Is there an electrochemistry formula used to calculate amount of time needed to reduce 1 mole?

[September24]

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EK 1001 #918 from chemistry

Suppoe ane elctrolytic cell utilizes the following cathode to silver plate a piece of jewelry:

(Ag+) + E- --> Ag E0=+0.8V


If 0.108 grams of silver is to be plated using 2.0 A current, for how long should the cell be run? (One faraday=96,485 C/mol*e-)
Answer is: 48 seconds

Explanation: use the following calculation:

(0.108 g Ag)(1 mol Ag/108g*Al)(1mol*e-/mol Ag)(96,485 C/mol e-)(1sec/2C)=48 seconds.

What is EK doing here?

 

[sps27]

.108 grams of silver is .108/108 moles of silver.
96,485 Colums is the charge of 1 mole of electrons.
We have 2 amperes supply i.e., 2 Columbs / sec

Now 1 mole of Silver needs 1 mole of electrons as per eqn....
So .001 mole of silver needs, .001 * 96,485 columbs
And we have a supply of 2 C for ever seconds.
So time required is,

(96,485 * .001 ) / 2 = which is roughly 96 /2 which is roughly 48 seconds.