Question

Chromium metal is electroplated using an acidic solution containing
CrOg according to the following equation:
CYO3(aq) + 6H* +6e - Cr(s) + 3H20
Calculate how many grams of chromium will be electroplated by
24,000 coulombs. How long will it take to electroplate 1.5 g chromium
using 12.5 A current ?
[Atomic mass of Cr = 52 g mol-1,1 F = 96500 C mol-1]​

Answer 1

1. charge Q = $(n\times n_f)F$

where n is number of moles of the species under consideration, $n_f$ is the n-factor [ simply n-factor means change in oxidation number ]

and F is charge required to deposit 1 mole of electron i.e., 96500C

here, n - factor = 6

Charge, Q = 24000C

so, 24000C = n × 6 × 96500C

or, n = 24000/(96500 × 6) = 0.04145mol

now, mole = weight in gm/atomic weight

so, 0.04145 = weight in gm/52g/mol

or, weight in gm of chromium = 0.04145 × 52 = 2.1554g

2. number of mole of Cr = 1.5/52

number of mole of electron passed = 6 × 1.5/52 [ because change in oxidation of Cr is +6 ]

we know, 1 mole of electron is equivalent to 96500C charges.

so, charge required = 6 × 1.5/52 × 96500

= 9 × 96500/52

= 16702C

time required = charge/current

= 16702/12.5

= 1336.16 sec.