Electrolysis

Curriculum Framework

Curriculum Framework

F= 96,500 C/mol of e C = amperes per second

Teaching Outline

Introduction—Venn DiagramElectrolysis Simulations—ConceptualFaraday’s Law—CalculationsWriting reaction for electrolysis Free Response Practice

Galvanic Electrolytic

Electrolysis—Simulation

Simulation Questions1. 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.

Simulation ExperimentDesign 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

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)

Electrolysis “Map” Helpful information

A = C per second

F=Faraday’s constant= 96500 C/mole e

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?

Example 2

If a constant current of 8.00 amperes is passed through a cell containing Zn2+ for 2.00 hours, how many grams of zinc will plate out onto the cathode?

Example 3

Calculate the amount of time required to produce 1000 grams of magnesium metal by electrolysis of molten MgCl2 using a current of 50 A.

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?

Example 5

Two cells, one containing aqueous AgNO3 and the other containing CuSO4 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?

Writing Electrolysis Reactions

Two types of situations:• Molten solutions—only two ions present

• Aqueous solutions—water possibly oxidized or reduced

Molten solutions

Cathode: Cation will be reducedAnode: Anion will be oxided

Aqueous solutionsCathode: Cation will be reduced OR

water will be reducedReduction of water: 2 H2O (l) + 2 e- --> H2 (g) + 2OH

(aq) Eo

red = -0.83 V

(The one with the higher reduction potential)

Anode: Anion will be oxidized OR water will be oxidizedOxidation of water: 2 H2O (l) --> 4 H+(aq) + O2(g) + 4 e- Eo

red = 1.23 V

(The one with the lower, more negative, reduction potential)

Copper (II) chlorideWriting REDOX reactions for Electrolytic CellsFor the electrolysis of aqueous CuCl2 using platinum (inert) electrodes. Find:The half-reaction at the Cathode: _______________________________________________ Eo = ________The half-reaction at the Anode: _______________________________________________ Eo = ________The overall redox reaction: _______________________________________________ Eo = ________Product(s) at the Cathode:________________ Product(s) at the Anode _________________The minimum voltage required: ________________ V

Sodium sulfate

For the electrolysis of Na2SO4(aq) using carbon (inert) electrodes. Find:

The half-reaction at the Cathode: ___________________________________________ Eo = ________The half-reaction at the Anode: ____________________________________________Eo = ________The overall redox reaction: ____________________________________________Eo = _______Product(s) at the Cathode:_________________ Product(s) at the Anode _________________The minimum voltage required: ________________ V

Copper (II) sulfate

For the electrolysis of CuSO4(aq) using inert electrodes. Find:The half-reaction at the Cathode: _____________________________________________ Eo = ________The half-reaction at the Anode: _____________________________________________ Eo = ________The overall redox reaction: _____________________________________________ Eo = ________Product(s) at the Cathode:_________________ Product(s) at the Anode _________________The minimum voltage required: ________________V

Potassium Iodide

For the electrolysis of KI in water. The half-reaction at the Cathode: ____________________________________________ Eo = ________The half-reaction at the Anode: ____________________________________________ Eo = ________The overall redox reaction: _____________________________________________ Eo =

________Product(s) at the Cathode:_________________ Product(s) at the Anode _________________The minimum voltage required: ________________ V