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23) It took 116 hours to produce 603 g of metal X by performing electrolysis on molten XCI3 with a current of 2.00 A. Calculate the molar mass of X. a) 55.8 g/mol b) 72.6 g/mol c) 27.0 g/mol d) 204 g.mol e) 209 g/mol

23) It took 116 hours to produce 603 g of metal X by performing electrolysis on molten XCI3 with a current of 2.00 A. Calculate the molar mass of X. a) 55.8 g/mol b) 72.6 g/mol c) 27.0 g/mol d) 204 g.mol e) 209 g/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
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Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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# Standard Reduction Potentials The following table lists standard reduction potentials (E°) for various redox couples in volts (V). Standard reduction potentials are used to predict the direction of electron flow in electrochemical cells. ## Reduction Potential Table | Couple | E° (Volts) | Couple | E° (Volts) | |-----------------|------------|----------------|------------| | F₂ | HF (H⁺) | +3.03 | SO₄²⁻ | H₂SO₃ (H⁺) | +0.20 | | F₂ | F⁻ | +2.87 | Sn⁴⁺ | Sn²⁺ | +0.15 | | S₂O₈²⁻ | SO₄²⁻ | +2.05 | S | H₂S (H⁺) | +0.141 | | BiO₃⁻ | Bi³⁺ | +2.0 | Hg₂Br₂ | Hg (Br⁻) | +0.140 | | H₂O₂ | H₂O (H⁺) | +1.78 | AgBr | Ag (Br⁻) | +0.0713 | | PbO₂ | PbSO₄ (H⁺, SO₄²⁻) | +1.685 | H⁺ | H₂ | +0.0000 | | Ce⁴⁺ | Ce³⁺ | +1.61 | Pb²⁺ | Pb | -0.126 | | MnO₄⁻ | Mn²⁺ (H⁺) | +1.491 | Sn²⁺ | Sn | -0.136 | | ClO₃⁻ | Cl⁻ | +1.47 | AgI | Ag (I⁻) | -0.152 | | PbO₂ | Pb²⁺ (H⁺) | +1.46 | Ni²⁺ | Ni | -0.250 | | Au³⁺ | Au | +1.42 | Co²
Transcribed Image Text:# Standard Reduction Potentials The following table lists standard reduction potentials (E°) for various redox couples in volts (V). Standard reduction potentials are used to predict the direction of electron flow in electrochemical cells. ## Reduction Potential Table | Couple | E° (Volts) | Couple | E° (Volts) | |-----------------|------------|----------------|------------| | F₂ | HF (H⁺) | +3.03 | SO₄²⁻ | H₂SO₃ (H⁺) | +0.20 | | F₂ | F⁻ | +2.87 | Sn⁴⁺ | Sn²⁺ | +0.15 | | S₂O₈²⁻ | SO₄²⁻ | +2.05 | S | H₂S (H⁺) | +0.141 | | BiO₃⁻ | Bi³⁺ | +2.0 | Hg₂Br₂ | Hg (Br⁻) | +0.140 | | H₂O₂ | H₂O (H⁺) | +1.78 | AgBr | Ag (Br⁻) | +0.0713 | | PbO₂ | PbSO₄ (H⁺, SO₄²⁻) | +1.685 | H⁺ | H₂ | +0.0000 | | Ce⁴⁺ | Ce³⁺ | +1.61 | Pb²⁺ | Pb | -0.126 | | MnO₄⁻ | Mn²⁺ (H⁺) | +1.491 | Sn²⁺ | Sn | -0.136 | | ClO₃⁻ | Cl⁻ | +1.47 | AgI | Ag (I⁻) | -0.152 | | PbO₂ | Pb²⁺ (H⁺) | +1.46 | Ni²⁺ | Ni | -0.250 | | Au³⁺ | Au | +1.42 | Co²
**Electrolysis Calculation Example** **Problem Statement:** It took 116 hours to produce 603 g of metal X by performing electrolysis on molten XCl₃ with a current of 2.00 A. Calculate the molar mass of X. **Multiple Choice Answers:** a) 55.8 g/mol b) 72.6 g/mol c) 27.0 g/mol d) 204 g/mol e) 209 g/mol **Solution Explanation:** To solve this problem, follow these steps: 1. **Calculate Total Charge (Q):** Using the formula: \( Q = I \times t \) Where: - \( I = 2.00 \) A (Current) - \( t = 116 \) hours \( = 116 \times 3600 \) seconds (time converted to seconds) \[ Q = 2.00 \, \text{A} \times 116 \times 3600 \, \text{s} \] \[ Q = 835200 \, \text{Coulombs (C)} \] 2. **Calculate Moles of Electrons:** Using Faraday's constant (\( F = 96485 \) C/mol) \[ \text{Moles of electrons} = \frac{Q}{F} \] \[ \text{Moles of electrons} = \frac{835200 \, \text{C}}{96485 \, \text{C/mol}} \] \[ \text{Moles of electrons} \approx 8.65 \, \text{mol} \] 3. **Relate Moles of Electrons to Moles of Metal X:** From the formula XCl₃, each mole of metal X requires 3 moles of electrons (since X has a +3 charge). \[ \text{Moles of X} = \frac{\text{Moles of electrons}}{3} \] \[ \text{Moles of X} = \frac{8.65 \, \text{mol}}{3} \] \[ \text{Moles of X} \approx 2.88 \, \text{mol} \] 4. **Calculate Molar Mass of Metal X:** \[ \text{
Transcribed Image Text:**Electrolysis Calculation Example** **Problem Statement:** It took 116 hours to produce 603 g of metal X by performing electrolysis on molten XCl₃ with a current of 2.00 A. Calculate the molar mass of X. **Multiple Choice Answers:** a) 55.8 g/mol b) 72.6 g/mol c) 27.0 g/mol d) 204 g/mol e) 209 g/mol **Solution Explanation:** To solve this problem, follow these steps: 1. **Calculate Total Charge (Q):** Using the formula: \( Q = I \times t \) Where: - \( I = 2.00 \) A (Current) - \( t = 116 \) hours \( = 116 \times 3600 \) seconds (time converted to seconds) \[ Q = 2.00 \, \text{A} \times 116 \times 3600 \, \text{s} \] \[ Q = 835200 \, \text{Coulombs (C)} \] 2. **Calculate Moles of Electrons:** Using Faraday's constant (\( F = 96485 \) C/mol) \[ \text{Moles of electrons} = \frac{Q}{F} \] \[ \text{Moles of electrons} = \frac{835200 \, \text{C}}{96485 \, \text{C/mol}} \] \[ \text{Moles of electrons} \approx 8.65 \, \text{mol} \] 3. **Relate Moles of Electrons to Moles of Metal X:** From the formula XCl₃, each mole of metal X requires 3 moles of electrons (since X has a +3 charge). \[ \text{Moles of X} = \frac{\text{Moles of electrons}}{3} \] \[ \text{Moles of X} = \frac{8.65 \, \text{mol}}{3} \] \[ \text{Moles of X} \approx 2.88 \, \text{mol} \] 4. **Calculate Molar Mass of Metal X:** \[ \text{
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