18.3 Electrolysis. Dr. Fred Omega Garces. Chemistry 201. Driving a non-spontaneous Oxidation-Reduction Reaction. Miramar College.

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1 18.3 Electrolysis Driving a non-spontaneous Oxidation-Reduction Reaction Dr. Fred Omega Garces Chemistry 201 Miramar College 1 Electrolysis

2 Voltaic Vs. Electrolytic Cells Voltaic Cell Energy is released from spontaneous redox reaction System does work on load (surroundings) Oxidation Reaction X g X + + e- Reduction Reaction e- + Y + g Y Overall (Cell) Reaction X + Y + g X + + Y, ΔG = 0 Electrolytic Cell Energy is absorbed to drive nonspontaneous redox reaction Surrounding (power supply) do work on system (cell) Oxidation Reaction A - g A + e- Reduction Reaction e- + B + g B Overall (Cell) Reaction A - + B + g A + B, Δ G> 0 General Characteristics of voltaic and electrolytic cells. A voltaic cell generates energy from a spontaneous reaction (ΔG < 0), whereas an electrolytic cell requires energy to drive a nonspontaneous reaction (ΔG > 0). In both types of cell two electrodes dip into electrolyte solution, and an external circuit provides the means for electrons to flow. Oxidation takes place at the anode, and reduction takes place at the cathode, but the relative electrode changes are opposite in the two cells. The anode in the electrolytic cell is now the positive (+) electrode and the cathode in the electrolytic cell is now the negative (-) electrode. 2 Electrolysis

Electrolytic Cell Two electrodes (inert or non inert) External potential (power) Source of electricity, battery (e- pump) g Drive a non

3 Electrolysis What is the reaction which occurs in an electrolysis process? Voltaic/Galvanic Cell - Produce electricity Electrolysis - Causes nonspontaneous reaction to occur by using external energy source. Electrolytic Cell Two electrodes (inert or non inert) External potential (power) Source of electricity, battery (e- pump) g Drive a non spontaneous reaction e- g + - e- g pulls electron Draws electron to positive (Oxidation) Anode (+) A g A - + e - push electron Pumps electron to reaction (Reduction) Cathode (-) e - + B + g B 3 Electrolysis

4 Electrolysis: Molten Salt Question: Electrolysis NaCl What is the direction of e- flow? Which is the Anode/ Cathode Which is the Oxidation / Reduction? Which is the Positive / Negative terminal? What are the Half rxns. With NaCl Electrode Ions... Anode Rxn Cathode Rxn E Pt(inert) Pt (inert) (inert) NaCl Na + Na + +e - g Na (s) V Cl- 2Cl - g Cl 2 + 2e V To drive rxn, Battery requires 4.07 V Complications due to H 2 O undergoing electrolysis process. This reaction must be considered. 4 Electrolysis

5 Water Complications in Electrolysis In an electrolysis, the most easily oxidized and most easily reduced reaction occurs. When water is present in an electrolysis reaction, then water (H 2 O) can be oxidized or reduced according to the reaction shown. Electrode Ions... Anode Rxn Cathode Rxn E Pt (inert) H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V H 2 O 2 H 2 O (l) g 4e - + 4H + (g) + O 2(g) V Net Rxn Occurring: 2 H 2 O g 2 H 2(g) + O 2 (g) E = V 5 Electrolysis

6 Electrolysis of NaCl Question: i) What is the direction of e- flow? ii) Which is the Anode/ Cathode iii) Which is the Oxidation / Reduction? iv Which is the Positive / Negative terminal? v) What are the Half rxns. Reaction which occurs is the reaction requiring least E potential Specie Ions... Anode Rxn Cathode Rxn E Pt (inert) (inert) NaCl Na + Na + + e - g Na (s) V Cl- 2Cl - g Cl 2 + 2e V H 2 O H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V H 2 O H 2 O (l) g 4e - + 4H + (aq) + O 2(g) V Overall Rxn: Anode: 2 H 2 O (l) g 4e - + 4H + (aq) + O 2(g) V Cathode: H 2 O (l) + 2e - g H 2(g) + 2OH - (aq) V E Cell = V Electrolysis of water is the reaction occurring. 6 Electrolysis

Electrolysis of MgCl 2 What is the reaction in an electrolysis reaction of MgCl 2. Reaction which occurs is the reaction requiring least E potential Specie Ions.

7 Electrolysis of MgCl 2 What is the reaction in an electrolysis reaction of MgCl 2. Reaction which occurs is the reaction requiring least E potential Specie Ions... Anode Rxn Cathode Rxn E Pt (inert) (inert) MgCl 2 Mg 2+ Mg 2+ +2e - g Mg (s) V Cl- 2Cl - g Cl 2 + 2e V H 2 O H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V H 2 O 2 H 2 O (l) g 4e - + 4H + (g) + O 2(g) V Overall Rxn: Anode: 2 H 2 O (l) g 4e - + 4H + (aq) + O 2(g) V Cathode: H 2 O (l) + 2e - g H 2(g) + 2OH - (aq) V E Cell = V MgCl 2 undergoes no chemical change. Electrolysis of water is the reaction occurring. 7 Electrolysis

36 V H 2 O H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) -0.83 V H 2 O 2 H 2 O (l) g 4e - + 4H + (g) + O 2(g) -1.

8 Electrolysis with non-inert Electrode What is the reaction in an electrolysis reaction of NaCl with iron nails Specie Ions... Anode Rxn Cathode Rxn E Fe Fe g Fe e V NaCl Na + Na + + e - g Na (s) V Cl- 2Cl - g Cl 2 + 2e V H 2 O H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V H 2 O 2 H 2 O (l) g 4e - + 4H + (g) + O 2(g) V Overall Rxn: Anode: Fe g Fe e V Cathode: H 2 O (l) + 2e - g H 2(g) + 2OH - (aq) V E Cell = -.39 V Iron oxidize at the anode and water is reduced at the cathode. A potential of 0.39 V is used to drive this reaction. 8 Electrolysis

A, Even though the half-reactions involve the same components, the cell operates because the half-cell

9 Electrolysis of Copper A net reaction of zero, yet a process does take place. A concentration cell based on the Cu/ Cu2+ half-reaction. A, Even though the half-reactions involve the same components, the cell operates because the half-cell concentrations are different. B, The cell operates spontaneously until the half-cell concentrations are equal. Note the change in electrodes (exaggerated here for clarity) and the equal color of solutions. 9 Electrolysis

10 Electro-Plating Example BLB: Calculate the minimum applied voltage required to cause the following electrolysis reaction to occur, assuming that the anode is platinum and the cathode is nickel: Ni 2+ (aq) + 2Br- (aq) g Ni (s) + Br 2(l) Specie Ions... Anode Rxn Cathode Rxn E Pt (Inert) (Inert) Ni +2 Ni 2+ Br- H 2 O H 2 O H 2 O 10 Electrolysis

11 Electro-Plating Example BLB: Calculate the minimum applied voltage required to cause the following electrolysis reaction to occur, assuming that the anode is platinum and the cathode is nickel: Ni 2+ (aq) + 2Br- (aq) g Ni (s) + Br 2(l) Specie Ions... Anode Rxn Cathode Rxn E Pt (Inert) (Inert) Ni +2 Ni 2+ Ni e - g Ni (s) V Br- 2Br - g Br 2 + 2e V H 2 O H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V H 2 O 2 H 2 O (l) g 4e - + 4H + (g) + O 2(g) V 11 Electrolysis

12 Electro-Plating Example BLB: Calculate the minimum applied voltage required to cause the following electrolysis reaction to occur, assuming that the anode is platinum and the cathode is nickel: Ni 2+ (aq) + 2Br- (aq) g Ni (s) + Br 2(l) Specie Ions... Anode Rxn Cathode Rxn E Pt (Inert) (Inert) Ni +2 Ni 2+ Ni e - g Ni (s) V Br- 2Br - g Br 2 + 2e V H 2 O H 2 O H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V H 2 O 2 H 2 O (l) g 4e - + 4H + (g) + O 2(g) V Overall Rxn: Cathode: Ni e - g Ni (s) V Anode: 2Br - g Br 2 + 2e V E Cell = V A larger voltage than the E Cell = 1.35 V is required to overcome cell resistance (overvoltage) if rxn is to occur at reasonable rate. 12 Electrolysis

Overvoltage At anode (Oxidation) E 2H 2 O g 4H + + O 2 + 4e - - 1.23 V Actually the voltage necessary to drive this reaction is greater than the theoretical value.

13 Overvoltage At anode (Oxidation) E 2H 2 O g 4H + + O 2 + 4e V Actually the voltage necessary to drive this reaction is greater than the theoretical value. Overvoltage - Voltage necessary to overcome activation energy Reason: i) Slow e- transfer at electrode surface ii) Process (Kinetics) of bond breaking. The actual overvoltage depends on the type of electrode and can only be determine experimentally. For water an additional 0.4V is necessary for oxidation, therefore, instead of V for water reduction, -1.6 V is necessary. 13 Electrolysis

14 Electrolysis Reactions order A Solution contains the ions H +, Cu 2+, Ca 2+ and Ni 2+, each at a concentration of 1.0 M. Determine the sequence of ions to be reduce at the Cathode? Reduction at Cathode E 1st) 2H + g H V 2nd) Cu +2 g Cu 0.34 V 3rd) Ca 2+ g Ca V 4th) Ni 2+ g Ni 0.25 V 14 Electrolysis

15 Electrolysis Reactions A Solution contains the ions H +, Cu 2+, Ca 2+ and Ni 2+, each at a concentration of 1.0 M. Determine the sequence of ions to be reduce at the Cathode? Reduction at Cathode E 1st) Cu +2 g Cu 0.34 V 2nd) Ni 2+ g Ni 0.25 V 3rd) 2H + g H V 4th) Ca 2+ g Ca V * * This reaction cannot be reduce in aqueous solution since H 2 O will be reduced first H 2 O reduction: H 2 O (l) + 2e- g H 2(g) + 2OH - (aq) V 15 Electrolysis

16 Electrolysis Stoichiometry The amount of metal plated can be calculated based on stoichiometry. 1 Ampere = Coulomb per time (C/s) 1 Faraday = C / mol e- Q 100: Silver is electroplated. i) What mass of silver plates if a current of 6.8A flows through the cell in 72 min. ii) What electrode will the silver plate-out. Reaction: Ag + (aq) + e- g Ag (s) (Reduction occurs at the cathode for electrolysis) Given: Goal: 6.8C = 1s Mass Ag (g) C = 1 mol e- 72 min = 432 sec g Ag = 1 mol Ag 60 sec 72 min 1 min 6.8 C 1 sec 1 mol e C mol Ag g Ag mol e- 1 mol Ag = Mass Ag = g of Silver 16 Electrolysis

17 Electrolysis Stoichiometry The amount of metal plated can be calculated based on stoichiometry. 1 Ampere = Coulomb per time (C/s) 1 Faraday = C / mol e- Q 99: Copper can be electroplated at the cathode of an electrolysis cell by the half-reaction: Reaction: Cu +2 (aq) + 2e- g Cu (s) (Reduction occurs at the cathode for electrolysis) How much time would it take to plate 325 mg of copper with 5.6 A? Given: Goal: 5.6 C = 1s Time for electrolysis C = 1 mol e- 325 mg =.325 g g Cu = 1 mol Cu 1 mol Cu g Cu g 2 mol e- 1 mol Cu C 1 mol e- 1 sec 5.6 C = Time electrolysis = 176 sec = 180 sec 17 Electrolysis

18 Voltaic Vs. Electrolytic Cells Voltaic: Generate E (+) Conductivity e- flow (external circuit) migration of ions (in cell) complete circuit (bridge) Anode (-) Oxidation Cathode (+) Reduction Electrolytic: Consumes E (-) Conductivity e- flow (pump by battery) migration of ions (in cell) Lowest E reaction occurs Anode (+) Oxidation Cathode (-) Reduction Oxidation Anode Reduction Cathode M g M + + e- M + + e - g M Neg - Terminal Pos - Terminal Reduction Cathode Oxidation Anode M + + e - g M M g M + + e- Neg - Terminal Pos - Terminal 18 Electrolysis

19 Summary General characteristics of voltaic and electrolytic cells. A voltaic cell generates energy from a spontaneous reaction (ΔG < 0), whereas an electrolytic cell requires energy to drive a nonspontaneous reaction (Δ G>0). In both types of cell, two external circuits provides the means or electrons to flow. Oxidation takes place all the anode, and reduction takes place at the cathode, but the relative electrode changes are opposite in the two cells. Voltaic Cell Energy is released from spontaneous redox reaction System does work on load (surroundings) Oxidation Reaction X g X + + e- Reduction Reaction e- + Y + g Y Overall (Cell) Reaction X + Y + g X + + Y, ΔG > 0 Electrolytic Cell Energy is absorbed to drive nonspontaneous redox reaction Surrounding (power supply) do work on system (cell) Oxidation Reaction A - g A + e- Reduction Reaction e- + B + g B Overall (Cell) Reaction A - + B + g A + B, Δ G> 0 19 Electrolysis

Electrochemical Cells

Electrochemistry Electrochemical Cells The Voltaic Cell Electrochemical Cell = device that generates electricity through redox rxns 1 Voltaic (Galvanic) Cell An electrochemical cell that produces an electrical