Chromium metal is electroplated from acidic aqueous solutions containing the dichromate ion, Cr₂O₇^2-. What is the minimum time needed to plate out 10.0 g of chromium metal from such a solution, if the current is 50.0 A?

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Quizplus · Accepted Answer

To calculate the time, use Faraday's laws of electrolysis. First, determine the moles of chromium: 10.0 g Cr / 51.996 g/mol = 0.1923 mol Cr. The reduction of Cr2O7^2- to Cr involves 6 electrons per Cr atom. Therefore, 0.1923 mol Cr requires 0.1923 mol x 6 mol e^- = 1.154 mol e^-. Using Faraday's constant (96485 C/mol e^-), the total charge needed is 1.154 mol e^- x 96485 C/mol e^- = 111,333 C. With a current of 50.0 A, time = charge/current = 111,333 C / 50.0 A = 2226.66 seconds = 37.1 minutes.

Chromium Metal Is Electroplated from Acidic Aqueous Solutions Containing the Dichromate