14.3 Electrolytic Cells  Electrolysis is the process of converting electrical energy into chemical energy.  Voltaic cells produce electrical energy and are viewed as exothermic reactions.  Electrolytic cells use electrical energy and are viewed as endothermic reactions.

Electrolytic cells require a flow of electrons from an outside energy source such as from a power supply or battery. Positive E  net values for electrochemical cells indicate that those reactions occur spontaneously. Electrolytic cells have negative E  net values which indicate that a minimum voltage must be applied to force a non-spontaneous reaction to occur. In electrolytic cells the cathode is negative and the anode is positive. (L.E.O.P.A. / G.E.R.N.C.)

Rechargeable Batteries Nickel-Cadmium (Ni-Cad) Batteries By using specific chemicals one can create a battery that can be easily recharged. At the cathode, the SOA is Ni(OH)(s). At the anode, the SRA is Cd(s). The net voltage of each cell is 1.25 V. Acts as a voltaic cell when it discharges. Acts as an electrolytic cell when it recharges.

Example 1: Aqueous Solutions Write half reactions and the net reaction. Compute the minimum voltage for the electrolysis of Pb(NO 3 ) 2 (aq).

Example 1: Aqueous Solutions Write half reactions and the net reaction. Compute the minimum voltage for the electrolysis of Pb(NO 3 ) 2 (aq). SOA OA Pb 2+ (aq) / NO 3 – (aq) / H 2 O (l) not on tables SRA GERNC: 2 ( Pb 2+ (aq) + 2 e –  Pb (s) ) LEOPA: 2 H 2 O (l)  O 2(g) + 4 H + (aq) + 4 e – Net: 2 Pb 2+ (aq) + 2 H 2 O (l)  2 Pb (s) + O 2(g) + 4 H + (aq) E° cell = E° r - E° r cathode anode = (-0.13) - (+1.23) = E  net = – 1.36 V  V min = V

Apparatus:

Power Supply C (s) Pb (s) C (s) e – V min = 1.36 V Pb 2+ (aq) NO 3 – (aq) Solution becomes acidic. H + (aq) O 2(g) Apparatus:

Example 2: Aqueous Electrolysis Draw and label an electrolytic cell for a solution of nickel (II) chloride.

Example 2: Aqueous Electrolysis Draw and label an electrolytic cell for a solution of nickel (II) chloride. SOA OA Ni 2+ (aq) / Cl – (aq) / H 2 O (l) RA SRA ***But, Cl – (aq) oxidizes faster than H 2 O (l) *** GERNC:Ni 2+ (aq) + 2 e –  Ni (s) LEOPA:2 Cl – (aq)  Cl 2(g) + 2 e – Net:Ni 2+ (aq) + 2 Cl – (aq)  Ni (s) + Cl 2(g)

E° cell = E° r - E° r cathode anode = (-0.26) - (+1.36) E  net = – 1.62 V  V min = V

Power Supply

C (s) Cl 2(g) bubbles produced C (s) e – V min = 1.62 V Ni 2+ (aq) Cl – (aq) Ni (s) electroplating Anode Cathode Power Supply

Example 3: Electrolysis of Water

SOA = H 2 O (l) SRA = H 2 O (l) GERNC:2 (2 H 2 O (l) + 2 e –  H 2(g) + 2 OH – (aq) ) LEOPA:2 H 2 O (l)  O 2(g) + 4 H + (aq) + 4 e – Redox: 6 H 2 O (l)  2 H 2(g) + O 2(g) + 4 H + (aq) + 4 OH – (aq) 4 H 2 O (l) Simplify the waters: 2 H 2 O(l)  2 H 2 (g) + O 2 (g)

E° cell = E° r - E° r cathode anode = (-0.83) - (+1.23) E  net = – 2.06V Therefore the minimum voltage needed to force this non-spontaneous reaction to occur is V.

Electrolysis of Molten Compounds When ionic compounds are melted the "melt" contains liquid ions. The "melt" is a good electrical conductor and can undergo electrolysis. No water is present.

Example 1: Write half and net equations and draw a labeled diagram for the electrolysis of molten zinc chloride. Species present: Zn 2+ (l) / Cl – (l) cation anion GERNC:Zn 2+ (l) + 2 e –  Zn (l) LEOPA:2 Cl – (l)  Cl 2(g) + 2 e – Net:Zn 2+ (l) + 2 Cl – (l)  Zn (l) + Cl 2(g)

Anode Power Supply The "Melt" Zn 2+ (l) Cl – (l) Zn (l) Cathode e – High Temp Source Cl 2(g)

Example 2: Describe the electrolysis of molten gallium oxide. Species present: Ga 3+ (l) / O 2– (l) GERNC:4 ( Ga 3+ (l) + 3 e –  Ga (l) ) LEOPA:3 ( 2 O 2– (l)  O 2(g) + 4 e – ) Net: 4 Ga 3+ (l) + 6 O 2– (l)  4 Ga (l) + 3 O 2(g)

initial volume Ga 3+ (l) O 2– (l) Ga (l) Cathode e – High Temp Source O 2(g) Anode Power Supply

Why is the electrolysis of molten compounds so important? The technique discovered by Sir Humphrey Davy became efficient in the late 1800's. If an aqueous solution of a salt containing an oxidizing agent that is a metallic ion weaker than water is electrolyzed, the water will be reduced at the cathode to form hydrogen gas. If a molten salt is electrolyzed, the metallic ion is reduced. Electrolysis permits production of active metals such as:  Cs + (l) + e –  Cs (l)  Li + (l) + e –  Li (l)  Al 3+ (l) + 3 e –  Al (l) – Ca 2+ (l) + 2 e –  Ca (l)

Industrial Electrolysis 1) Chlor-Alkali Electrolytic Cells NaCl(aq) saturated brine is pumped from an underground salt dome at Dow Chemical in Fort Saskatchewan. NaCl(aq) is pumped between huge electrodes. Cathode: 2 H 2 O (l) + 2 e –  H 2(g) + 2 OH – (aq) Anode: 2 Cl – (aq)  Cl 2(g) + 2 e – Net: 2 H 2 O (l) + 2 Cl – (aq)  H 2(g) + 2 OH – (aq) + Cl 2(g)  Na + (aq) remains in solution with the OH – (aq) : NaOH (aq)

Uses a) Cl 2(g) is used to make  disinfectant for drinking water  bleaches ( NaOCl (aq), Ca(OCl) 2(s)  plastics (polyvinyl chloride)  pesticides (2,4 - D)  solvents (C 2 Cl 4 - dry-cleaning)

b) H 2 (g) is used to make  ammonia  hydrogen peroxide  Margarine c) NaOH (s) - Caustic Soda ( lye ) is used to make  cellophane  pulp and paper  aluminum  detergents

2) Downs Process Molten NaCl (l) is electrolyzed. Cathode: Na + (l) + e –  Na (l) Anode:2 Cl – (l)  Cl 2(g) + 2 e – Note: NaCl (s) is dissolved in molten CaCl 2(l) to reduce NaCl's melting point (805  C). Cl 2(g) NaCl (l) Na (l) Carbon anode Cathode

Electrolysis of Molten Sodium Chloride

Uses Sodium is cooled to form a solid and used in sodium vapour lamps and as a coolant in some nuclear power reactors. Chlorine is sold for commercial use already discussed.

3) Hall-Heroult Process  Bauxite (Al 2 O 3(s) ) is dissolved in molten cryolite (Na 3 AlF 6(l) ) at about 970  C. Cathode: 4(Al 3+ (l) (cryolite) + 3 e –  Al (l) ) Anode: 3(2 O 2– (l) (cryolite)  O 2(g) + 4 e – ) Net: 4 Al 3+ (cryolite) + 6 O 2- (cryolite) → 4 Al(l) + 3 O 2 (g) Overall the reaction is a decomposition of aluminum oxide: 2 Al 2 O 3 (s) → 4 Al(s) + 3 O 2 (g)

Refining of Metals Electrorefining is the process of using an electrolytic cell to obtain high-grade metals at the cathode from an impure metal at the anode. Electrowinning is the process of using an electrolytic cell to reduce metal cations from a molten or aqueous electrolyte at the cathode.

Electroplating It is plating of a metal at the cathode of an electrolytic cell. Example: Cathode (spoon): Ag + (aq) + 1e - → Ag(s) Anode: Ag(s) → 1e - + Ag + (aq)

Voltaic CellsElectrolytic Cells 1) 2) 3) 4)

Porous Barrier anions (–) cations (+) e – Cathode SOA is Reduced GERPC Anode SRA is Oxidized LEONA e – Anode SRA is Oxidized LEOPA Cathode SOA is Reduced GERNC anions (–) cations (+) D.C. Power Supply 5) 6)

Voltaic CellsElectrolytic Cells 1)Energy Conversion: Chemical Electrical Energy Conversion: Electrical Chemical 2)SpontaneousNon Spontaneous 3)Exothermic Produces electricity Endothermic Absorbs electrical energy from a "power supply" 4) E  net = "+" - internally driven E  net = "–"  V min = + E  - externally driven

Porous Barrier anions (–) cations (+) e – Cathode SOA is Reduced GERPC Anode SRA is Oxidized LEONA e – Anode SRA is Oxidized LEOPA Cathode SOA is Reduced GERNC anions (–) cations (+) D.C. Power Supply 5) 6)Examples: a) alkaline cell b) lead - acid car battery Examples: a) at DOW: NaCl (aq) Electrolysis b) Electroplating

Read pgs. 639 – 650 pgs. 640, 644, 645, 649 Practice #'s 1, 2, 3, 5, 6, 12, 13, 14, 16 pg. 651 Section 14.3 Questions #’s 4, 9, 15