Siyapath
Sign up

Lesson: Chapter 12 — Electrochemistry

Putting the chapter together 1 of 3

Chapter summary

The whole chapter in one list.

Electrochemical cells

  • Electrochemical cells are used to turn chemical energy into electrical energy.
  • A simple cell can be made by joining two rods of different metals with a conducting wire and dipping them in an acid solution.
  • In a simple electrochemical cell the more active metal acts as the anode and the less active metal as the cathode.
  • An oxidation half reaction happens at the anode and a reduction half reaction at the cathode.
  • In an electrochemical cell the anode is the negative terminal and the cathode the positive terminal.
  • Electrons flow along the wire from the anode to the cathode.
  • Conventional current is taken as flowing from the positive terminal (cathode) to the negative terminal (anode).

Electrolysis

  • Bringing about chemical changes in substances by passing an electric current through a conducting solution or liquid is electrolysis.
  • For this, carbon or metal electrodes are connected to an outside supply and dipped in the solution, so electricity passes through the solution or liquid.
  • A liquid or solution that carries electricity is an electrolyte. To conduct, an electrolyte must contain ions that can move.
  • The positive electrode of an electrolytic cell acts as the anode, so an oxidation half reaction happens at it.
  • Making useful products from what forms at the electrodes is an industrial use of electrolysis.
  • Sodium metal is obtained industrially by electrolysing molten sodium chloride. The by-product, chlorine, is used for other useful purposes.

Corrosion and rusting

  • A metal's surface undergoing chemical change through exposure to the atmosphere and moisture is corrosion of metals; iron and steel corroding like this is called, in particular, rusting.
  • Oxygen and moisture are essential for iron to rust.
  • The corrosion of iron is an electrochemical process. Its anode reaction is Fe (s) → Fe2+ (aq) + 2e, and its cathode reaction 2H2O (l) + O2 (g) + 4e → 4OH- (aq).
  • The overall corrosion reaction comes from these: 2H2O (l) + O2 (g) + 2Fe → 2Fe(OH)2 (s). Fe(OH)2 is oxidised further to hydrated ferric oxide (Fe2O3.H2O), or rust.
  • Acids and salts such as sodium chloride increase the rate of rusting; bases reduce it.
  • Rusting can be prevented by keeping iron from contact with the essentials, oxygen and moisture. For this, paint, grease or tin can be coated on the iron as a protective layer.
  • When a metal more active than iron is joined to it, the active metal acts as the anode and the iron as the cathode, so rusting is prevented. This is sacrificial protection; galvanising iron is an example.

The textbook's summary writes the rust formula Fe2O3.H2O; its body text explains that Fe2O3.xH2O is more accurate, because the amount of water varies.