Chemistry

Problem

An industrial process electrolyzes molten

\mathrm{CrCl}_3

in a fused-salt cell. A current of 2.00 A is maintained for exactly 2.00 h. Molar mass of Cr: 52.00 g

\text{mol}^{-1}

Faraday constant:

F = 9.6485 \times 10^{4}\,\text{C}\,\text{mol}^{-1}

The only cathode reaction is

\mathrm{Cr}^{3+} + 3\,e^{-} \rightarrow \mathrm{Cr(l)}

. a. Part 1: a) Calculate the total charge delivered to the cell over the 2.00 h run. b. Part 2: b) Determine the number of moles of chromium metal produced. c. Part 3: c) Compute the mass of chromium collected at the cathode. d. Part 4: d) Suppose the current efficiency for chromium deposition is later found to be 88.0 %. Re-evaluate the actual mass of Cr obtained under this efficiency.

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