A current of 3.58 A is passed through a Ni(NO,), solution. How long, in hours, would this current have to be applied to plate out 6.40 g of nickel?
A current of 3.58 A is passed through a Ni(NO,), solution. How long, in hours, would this current have to be applied to plate out 6.40 g of nickel?
General Chemistry - Standalone book (MindTap Course List)
11th Edition
ISBN:9781305580343
Author:Steven D. Gammon, Ebbing, Darrell Ebbing, Steven D., Darrell; Gammon, Darrell Ebbing; Steven D. Gammon, Darrell D.; Gammon, Ebbing; Steven D. Gammon; Darrell
Publisher:Steven D. Gammon, Ebbing, Darrell Ebbing, Steven D., Darrell; Gammon, Darrell Ebbing; Steven D. Gammon, Darrell D.; Gammon, Ebbing; Steven D. Gammon; Darrell
Chapter19: Electrochemistry
Section: Chapter Questions
Problem 19.124QP: An aqueous solution of an unknown salt of vanadium is electrolyzed by a current of 2.50 amps for...
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![**Educational Text: Electroplating Nickel from Ni(NO₃)₂ Solution**
A current of 3.58 A is passed through a Ni(NO₃)₂ solution. How long, in hours, would this current have to be applied to plate out 6.40 g of nickel?
**Time required: [_____________] h**
**Explanation:**
In this setup, you are tasked with calculating how long it will take to electroplate a specified amount of nickel using a given electrical current. The calculation involves principles of electrochemistry, specifically relating to Faraday's laws of electrolysis.
1. **Molar Mass and Charge:**
- First, calculate the moles of nickel needed using the atomic mass of nickel (approximately 58.69 g/mol).
- Nickel has an ion charge of 2+ in the solution, meaning 2 moles of electrons are required to deposit 1 mole of nickel.
2. **Faraday’s Constant:**
- The constant (approximately 96485 C/mol e⁻) is used to relate charge and moles of electrons.
3. **Calculate Total Charge Needed:**
- Use the moles of nickel and multiply by the charge per mole to find the total charge required for the electroplating.
4. **Determine Time:**
- Finally, use the formula \( \text{Time} = \frac{\text{Total Charge}}{\text{Current}} \) to find the time in seconds, and convert to hours by dividing by 3600.
This practical scenario highlights the application of theoretical chemistry in industrial and laboratory settings, demonstrating the use of electrolytic processes for metal deposition.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fd891009c-d394-441a-930b-8d5c1fe8da7b%2Fb92cff38-f0d9-4ef8-8c0d-39ec5e3c3434%2Fmccgnz.jpeg&w=3840&q=75)
Transcribed Image Text:**Educational Text: Electroplating Nickel from Ni(NO₃)₂ Solution**
A current of 3.58 A is passed through a Ni(NO₃)₂ solution. How long, in hours, would this current have to be applied to plate out 6.40 g of nickel?
**Time required: [_____________] h**
**Explanation:**
In this setup, you are tasked with calculating how long it will take to electroplate a specified amount of nickel using a given electrical current. The calculation involves principles of electrochemistry, specifically relating to Faraday's laws of electrolysis.
1. **Molar Mass and Charge:**
- First, calculate the moles of nickel needed using the atomic mass of nickel (approximately 58.69 g/mol).
- Nickel has an ion charge of 2+ in the solution, meaning 2 moles of electrons are required to deposit 1 mole of nickel.
2. **Faraday’s Constant:**
- The constant (approximately 96485 C/mol e⁻) is used to relate charge and moles of electrons.
3. **Calculate Total Charge Needed:**
- Use the moles of nickel and multiply by the charge per mole to find the total charge required for the electroplating.
4. **Determine Time:**
- Finally, use the formula \( \text{Time} = \frac{\text{Total Charge}}{\text{Current}} \) to find the time in seconds, and convert to hours by dividing by 3600.
This practical scenario highlights the application of theoretical chemistry in industrial and laboratory settings, demonstrating the use of electrolytic processes for metal deposition.
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