Redox I

• oxidation - the gain of oxygen, but loss of electrons
• reduction - the loss of oxygen, but gain of electrons
• oxidising agent - they oxidise another species, and undergo reduction themselves (they accept an electron from the species being oxidised).
• reducing agent -they reduce another species, and undergo oxidation themselves (they donate an electron to the species being reduced)
• atoms in an element have an oxidation number of 0
• simple ions have an oxidation state which matches the charge on the ion.
• set oxidation states:

  F = always -1

  O = -2 (except +2 when reacted with F and -1 in peroxides (H2O2)

  Cl, Br, I = -1 (can be positive when bonded to F or O)

  H = +1 (except -1 in metal hydrides)
• set oxidation states:

  group 1 metals - always +1

  group 2 metals - always +2

  aluminium - always +3
• the sum of oxidation states of elements in a compound is equal to the charge on the compound.
• oxidation half equation (for Iodine)

  2I- (aq) = I2 (aq)+ 2e-
• reduction half equation (for bromine)

  Br2 (aq) + 2e- = 2Br- (aq)
• Balancing complicated half equations steps (acidic conditions)

  E: Balance the element being oxidised/reduced

  e: add electrons to balance the oxidation state.

  O: add H2O to balance oxygen

  H: add H+ to balance hydrogen (so number of hydrogens is the same on both sides)

  c: check charges balance.
• when combing half equations to make ionic/normal equations, remember:

  • electrons need to be on opposite sides of the reactions.
  • balance so the number of electrons is the same.
• balancing complicated half equations (alkaline conditions)

  1. balance half equation normally
  2. For every$H^+$ion present, add the same number of $OH^-$ions to both sides of the reaction. When both ions are on the same side, they combine to form water.
  3. Cancel out the equation as normal.