AQA Chemistry Unit 2: 11 Group 7, Halogens

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Basic Information

  • Halogen has 7 electrons in its highest energy level, reacts to gain electron and causes the loss of electron so is a oxidising agent
  • Halide has full highest energy level, reacts to lose electron and causes the gain of electron so is a reducing agent

The Halogens

  • Down the group atomic radius increases, boiling point increases due to bigger van der Waals due to more electrons meaning element has larger temporary dipoles and the electronegativity decreases as the larger the atom so the more electrons in energy levels meaning more shielding
  • Chemical Trend: Halogens have high oxidising power so cause oxidation. They themselves are reduced so gain an electron to have a full energy level
  • Down the group oxidation power decreases since more shielding from the nucleus and there is more distance from the nucleus
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Chemical Reactions of the Halogens

  • The oxidising ability of the halogens increases up the group
  • F2>Cl2>Br2>I2

Displacement Reactions

  • Halogens react with metal halides to displace them (If the halogen is more reactive)
  • The halides are colourless solutions
  • Fluorine oxidises Chloride, Bromide and Iodide (Cl2 is faint yellow, Br2 is orange, I2 is red)
  • Chlorine oxidises Bromide and Iodide (Br2 is yellow, I2 is orange/brown)
  • Bromine oxidises Iodide (red/brown)
  • Fluorine cannot be tested in an aqueous solution as it reacts with water
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Reactions of Halide Ions

  • Halides can lose their extra electrons
  • They act as reducing agents
  • Reducing power increases down the group as:
  • Larger atomic radius
  • Less attraction between nucleus and highest energy level
  • Due to increased distance there is more shielding

Sodium Chloride and Concentrated Sulfuric Acid

  • Chloride: White steamy fumes of hydrogen chloride, solid sodium hydrogensulfate formed
  • Isn't a redox reaction as there is no oxidation state change it is a acid-base reaction
  • The chloride ion is too weak a reducing agent to reduce the sulfur
  • NaCl(s) + H2SO4(l) -> NaHSO4(s) + HCl(g)
  • Similar reaction with sodium fluorine as the fluorine ion is an even weaker reducing agent
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Reactions of Halide Ions Cont

Sodium Bromide and Concentrated Sulfuric Acid

  • Steamy white fumes of hydrogen bromide formed
  • Brown fumes of bromine and colourless sulfur dioxide formed
  • Two reactions occur:
    • Sodium hydrogensulfate and hydrogen bromide are produced in a acid base reaction
    • NaBr(s) + H2SO4(l) -> NaHSO4(s) +HBr(g)
    • Bromide ions are strong enough reducing agents to reduce sulfuric acid to SO2
    • 2H+ + 2Br- + H2SO4(l) -> SO2(g) + 2H2O(l) +Br2(l)
  • This is a redox reaction, the reactions are exothermic.
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Reactions of Halide Ions Cont

Sodium Iodide and Sulfuric Acid

  • White steamy fumes of hydrogen iodide
  • Black solid of iodine, bad egg smell of hydrogen sulfide
  • Yellow solid sulfur maybe seen
  • Hydrogen Iodide produced in an acid-base reaction:
  • NaI(s) + H2SO4 -> NaHSO4(s) + HI(g)
  • As Iodide is a better reducing agent it can reduce the sulfur further:
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Metal Halides with Silver Ions

All metal halides react with silver ions in aqueous solution

Silver fluoride doesn't form a precipitate as it is soluble in water

  • Dilute nitric acid is added to the halide solution to remove any soluble carbonate or hydroxide impurities that would interfere with the test as they would form an insoluble silver carbonate or silver hydroxide
  • A few drops of sliver nitrate solution are added and the halide precipitate forms
  • The reaction can test for halides as they form different coloured precipitates
  • As silver bromide and silver iodide have similar colours concentrated ammonia solution is added; silver bromine dissolves and silver iodide doesn't

Test for Halides

  • Silver fluoride forms no precipitate, chloride forms white precipitate, bromide forms cream precipitate and iodide a pale yellow precipitate
  • Chloride dissolves in dilute ammonia, bromide in concentrated ammonia and iodide is insoluble in ammonia
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Uses of Chlorine

  • Uses: Kill bacteria e.g. in drink water or swimming pools

Reaction with Chlorine in Cold Water

  • The reaction is disproportionation as the same element is oxidised and reduced
  • When UI is added to the mixture it turns red first then is bleached white

Cl2 + H2O -> HCl + HClO

  • It is a reversible reaction
  • The HClO (chloric I acid) and HCL kill the bacteria
  • HClO acts as bleach and is also an oxidising agent so kills the bacteria by oxidation
  • The other halogens react much more slowly as you go down the group
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Uses of Chlorine Cont

Reaction of Chlorine with Water in Sunlight

2Cl2 + 2H2O -> 4HCl + O2

  • In shallow pools chlorine is rapidly lost due to the sunlight
  • Solutions:
    •  Frequent testing of chlorine levels and adding more chlorine
    • Avoid direct chlorination and use solid sodium/calcium chlorate (NaClO or Ca(ClO)2)
  • Benefits:
    •  Avoiding direct chlorination is easier to measure, dissolve and handle
    • Both salts are less hazardous than toxic Cl2
  • It is a reversible reaction

NaClO + H2O -> Na+ + OH- + HClO

  • The pools are kept slightly acidic since H+ ions neutralise OH-, the equilibrium position shifts to replace the hydroxide ions so moves right making more chloric I acid
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Uses of Chlorine Cont

Reaction of Chlorine with Alkali

Cl2 + 2NaOH -> NaCl + NaClO + H2O

  • It is a reversible reaction
  • NaClO is an active ingredient in bleach
  • The reaction is disproportionation
  • The Sodium Chlorate is an oxidising agent so it is reduced but by doing this it oxidises the bacteria killing it
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