Group 7, the halogens

• Halogens

  • Have 7 electrons in the outermost shell
  • Are non-metals
  • Exist as diatomic molecules and the two atoms are covalently bonded
• Van der Waals' forces

  Weak intermolecular forces that hold the diatomic halogen molecules together
• Chlorine is a green/yellow gas at room temperature
• Bromine is a brown liquid at room temperature
• Iodine is a grey/black solid at room temperature
• Reaction of halogens with metallic elements

  Metallic elements lose electron and halogens gain electron

  Oxidation number of chlorine changes from 0 to -1
• Displacement reaction

  A more reactive halogen replaces a less reactive halogen from its halide solution

  Chlorine replaces bromine from sodium bromide solution

  Chlorine acts as an oxidising agent and the bromide ions are oxidised
• Extraction of halogens

  Chlorine can be used to extract bromine from seawater

  Iodine can be extracted from kelp (seaweed) by heating with manganese dioxide and concentrated sulphuric acid
• Reaction of bromide ions with sulphuric acid

  Produces hydrogen bromide gas in the first stage

  In the second stage, bromide ions act as a reducing agent, reducing the sulfuric acid to sulphur dioxide
• Reaction of iodide ions with sulphuric acid

  Produces hydrogen iodide gas in the first stage

  In the second and third stages, iodide ions act as a reducing agent, reducing the sulfuric acid to sulphur dioxide, sulphur and hydrogen sulphide
• Reaction of chlorine with water in sunlight

  2Cl2(g) + 2H2O(l) → 4HCl(aq) + O2(g)

  Chlorine is rapidly lost from swimming pools in direct sunlight
• Reaction of chlorine with cold sodium hydroxide

  Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)

  Sodium chlorate (I) is an oxidising agent and the active ingredient in household bleach
• Reaction of chlorine with hot sodium hydroxide

  3Cl2(aq) + 6NaOH(aq) → 5NaCl(aq) + NaClO3(aq) + 3H2O(l)

  Produces sodium chlorate (V)
• The halide ion test does not detect fluoride ions as silver fluoride is soluble and will not form a precipitate
• The nitric acid is added to remove any carbonate ions or hydroxide ions that would also form precipitates
• Trend in melting/boiling points in group 7

  • increases down the group
  • number of electrons in molecules increases down group = stronger Van der Waals forces
  • this increases chances of dipoles arising within a molecule which would induce dipoles in neighbouring molecules = strengthens intermolecular bond
• Chlorine has a low boiling point becuase the forces between the molecules are weak.

  Explain how these forces arise between molecules of chlorine. (3)

  • random movement of electrons in one molecule creates a dipole
  • induces a dipole in a neighbouring molecule
  • these temporary dipoles attract
• Trend in bond enthalpies in group 7

  • decrease going down the group
  • due to the increasing atomic radius, increased shielding and reduced attraction between the nucleus and bonding pair of electrons
  • exception to this is fluorine.
  • it has such a small atomic radius that bonding and non-bonding electrons repel each other, reducing the bond enthalpy.
• Trend in oxidising ability of halogens

  • oxidising ability decreases down group 7
  • atomic radius increases down group = weaker force of attraction between nucleus and outer shell of electrons
  • increased shielding effect down group
  • means it’s harder to gain an electron as oxidising agents are electron acceptors
• Fluorine is the strongest oxidising agent of group 7 because it has least number of electrons, so very small atomic radius so the attraction between nucleus and outer electrons is strong. This and the least shielding effect means it’s easier to accept an electron into its outermost shell.
• Trend in electronegativity in group 7

  • decreases going down the group
  • as the atomic radius increases, the shared pair of electrons in the covalent bond are further from the nucleus, reducing the strength of the attraction.
  • increasing atomic radius also means increasing shielding effect, again reducing the electronegativity.
• Trend in reducing ability of halide ions

  • Going down group 7, the halide ions become stronger reducing agents.
  • This is because the ionic radius increases so the outer electrons are less strongly attracted.
  • Further down the group electrons are more easily lost, making them better reducing agents.
• Chloride ions with sulphuric acid

  • Solid sodium chloride reacts with concentrated sulphuric acid. This reaction is performed in a fume cupboard due to the HCl fumes produced. Sodium hydrogensulfate is also formed.

  NaCl (s) + H2SO4 (l)→ NaHSO4 (s) + HCl (g)

  • This is not a redox reaction as no oxidation states change in the process.
  • The chloride ion is too weak a reducing agent to change the oxidation state of sulphur.
• Give an equation for the reaction of solid sodium bromide with concentrated sulfuric acid to form bromine.

  State one observation made during this reaction.

  • 2H2SO4 + 2NaBr → Na2SO4 + SO2 + Br2 + 2H2O
  • orange/brown fumes/solution
• Solid sodium iodide reacts with concentrated sulfuric acid to form iodide and sulfur in a redox reaction

  1. Give a half equation to show the conversion of iodide ions to iodine.
  2. Give a half-equation to show the conversion of sulfuric acid to sulfur.
  3. Give an overall equation for this redox reaction.
  4. Identify one other sulfur-containing reduction product formed when solid sodium iodide reacts with concentrated sulfuric acid.
  5. 2I- → I2 + 2e-
  6. H2SO4 + 6H+ + 6e- → S + 4H2O
  7. 6H+ + 6I- + H2SO4 → 3I2 + S + 4H2O
  8. SO2 or H2S
• State one observation when solid sodium chloride reacts with concentrated sulfuric acid.

  Give an equation for the reaction.

  State the role of chloride ions in the reaction.

  • Misty or steamy or white fumes/gas
  • NaCl + H2SO4 → NaHSO4 + HCI

  Base OR proton acceptor
• Give an equation for the redox reaction between the solid sodium bromide and concentrated sulfuric acid.

  Explain, using oxidation states, why this is a redox reaction.

  • 2NaBr + 2H2SO4 → Na2SO4 + Br2 + SO2 + 2H2O
  • Br changes oxidation state from -1 to 0 and is oxidised
  • S changes oxidation state from +6 to +4 and is reduced
• State what is observed when aqueous chlorine is added to sodium bromide solution.

  Give an ionic equation for the reaction.

  • Yellow or orange solution
  • Cl2 + 2Br- → 2Cl- + Br2
• Solid sodium chloride reacts with concentrated sulfuric acid.

  Give an equation for this role.

  State the role of sulfuric acid in this reaction.

  • NaCl + H2SO4 → NaHSO4 + HCl
  • Proton donor
• Fumes of sulfur dioxide are formed when sodium bromide reacts with concentrated sulfuric acid.

  For this reaction

  • give an equation
  • give one other observation
  • state the role of the sulfuric acid.
  • 2NaBr + 2H2SO4 → Na2SO4 + SO2 + Br2 + 2H2O
  • brown gas/fumes or orange gas/fumes
  • Oxidising agent
• Reaction of iodide ions with sulphuric acid

  1. NaI (s) + H2SO4 (l)→ NaHSO4 (s) + HI(g)
  2. 2H+ + 2I- + H2SO4 (l)→ I2(g) + SO2(g) + 2H2O (l)
  3. 6H+ + 6I- + H2SO4 (l)→ 3I2(g) + S(s) + 4H2O (l)

  Overall equation for steps 2 and 3:        

  8H+ + 8I- + H2SO4 (l)→ 4I2(g) + H2S(s) + 4H2O (l)

  • Fumes of HI are white in colour.
  • Iodine fumes are purple in colour.
  • Hydrogen sulphide has a rotten egg smell.
• Chlorine undergoes a redox reaction called as disproportionation, where chlorine is both oxidised and reduced.


  • Reaction of chlorine with water for water treatment
  • Chlorine is used in water treatment to kill bacteria and make water safe for drinking.
  • Chlorine reacts with water to form hydrochloric acid and chloric (I) acid.

  Cl2 (aq) + H2O (l) → HCl (aq) + HClO (aq)
• Chloric (I) acid decomposes slowly in water to form reactive oxygen molecule that kills the bacteria in water.

  HClO → HCl +[O]
• Chlorine may react with organic compounds in water to form products that could be hazardous to health. It is important to use sodium hydroxide to remove chlorine in water before using it in aquariums as chlorine could be harmful for aquatic organisms.
• An alternative to direct chlorination is to add solid sodium or calcium chlorate (I) which will dissolve in water to form the chlorate ions.

  NaClO (s) + H2O (l) ↔ Na+ (aq) + OH- (aq) + HClO (aq)

  • In order to shift the equilibrium to the right and maintain the concentration of HClO, the water must be slightly acidic, but not so much as to endanger people.
  • Under alkaline conditions, the equilibrium will shift to the left.
• Reaction of chlorine with sodium hydroxide

  Chlorine reacts with cold sodium hydroxide to form sodium chloride and sodium chlorate (I) in a redox reaction

  Cl2 (aq) + 2NaOH (aq) → NaCl (aq) + NaClO (aq) + H2O (l)

  • Sodium chlorate (I) is an oxidising agent, and is the active ingredient in household bleach and kills bacteria.
• Give an equation for the reaction of chlorine with water.

  Give a reason why chlorine is added to drinking water.

  1. Cl2 + H2O ↔ HCI + HClO
  2. Kills bacteria
• Chlorine reacts with cold, aqueous sodium hydroxide in the manufacture of bleach.

  Give an equation for this reaction.

  • Cl2 + 2NaOH → NaCl + NaCIO + H2O
• Chlorine is used to treat water even though it is toxic to humans.

  Give one reason why water is treated with chlorine.

  Explain why chlorine is added to water even though it is toxic.

  Give an equation for the reaction of chlorine with cold water.

  1. Reason: kill bacteria
  2. Explanation: health benefit outweighs risk
  3. Equation: Cl2 + H2O ↔ HCl + HClO
• Chlorine reacts with hot aqueous sodium hydroxide as shown in the equation.

  3Cl + 6NaOH → NaClO3 + 5NaCl + 3H2O

  Give the oxidation state of chlorine in NaClO3 and in NaCl

  State, in terms of redox, what happens to chlorine in the reaction.

  1. NaCIO3 : +5 and NaCl: -1
  2. Chlorine is oxidised and reduced