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The Halogens

The Halogens. Chapter 40. Gp VII. F. Cl. Br. I. At. Position in Periodic Table. General information. Characteristic properties. Very reactive non-metals High electronegativity High electron affinity Bonding and Oxidation states Colour. Electronegativity.

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The Halogens

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  1. The Halogens Chapter 40

  2. Gp VII F Cl Br I At Position in Periodic Table

  3. General information

  4. Characteristic properties • Very reactive non-metals • High electronegativity • High electron affinity • Bonding and Oxidation states • Colour

  5. Electronegativity A measure of the attracting ability of bonding electrons. Due to their large effective nuclear charge, halogens are the most electronegative element in the periodic table.

  6. Electron Affinity Generally high electron affinity. All are exothermic. X(g) + e- X-(aq)

  7. Bonding and Oxidation State • ns2np5 • Ionic or covalent bond with oxidation state –1 or +1 (F shows –1 only) • Except F, all other halogens can expand their octet by using the low-lying, vacant d-orbitals to form bonding. Their oxidation states range from –1 to +7.

  8. nd nd np np ns (+3 states) ns   (0, +1, -1 states) HClO2, ClO2-       Cl2, Cl-, OCl2 Bonding and Oxidation State

  9.   nd   nd np   np ns  (+7 state) ns   (+5 state) Cl2O7, HIO4    HClO3, HBrO3, I2O5 Bonding and Oxidation State

  10. Colour

  11. White light Iodine Colour

  12. E2 E1 Iodine Fluorine Check point 40-1

  13. b.p. 400 m.p. (Temp. oC) 0 F Cl Br I At -300 Variation in properties of Halogens m.p./b.p. increase as Mr increase because Van der Waals’ Force increase with Mr

  14. 4.0 4 3.0 3 2.8 2.5 2 1 F Cl Br I Electronegativity F form ionic compound with –1 oxidation state. I shows highly positive oxidation states with more electronegative element. e.g. KIO4

  15. 340 320 280 F Cl Br I Electron affinity EA decrease from Cl to I due to increasing atomic size which lowers the nuclear attraction to the added electron. F has exceptional lower EA because its 2nd shell is already crowded with 7 e-, the additional e- will experienced a greater e-e repulsion than other halogens.

  16. H ½ X2(std.state)  X-(aq) Hatom Hhyd X(g) HEA X-(g) H = Hatom + HEA + Hhyd Variation in Chemical Properties of elements All are good oxidizing agents. (F2>Cl2>Br2>I2)

  17. Oxidizing strength Oxidizing power Hatom HEA Hhyd H

  18. Oxidizing properties • Oxidation of metals • a. F2 oxidizes all metals including Au and Ag. • b. Cl2 oxidizes common metals such as Na and Cu. • Reactions with Fe2+(aq) • 2Fe2+ + Cl22Fe3+ + 2Cl- • 2Fe2+ + Br22Fe3+ + 2Br- • 2Fe3+ + 2I-2Fe2+ + I2

  19. Oxidizing properties 3. Reaction with phosphorus 2P + 3X2 2PX3 2P + 5X2  2PX5 F2 and Cl2 form PX5 as they have high oxidizing Strength. They combine with P to exhibit its maximum oxidation state. Br2 and I2 form PX3 only.

  20. Reaction with water Fluorine oxidizes water to form HF and O2 2F2 + 2H2O  2HF + O2 Chlorine undergoes disproportionation(self oxidation and reduction) to form HCl and HOCl. Cl2 + H2O  HCl + HOCl

  21. Reaction with water A mixture of Cl2(aq), HCl(aq) and HOCl(aq) is called chlorine water. • OCl- , chloric(I) ion • Strong oxidizing agent with bleaching property • OCl- + dye (coloured)  Cl- + dye (colourless) • Unstable to heat and light • 2OCl- O2 + 2Cl-

  22. Reaction with water Br2 is only slightly soluble in water. Br2 (aq) + H2O (l)  HBr(aq) + HOBr(aq) • OBr-(aq) is unstable and with bleaching property: • 2OBr- O2 + 2Br- • 2. OBr- + dye(coloured)  Br- + dye(colourless)

  23. Reaction with water I2 does not react with water and only very slightly soluble in water. I2 is more soluble in ethanol or in KI(aq). I2(aq) + I-(aq) I3-(aq)

  24. Reaction with alkalis All halogens react with aqueous alkalis and disproportionate in it. The products depend on temperature and concentration of the alkalis.

  25. Reaction with alkalis

  26. Reaction with alkalis

  27. Reaction with alkalis The reverse reaction of I2 and KOH is used to prepare standard iodine solution for iodometric titration. Know amount of iodine is generated by dissolving known quantity of KIO3 in excess KI and H+. 5KI + KIO3 + 3H+  3I2 + 6K+ + 3H2O

  28. Variation in properties of halides Halogen Displacement reactions Oxidizing strength : F2 >Cl2 >Br2 >I2 Cl2 + 2Br- 2Cl- + Br2 Cl2 + 2I- 2Cl- + I2 Br2 + 2I- 2Br- + I2

  29. Halide + c. H2SO4 NaCl + H2SO4 NaHSO4 + HCl 2NaCl + H2SO4 Na2SO4 + HCl (500oC) 2NaBr + 2H2SO4 2HBr + Na2SO4 2HI + H2SO4  SO2 + Br2 + 2H2O 2NaI + 2H2SO4 2HI + Na2SO4 8HI + H2SO4  H2S + 4I2 + 4H2O Relative reducing power: HCl < HBr < HI

  30. Halide + c. H3PO4 3NaCl + H3PO4 Na3PO4 + 3HCl 3NaBr + H3PO4 Na3PO4 + 3HBr 3NaI + H3PO4 Na3PO4 + 3HI H3PO4 is NOT a strong oxidizing agent. It reacts with halides to form HX.

  31. Halides + Ag+(aq) Ag+(aq) + Cl-(aq)  AgCl(s) , white precipitate Ag+(aq) + Br-(aq)  AgBr(s) , pale yellow precipitate Ag+(aq) + I-(aq)  AgI(s) , yellow precipitate

  32. Halides + Ag+(aq)

  33. Relative acidity of HX

  34. Relative Acidity of HX H HX(aq) H+(aq) + X-(aq) H5 H6 H1 H+(g) + X-(g) H3 H4 H2 HX(g) H(g) + X(g)

  35. Explanation of weak acid strength of HF: • HF has the greatest bond dissociation • energy and exceptionally small • electron affinity. It has the least • exothermic H. • Due to formation of strong hydrogen • bond and greatest degree of hydration, • HF has the smallest decrease of TS. Relative Acidity of HX H1 H2 H3 H4 H5 H6 H TS G

  36. Anomalous behaviour of HF According to La Chatelier’s Principle, the degree of dissociation of a weak acid increases with dilution. HA H+ + X- However, the acidity of conc. HF is greater than dilute HF. Why? In conc. HF, HF molecules form dimers H2F2 by hydrogen bond. The acid strength of H2F2 is stronger than HF. H2F2 H+ + HF2-

  37. Uses of Halogens and Halides • Teflon • Na2SiF6 or NaF is used to fluoridate drinking water. • Chlorine bleach, sterilization of water, PVC. • Anti-knocking agent (leaded petrol) • AgBr coating in film • Iodine tincture • AgI in coating film

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