2(g) + 2H2O → O2 + 4H+ + 2Cl-

Disproportionation of Cu(I) Use standard cell potentials to predict whether Cu+ will be stable with respect to disproportionation to Cu(s) and Cu2+. Calculate the equilibrium constant for this process.

Cell reactions Write the cell reaction corresponding to each of the following cells, and predict the reversible potential between the electrodes. Indicate whether this reaction will proceed to the right, and state the direction of electron flow through the external circuit when the two electrodes are connected.

  1. Fe(s) | Fe2+(1.0M ) || H+(1.0M ) | H2(1.0 atm) | Pt(s)
  2. Pt(s) | Cl2(g, 10¡ 4 atm) | Cl- (.01M ) || Cl- (.01M ) | AgCl(s) | Ag(s)
  3. Cu(s) | Cu(NO3)2(0.01M ) || Cu(NO3)2(.10M ) | Cu(s)

Equilibrium compositions An electrochemical cell consists of a silver electrode dipping into 0.10M AgNO3 and a nickel electrode dipping into 0.10 M Ni(NO3)2.

  1. Write the net reaction that occurs as this cell generates electric current, and show the reaction that occurs at each electrode. Indicate the EMF and free energy contributed by each half cell, and then calculate the same quantities for the net cell reaction.
  2. If a steady current of 0.10 A is supplied by this cell for 1.0 hour, what will be the change in mass of the nickel electrode?
  3. If the volumes of the Ni2+ and Ag+ solutions are equal and the reaction is allowed to proceed until no more current flows, what will be the approximate final concentrations of these two ionic species?

Solubility product The standard half-cell potential for the electrode Cu2+ + I- + e- → CuI is +.860 volt. Use this value, together with any other required EMF values, to calculate the solubility product constant for CuI(s).