A classroom experiment was set up to simulate the industrial extraction of zinc metal from an aqueous solution of zinc ions by electrolysis - VCE - SSCE Chemistry - Question 7 - 2009 - Paper 1

First, we need to calculate the number of moles of zinc produced:

$n = \frac{0.900 \text{ g}}{65.4 \text{ g/mol}} = 0.01376 \text{ mol}$

Since it takes 2 moles of electrons to produce 1 mole of zinc, the moles of electrons transferred, (n(e^-)), can be calculated as:

$n(e^-) = 2 \times n = 2 \times 0.01376 \text{ mol} = 0.02752 \text{ mol}$

Next, we calculate the charge using the formula:

$Q = n(e^-) \times F$

Where (F = 96500 \text{ C/mol}) is Faraday's constant:

$Q = 0.02752 \text{ mol} \times 96500 \text{ C/mol} = 2656.0 \text{ C}$

Now, with electrolysis time of 30 minutes (or 1800 seconds), the current can be calculated using:

$I = \frac{Q}{t} = \frac{2656.0 \text{ C}}{1800 \text{ s}} \approx 1.47 \text{ A}$

Expressing with significant figures:

$I \approx 1.48 \text{ A}$