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Halogens

Halogens. AS. Generally:. Oxidising agents Germicides Note: Atoms are halogens Ions are halides Ions have 8 electrons by borrowing one, so single negative charge (F¯ etc.) Not very soluble in water, more soluble in organic solvent, usually cyclohexane. Electronic configurations.

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Halogens

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  1. Halogens AS

  2. Generally: • Oxidising agents • Germicides • Note: Atoms are halogens • Ions are halides • Ions have 8 electrons by borrowing one, so single negative charge (F¯ etc.) • Not very soluble in water, more soluble in organic solvent, usually cyclohexane

  3. Electronic configurations • 2s22p5 • 3s23p5 • 4s2(3d10)4p5 • 5s2(4d10)5p5

  4. Atomic radius increases down group because of: • More shielding / repulsion • More shells • (Even though) Bigger nucleus • Ionic radius matches atomic radius

  5. Boiling point • (X2 molecules) • Increases down G7 • Because of increased Van der Waal’s forces • F & Cl gases • Br liquid • I & At solid

  6. Reactivities (oxidising ability) • ½X2 + eˉ Xˉ • Reactivity of atoms decreasesdown group because of: • Stronger X-X bond • Electron affinity not changing • Hydration / Lattice energy decreasing – making, say, NaI releases less energy than making, say, NaF

  7. Displacement • Higher-reactivity halogens will displace lower-reactivity halogens • F > Cl > Br > I > (At) • E.g. Cl2 + 2NaBr  Br2 + 2NaCl • Or Cl2 + 2Br- 2Cl- + Br2 • Displaced bromine – yellow/brown colour • Displaced iodine – brown and/or black precipitate • Fluorine too dangerous

  8. Electronegativity and polarisation • Electronegativity decreases down group • More shells so more shielding • Radius increases so less nuclear attraction for shared electrons • Bigger nucleus should attract electrons but it has less effect than other two • Therefore hydrogen halides are decreasingly polar

  9. Hydrogen halides (HCl etc.) • All colourless gases, very soluble in water, dissociate well,  strong acids • E.g. HCl(aq) H+(aq) + Cl-(aq) • Bond enthalpy decreases down group, easier to separate H+ and X- • So  strength of acid increases •  HI > HBr > HCl • These acids are oxidising agents – the hydrogen ion (H+) can take an electron • Note: conc. HF not a strong acid – bond dissociation enthalpy too high, diluted is stronger

  10. Tests for ions: • Test solutions with acidified silver nitrate to make a silver halide: • (acidified to avoid formation of carbonates) • F – no change - AgF is soluble • Cl – white AgCl precipitate • Br – cream AgBr precipitate • I – yellow AgI precipitate

  11. General reaction: • AgNO3(aq) + Xˉ(aq) AgX(s) + NO3ˉ(aq) • or Ag+(aq) + Xˉ(aq) AgX (s) • Check with dilute and conc. ammonia solution (NH4OH): • Fˉ no reaction anyway • AgCl precipitate soluble in dilute or concentrated ammonia • AgBr precipitate soluble in concentrated ammonia only • AgI precipitate insoluble, even in concentrated ammonia

  12. A redox reaction: • Cl2 + H2O  HCl + HClO • Note oxidation states: • One chlorine goes from 0 to -1 in HCl • Other goes 0 to +1 in HClO • This is a disproportionation – one element is simultaneously oxidised and reduced.

  13. Reduction by halideions • Iodide ion I- is very reducing – • “happy” to lose electron • didn’t really want it • bromide less so, etc. • Decreasing reducing power up the group • Electrons lost particularly to strong oxidising agents (H2SO4, F2 etc.)

  14. Homework • Find out the reactions of H2SO4 with F- to I- and memorise

  15. Halide ion reactions with concentrated H2SO4 MEMORISE!! • Sulphur has an oxidation number of +6 in SO42ˉ • Fˉ and Clˉ cannot reduce sulphur – only white fumes of HF or HCl seen when NaF or NaCl added to conc. sulphuric acid • Fˉ + H2SO4 HF + HSO4ˉ • Clˉ + H2SO4 HCl + HSO4ˉ

  16. (Brˉ) NaBr + H2SO4 • Brˉ can reduce sulphur from +6 to +4, producing HBr, SO2, Br2 • Observe white fumes of HBr, invisible SO2 (bubbles) and brown fumes of Br2 (or possibly liquid Br2) • Brˉ + H2SO4 HBr + HSO4ˉ • Then • 2HBr + H2SO4 Br2 + SO2 + 2H2O

  17. (Iˉ) NaI + H2SO4 • Iˉ can reduce sulphur from +6 to +4, then 0, then -2, producing HI, SO2, S, H2S, I2 • White fumes of HI, invisible SO2 and H2S (bubbles, distinctive smell), yellow sulphur, and purple fumes of I2 • Iˉ + H2SO4 HI + HSO4ˉ Then • 2HI + H2SO4 I2 + SO2 + 2H2O then

  18. General reaction: • NaX(s) + H2SO4(aq) NaHSO4(aq) + HX(g) • fumes in moist air • Or Xˉ + H2SO4 HSO4ˉ + HX • In these, the halide is acting as a base (proton acceptor), accepting H+ • E.g. • NaF(s) + H2SO4(aq) NaHSO4(aq) + HF(g)

  19. Uses of chlorine and chlorate(I) compounds • Under most conditions: • H2O + Cl2Ý HCl + HClO • This is a disproportionation, one chlorine atom is reduced, the other is oxidised. • This makes “chlorine water” which can decompose photolytically (in sunlight): • 2Cl2 + 2H2O  O2 + 4HCl

  20. Chlorination • Water treatment • Drinking water – 0.7mg dm-3 • Swimming pools more concentrated • Kills bacteria, especially E.coli from bottoms.

  21. Reaction with NaOH Cl2 + 2NaOH Ý NaCl + NaClO + H2O • This mixture of sodium chloride and sodium chlorate(I) is used as a bleach.

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