Group 7 - Halogens

• Physical properties of the halogens

  • Fluorine is a highly reactive pale yellow gas
  • Chlorine is a pale green gas, toxic in high concentrations
  • Bromine is a red liquid that gives off dense brown/orange poisonous fumes
  • Iodine is a shiny grey solid that sublimes to form a purple gas
• As you go down the group

  The melting and boiling point increases
• Reason for trend in melting and boiling point

  • The halogens are covalent simple molecular substances
  • There are weak Van der Waal's forces between the molecules which increase as molecular size increases
  • Van der Waal's forces are greater for larger molecules because they contain more electrons making them more polarisable
  • More energy is needed to melt and boil the halogens as you go down the group

  
• As you go down the group

  The electronegativity decreases
• Reason for trend in electronegativity

  • The atomic radius increases as you go down the group because there are more full shells of electrons
  • There are more full shells shielding the bonding pair of electrons from the attractive force of the nucleus
  • The force of attraction from the nucleus for the bonding pair of electrons decreases as you go down the group, so electronegativity decreases

  
• Oxidising ability of the halogens

  Compares their relative tendency to accept electrons
• Oxidising ability of the halogens

  • Fluorine is the most powerful oxidising agent
  • Iodine is the least powerful oxidising agent
• Halogens acting as oxidising agents


  1. F2 + 2e- → 2F-

  2. Cl2 + 2e- → 2Cl-

  3. Br2 + 2e- → 2Br-
• Halogen displacement reactions

  • Chlorine will displace bromide and iodide ions from solution
  • Bromine will displace iodide ions from solution
• Halogen displacement reactions


  • Cl2(aq) + 2Br-(aq) → Br2(aq) + 2Cl-(aq)
  • Cl2(aq) + 2I-(aq) → I2(aq) + 2Cl-(aq)
  • Br2(aq) + 2I-(aq) → I2(aq) + 2Br-(aq)
• Observations of halogen displacement reactions

  • Formation of pale-yellow solution showing chlorine has displaced bromide
  • Formation of a brown solution showing chlorine has displaced iodide
  • Formation of a brown solution showing bromine has displaced iodide
• Distinguishing bromine and iodine in solution


  Shaking with cyclohexane causes an orange layer of bromine or purple iodine as the halogen is more soluble in this non-polar solvent
• Reducing ability of the halide ions

  Compares their relative tendency to donate electrons
• As you go down the group


  The halide ions become increasingly powerful reducing agents
• Reason for trend in reducing ability of halide ions

  • The ions get bigger as you go down the group
  • There are more full shells shielding the outmost electrons from the attractive force of the nucleus
• Iodide ions are stronger reducing agents than chloride ions
• Fluoride F- and Chloride Cl-

  These are the weakest reducing agents and are not powerful enough to reduce sulfur from its +6 oxidation state
• Excess silver nitrate


  Used to ensure all chloride and iodide ions are precipitated
• Half equations for oxidation

  • 2Br- → Br2 + 2e-
  • 2I- → I2 + 2e-
• Half equations for reduction

  • H2SO4 + 2H+ + 2e- → SO2 + 2H2O
  • H2SO4 + 6H+ + 6e- → S + 4H2O
  • H2SO4 + 8H+ + 8e- → H2S + 4H2O
• Reaction of chlorine with water

  Cl2(g) + H2O(l) ⇌ HClO(aq) + HCl(aq)

  
  

  This is a disproportionation reaction because in Chlorine goes from an oxidation state of 0 in Cl2 to -1 in HCl and +1 in HClO.
• A disproportionation reaction is one where the same species is simultaneously oxidised and reduced
• Chlorine is bubbled through water in bright sunlight

  2Cl2(g) + 2H2O(l) → 4H+(aq) + 4Cl-(aq) + O2(g)
• Reason for treating water with chlorine

  To kill bacteria
• Despite its toxicity, the benefits of using chlorine outweigh any negative effects
• Reaction of Cl2, Br2 and I2 with cold dilute NaOH

  Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)
• Chlorine is oxidised in the reaction with hot aqueous sodium hydroxide
• Testing for halide ions


  1. Add HNO3 to the solution

  2. Add AgNO3

  3. Observe for precipitate formation

  4. Cl- gives white, Br- gives cream, I- gives yellow
• Telling AgCl and AgBr apart

  1. AgCl redissolves in dilute ammonia

  2. AgBr redissolves in concentrated ammonia

  3. AgI does not redissolve
• Testing for sulfate ions

  1. Add HCl to the solution

  2. Add BaCl2

  3. Observe for white precipitate formation

  
  

  Ba2+(aq) + SO42-(aq) -> BaSO4(s)

  
  

  It is important to first acidify the solution we want to test with hydrochloric acid to remove any carbonate ions which would also form a white precipitate of barium carbonate and would give a false positive test masking the desired observations.
• Testing for carbonate ions

  1. Add dilute acid to the solid or solution

  2. Observe for fizz/bubbles due to CO2 production

  
  

  • 2H+ + CO32- -> CO2 + H2O
  • H+ + HCO3- -> CO2 + H2O
• Testing for hydroxide ions


  1. Forms insoluble precipitates with metal ions

  2. Fe2+(green precipitate), Fe3+ (brown precipitate) and Cu2+ (blue precipitate).

  3. Turns litmus/universal indicator blue/purple
• Testing to show solution C contains ammonium sulfate


  1. Add NaOH, observe for ammonia gas

  2. Add BaCl2, observe for white precipitate

  3. Confirm presence of sulfate and ammonium ions
• Identifying halide ions in solution D

  1. Add excess AgNO3, filter and weigh precipitate

  2. Add excess dilute ammonia, filter and weigh remaining solid

  3. Use masses to determine mixture of halide ions
• Because the reaction is reversible, at equilibrium all three species are present, Cl2, HCl and HClO. Chlorine water is pale green owing to the presence of Cl2. A solution of chlorine water will first turn universal indicator red due to the acidity of HClO and HCl, it will then turn colourless as HClO bleaches the colour.
• NaCl and NaClO


  • The products NaCl and NaClO are used as bleach and to disinfect/kill bacteria.
  • NaClO is named sodium chlorate (I)
  • NaClO3 a species formed when chlorine reacts with hot concentrated NaOH is name sodium chlorate (V).
  • The roman numeral to show the oxidation state is an important addition to the IUPAC name.
• Halide Precipitation Reactions 


  • Ag+(aq) + Cl-(aq) -> AgCl(s)
  • Ag+(aq) + Br-(aq) -> AgBr(s)
  • Ag+(aq) + I-(aq) -> AgI(s)
• Exam questions might ask...

  • It is important to first acidify the solution toil want to test with nitric acid to remove any carbonate ions which would also form a white precipitate of silver carbonate and would give a false positive test masking the desired observations.
  • We must acidify with nitric and not hydrochloric acid, this is because the Cl- from the fully dissociated HCl in solution would also give a false positive result.
• Reactions of F- and Cl-

  • These are the weakest reducing agents and are not powerful enough to reduce sulfur from its +6 oxidation state.
  • Only acid base reactions occur to produce acid gases which turn moist blue litmus paper red.
  • NaF(s) + H2SO4(l) -> NaHSO4(s) + HF(g)
  • Observations: White steamy fumes of HF are formed and moist blue litmus turns red.
  • NaCl(s) + H2SO4(l) -> NaHSO4(s) + HCl(g)
  • Observations: White steamy fumes of HCl are formed and moist blue litmus turns red.
• Reactions of Br-


  • Reaction 1 – acid-base
  • NaBr(s) + H2SO4(l) -> NaHSO4(s) + HBr(g)
  • Observations: misty fumes of HBr that turn damp blue litmus paper red

  
  

  • Reaction 2 – reduction of sulfur from +6 in H2SO4 to +4 in SO2
  • 2NaBr(s) + 3H2SO4(l) -> 2NaHSO4(s) + SO2(g) + Br2(g)+ 2H2O(l)
  • Observations: orange fumes of bromine and a gas that smells of rotten eggs and turns acidified dichromate paper from orange to green.

  
  

  • Half equations for the redox reaction:
  • 2Br- -> Br2 + 2e-
  • 2H+ + H2SO4 + 2e- -> SO2 + 2H2O