Question

How many seconds are required to deposit 0.161 grams of iron metal from a solution that contains Fe₂ ions, if a current of 0.651 A is applied?

Answer 1

To calculate the time required to deposit 0.161 grams of iron metal from a solution containing Fe₂ ions with a current of 0.651 A, we can use Faraday's law of electrolysis and calculate the amount of charge required to deposit the given amount of iron metal.

Step-by-step explanation:

To calculate the time required to deposit 0.161 grams of iron metal from a solution containing Fe₂ ions with a current of 0.651 A, we can use the concept of electroplating and Faraday's law of electrolysis.

Using Faraday's law, we can calculate the amount of charge required to deposit the given amount of iron:

1. First, convert the mass of iron to moles by dividing it by the molar mass of iron (55.85 g/mol).
2. Next, use the balanced equation of the redox reaction to determine the moles of electrons involved in the deposition of one mole of iron.
3. Finally, use the equation Q = It, where Q is the charge in coulombs, I is the current in amperes, and t is the time in seconds, to find the time required to deposit the given amount of iron.

Substitute the values into the equation Q = It to find the time required in seconds.