(may be broke/outdated!)
An antique automobile bumper is to be chrome plated. The bumper, which is dipped into an acidic Cr2O72- solution, serves as a cathode of an electrolytic cell. The atomic mass of Cr is 51.996; 1 faraday = 96,485 coulombs. If oxidation of H2O occurs at the anode, how many moles of oxygen gas will evolve for every 1.00 x 102 grams of Cr(s) deposited?
How do you set this problem up? Any help would be great thanks!
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One Response
First you need a balanced chemical equation.
Presumably your reduction is
14 H(+) + 12 e(-) + Cr2O7(2-) = 2 Cr + 7 H2O
And presumably your oxidation is
2 H2O = O2 + 4 H(+) + 4 e(-)
To balance this you need three of the latter equation, and when you add together you get:
2 H(+) + Cr2O7(2-) = 2 Cr + 3 O2 + H2O
Yes that seems to balance.
Now you can forget the electricity part. Just take 100 grams, convert it to moles of chromium, multiply by 3/2 to get the moles of oxygen gas produced.