Question 13.4: In a process called flash electroplating, a current of 2.50 ......
In a process called flash electroplating, a current of 2.50 \times 10^3\ A passes through an electrolytic cell for 5.00 minutes. How many moles of electrons are driven through the cell?
Strategy Because we know both the current and the time for which it was applied, we can use the relationship between charge and current in Equation 13.8 to obtain the charge. Then Faraday’s constant will let us convert that charge into moles of electrons. We must be careful with units, however, because the ampere is charge per second and time is given here in minutes.
\begin{gathered}\text { Charge }=\text { current } \times \text { time } \\\qquad Q=I \times t\end{gathered} (13.8)
Q=2500\ A \times 300\ s =7.50 \times 10^5\ C
Now use Faraday’s constant:
7.50 \times 10^5\ C \times\left(\frac{1\text{ mol e }^{-}}{96,485\ C }\right)=7.77\text{ mol e }^{-}
Discussion This two-step manipulation is really a stoichiometry problem, as it allows us to find the number of moles of something (electrons, in this case) in a chemical reaction. The central importance of moles in stoichiometry problems is well established at this point in our study of chemistry. Just as for other types of stoichiometry problems that we’ve considered, we sometimes need to use these relationships in reverse. For example, if you need to know how long to run a given current to obtain a desired amount of plated material, these relationships would still be used, only the known quantity would be moles of electrons.
Check Your Understanding You need 0.56 mol of electrons to deposit a thin covering of silver on a part that you are using for a prototype machine. How long would you need to run a current of 5.0 A to obtain this number of moles?