2.6.3 redox reactions_of_the_halogens

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2.6.3 redox reactions_of_the_halogens

  1. 1. 2.6.3 Redox Reactions of the Halogens i.e. Group VII elements
  2. 2. Introduction • • • • • • Redox comes from reduction and oxidation Oxidation Is Loss of electrons Reduction Is Gain of electrons Remember this by OIL RIG Halogens are powerful oxidising agents They remove [take] electrons from other elements to become negative ions
  3. 3. Experiment 1 Reaction of chlorine gas with potassium iodide
  4. 4. Safety Precautions • KMnO4 – powerful oxidising agent and irritant • HCl - corrosive • Cl2 - poisonous • KI – irritant • Wear safety glasses and rubber gloves
  5. 5. Apparatus • Test tubes • Potassium manganate(VII) crystals • Conc. Hydrochloric acid • Potassium iodide solution • Litmus paper [blue]
  6. 6. Method • Pour some KI solution into a test tube. • Place some potassium manganate(VII) crystals [KMnO4] in a test tube • Add some conc. HCl • Note what happens. • Pour the chlorine gas into the potassium iodide and shake. • Note what happens.
  7. 7. Results • A greenish yellow gas is produced when the conc. HCl is dropped onto the KMnO4 • this gas is chlorine [Cl2] • MnO4- + HCl = Mn2+ + H2O + Cl2 • It is heavier than air • It has a characteristic smell [swimming pool] • It turns blue litmus red - then bleaches it • When chlorine is poured into KI solution turns from colourless to brown as iodine is released.
  8. 8. • • • • • Cl2(g) + 2 KI(aq) = 2 KCl(aq) + I2(aq) Another way of writing the equation is Cl2(g) + 2 I-(aq) = 2 Cl- (aq) + I2(aq) the chlorine has displaced the Iodine and so is more reactive The chlorine has been reduced [gained an electron to become Cl -] The iodide ion I- in KI has been oxidised [lost an electron to become I2 ] The chlorine is an oxidising agent.
  9. 9. Experiment 2 Reaction of Cl2 with KBr
  10. 10. • Make chlorine as in last experiment MnO4- + HCl = Mn2+ + H2O + Cl2 • Put some KBr solution into a test tube • Pour the Cl2 into the test tube of KBr • Note what happens • The KBr solution turns from colourless to red/brown as bromine is released
  11. 11. • • • • • Cl2(g) + 2 KBr(aq) = 2 KCl(aq) + Br2(aq) the chlorine has displaced the Bromide ion and so is more reactive The chlorine has been reduced [gained electron to become Cl-] The bromide ion Br- in KBr has been oxidised [lost a electron to become Br2.] The chlorine is an oxidising agent. Iodine is darker than Bromine
  12. 12. Conclusions • Non-metals can also be arranged in a reactivity series. • F• Cl• Br• [O2-] • I• [SO42-] More Reactive Less Reactive
  13. 13. Experiment 3 Reaction of 2+ Cl2 with Fe
  14. 14. Apparatus • KMnO4 crystals • Conc. HCl • Iron (II) sulphate solution FeSO4
  15. 15. Method • • • • • • • Make up a solution of iron(II)sulphate it is pale green in colour Make some chlorine gas mix the gas and the iron(II)sulphate solution Note what happens the solution turns brown 2 Fe2+ + Cl2 = 2 Fe3+ + 2 ClGreen Brown • the Fe2+ has been oxidised to Fe3+
  16. 16. Key Points to remember • OIL • RIG • F most reactive • More reactive element displaces less reactive from solution
  17. 17. Experiment 4 Reaction of Cl2 with Na2SO3
  18. 18. Apparatus • • • • KMnO4(s) Conc. HCl Sodium sulphite solution [Na2SO3] Barium Nitrate Solution
  19. 19. • • • • • • • • • Make up a solution of sodium sulphite Test it with Barium nitrate solution A white precipitate forms Add HCl and the precipitate re-dissolves This tells us the solution is a sulphite Take a fresh sample of the solution react it with some chlorine Add barium nitrate solution Note what happens
  20. 20. • A white precipitate forms • Add Hydrochloric acid to see what happens to the precipitate • The precipitate does not re-dissolve • the sulphite has been oxidised to a sulphate Cl2 + SO32- + H2O = 2 Cl- + SO42- + 2H+ • look at oxidation numbers • S goes from +4 to +6 so oxidised by Cl2

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