Halogens

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  • Created by: JasmineR
  • Created on: 30-04-16 16:10
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  • Halogens
    • 1) Trends
      • Boiling and melting point
        • Increase down the group
        • As the molecules become larger they have more electrons and so larger van der Waal forces.
          • Therefore, more energy needed to break the forces.
      • Electronegativity
        • Increases down the group
          • The atomic radii increases due to the increasing number of shells.
            • The nucleus is therefore less able to attract the bonding pairs of electrons
        • The relative tendency of an atom in a molecule to attract electrons in a covalent bond.
    • 2) Displacement reactions of halide ions by halogens
      • A halogen that is a strong oxidising agent will displace a halogen with a lower oxidising power from one of its compounds
      • The oxidising strength decreases down the group
        • Oxidising agents are electron acceptors
      • Chlorine will displace both bromide and iodide ions
      • Bromine will displace iodide ions
      • Iodine won't displace anything
      • The colour of the solution in the test tube shows which free halogen is present in the solution
        • Chlorine = very pale green solution (often colourless)
        • Bromine = Yellow solution
        • Iodine = brown solution (sometimes black solid is present)
      • Be able to write these reactions as two half equations showing oxidation or reduction
        • E.g.
          • 2Br- (aq) --> Br2 (aq) + 2e-
            • Cl2(aq) + 2e-  --> 2Cl- (aq)
    • 3) The reactions of halide ions with silver nitrate
      • This reaction is used as a test to identify which halide ion is present
      • The test solution is made with nitric acid, and then silver nitrate soltution
        • The nitric acid will react with any carbonates present to prevent the formation of Ag2CO3
          • Ag2CO3 would mask the desired observations
        • Results of this:
          • Fluorides produce no precipitate
          • Chlorides produce a white precipitate
          • Bromides produce a cream precipitate
          • Iodides produce a pale yellow precipitate
      • The silver halide precipitates can be treated with ammonia solution to help differentiate between them if colours look similar
        • Silver chloride dissolves in dilute ammonia to form a complex ion
        • Silver bromide dissolves in concentrated ammonia to from a complex ion
        • Silver iodide doesn't react with ammonia as it is too insoluble
    • 4) The reaction of halide salts with concentrated sulphuric acid
      • The halides show increasing power as reducing agents going down the group
        • A reducing agent donates electrons
        • This is because as the ions get bigger it is easier for the outer electrons to be given away as the pull from the nucleus on them becomes smaller
      • Fluoride and Chloride
        • Fluoride and chloride ions aren't strong enough reducing agents to reduce S in H2SO4
          • No redox reactions occur
            • Only acid-base reactions occur
              • H2SO4 plays the role of an acid (proton donor)
        • 1)   NaF(s) + H2SO4(l)    --> NaHSO4 (s)  + HF(g)
          • Observations: White steamy fumes of HF are evolved
        • 2)   NaCl (s) + H2SO4 (l) --> NaHSO4 (s) + HCl (g)
          • Observations: White steamy fumes of HCl are evolved
      • Bromide
        • Bromide ions are stronger reducing agents than Cl- and F-
          • After the initial acid-base reaction they reduce the S in H2SO4 from +6 to +4 in SO2
        • 1)    Acid-base step: NaBr (s) + H2SO4 (l)    --> NaHSO4(s) + HBr (g)
          • Observations: White steamy fumes of HBr evolved.
        • 2)   Redox step:   2HBr + H2SO4       --> Br2 (g) + SO2(g) + 2H2O (l)
          • Observations: Red fumes of Bromine evolved and a colourless, acidic gas SO2
        • Reduction production: Sulphur dioxide
      • Iodide
        • Iodide ions are the strongest halide reducing agents
          • They can reduce S from +6 in H2SO4 to +4 in SO2, to 0 in S and -2 in H2S
        • 1)  NaI (s) + H2SO4(l)   --> NaHSO4 (s) + HI (g)
          • Observations: White steamy fumes of Hi are evolved
        • 2)    2HI + H2SO4   --> I2(s)  + SO2(g) + 2H2O (l)
          • Observations: Black solid and purple fumes of iodine are evolved. A colourless, acidic gas of SO2
        • 3)    6HI + H2SO4   --> 3I2 + S (s) + 4H2O (l)
          • Observations: A yellow solid of sulphur
        • 4)   8HI + H2SO4   --> 4I2 (s) + H2S(g) + 4H2O (l)
          • Observations: H2S (hydrogen sulphide), a gas with a bad egg smell
        • Reduction products: sulphur dioxide, sulphur and hydrogen sulphide
        • Note that the H2SO4 plays the role of acid in the first step producing HI and then acts as an oxidising agent in the 3 redox steps
    • 5) The disproportionation reactions of chlorine and chlorate(I)
      • Disproportionation is a reaction where an element simultaneously oxidises and reduces
      • Chlorine with water:
        • Cl2(aq) + H2O(l) --> HClO(aq) + HCl(aq)
          • Chlorine is both simultaneously reducing and oxidising
          • If universal indicator is added it will first turn red due to the acidity of both products. The it will turn colourless as the HClO bleaches the colour
      • Chlorine with water in sunlight
        • 2Cl2 + 2H2O  -->   4H+         + 4Cl- + O2
          • The greenish colour of these solutions is due to the Cl2
            • Cl2(aq) + H2O(l) --> HClO(aq) + HCl(aq)
              • Chlorine is both simultaneously reducing and oxidising
              • If universal indicator is added it will first turn red due to the acidity of both products. The it will turn colourless as the HClO bleaches the colour
        • The same reaction occurs to the equilibrium mixture of chlorine water.
          • The greenish colour of chlorine water fades as the Cl2 reacts and a colourless gas (O2) is produced
      • Chlorine is used in water treatment to kill bacteria
        • It has been used to treat drinking water and the water in swimming pools
        • The benefits to health of water treatment by chlorine outweigh its toxic effects
      • Reaction of chlorine with cold dilute sodium hydroxide solution
        • Cl2 (and Br2 and I2) in aqueous solutions will react with cold sodium hydroxide
        • The colour of the halogen solution will fade to colourless
        • Cl2 (aq) + 2NaOH (aq) -->        NaCl (aq) + NaClO (aq) + H2O (l)
          • The mixture of NaCl and NaClO is used as bleach and to disinfect/kill bacteria
    • 6) Naming chlorate and sulphates
      • The various forms of sulfur and chlorine compounds where oxygen is combined are all called sulfates and chlorates
      • The relevant oxidation number is given in roman numberals
        • If asked to name these compounds remember to add the oxidation number
      • E.g.
        • NaClO: sodium chlorate(I)
        • NaClO3: sodium chlorate(V)
        • K2SO4: potassium sulfate(VI)
        • K2SO3: potassium sulfate(IV)
  • Electronegativity
    • Increases down the group
      • The atomic radii increases due to the increasing number of shells.
        • The nucleus is therefore less able to attract the bonding pairs of electrons
    • The relative tendency of an atom in a molecule to attract electrons in a covalent bond.

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