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How many grams of copper metal can be deposited from Cu²+ (aq) when a current of 2.50 A is run for 2.55 h? (F= 96,500 C/ mol)

How many grams of copper metal can be deposited from Cu²+ (aq) when a current of 2.50 A is run for 2.55 h? (F= 96,500 C/ mol)

Chemistry
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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**Question 19 of 20** **Electrochemistry Calculation** How many grams of copper metal can be deposited from Cu²⁺(aq) when a current of 2.50 A is run for 2.55 hours? (F = 96,500 C/mol) --- **Explanation and Calculation**: To solve this problem, we will use Faraday's laws of electrolysis, which relate the amount of substance deposited at an electrode to the total charge passed through the electrolyte. 1. **Calculate Total Charge (Q):** - Current (I) = 2.50 A - Time (t) = 2.55 hours = 2.55 x 3600 seconds (since 1 hour = 3600 seconds) - \( Q = I \times t \) 2. **Calculate Moles of Electrons (n):** - Given: Faraday's constant (F) = 96,500 C/mol - Number of moles of electrons, \( n = \frac{Q}{F} \) 3. **Determine Moles of Copper Deposited:** - The reaction involves the reduction of Cu²⁺ ions to Cu. - Each Cu requires 2 moles of electrons to be deposited as metal. 4. **Calculate Mass of Copper Deposited:** - Atomic mass of copper (Cu) ≈ 63.55 g/mol - Use the formula: mass = moles of copper \(\times\) atomic mass 5. **Use the keypad to input the value:** - Enter the final calculated mass in grams (g) using the numeric keypad displayed. By following these steps, you can find the mass of copper metal that can be deposited under the given conditions.
Transcribed Image Text:**Question 19 of 20** **Electrochemistry Calculation** How many grams of copper metal can be deposited from Cu²⁺(aq) when a current of 2.50 A is run for 2.55 hours? (F = 96,500 C/mol) --- **Explanation and Calculation**: To solve this problem, we will use Faraday's laws of electrolysis, which relate the amount of substance deposited at an electrode to the total charge passed through the electrolyte. 1. **Calculate Total Charge (Q):** - Current (I) = 2.50 A - Time (t) = 2.55 hours = 2.55 x 3600 seconds (since 1 hour = 3600 seconds) - \( Q = I \times t \) 2. **Calculate Moles of Electrons (n):** - Given: Faraday's constant (F) = 96,500 C/mol - Number of moles of electrons, \( n = \frac{Q}{F} \) 3. **Determine Moles of Copper Deposited:** - The reaction involves the reduction of Cu²⁺ ions to Cu. - Each Cu requires 2 moles of electrons to be deposited as metal. 4. **Calculate Mass of Copper Deposited:** - Atomic mass of copper (Cu) ≈ 63.55 g/mol - Use the formula: mass = moles of copper \(\times\) atomic mass 5. **Use the keypad to input the value:** - Enter the final calculated mass in grams (g) using the numeric keypad displayed. By following these steps, you can find the mass of copper metal that can be deposited under the given conditions.
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