91Ó°ÊÓ

Identify each statement as true or false. Rewrite each false statement to make it true. (a) Oxidation always occurs at the anode of an electrochemical cell. (b) The anode of a discharging voltaic cell is the site of reduction and is negative. (c) Standard-state conditions for electrochemical cells are a concentration of \(1.0 \mathrm{M}\) for dissolved species and a pressure of 1 bar for gases. (d) The potential of a voltaic cell does not change with temperature. (e) All product-favored oxidation-reduction reactions have a standard cell potential \(E_{\text {cell }}^{\circ}\), with a negative sign.

Draw a diagram of each cell. Label the anode, the cathode, the species in each half-cell solution, the direction of electron movement in an external circuit, and the direction of movement of ions within the cell. (a) \(\mathrm{Cu}(\mathrm{s})\left|\mathrm{Cu}^{2+}(\mathrm{aq}) \| \mathrm{Fe}^{2+}(\mathrm{aq})\right| \mathrm{Fe}(\mathrm{s})\) (b) \(\mathrm{Pt}(\mathrm{s})\left|\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}), \mathrm{H}^{+}(\mathrm{aq}) \| \mathrm{Fe}^{2+}(\mathrm{aq}), \mathrm{Fe}^{3+}(\mathrm{aq})\right| \mathrm{Pt}(\mathrm{s})\)

In principle, a battery could be made from aluminum metal and chlorine gas. (a) Write a balanced equation for the reaction that would occur in a battery using \(\mathrm{Al}^{3+}(\mathrm{aq}) \mid \mathrm{Al}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{~g}) \mid \mathrm{Cl}^{-}(\) aq \()\) half-cells. (b) Identify the half-reaction at the anode and at the cathode. Do electrons flow from the \(\mathrm{Al}\) electrode when the cell does work? Explain. (c) Calculate the standard potential, \(E_{\text {cell }}^{\circ}\), for the battery.

Calculate the cell potential of a concentration cell that contains two hydrogen electrodes if the cathode contacts a solution with \(\mathrm{pH}=7.8\) and the anode contacts a solution with (conc. \(\left.\mathrm{H}^{+}\right)=0.05 \mathrm{M}\).

Consider the electrolysis of water in the presence of very dilute \(\mathrm{H}_{2} \mathrm{SO}_{4}\). What species is produced at the anode? At the cathode? What are the relative amounts of the species produced at the two electrodes?

A student wanted to measure the copper(II) concentration in aqueous solution. For the cathode half-cell she used a silver electrode with a 1.00 -M solution of \(\mathrm{AgNO}_{3}\). For the anode half-cell she used a copper electrode dipped into the aqueous sample. The cell gave \(E_{\text {cell }}=\) \(0.62 \mathrm{~V}\) at \(25^{\circ} \mathrm{C}\). Calculate the copper(II) ion concentration of the solution.