During electrorefining of Cu from CUSO4, how much time (in hours) is needed to produce 250 g Cu on the cathode if the current is kept at 11 A?
During electrorefining of Cu from CUSO4, how much time (in hours) is needed to produce 250 g Cu on the cathode if the current is kept at 11 A?
Chemistry
10th Edition
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Chapter1: Chemical Foundations
Section: Chapter Questions
Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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![**Transcription for Educational Website:**
### Electrorefining of Copper
#### Question:
During electrorefining of Cu from CuSO₄, how much time (in hours) is needed to produce 250 g of Cu on the cathode if the current is kept at 11 A?
#### Explanation:
This problem involves using Faraday’s laws of electrolysis to calculate the time required for electrorefining copper from copper sulfate.
**Key Concepts:**
1. **Faraday’s First Law of Electrolysis**: The mass of a substance deposited at an electrode is directly proportional to the charge passed through the electrolyte.
2. **Faraday’s Second Law of Electrolysis**: The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their equivalent weights.
**Calculations:**
- The molar mass of copper (Cu) is approximately 63.5 g/mol.
- Assuming a valency of +2 for Cu in copper sulfate, each mole of copper requires 2 faradays (F) of charge to be deposited.
- 1 Faraday (F) = 96,485 Coulombs (C).
**Steps to calculate time:**
1. Calculate the moles of copper:
\[
\text{Moles of Cu} = \frac{250 \, \text{g}}{63.5 \, \text{g/mol}} \approx 3.937 \, \text{mol}
\]
2. Calculate the total charge needed using Faraday’s laws:
\[
\text{Total Charge} = \text{Moles of Cu} \times 2 \, F \, \times 96,485 \, C/mol
\]
3. Calculate the time (t) using the formula:
\[
t = \frac{\text{Total Charge}}{\text{Current (I)}}
\]
Convert the result to hours for the final answer.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fd2691140-3c8d-4e92-bf51-05dba936fe25%2Fd0921132-343f-45fa-a4be-948bcbd1c5f3%2Fz4znxzr.jpeg&w=3840&q=75)
Transcribed Image Text:**Transcription for Educational Website:**
### Electrorefining of Copper
#### Question:
During electrorefining of Cu from CuSO₄, how much time (in hours) is needed to produce 250 g of Cu on the cathode if the current is kept at 11 A?
#### Explanation:
This problem involves using Faraday’s laws of electrolysis to calculate the time required for electrorefining copper from copper sulfate.
**Key Concepts:**
1. **Faraday’s First Law of Electrolysis**: The mass of a substance deposited at an electrode is directly proportional to the charge passed through the electrolyte.
2. **Faraday’s Second Law of Electrolysis**: The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their equivalent weights.
**Calculations:**
- The molar mass of copper (Cu) is approximately 63.5 g/mol.
- Assuming a valency of +2 for Cu in copper sulfate, each mole of copper requires 2 faradays (F) of charge to be deposited.
- 1 Faraday (F) = 96,485 Coulombs (C).
**Steps to calculate time:**
1. Calculate the moles of copper:
\[
\text{Moles of Cu} = \frac{250 \, \text{g}}{63.5 \, \text{g/mol}} \approx 3.937 \, \text{mol}
\]
2. Calculate the total charge needed using Faraday’s laws:
\[
\text{Total Charge} = \text{Moles of Cu} \times 2 \, F \, \times 96,485 \, C/mol
\]
3. Calculate the time (t) using the formula:
\[
t = \frac{\text{Total Charge}}{\text{Current (I)}}
\]
Convert the result to hours for the final answer.
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