Topic 8

• oxidation


  loss of electrons
• reduction


  gain of electrons
• oxidising agents


  electron acceptors (they get reduced)
• reducing agents


  electron donors (they get oxidised)
• redox reaction


  a reaction that involves reduction and oxidation
• disproportionation reaction


  a reaction in which the same species is simultaneously oxidised and reduced
• oxidation number


  the charge that an element in a compound or ion would have if the compound was fully ionic
• Oxidation number of an element is 0
• Oxidation number of an element in a simple ion is the charge of the ion
• Oxidation number of F is -1
• Oxidation number of H is +1 with non-metals, but -1 with metals
• Oxidation number of O is -2, apart from the element, in a peroxide or with fluorine
• Oxidation numbers in a neutral compound must add to 0
• Oxidation numbers in an ion must add to the charge of the ion
• The names of compounds or ions that contain an element that can have multiple oxidation numbers have the oxidation number in Roman numerals.
• In a half equation showing reduction, the electrons are on the left of the arrow.
• In a half equation showing oxidation, the electrons are on the right of the arrow.
• Half equations rules:

  • balance atoms
  • balance oxygen with H2O
  • balance hydrogen with H+
  • balance electrons
• Half equations can be combined to give the overall equation:

  • number of electrons in each half equation must be the same, so multiply one equation
  • combine reactants and products
  • remove electrons
  • cancel out any species that are on both sides of the arrow
• Halogens are weaker oxidising agents as the group is descended:

  • outer electron is in a higher energy level, further from the nucleus, with more shielding
  • despite the additional protons, the attraction between the nucleus and the electron being gained is weaker
  • harder for the halogen to gain electrons

  e.g. Cl2 + 2e- -> 2Cl-

  (Test for this is displacement reactions by adding a halogen to a halide)
• Halides are stronger reducing agents as the group is descended:

  • outer electron is in a higher energy level, further from the nucleus, with more shielding
  • despite the additional protons, the outer electrons are less attracted to the nucleus
  • halide ions lose electrons more easily

  e.g. 2Cl- -> Cl2 + 2e-

  (Test for this is halides ions with concentrated sulphuric acid)
• Concentrated sulphuric acid with F- and Cl-:

  • acid base reaction where H2SO4 is a proton donor, halide ion is a base
  • not a redox reaction
  • halide ion does not reduce H2SO4

  1. NaF + H2SO4 -> NaHSO4 + HF
  2. NaCl + H2SO4 -> NaHSO4 + HCl

  • HCl turns damp blue litmus paper red
  • forms white clouds with ammonia
• Concentrated sulphuric acid with Br-:

  • 2 reactions
  • acid base reaction where H2SO4 is a proton donor, and Br- is a base

  1. NaBr + H2SO4 -> NaHSO4 + HBr

  • redox reaction where bromide ions are oxidised to bromine, and H2SO4 is reduced to sulphur dioxide
  • Br- can reduce H2SO4, and reduces S from +6 to +4

  2. H2SO4 + 2H+ + 2Br- -> Br2 + SO2 + 2H2O

  • HBr forms steamy fumes
  • Br2 forms brown fumes
  • SO2 is a colourless choking gas
• Concentrated sulphuric acid with I-:

  • 2 reactions
  • acid base reaction where H2SO4 is a proton donor, and I- is a base

  1. NaI + H2SO4 -> NaHSO4 + HI

  • redox reaction where iodide ions are oxidised to iodine, and H2SO4 is reduced to sulphur dioxide, sulphur, and hydrogen sulphide
  • I- can reduce H2SO4, and reduces S from +6 to +4, 0, and -2

  2. 2I- -> I2 + 2e-

  • I2 forms purple fumes or a black solid

  3. H2SO4 + 2H+ + 2e- -> SO2 + 2H2O

  • SO2 is a colourless choking gas

  4. H2SO4 + 6H+ + 6e- -> S + 4H2O

  • S is a yellow solid

  5. H2SO4 + 8H+ + 8e- -> H2S + 4H2O

  • H2S has a bad egg smell
• Solubility of halogens in water decreases down the group. Halogens are more soluble in non-polar solvents e.g. cyclohexane.
• When halogens are added to water and cyclohexane, 2 layers of different colours form:

  • Cl2 -> pale green in cyclohexane, pale green in water
  • Br2 -> orange in cyclohexane, yellow-orange in water
  • I2 -> purple in cyclohexane, pale brown in water

  The colour of the water becomes paler when the cyclohexane is added/
• G1 with oxygen:

  4X(s) + O2(g) -> 2X2O
• G1 with chlorine:

  2X(s) + Cl2(g) -> 2XCl
• G1 with water:

  2X(s) + 2H2O -> 2XOH(aq) + H2(g)
• G2 with oxygen:

  2Y(s) + O2(g) -> 2YO(s)
• G2 with chlorine:

  Y(s) + Cl2(g) -> YCl2(s)
• Mg with steam:

  Mg(s) + H2O(g) -> MgO(s) + H2(g)
• G2 with water:

  Y(s) + 2H2O(l) -> Y(OH)2(aq) + H2(g)
• Chlorine with water:

  Cl2 + H2O -> HCl + HClO

  Cl is reduced to -1 in HCl, and oxidised to +1 in HClO
• Chlorine with cold alkali:

  Cl2 + 2NaOH -> NaCl + NaClO + H2O

  Cl is reduced to -1 in NaCl, and oxidised to +1 in NaClO
• Chlorine with hot alkali:

  3Cl2 + 6NaOH -> 5NaCl + NaClO3 + 3H2O

  Cl is reduced to -1 in NaCl, and oxidised to +5 in NaClO3