1. Electrolysis

2. Curriculum Framework

3. F= 96,500 C/mol of eC = amperes per second

4. Teaching Outline Introduction—Venn Diagram Electrolysis Simulations—Conceptual Faraday’s Law—Calculations Writing reaction for electrolysis Free Response Practice

5. GalvanicElectrolytic

6. Electrolysis—Simulation

7. Simulation Questions 1. Describe the change in amount of metal on each electrode. How are these changes related? 2. What flows from the + electrode in the external circuit via the wire? 3. What causes the direction of the flow? 4. What flows from the + electrode in the solution? 5. Describe the action that causes the metal ions to plate on to the electrode. Write a chemical equation that summarizes your explanation.

8. Simulation Experiment Design an experiment to answer each of the following experimental questions: How does the amount of time affect the change in mass on the two electrodes? How does the number of amps affect the change in mass on the two electrodes? How does the type of metal affect the change in mass on the two electrodes? For each experiment, include Independent variable Dependent variable Data table Summary of results

9. Faraday’s Law The quantity of metal produced and/or consumed in an electrolytic cell is dependent upon Type of metal (molar mass) Oxidation state of the metal Time Amperage (current)

10. Electrolysis “Map” Helpful information A = C per second F=Faraday’s constant= 96500 C/mole e

11. Example 1 The current in a given wire is 1.80 amp. How many coulombs will pass a given point on the wire in 1.36 minutes?

12. Example 2 If a constant current of 8.00 amperes is passed through a cell containing Zn 2+ for 2.00 hours, how many grams of zinc will plate out onto the cathode?

13. Example 3 Calculate the amount of time required to produce 1000 grams of magnesium metal by electrolysis of molten MgCl 2 using a current of 50 A.

14. Example 4 What amperage is required to plate out 50.00 grams of Cr from a Cr +3 solution in a period of 8.00 hours?

15. Example 5 Two cells, one containing aqueous AgNO 3 and the other containing CuSO 4 are set up in series. In a given electrolysis that results in depositing 1.25 g of silver in the first cell, how much copper should deposit simultaneously in the second cell?

16. Writing Electrolysis Reactions Two types of situations: Molten solutions—only two ions present Aqueous solutions—water possibly oxidized or reduced

17. Molten solutions Cathode: Cation will be reduced Anode: Anion will be oxided

18. Aqueous solutions C athode: Cation will be reduced OR water will be reduced Reduction of water: 2 H 2 O (l) + 2 e - --> H 2 (g) + 2OH (aq) E o red = -0.83 V (The one with the higher reduction potential) Anode: Anion will be oxidized OR water will be oxidized Oxidation of water: 2 H 2 O (l) --> 4 H + (aq) + O 2 (g) + 4 e - E o red = 1.23 V (The one with the lower, more negative, reduction potential)

19. Copper (II) chloride Writing REDOX reactions for Electrolytic Cells For the electrolysis of aqueous CuCl2 using platinum (inert) electrodes. Find: The half-reaction at the Cathode: _______________________________________________ E o = ________ The half-reaction at the Anode: _______________________________________________ E o = ________ The overall redox reaction: _______________________________________________ E o = ________ Product(s) at the Cathode:________________ Product(s) at the Anode _________________ The minimum voltage required: ________________ V

20. Sodium sulfate For the electrolysis of Na 2 SO 4 (aq) using carbon (inert) electrodes. Find: The half-reaction at the Cathode: ___________________________________________ E o = ________ The half-reaction at the Anode: ____________________________________________E o = ________ The overall redox reaction: ____________________________________________E o = _______ Product(s) at the Cathode:_________________ Product(s) at the Anode _________________ The minimum voltage required: ________________ V

21. Copper (II) sulfate For the electrolysis of CuSO4(aq) using inert electrodes. Find: The half-reaction at the Cathode: _____________________________________________ E o = ________ The half-reaction at the Anode: _____________________________________________ E o = ________ The overall redox reaction: _____________________________________________ E o = ________ Product(s) at the Cathode:_________________ Product(s) at the Anode _________________ The minimum voltage required: ________________V

22. Potassium Iodide For the electrolysis of KI in water. The half-reaction at the Cathode: ____________________________________________ E o = ________ The half-reaction at the Anode: ____________________________________________ E o = ________ The overall redox reaction: _____________________________________________ E o = ________ Product(s) at the Cathode:_________________ Product(s) at the Anode _________________ The minimum voltage required: ________________ V