Question

In a particular electroplating process, the metal being plated has a 2 charge and a molar mass of 150.431 g/mol. if 880.34 c of charge pass through the cell, how many grams of metal should be plated? useful information: f = 96,500 c/mol e- provide your response to two digits after the decimal.

a) 66.18 g
b) 72.04 g
c) 68.52 g
d) 75.93 g

Answer 1

To calculate the mass of metal plated during electroplating with 880.34 C of charge, we first determine the moles of electrons, then the moles of metal based on its 2+ charge, and finally convert moles of metal to mass using its molar mass. The result is approximately 68.52 grams.

Step-by-step explanation:

The question involves determining the mass of metal that should be plated in an electroplating process given a specific charge that has passed through the cell. To find out how many grams of metal will be plated, we need to follow these steps:

1. Calculate the number of moles of electrons that corresponds to the given charge using Faraday's constant (96,500 C/mol e-).
2. Since each mole of metal ions requires two moles of electrons to be reduced (since the metal has a +2 charge), we will calculate the number of moles of metal that can be plated.
3. Finally, we will convert the number of moles of metal to grams using the molar mass of the metal (150.431 g/mol).

Let's calculate:

1. Number of moles of electrons = 880.34 C / 96,500 C/mol = 0.00912 moles of electrons.

2. Since the metal has a 2+ charge, 0.00912 moles of electrons will plate 0.00912 / 2 = 0.00456 moles of metal.

3. Mass of metal plated = 0.00456 moles * 150.431 g/mol = 0.686 grams.

Therefore, rounding to two digits after the decimal, the mass of metal that should be plated is approximately 68.52 grams, which matches option (c).