Electrochemistry is the chemistry of reactions which involve electron transfer. In spontaneous reactions electrons are released with energy which can be used in electrochemical cells. In non-spontaneous reactions electrons have to be supplied with energy in order to produce chemicals that are wanted in electrolytic cells or electrolysis. In both electrochemical and electrolytic cells the terms anode and cathode have the following meanings: Anode: The electrode where oxidation is occurring Cathode: The electrode where reduction is occurring

This is the use of a oxidation-reduction reaction to produce an electric current (electricity) To do this you have to separate the two half equations from one another and provide a path for the electrons to flow Half-cell where occurs Half-cell where occurs reduction oxidation But there is a problem…….

Oxidizing agent Reducing agent Becomes negative as electrons arrive Becomes positive as electrons leave e- Electrons should move through the wire from Fe 2+ to MnO 4 -. But they wont

This problem can be solved The solutions need to be linked so that ions can can also flow to keep the net charge in each beaker equal. This can be done by using a salt bridge Or a porous disk in a tube linking the two solutions.

This is called an electrochemical battery or a galvanic cell The electrons are forced to flow through a wire by separating the reducing agent from the oxidizing agent.

Oxidation occurs at the anode Reduction occurs at the cathode

Lead storage batteries are used in all automobiles. It uses lead metal, Pb as the reducing agent and lead (IV) oxide, PbO 2 as the oxidizing agent Sulfuric acid, H 2 SO 4 provides the H + needed for the reaction. It also provides SO 4 2- ions that react with the Pb 2+ ions to form PbSO 4 (s) The PbSO 4 (s) coats the lead electrodes and will cause the battery to go flat. The battery can be recharged by forcing current in the opposite direction to reverse the cell reaction.

Lead Acid Battery (Car Battery): Anode: Pb (s)  Pb 2+ (aq) + 2e- Cathode: PbO 2 (s) + 2e- + 4H + (aq)  Pb 2+ (aq) +2H 2 O (l) The overall equation is PbO 2 (s) + Pb (s) + 2H 2 SO 4(aq)  Pb SO 4(aq) + 2H 2 O (l)

You can reverse the redox reaction to recharge the car battery. Your car uses an alternator which generates electricity which is used to recharge the battery You can do this by passing an electric current back into the battery. This turns the Pb 2+ back into PbO 2 (s) + Pb Pb SO 4(s) + 2H 2 O (l)  PbO 2 (s) + Pb (s) + 2H 2 SO 4(aq)

Dry Cell Batteries These are small efficient batteries used in calculators, electronic watches, portable radios and tape players. They are called dry cells because they don’t contain a liquid electrolyte.

The Common Dry Cell Battery Zn  Zn 2+ (aq) + 2e- Anode: 2 NH MnO 2 + 2e-  Mn 2 O 3 + NH 3 + H 2 O Cathode:

H 2 O + 2MnO 2 + 2e-  Mn 2 O 3 + 2OH - Alkaline dry cell battery Anode: Cathode: Zn + 2OH -  ZnO (s) + H 2 O + 2e-

Mercury Battery This is an example of a dry battery. It is the battery found in calculators

This type of battery can’t be recharged. Anode: Zn (s)  Zn 2+ (aq) + 2e- Cathode: HgO (s) + 2e- + 2H +  Hg (aq) + H 2 O

4Fe (s) + 3O 2 (g)  2Fe 2 O 3 (s) Fe (s)  Fe e- O 2 (g) + 4e-  2O 2- Oxidation: Reduction:

How do you stop corrosion? You need to prevent water and oxygen from coming into contact with metals that will oxidize by: Coating metals in oil, paint, plastic or another metal Cathodic protection Or create alloys of iron that resist corrosion like stainless steel Corrosion destroys more than $7 billion worth of equipment in the USA each year. 20% of tha annual steel and iron production in the USA goes to replacing corroded materials

This is a way of protecting steel on ships, buried fuel tanks and pipelines A metal that is a better reducing agent than steel is attached to the steel

The magnesium is a stronger reducing agent than iron. So electrons flow from the magnesium into the iron, rather than from out of the iron. This stops the iron from being oxidized. As oxidation of the magnesium occurs, the magnesium dissolves and needs to be replaced periodically.

This is the process of forcing a current through a cell to produce a chemical change that would not otherwise occur. Uses: Recharging batteries Electroylysis of water to produce O 2 and H 2 (g) Another important use of electrolysis is the production of metals. Aluminum is an example of a metal produced by this process.

Electrolysis of water: H 2 O (l)  O 2 (g) + H 2 (g) O 2 (g) H 2 (g)