Halogens: Group 7 Trends and Redox Reactions

Group 7(17), The Halogens

Trends in Properties

• Group 7 elements are known as the halogens, they include Fluorine, Chlorine, Bromine, Iodine and Astatine.
• Electronegativity is how well an atom attracts electrons to itself in a covalent bond.
• Electronegativity decreases going down Group 7 because:
  • atomic radii increases
  • more shielding by inner shell electrons
  • incoming electrons are further from the positive nucleus
  • increased shielding outweighs the increased nuclear charge
• Going down Group 7, the state of the halogens at room temperature changes.
  • Fluorine and chlorine are gases
  • Bromine is a liquid
  • Iodine and astatine are solids
• The halogens get darker in colour going down the group
• Going down Group 7, the boiling points increase because:
  • atoms have more electrons
  • stronger van der Waals forces between the molecules
  • more energy is needed to overcome these forces

Group 7(17) Redox Reactions

• Halogens react with metals by accepting an electron from the metal atom and becoming an ion with a 1- charge.
  • E.g. Ca (s) + Cl2 (g) → CaCl2 (s)
  • Calcium loses electrons, Chlorine gains electrons
• Halogens are oxidising agents that cause another substance to lose electrons because they are reduced themselves as they gain electrons.
• The oxidising ability of the halogens decreases going down the group.
• Halogen displacement reactions involve a halogen displacing a less reactive halide in a metal halide.
  • E.g. Cl2 (aq) + 2NaBr (aq) → 2NaCl (aq) + Br2 (aq)
  • Here, chlorine has displaced the bromide because it is more reactive
  • Ionic equation: Cl2 (aq) + 2Br¯ (aq) → 2Cl-(aq) + Br2 (aq)
• In this example, bromine has been oxidised and chlorine has been reduced, with chlorine acting as an oxidising agent.
• Halide ions can act as reducing agents to donate electrons to another atom, getting oxidised themselves in the process.
• The reducing ability of the halide ions increases going down the group as the ions become larger.
• Outer electrons get further from the nucleus and experience more shielding, being held less tightly and more easily lost.

Sodium halides with concentrated sulfuric acid

• The trend in reducing ability of the halides can be demonstrated by reacting solid sodium halides with concentrated sulfuric acid
• The ionic equation for these reactions is: H2SO4(I) + X¯(aq) →HX(g) + HSO4 (aq)
• NaCl with conc. Sulfuric acid reaction takes place as: H2SO4 (I) + NaCl (s) → HCl (g) + NaHSO4 (S)
  • HCl gas produces steamy fumes, with solid sodium hydrogensulfate is also observed
  • No redox reaction takes place with chlorine as it is a weak reducing agent; the oxidation state of each substance remains the same
• NaBr with conc. sulfuric acid initially: H2SO4 (I) + NaBr (s) → HBr (g) + NaHSO4 (S)
  • HBr is produced as steamy fumes, solid sodium hydrogensulfate is observed
  • Bromide ions then reduce sulfuric acid to sulfur dioxide gas: 2HBr (g) + H2SO4 (I) → Br2 (g) + SO2 (g) + 2H2O (1)
  • Brown fumes of bromine are observed with colourless sulfur dioxide
  • Sulfur is reduced, with the oxidation state decreasing, and bromine is oxidised so the oxidation state increases
  • A redox reaction occurred demonstrating bromide ions are stronger reducing agents than chloride ions
• NaI with conc. sulfuric acid initially: H2SO4 (I) + Nal (s) → HI (g) + NaHSO4 (s)
  • HI gas is produced as steamy fumes, solid sodium hydrogensulfate is also observed
  • Iodide ions are much stronger reducing agents reducing sulfuric acid to sulfur dioxide, then to sulfur, and finally to hydrogen sulfide
  • Formation of hydrogen sulfide: 8HI (g) + H2SO4 (l) → 4I2 (s) + H2S (g) + 4H2O (1)
    • Iodine is seen as a black solid, hydrogen sulfide has a strong smell of bad eggs, and sulfur is a yellow solid

Testing for Halide Ions

• Halide ions can be identified in an unknown solution by dissolving the solution in nitric acid and then adding silver nitrate solution dropwise.
• Nitric acid is used to prevent false positives and silver nitrate will react as: Ag+ (aq) + X-(aq) → AgX (s)
• A silver halide forms upon addition of silver nitrate solution, its colour depends on the halide ion present:
  • Silver chloride (AgCl) is a white precipitate
  • Silver bromide (AgBr) is a cream precipitate
  • Silver iodide (Agl) is a yellow precipitate
• Silver fluoride is soluble
• White, cream and yellow precipitates look very similar so ammonia is sometimes added to confirm the identity of the halide ion
• Dilute followed by concentrated ammonia is added to the silver halide solution:
  • If the precipitate dissolves in dilute ammonia, the unknown halide is chloride
  • If the precipitate does not dissolve in dilute but does dissolve in concentrated ammonia, the unknown halide is bromide
  • If the precipitate does not dissolve in dilute or concentrated ammonia, the unknown halide is iodide

Uses of Chlorine & Chlorate(I)

• A disproportionation reaction is a reaction in which the same species is both oxidised and reduced, e.g. chlorine with water or cold, dilute alkali.
• When chlorine reacts with water, both chloric (I) acid, and hydrochloric acid are formed:
  • Cl2(aq) + H2O (1) → HCl(aq) + HCLO(aq)
  • Chloric (1) acid, HCIO, sterilises water by killing bacteria
  • Chlorine is oxidised from 0 to +1, and reduced from 0 to -1 making it a disproportionation reaction
• In sunlight, the following reaction occurs: 2Cl2 (g) + 2H2O (l) → 4HCl (aq) + O2 (g); the chlorine has to be topped up on a regular basis increasing costs
• An alternative to directly adding chlorine, is to add chlorate ions, CIO- in the form of sodium or calcium chlorate(I):
  • NaClO (aq) + H2O (l) → Na+ (aq) + OH- (aq) + HCIO-(aq)
  • CIO-(aq) acts as a sterilising agent cleaning the water
• Advantages and disadvantages are assessed as chlorine is a toxic substance; levels of chlorine in water need to be carefully monitored

Simple test tube reactions

• Required Practical 3 identifying anions and cations; 0.1 mol dm-3 solutions are used to investigate Group 2 cations, ammonium ions, Group 7 halide ions, hydroxide ions, carbonate ions and sulfate ions

• To test for Group 2 metal cations using NaOH:

  1. Place 10 drops of 0.1 mol dm-3 barium chloride solution in a clean test tube
  2. Add 10 drops of 0.6 mol dm-3 sodium hydroxide solution, mixing well.
  3. Continue to add sodium hydroxide solution, dropwise until in excess

• To test for Group 2 metal cations using dilute sulfuric acid:

  1. Place 10 drops of 0.1 mol dm-3 barium chloride solution in a clean test tube
  2. Add 10 drops of 1.0 mol dm-3 sulfuric acid, mixing well
  3. Continue to add sulfuric acid, dropwise until in excess

• To test for ammonium ions:

  1. Place 10 drops of ammonium chloride in a clean test tube
  2. Add 10 drops of 0.4 mol dm-³ sodium hydroxide solution and shake the mixture
  3. Warm the mixture in the test tube gently using a water bath
  4. Test the fumes released from the mixture with litmus paper

• To test for hydroxide ions:

  1. Add 1 cm³ of 0.4 mol dm-3 sodium hydroxide solution to a test tube with red litmus paper/universal indicator
  2. Test for alkaline gas using ammonia solution, which forms hydroxide ions when it comes into contact with water
     1. Take 5 drops of 1.0 mol dm-3 ammonia solution and place on a filter paper inside a petri dish with lid
     2. Dampen a piece of red litmus paper with deionised water and place on the other side of the petri dish; replace the lid and observe for a few minutes

• To test for carbonate ions:

  1. Add 2 cm³ of calcium hydroxide (limewater) to a test tube
  2. Add 3 cm³ of 0.5 mol dm-3 sodium carbonate solution to a separate test tube and add an equal volume of dilute hydrochloric acid
  3. Immediately connect a delivery tube with the open end into the limewater test tube

• To test for sulfate ions:

  1. Add 1 cm³ of 0.1 mol dm³ magnesium sulfate solution into a test tube
  2. Add an equal volume of dilute hydrochloric acid
  3. Add an equal volume of 0.1 mol dm¯³ barium chloride solution

• To test for halide solutions:

  1. Add 10 drops of 0.1 mol dm-3 potassium chloride in a clean test tube
  2. Add 5 drops of dilute nitric acid and shake
  3. Add 10 drops of 0.05 mol dm-3 silver nitrate solution to the resulting mixture
  4. Add an excess of 2.0 mol dm-3 ammonia solution and shake to mix thoroughly
  5. Repeat steps with concentrated ammonia
  6. Repeat steps replacing potassium chloride with bromide and then potassium iodide

• To test for solid halide solutions:

  1. Place a small spatula of solid potassium chloride in a clean dry test tube
  2. Slowly add a few drops of concentrated sulfuric acid
  3. Test the gas evolved with moist blue litmus paper
  4. Repeat with solid potassium bromide, but test the gas produced using a narrow strip of filter paper that has been dipped in acidified potassium dichromate solution
  5. Repeat with potassium iodide, but test the gas produced using a narrow strip of filter paper that has been dipped in lead nitrate solution

• Sulfuric acid can be used to test for Group 2 metal ions

  • Ba2+= white precipitate formed which does not dissolve in excess
  • Ca2+= slight white precipitate which does not dissolve in excess
  • Mg2+= slight white precipitate will redissolve to give a colourless solution in excess
  • Sr2+= white precipitate which does not dissolve in excess

• Silver nitrate solution is used to test for the presence of halide ions:

  • Chloride ions = white precipitate
  • Bromide ions = cream precipitate
  • Iodide ions = yellow precipitate