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: cathode:

› Cell voltage: Mg Mg2+ + 2e- Eº = 2.37 V

Cu2+ + 2e- Cu Eº = 0.34 V Mg + Cu2+ Mg2+ + Cu Eº =

2.71 V

› Cell notation:Mg | Mg2+ | | Cu2+ | Cu Eº = 2.71

Mg Cu

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) + 4e-

O2(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.