Final answer:
Using Faraday's laws, we calculate that a current of 1.115 A is required to plate out 1.22 g of nickel from a Ni2+ solution in one hour.
Step-by-step explanation:
To determine the current required to plate out 1.22 g of nickel from a solution of Ni2+, we can use Faraday's laws of electrolysis which states that the amount of substance deposited at an electrode is proportional to the quantity of electric charge passed through the electrolyte.
Nickel has a molar mass of 58.69 g/mol and is divalent (Ni2+), meaning two moles of electrons are needed to deposit one mole of nickel. The constant we need is Faraday's constant, which is approximately 96500 C/mol of electrons.
If 1.22 g of nickel is to be plated in 1.0 hour (3600 seconds), we first convert the mass of nickel to moles:
(1.22 g) / (58.69 g/mol) = 0.0208 mol Ni
Since 2 mol of electrons are required for every mol of Ni, we need:
0.0208 mol Ni x (2 mol e-/mol Ni) = 0.0416 mol e-
Multiplying this number of moles by Faraday's constant gives us the total charge needed:
0.0416 mol e- x 96500 C/mol e- = 4014 C
Finally, we find the current by dividing the charge by the time in seconds:
Current (I) = Total charge (Q) / Time (t)
I = 4014 C / 3600 s = 1.115 A
Therefore, a current of 1.115 A is required to plate out 1.22 g of nickel in one hour.