A voltaic cell using Cu²+/Cu and Al³+/Al half-cells is set up at standard conditions, and each compartment has a volume of 225 mL. What is the [Al³+] after the cell has delivered 0.120 A for 29.0 hours at 25°C? (E° for Cu²+/ Cu = 0.340 V and Eº for Al³+/AI = -1.660 V.)

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### Electrochemistry Problem on Voltaic Cells

**Question Overview:**

A voltaic cell using Cu²⁺/Cu and Al³⁺/Al half-cells is set up at standard conditions. Each compartment has a volume of 225 mL. Calculate the concentration of Al³⁺ ions after the cell has delivered 0.120 A for 29.0 hours at 25°C. Use the following standard electrode potentials:
- \( E^° \) for Cu²⁺/Cu = 0.340 V
- \( E^° \) for Al³⁺/Al = -1.660 V

---

**Solution Steps:**

1. **Calculate Total Charge Delivered:**
   Convert the current and time into Coulombs:
   \[
   \text{Charge (Q)} = \text{Current (I)} \times \text{Time (t)}
   \]
   \[
   Q = 0.120 \, \text{A} \times (29.0 \, \text{hours} \times 3600 \, \text{seconds/hour})
   \]

2. **Calculate Moles of Electrons:**
   Using the formula:
   \[
   \text{Moles of electrons} = \frac{Q}{F}
   \]
   where \( F \) is Faraday's constant (96485 C/mol).

3. **Determine Moles of Al³⁺ Reacted:**
   Balance the reaction and calculate moles of Al³⁺ formed or reacted.

4. **Calculate Final Concentration of Al³⁺:**
   Using:
   \[
   \text{Concentration} = \frac{\text{Moles}}{\text{Volume (L)}}
   \]

**Inputs and Controls:**

- Below the question is a keypad that can be used for inputting answers.
- Additional resources and tips can be accessed by tapping on the designated area.

**Note:**

This problem exemplifies a real-world application of electrochemical principles, combining theoretical knowledge with practical calculations commonly encountered in chemistry and physics studies.
Transcribed Image Text:### Electrochemistry Problem on Voltaic Cells **Question Overview:** A voltaic cell using Cu²⁺/Cu and Al³⁺/Al half-cells is set up at standard conditions. Each compartment has a volume of 225 mL. Calculate the concentration of Al³⁺ ions after the cell has delivered 0.120 A for 29.0 hours at 25°C. Use the following standard electrode potentials: - \( E^° \) for Cu²⁺/Cu = 0.340 V - \( E^° \) for Al³⁺/Al = -1.660 V --- **Solution Steps:** 1. **Calculate Total Charge Delivered:** Convert the current and time into Coulombs: \[ \text{Charge (Q)} = \text{Current (I)} \times \text{Time (t)} \] \[ Q = 0.120 \, \text{A} \times (29.0 \, \text{hours} \times 3600 \, \text{seconds/hour}) \] 2. **Calculate Moles of Electrons:** Using the formula: \[ \text{Moles of electrons} = \frac{Q}{F} \] where \( F \) is Faraday's constant (96485 C/mol). 3. **Determine Moles of Al³⁺ Reacted:** Balance the reaction and calculate moles of Al³⁺ formed or reacted. 4. **Calculate Final Concentration of Al³⁺:** Using: \[ \text{Concentration} = \frac{\text{Moles}}{\text{Volume (L)}} \] **Inputs and Controls:** - Below the question is a keypad that can be used for inputting answers. - Additional resources and tips can be accessed by tapping on the designated area. **Note:** This problem exemplifies a real-world application of electrochemical principles, combining theoretical knowledge with practical calculations commonly encountered in chemistry and physics studies.
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