How much time is required to electroplate 6.530 µg of chromium from a potassium dichromate solution, by using a current of 118 mA?
I thought the ratio of e- to moles chromium would be 6/2 but I guess that is wrong? I am using t= nF/A
I already made the proper conversions to grams and A's....just confused on the potassium dichromate to chromium. I have it ias KCr2O7 + 6e-==> 2CrO4^2- which is where i got the ratio to be 6/2
To determine the time required to electroplate 6.530 µg of chromium from a potassium dichromate solution using a current of 118 mA, you need to consider the stoichiometry of the electroplating reaction.
The balanced equation for the electroplating reaction is:
2CrO4^2- + 10e- + 16H+ → 2Cr3+ + 8H2O
From this equation, you can see that it takes 10 electrons to produce 2 moles of chromium (Cr3+). Therefore, the ratio of electrons to moles of chromium is 10/2 = 5/1.
To calculate the time (t) required for electroplating, you can use the formula:
t = (n * F) / I
Where:
t = time (s)
n = moles of chromium
F = Faraday's constant (96485 C/mol)
I = current (A)
First, let's find the number of moles of chromium:
Mass of chromium = 6.530 µg = 6.530 × 10^-6 g
Molar mass of chromium (Cr) = 52 g/mol
Moles of chromium = (mass of chromium) / (molar mass of chromium)
= (6.530 × 10^-6 g) / (52 g/mol)
Next, calculate the time required for electroplating:
t = (n * F) / I
= [(6.530 × 10^-6 g) / (52 g/mol)] * (96485 C/mol) / (118 × 10^-3 A)
After performing the calculations, you will obtain the time required to electroplate 6.530 µg of chromium from the potassium dichromate solution with the given current.