The halogens

• The halogens are the most reactive non-metallic group. The elements do not occur in their elemental form in nature, they occur as stable halide ion dissolved in sea water or combines with sodium or potassium as solid deposits, such as in salt mines containing common salt, NaCl
• At RTP, all the halogens exist as diatomic molecules, X2. The group contains elements in all three physical states at RTP, changing from gas, to liquid, to solid down the group.
• In their solid states the halogens form lattices with simple molecular structures
• Trend in boiling point down the group:

  1. more electrons
  2. stronger London forces
  3. more energy required to break the IMF
  4. boiling point inccreases
• REDOX reaction: each halogen has 7 outer-shell electron, one short of the electronic configuration of a noble gas. Two electrons are in the outer s sub-shell and 5 in the outer p sub-shell
• REDOX reactions are the most common reactions of the halogens. Each halogen atom is reduced, gaining one electron to form a 1- halide ion

  Cl2 + 2e- = 2Cl- - chlorine is reduced
• Another species loses electrons to halogen atoms - it is oxidised. The halogen is called an oxidising agent, because it has oxidised another species
• Halogen-Halide Displacement reaction: displacement reactions of halogens with halide ions can be carried out on a test-tube scale. The results of the displacement reactions show that reactivity decreases down the group
• A solution of each halogen is added to aqueous solutions of the other halides, e.g. solution of Cl2 is added to 2 aqueous solutions containing Br- and I- ions.
• If the halogen added is more reactive than the halide present:

  1. a reaction takes place - the halogen displacing the halide from solution
  2. the solution changes colour
• Solutions of iodine and bromine in water can appear a similar brown colour, depending on the concentration - to tell them apart an organic non-polar solvent such as cyclohexane can be added and the mixture shaken
• Non-polar halogens dissolve more readily in cyclohexane than i water. In cyclohexane their colour are much easier to tell apart, with iodine being a deep violet
• Solutions in water:

  Cl2 - pale green

  Br2 - orange

  I2 - brown
• Solution in cyclohexane (top layer):

  Cl2 - pale green

  Br2 - orange

  I2 - violet
• orange colour from Br2 formation:

  Cl2 (aq) + 2Br- (aq) = 2Cl- (aq) + Br2 (aq)
• Violet colour from I2 formation:

  Cl2 (aq) + 2I- (aq) = 2Cl- (aq) + I2 (aq)
• Violet colour from I2 formation:

  Br2 (aq) + 2I- (aq) = 2Br- (aq) + I2 (aq)
• Chlorine reacts with both Br- and I-

  Bromine reacts with I- only

  Iodine does not react at all
• Reaction of chlorine with bromide ions is a redox reaction:

  full equation: Cl2 (aq) + 2NaBr (aq) = 2NaCl (aq) + Br2 (aq)

  ionic equation: Cl2 (aq) + 2Br- (aq) = 2Cl- (aq) + Br2 (aq)

  2 Br in 2Br- = each increases by +1 (total increase +2) - oxidation

  2 Cl in Cl2 = each Cl decreases by -1 (total decrease -2) - reduction
• Fluorine is a pale yellow gas and reacts with almost an substance it comes in contact with
• Astatine is extremely rare because it is radioactive and decays rapidly and the elements has never actually been seen - predicted to be the least reactive halogen
• In redox reactions, halogen react by gaining electrons. Down the group, the tendency to gain an electron decreases and the halogens become less reactive
• Down the group:

  1. atomic radius increases
  2. more inner shells so shielding increases
  3. less nuclear attraction to capture an electron from another species
  4. reactivity decreases
• Fluorine is the strongest oxidising agent, gaining electrons from other species more readily than the others. The halogens become weaker oxidising agents down the group
• Disproportionation: a redox reaction in which the same element is both oxidised and reduced. The reaction of chlorine with water and with cold, dilute sodium hydroxide are examples of this
• Cl used in water purification. Began to be widely used as a disinfectant for drinking water treatment over 100 years ago, revolutionised public health by reducing the incidence of waterborne diseases by killing harmful bacteria
• When small amounts of chlorine are added to water, a disproportionation reaction takes place. For each chlorine molecule, one atom is oxidised and the other is reduced
• Cl2 (aq) + H2O (L) = HClO (aq) + HCl (aq)

  0 = -1 reduction

  0 = +1 oxidation
• The two products are both acids (chloric (I) acid and HCl. The bacteria are killed by chloric (I) acid and chlorate (I) ions ClO- rather than by chlorine. Chloric (I) acid alist acts as a weak bleach - can demonstrate this by adding some indicator solution to a solution of chlorine in water. The indicator first turns red, from the presence of the 2 acids. The colour then disappears as the bleaching action of chloric (I) acid takes effect
• The reaction of chlorine with water is limited by the low solubility of chlorine in water.
• If the water contains dissolved sodium hydroxide, much more chlorine dissolves and another disproportionation reaction takes place
• The reaction of chlorine with cold, dilute aqueous sodium hydroxide:

  Cl2 (aq) + 2NaOh (aq) = NaClO (aq) + NaCl (aq) + H2O (l)

  0 = -1 reduction

  0 = +1 oxidation
• The resulting solution contains a large concentration of chlorate (I), ClO- ions from the sodium chlorate (I), NaClO, that is formed. This solution finds a use as household bleach, made by reacting chlorine with cold dilute aqueous sodium hydroxide
• Although chlorine is beneficial in ensuring that our water is fit to drink and that bacteria are killed, it is also an extremely toxic gas. Chlorine is a respiratory irritant in small concentrations, and large concentrations can be fatal