 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 | Zn 2+ | | Cu 2+ | Cu › “|” separation of electrode and ions › “| |” salt bridge

 Electrons are produced by oxidation of Zinc at anode. The electrons are used by Cu 2+ 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  Mg e - Eº = 2.37 V Cu e -  CuEº = 0.34 V Mg + Cu 2+  Mg 2+ + CuEº = 2.71 V › Cell notation:  Mg | Mg 2+ | | Cu 2+ | CuEº = 2.71 MgCu

 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) + SO 4 2- (aq) → PbSO 4(s) + 2e - (oxidation) PbO 2(s) +4 H + (aq) + SO 4 2- (aq) + 2e - → PbSO 4(s) + 2H 2 O (l) (reduction )

The anode and cathode are immersed in H 2 SO 4 and are coated with PbSO 4 as the electrical current is drawn. The battery goes dead when too much PbSO 4 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 Fe 2+ (aq) + 4e - O 2(g) + 2 H 2 O (l) + 4e - → 4 OH - (aq) 2 Fe (s) + O 2(g) + 2 H 2 O (l) → 2 Fe(OH) 2(s)  The Fe(OH) 2 undergoes several additional reactions to for Fe 2 O 3 (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.