A galvanic cell is based on the following half-reactions:
$$\mathrm{Cu}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu}(s) \quad
\mathscr{E}^{\circ}=0.34 \mathrm{V}$$
$$\mathrm{V}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{V}(s) \quad
\mathscr{E}^{\circ}=-1.20 \mathrm{V}$$
In this cell, the copper compartment contains a copper electrode and
\(\left[\mathrm{Cu}^{2+}\right]=1.00 M,\) and the vanadium compartment contains
a vanadium electrode and \(\mathrm{V}^{2+}\) at an unknown concentration. The
compartment containing the vanadium \((1.00 \mathrm{L}\) of solution) was
titrated with 0.0800\(M \mathrm{H}_{2} \mathrm{EDTA}^{2-}\) , resulting in the
reaction
$$\mathrm{H}_{2} \mathrm{EDTA}^{2-}(a q)+\mathrm{V}^{2+}(a q)
\rightleftharpoons \mathrm{VEDTA}^{2-}(a q)+2 \mathrm{H}^{+}(a q)$$
$$\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad\quad
K=?$$
The potential of the cell was monitored to determine the stoichiometric point
for the process, which occurred at a volume of 500.0 \(\mathrm{mL}
\mathrm{H}_{2} \mathrm{EDTA}^{2-}\) solution added. At the stoichiometric
point, \(\mathscr{E}_{\text {cell}}\) was observed to be 1.98 \(\mathrm{V}\) . The
solution was buffered at a pH of 10.00 .
a. Calculate\(\mathscr{E}_{\text {cell}}\) before the titration was carried out.
b. Calculate the value of the equilibrium constant, \(K\), for the titration
reaction.
c. Calculate \(\mathscr{E}_{\text {cell}}\) at the halfway point in the
titration.
