group 7

• fluorine (F2) is a highly reactive pale yellow gas
• chlorine (Cl2) is a greeinish, reactive gas that is poisonous in high concentrations
• bromine (Br2) is a red liquid that gives off dense brown/orange poisonous fumes
• iodine (I2) is a shiny grey solid that sublimes to purple gas
• melting and boiling point increase down group 7

  molecules become larger and have more electrons so the van der waal forces between molecules get stronger and require more energy to overcome
• electronegativity is the relative tendency of an atom in a molecule to attract electrons in a covalent bond to itself
• electronegativity decreases down group 7 as the atomic radii increases due to the increasing number of shells

  nucleus is therefore less able to attract the bonding pair of electrons
• a halogen that is a strong oxidising agent will displace a halogen that has a lower oxidising power from one of its compounds
• oxidising strength decreases down group 7
• oxidising agents are electron acceptors
• chlorine will displace bromide and iodide ions
• bromine will displace iodide ions
• displacement reaction of potassium chloride and chlorine

  very pale green solution, no reaction
• displacement reaction of chlorine and potassium bromide

  yellow solution, Cl has displaced Br

  Cl2 (aq) + 2Br- (aq) --> 2Cl- (aq) + Br2 (aq)
• displacement reaction of chlorine and potassium iodide

  brown solution, Cl has displaced I

  Cl2 (aq) + 2I- (aq) --> 2Cl- (aq) + I2 (aq)
• displacement reaction bromine and potassium chloride

  yellow solution, no reaction
• displacement reaction bromine with potassium bromide

  yellow solution, no reaction
• displacement reaction of bromine with potassium iodide

  brown solution, Br has displaced I

  Br2 (aq) + 2I- --> 2Br- (aq) + I2 (aq)
• displacement reaction iodine and potassium chloride

  brown solution, no reaction
• displacement reaction iodine and potassium bromide

  brown solution, no reaction
• displacement reaction iodine and potassium iodide

  brown solution, no reaction
• test halide ions with silver nitrate acidified with nitric acid

  nitric acid reacts with any carbonates present to prevent Ag2CO3 precipitate forming that would mask observations
• reaction of halide ions with silver nitrate

  fluoride produces no precipitate

  chlorides produce a white precipitate

  bromides produce a cream precipitate

  iodides produce a pale yellow precipitate

  Ag+ (aq) + X- (aq) --> AgX (s)
• silver halide precipitates can be treated with ammonia to differentiate between similar colours
• Silver chloride dissolves in dilute ammonia to form a complex ion

  AgCl(s) + 2NH3 (aq) --> [Ag(NH3 )2 ]+ (aq) + Cl - (aq) Colourless solution
• Silver bromide dissolves in concentrated ammonia to form a complex ion

  AgBr(s) + 2NH3 (aq) --> [Ag(NH3 )2 ]+ (aq) + Br - (aq) Colourless solution
• Silver iodide does not react with ammonia – it is too insoluble
• reducing agent donates electrons

  halides react more vigorously as reducing agents going down the group as tendency to donate electrons increases when ions get bigger and nuclear attraction decreases

  we see this when reacting halides with concentrated sulfuric acid
• F- is not a strong enough reducing agent to reduce the S in H2SO4, so acid-base reaction occurs instead of redox

  NaF(s) + H2SO4 (l) --> NaHSO4 (s) + HF(g)

  white steamy fumes of HF
• Cl- ions are not strong enough reducing agents to reduce the S in H2SO4, so acid-base reactions occur instead of redox

  NaCl(s) + H2SO4 (l) --> NaHSO4 (s) + HCl (g)

  white steamy fumes of HCl observed
• Bromide ions reduce sulfur in H2SO4 from +6 to +4 in SO2

  In the 1st step H2SO4 acts as a acid and in the 2nd acts as an oxidising agent

  observe white steamy fumes of HBr, orange fumes of bromine and colourless acidic SO2 gas
• equations for bromine with sulfuric acid

  Acid- base step: NaBr(s) + H2SO4 (l) --> NaHSO4 (s) + HBr(g)

  Redox step: 2 H+ + 2 Br - + H2SO4 --> Br2 (g) + SO2 (g) + 2 H2O(l)

  Ox ½ equation 2Br - --> Br2 + 2e-

  Re ½ equation H2SO4 + 2 H+ + 2 e- --> SO2 + 2 H2O
• i- ions are the strongest halide reducing agents, reduce sulfur from +6 in H2SO4 to +4 in SO2, 0 in S and -2 in H2S

  H2SO4 acts as an acid in the first step and oxidising agent in the remaining steps

  observe white steam fumes of HI, black solid and purple fumes of iodine, colourless acidic gas SO2, yellow solid sulfur, H2S foul egg smelling gas
• equations for reactions of iodine with sulfuric acid

  NaI(s) + H2SO4 (l) -->NaHSO4 (s) + HI(g)

  2 H+ + 2 I- + H2SO4 --> I2 (s) + SO2 (g) + 2 H2O(l)

  6 H+ + 6 I- + H2SO4 --> 3 I2 + S (s) + 4 H2O (l)

  8 H+ + 8 I- + H2SO4 --> 4 I2 (s) + H2S(g) + 4 H2O(l)
• half equations for iodine and sulfuric acid

  Ox ½ equation 2I - --> I2 + 2e-

  Re ½ equation H2SO4 + 2 H+ + 2 e- --> SO2 + 2 H2O

  Re ½ equation H2SO4 + 6 H+ + 6 e- --> S + 4 H2O

  Re ½ equation H2SO4 + 8 H+ + 8 e- --> H2S + 4 H2O
• disproportionation is the name for a reaction where an element simultaneously oxidises and reduces
• chlorine reaction with water

  Cl2 (g) + H2O (l) ⇌ HClO (aq) + HCl (aq)
• equation when chlorine is bubbled through water WITH the presence of sunlight

  2Cl2 + 2H2O --> 4H+ + 4Cl - + O2

  green colour of chlorine water fades as Cl2 reacts and colourless O2 is produced
• chlorine reduces and oxidises simultaneously

  if universal indicator is added it will turn red due to acidity of both reaction products and then colourless as HClO bleaches colour
• chlorine is used in water treatments to kill bacteria

  treats drinking water and water in swimming pool