Chem 2 AP

UNIT 13 WORKSHEET #1 – Redox and Electrochemical Cells

1.In each of the following half-reactions, give the species being reduced and the number of electrons needed to balance the half-reactions:

a.AgBrO3 + ?e- → Ag + BrO3-

b.HCrO4- + 7H+ + ?e- → Cr3+ + 4H2O

c.WO3 + 6H+ + ?e- → W + 3H2O

2.Identify the species in each of the following reactions that would receive electrons from the cathode and that would lose electrons at the anode in each of the following galvanic cells:

a.Au3+(aq) + Zn(s)  Au+(aq) + Zn2+(aq)

b.3 Pu6+(aq) + 2 Cr3+(aq)  2 Cr6+(aq) + 3 Pu4+(aq)

3.Regarding the following reaction:F2(g) + 2 I-(aq) → 2 F-(aq) + I2(s)

a.List the species being oxidized.

b.List the species being reduced.

c.Calculate E0cell (see Appendix M page A-32 in your textbook)

d.Identify the species which receives electrons from the cathode.

e.Identify the species donates electrons to the anode.

4.Answer the same questions that were posed in problem #3 for the following reaction:

Hg2+(aq) + Zn(s) → 2 Hg(s) + Zn2+(aq);Hg22+ + 2e- → 2 HgE0 = +0.80V

5.Given the following half-reactions:

PtCl4- + 4e-  Pt + 4Cl-E0 = 0.73 V

Fe3+ + 3e-  FeE0 = -0.037 V

a.Write the overall equation for the cell.

b.Calculate E0cell

c.Draw a diagram for the galavanic cell and identify the anode and the cathode, the

direction of electron flow and the electrode in each half-cell.

6.A galvanic cell is constructed using a chromium electrode in a 1.00-molar solution of Cr(NO3)3 and a copper electrode in a 1.00-molar solution of Cu(NO3)2. Both solutions are at 25C.

(a) Write a balanced net ionic equation for the spontaneous reaction that occurs as the cell

operates. Identify the oxidizing agent and the reducing agent.

(b)A partial diagram of the cell is shown below.

(i)Which metal is the cathode?

(ii) What additional component is necessary to make the cell operate?

(iii)What function does the component in (ii) serve?

(c) How does the potential of this cell change if the concentration of Cr(NO3)3 is changed to 3.00-

molar at 250C? Explain.

7.An electrochemical cell consists of a tin electrode in an acidic solution of 1.00 molar Sn2+

connected by a salt bridge to a second compartment with a silver electrode in an acidic solution

of 1.00 molar Ag+.

(a)Write the equation for the half–cell reaction occurring at each electrode. Indicate which half–reaction occurs at the anode.

(b)Write the balanced chemical equation for the overall spontaneous cell reaction that occurs when the circuit is complete. Calculate the standard voltage, E, for this cell reaction.

(c)Calculate the equilibrium constant for this cell reaction at 298K.

(d)A cell similar to the one described above is constructed with solutions that have initial concentrations of 1.00 molar Sn2+ and 0.0200 molar Ag+. Will the initial voltage be higher or lower than E0cell?