Additional electrochemistry

History of Electrochemistry Key people: Luigi Galvani Alessandro Volta Humphry Davy Michael Faraday

Luigi Galvani Showed that electricity is generated when two different metals are placed in a conducting solution

Alessandro Volta Was the first to construct a battery using lead and zinc plates

Humphry Davy Used electrolysis to isolate reactive elements such as sodium, potassium and calcium

Michael Faraday Made many discoveries about electrolysis

Electrochemical Series The electrochemical series is an arrangement of metals in order of tendency to lose electrons Usually, the higher a metal is in the electrochemical series, the more reactive it is. The position of a metal in the series relative to another may be determined experimentally.

Determining the relative positions of zinc and copper

Electrolysis of molten lead bromide Electrolyte is composed of Pb2+ ions and Br- ions The Pb2+ ions move to the cathode, where they are reduced: Pb2+(l) + 2e- → Pb(l) The Br- ions move to the anode, where they are oxidised: 2Br-(l) - 2e- → Br2(g)

Electrolysis of molten lead bromide (continued) Inert electrodes used, to prevent the bromine formed from reacting Overall reaction for the electrolysis is: Pb2+(l) + 2Br-(l) → Pb(l) + Br2(g)

Electrolysis of molten lead bromide

Corrosion Corrosion of metals is caused by the action of air, water, and other chemicals, such as acids, on the metal surface

Corrosion of iron and steel An electrochemical cell is formed with a cathode and an anode. The anode gets eaten away as the iron is oxidised: Fe(s) → Fe2+(aq) + 2e- At the cathode, the dissolved oxygen in the water reacts: H2O(l) + 1/2O2(aq) + 2e- → 2OH-(aq)

Corrosion of iron and steel (continued) The Fe2+ and OH- ions diffuse away from the electrodes, and form a precipitate of iron(II) hydroxide: Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s) This is then oxidised by dissolved oxygen to form rust, Fe2O3.xH2O, where x is variable Hydrated iron(III) oxide (rust) flakes off and does not protect the metal from further corrosion

Rusting of iron

Acceleration of corrosion The presence of salt in water accelerates corrosion of iron and steel This is because it increases the conductivity of the electrochemical cell solution If iron is in contact with a metal such as lead, which is below it in the electrochemical series, as, for example, when lead piping is joined to an iron storage tank, the rate of corrosion is accelerated greatly

Prevention of corrosion Painting (to exclude air) Coating with tin, plastic, grease or oil (to exclude air) Protection against rust afforded by these methods lasts only until the coating is scratched

Galvanising Involves coating the metal with zinc The zinc corrodes and the iron is protected Zn(s) → Zn2+(aq) Cathodic protection, where the cathode is protected and the anode is eaten away

How zinc protects Fe from rusting, while tin does not

Sacrificial anodes A sacrificial anode made from a more reactive metal such as magnesium or zinc is connected to steel The anode is eaten away, and the steel, which is the cathode, is protected Sacrificial anodes are widely used to protect the steel of ships and underground or submerged oil pipelines and water pipes

Corrosion-resistant metals Some metals form protective oxide coatings when exposed to air. The oxide coatings formed by aluminium, nickel, chromium and magnesium do not flake off easily These metals are consequently protected from further corrosion.