Unit Fourteen: Redox Reactions

Unit Fourteen: Redox Reactions

Deals with the relation of the flow of electric current to chemical changes and the conversion of chemical to electrical energy (Electrochemical Cell) and electrical to chemical energy (Electrolysis) A device that can create electrical current from a

spontaneous redox reaction Electrodes: Anode oxidation occurs (- in chemistry) Cathode reduction occurs (+ in chemistry) Salt bridge saturated salt solution that connects the two half-cells

Cell notation shorthand form used to describe the cell Zn | Zn2+ | | Cu2+ | Cu | separation of electrode and ions | | salt bridge

Electrons are produced by oxidation of Zinc at anode. The electrons are used by Cu2+ for reduction at the cathode. The electrochemical cell dies when the anode is used up Cations migrate toward cathode and

anions migrate toward anode Mg and Cu Determine anode: Mg cathode: Cu Cell voltage: Mg Mg2+ + 2eE = 2.37 V Cu2+ + 2e- Cu E = 0.34 V

Mg + Cu2+ Mg2+ + Cu E = 2.71 V Cell notation: Mg | Mg2+ | | Cu2+ | Cu E = 2.71 Pb and Cu Ni and Fe

Found in most automobiles Consist of six electrochemical cells wired in series Each cell produces 2 volts for a total of 12 volts Each cell contains a porous lead anode where

oxidation occurs according to the following reaction Pb(s) + SO42-(aq) PbSO4(s) + 2e- (oxidation) PbO2(s) +4 H+(aq) + SO42-(aq) + 2e- PbSO4(s) + 2H2O(l) (reduction) The anode and cathode are immersed in H2SO4 and are coated with PbSO4 as the

electrical current is drawn. The battery goes dead when too much PbSO4 develops. Recharged by running the electrical current in reverse.

The reactants are constantly replenished Most common is the hydrogen-oxygen fuel cell Hydrogen gas flows past the anode and undergoes oxidation. Oxygen flows past the cathode and undergoes reduction. The sum of the two half reactions only product produced is water.

The fuel constantly flows trhrough the battery, generating electrical current as they undergo a redox reaction. Electrical current is used to drive an otherwise nonspontaneous redox reaction.

Electrolytic Cell an electrochemical cell used for electrolysis Used to produce metals from metal oxides and to plate metals onto other metals. Silver is being oxidized on the left side and reduced on the right. As it is reduced, it is deposited on the object to be plated.

Most common is the rusting of iron 2 Fe(s) 2 Fe2+(aq) + 4eO2(g) + 2 H2O(l) + 4e- 4 OH-(aq) 2 Fe(s) + O2(g) + 2 H2O(l) 2 Fe(OH)2(s) The Fe(OH)2 undergoes several additional reactions to for Fe2O3 (orange substance called rust).

Preventing rust Keep dry (rust cannot occur without moisture) Coat iron with substance impervious to water Sacrificial electrode Must be composed of metal above iron in activity series Sacrificial electrode oxidizes in place of iron Galvanized

Coat iron with a metal above itself on the activity series Zinc, for example, will oxidize before iron. Zinc oxide does not crumble, so it remains on the iron as a protective coating.

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