1. Additional electrochemistry

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

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

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

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

6. Michael Faraday
Made many discoveries about electrolysis

7. 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.

8. Determining the relative positions of zinc and copper

9. 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)

10. 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)

11. Electrolysis of molten lead bromide

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

13. 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)

14. 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) OH-(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

15. Rusting of iron

16. 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

17. 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

18. 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

19. How zinc protects Fe from rusting, while tin does not

20. 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

21. 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.