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1 Characteristic Properties of the Halogens 2 Group VIIA elements include  fluorine  chlorine  bromine  iodine  astatine Introduction halogens (Salt.

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Presentation on theme: "1 Characteristic Properties of the Halogens 2 Group VIIA elements include  fluorine  chlorine  bromine  iodine  astatine Introduction halogens (Salt."— Presentation transcript:

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2 1 Characteristic Properties of the Halogens

3 2 Group VIIA elements include  fluorine  chlorine  bromine  iodine  astatine Introduction halogens (Salt producers)

4 3 Astatine  chemistry not much known  radioactive  the total amount present in the Earth's crust is probably less than 30 g at any one time. Introduction

5 4 Halogens are p-block elements  outermost shell electronic configuration of ns 2 np 5

6 5  one electron short of the octet structure Halogens are p-block elements

7 6 In the free elemental state Introduction they complete their octets by sharing their single unpaired p-electrons

8 7 they either gain an additional electron to form halide ions or When halogens react with other elements share their single unpaired p- electrons to form single covalent bonds

9 8  highest among the elements in the same period  have a high tendency to attract electrons  strong oxidizing agents High Electronegativity / Electron Affinity

10 9  -1 is the most common oxidation state of halogens in their compounds Ionic : NaF, NaCl, NaBr, NaI Covalent :HF, HCl, HBr, HI High Electronegativity / Electron Affinity

11 10 Variable Oxidation State All halogens (except fluorine ) can expand their octet of electrons by utilizing the vacant, energetically low-lying d-orbitals.

12 11 “Electrons-in-boxes” diagrams of the electronic configuration of a halogen atom of the ground state and various excited states

13 12 The half-filled orbital(s) overlap(s) with those of more electronegative atoms (e.g. O )  positive oxidation state (+1, +3, +5, +7)

14 13 Oxidation state of halogen Ion / Compound –1 F – Cl – Br – I – HFHClHBrHI OF 2 0F 2 Cl 2 Br 2 I 2 +1 Cl 2 OBr 2 O HOClHOBr OCl – OBr – +3 HClO 2 ClO 2 – Various oxidation states of halogens in their ions or compounds

15 14 Oxidation state of halogen Ion / Compound +4ClO 2 BrO 2 +5 HClO 3 HBrO 3 I 2 O 5 ClO 3 – BrO 3 – HIO 3 IO 3 – +6Cl 2 O 6 BrO 3 +7 Cl 2 O 7 H 5 IO 6 HClO 4 HIO 4 ClO 4 – IO 4 – Various oxidation states of halogens in their ions or compounds

16 15 Fluorine (  1)  the most electronegative element  only one unpaired p electron available for bonding  oxidation state is limited to –1

17 16 Fluorine (  1)  cannot expand its octet  no low-lying empty d orbitals available  the energy required to promote electrons into the third quantum shell is very high Absence of HFO, HFO 2, HFO 3, HFO 4

18 17 Variation in Physical Properties 1. Melting point / boiling point  down the group Halogen Melting point (  C) Boiling point (  C) Fluorine Chlorine Bromine Iodine Astatine –220 –101 – –188 –

19 18 Variations in melting point and boiling point of the halogens

20 19 Variation in Physical Properties 1. Melting point / boiling point  down the group The molecular size  down the group  The electron cloud is more easily polarized  Induced dipoles are formed more easily  Stronger London dispersion forces

21 20 2. Colour becomes darker down the group HalogenF2F2 Cl 2 Br 2 I2I2 Colour Pale yellow Greenish yellow Reddish brown Violet black

22 21 Appearances of halogens at room temperature and pressure: chlorine chlorine

23 22 Appearances of halogens at room temperature and pressure: bromine bromine

24 23 Appearances of halogens at room temperature and pressure: iodine iodine

25 24 Colour All halogens  coloured  the absorption of radiation in the visible light region of the electromagnetic spectrum The colour is due to the unabsorbed radiation in the visible light region

26 25 Colour Fluorine atom  has the smallest size  absorbs the radiation of relatively high frequency (i.e. blue light )  appears yellow (the unabsorbed radiation)

27 26 Colour Atoms of other halogens  larger sizes  absorb radiation of lower frequency

28 27 Colour Iodine  absorbs the radiation of relatively low frequency (i.e. yellow light )  appears violet

29 28 Q.1The colour of astatine is black.

30 29 Colour Halogens  different colours when dissolved in different solvents

31 30 Halogen Colour in pure formin waterin 1,1,1-trichloroethane F2F2 Pale yellow Cl 2 Greenish yellowPale yellowYellow Br 2 Reddish brownYellowOrange I2I2 Violet black Yellow (only slightly soluble) Brown in KI(aq) Violet Colours of halogens in pure form and in solutions

32 31 Colour Halogens  non-polar molecules  not very soluble in polar solvents (such as water )  but very soluble in organic solvents (such as 1,1,1-trichloroethane )

33 32 Colours of halogens in water: (a) chlorine; (b) bromine; (c) iodine (a)(b)(c)

34 33 Colours of halogens in 1,1,1-trichloroethane: (a) chlorine; (b) bromine; (c) iodine (a)(b)(c)

35 34 3. Electron Affinity  down the group HalogenFClBrIAt E.A. kJ/mol 

36 35 The number of electron shells and size of atoms  down the group  The nuclear attraction for the additional electron  down the group  Electron affinity  from Cl to I

37 36 Atoms of fluorine have the smallest size among the halogens  The addition of an extra electron to the small quantum shell(n=2) results in great repulsion among the electrons.  Fluorine has a lower electron affinity than Cl and Br.

38 37 HalogenFClBrIAt Electronegativity Electronegativity  down the group

39 38 The number of electron shells and size of atoms  down the group  The nuclear attraction for the bonding electrons  down the group  Electronegativity  down the group

40 39 Fluorine has the highest electronegativity because it is the most reactive elements. The electronegativity of fluorine is arbitrarily assigned as 4.0.

41 40 Variation in Chemical Properties Reactivity : F 2 > Cl 2 > Br 2 > I 2 React by gaining electrons Oxidizing power : F 2 > Cl 2 > Br 2 > I 2

42 41 1. Reactions with Sodium All halogens  combine directly with sodium to form sodium halides  the reactivity decreases down the group from fluorine to iodine

43 42 Fluorine  react explosively to form sodium fluoride 2Na(s) + F 2 (g)  2NaF(s) 1. Reactions with Sodium

44 43 Chlorine  reacts violently to form sodium chloride 2Na(s) + Cl 2 (g)  2NaCl(s) 1. Reactions with Sodium

45 44 Bromine  burns steadily in bromine vapour to form sodium bromide 2Na(s) + Br 2 (g)  2NaBr(s) 1. Reactions with Sodium

46 45 Iodine  burns steadily in iodine vapour to form sodium iodide 2Na(s) + I 2 (g)  2NaI(s) 1. Reactions with Sodium

47 46 Na(s) + X 2 NaX(s) Na + (g) + X 2 (g) I.E. Na + (g) + X(g) B.E. Na + (g) + X  (g) E.A. Vigor of reaction depends on 1.The activation energy (endothermic) 2.The lattice energy (exothermic)  Activation energy

48 47 Na(s) + X 2 NaX(s) Na + (g) + X(g) Na + (g) + X 2 (g) I.E. B.E. Na + (g) + X  (g) E.A. F has an exceptionally low B.E. & zero F is the most reactive (g)

49 48 Na(s) + X 2 NaX(s) Na + (g) + X(g) Na + (g) + X 2 (g) I.E. B.E. Na + (g) + X  (g) E.A. The lattice enthalpy of NaF is most negative

50 49 Na(s) + X 2 NaX(s) Na + (g) + X(g) Na + (g) + X 2 (g) I.E. B.E. Na + (g) + X  (g) E.A. Cl has zero Cl is more reactive than Br & I (g)

51 50 Na(s) + X 2 NaX(s) Na + (g) + X(g) Na + (g) + X 2 (g) I.E. B.E. Na + (g) + X  (g) E.A. Lattice enthalpy : NaCl > NaBr > NaI

52 51 Na(s) + X 2 NaX(s) Na + (g) + X(g) Na + (g) + X 2 (g) I.E. B.E. Na + (g) + X  (g) E.A. (s)/(l) Br is more reactive than I : Br 2 (l) < I 2 (s)

53 52 Na(s) + X 2 NaX(s) Na + (g) + X(g) Na + (g) + X 2 (g) I.E. B.E. Na + (g) + X  (g) E.A. Lattice enthalpy : NaBr > NaI

54 53 Q.2(a) Variation:bond enthalpy decreases from Cl 2 to I 2 Reason : The size of atoms and thus the bond length between atoms increases down the group. The shared electron pair is getting further away from the bonding nuclei.  weaker bond and lower B.E. F 2 has an exceptionally small B.E. because the F atoms are so small that the repulsive forces between lone pairs on adjacent bonding atoms become significant.

55 54 Q.2(b) The lattice enthalpy becomes less negative down the group. It is because the anionic radius, r -, increases down the group.

56 55 F 2 reacts explosively even in the dark at  200  C Cl 2 reacts explosively in sunlight Br 2 reacts moderately on heating with a catalyst I 2 reacts slowly and reversibly even on heating 2.1Reactions with hydrogen X 2 + H 2 (g)  2HX(g)

57 56 Q.3 Explain the extreme reactivity of fluorine in terms of the bond enthalpies of F–F and H–F bonds. Fluorine has an exceptionally small F-F bond enthalpy. Thus, the activation energy of its reaction with hydrogen is also exceptionally small. Hydrogen fluoride has the highest bond enthalpy among the hydrogen halides. Thus, the formation of HF from H 2 and F 2 is the most exothermic. The energy released from the reaction further speeds up the reaction. F 2 + H 2 (g)  2HF(g)

58 57 Chlorine removes hydrogen completely from turpentine(C 10 H 16 ) C 10 H 16 (l) + 8Cl 2 (g)  10C(s) + 16HCl(g)

59 58 Q.4 The cotton wool bursts into flames and the gas jar is filled with dark smoke (of carbon) and white fumes (of HCl) HCl gives dense white fumes with ammonia.

60 59 2.2Reactions with phosphorus F 2 + P  PF 5 Cl 2 + P  PCl 3 + PCl 5 Br 2 + P  PBr 3 I 2 + P  PI 3 F 2 is the strongest oxidizing agent, it always oxidizes other elements to their highest possible oxidation states.

61 60 2.2Reactions with phosphorus F 2 + P  PF 5 Cl 2 + P  PCl 3 + PCl 5 Br 2 + P  PBr 3 I 2 + P  PI 3 Br 2 and I 2 are NOT strong enough to oxidize P to its highest possible oxidation state.

62 61 2.3Reactions with xenon Fluorine reacts directly with all non-metals except nitrogen, helium, neon and argon. It will even react with diamond and xenon on heating. C(diamond) + 2F 2  CF 4 Xe + F 2  XeF 2 Xe + 2F 2  XeF 4 Xe + 3F 2  XeF 6

63 62 2.3Reactions with xenon It is because (a)Xenon can expand its octet by utilizing vacant, low-lying d-orbitals.

64 63 By VB Theory, To form two Xe-F bonds in XeF 2, a 5p electron in Xe has to be promoted to a 5d orbital. Xe  5s5p  F 2s2p  Xe*  5s5p  5d

65 64 By VB Theory, To form four Xe-F bonds in XeF 4, two 5p electrons in Xe have to be promoted to two 5d orbitals. Xe  5s5p  5d Xe**  5s5p  5d 

66 65 By VB Theory, To form six Xe-F bonds in XeF 6, three 5p electrons in Xe have to be promoted to three 5d orbitals. Xe  5s5p  5d Xe***  5s5p  5d 

67 66 2.3Reactions with xenon The gap between np and nd sub-shells  down the group, thus, Tendency to form bonds  down the group : - Xe > Kr > Ar > Ne > He the promotion of electrons from np sub-shell to nd sub-shell becomes easier down the group.

68 67 Xe  5s5p  5d Xe***  5s5p  5d  Also, the energy released by forming more single bonds outweighs the energy required for promoting 5p electrons to 5d orbitals.

69 68 3Reactions with other reducing agents I 2 is the weakest oxidizing agents among the halogens.

70 69 3.1All halogens(except I 2 ) oxidize Fe 2+ to Fe 3+ Half reaction Standard electrode potential (V) Cl 2 (aq) + 2e –  2Cl – (aq) Br 2 (aq) + 2e –  2Br – (aq) Fe 3+ (aq) + e –  Fe 2+ (aq) I 2 (aq) + 2e –  2I – (aq) X 2 (aq) + 2Fe 2+ (aq)  2X  (aq) + 2Fe 3+ (aq) ( X = F, Cl, Br)

71 70 3.1All halogens(except I 2 ) oxidize Fe 2+ to Fe 3+ Half reaction Standard electrode potential (V) Cl 2 (aq) + 2e –  2Cl – (aq) Br 2 (aq) + 2e –  2Br – (aq) Fe 3+ (aq) + e –  Fe 2+ (aq) I 2 (aq) + 2e –  2I – (aq) I 2 (aq) + 2Fe 2+ (aq)  No reaction

72 71 3.2All halogens(except I 2 ) oxidize S 2 O 3 2  to SO 4 2  4X 2 (aq) + S 2 O 3 2  (aq) + 5H 2 O(l)  8X  (aq) + 10H + (aq) + 2SO 4 2  (aq) (X = F, Cl, Br) I 2 (aq) + 2S 2 O 3 2  (aq)  2I  (aq) + S 4 O 6 2  (aq) Used in iodometric titration

73 72 (i)2I  (aq) + 2Fe 3+ (aq)  I 2 (aq) + 2Fe 2+ (aq) (excess) (unknown) Determination of [Fe 3+ (aq)] by iodometric titration Using starch as indicator (ii) I 2 (aq) + 2S 2 O 3 2  (aq)  2I  (aq) + S 4 O 6 2  (aq) (standard solution)

74 73 4Displacement reactions Cl 2 (aq) + 2Br  (aq)  2Cl  (aq) + Br 2 (aq) Cl 2 (aq) + 2I  (aq)  2Cl  (aq) + I 2 (aq) Br 2 (aq) + 2I  (aq)  2Br  (aq) + I 2 (aq) I 2 (aq) + I  (aq)  I 3  (aq) (yellow) (brown) More reactiveLess reactive

75 74 4Displacement reactions Cl 2 (aq) + 2I  (aq)  2Cl  (aq) + I 2 (aq) Br 2 (aq) + 2I  (aq)  2Br  (aq) + I 2 (aq) I 2 (aq) + I  (aq)  I 3  (aq) (yellow) (brown) What would be observed if an excess of Cl 2 (aq) or Br 2 (aq) is added to I  (aq)? The solution turns cloudy and a black solid settles at the bottom

76 75 Aqueous solution Halogen added F2F2 Cl 2 Br 2 I2I2 F–F– No reaction Cl – A pale yellow solution is formed (Cl 2 is formed) No reaction Reactions of halide ions with halogens

77 76 Aqueous solution Halogen added F2F2 Cl 2 Br 2 I2I2 Br – A yellow solution is formed (Br 2 is formed) No reaction I–I– A yellowish brown solution is formed (I 3  is formed) No reaction Reactions of halide ions with halogens

78 77 Q.5 Shake hexane or 1,1,1-trichloroethane with the two solutions respectively. The one that turns the organic layer violet is I 3  (aq). The one that turns the organic layer orange or brown is Br 2 (aq).

79 78 1,1,1-trichloroethane Br 2 I2I2 Br 2 (aq) I 3  (aq) If hexane is used, the upper layer will be the organic layer

80 79 Disproportionation is a chemical change in which oxidation and reduction of the same species (which may be a molecule, atom or ion) take place at the same time. 5.Disproportionation

81 80 A.Reactions with Water HOCl : chloric(I) acid or hypochlorous acid Chlorine water  a mixture of hydrochloric acid and chloric(I) acid

82 81 Chlorate(I) ion, OCl  is also known as hypochlorite ion  unstable  decomposes when exposed to sunlight or high temperatures to give chloride ions and oxygen 2OCl – (aq)  2Cl – (aq) + O 2 (g) A.Reactions with Water

83 82 Chlorate(I) ion  bleaches by oxidation A.Reactions with Water Cl 2 (aq) + H 2 O(l) 2H + (aq) + Cl – (aq) + OCl – (aq) OCl – (aq) + dye  Cl – (aq) + (dye + O) coloured colourless

84 83 Bromine  only slightly soluble in water  mainly exists as molecules in saturated bromine water A.Reactions with Water

85 84 When the solution is diluted  hydrolysis takes place  hydrobromic acid and bromic(I) acid (hydrobromous acid) are formed Br 2 (l) + H 2 O(l) HBr(aq) + HOBr(aq) A.Reactions with Water

86 85 Bromate(I) ion, OBr   also unstable  bleaches dyes by oxidation OBr – (aq) + dye coloured  Br – (aq) + (dye + O) colourless A.Reactions with Water

87 86 Iodine  does not react with water  only slightly soluble in water A.Reactions with Water

88 87 Fluorine reacts vigorously with water to form hydrogen fluoride and oxygen A.Reactions with Water Being the strongest oxidizing agent, F 2 undergoes reduction rather than disproportionation with water. 2F 2 (g) + 2H 2 O(l)  4HF(aq) + O 2 (g) 0 11

89 88 Chlorine reacts similarly at high temperature or when exposed to light 2Cl 2 (aq) + 2H 2 O(l)  2HCl(aq) + 2HOCl(aq) 2HOCl(aq) 2HCl(aq) + O 2 (g) Heat or light Overall : 2Cl 2 (aq) + 2H 2 O(l) 4HCl(aq) + O 2 (g) Heat or light A.Reactions with Water

90 89 All halogens react with aqueous alkalis All halogens ( except F 2 ) undergoes disproportionation with alkalis In general, Reactivity decreases down the group B.Reactions with Alkalis

91 90 The products formed depend on 1.Temperature 2.The type of halogen reacted 3.The concentration of alkali used B.Reactions with Alkalis

92 91 Effect of temperature (a)At lower temperatures, X2X2 Cl 2 Br 2 I2I2 T 1 /  C  20 00<0<0 X 2 (aq) + 2OH  (aq) XO  (aq) + X  (aq) + H 2 O(l) T1T1 0+1 11 B.Reactions with Alkalis

93 92 Effect of temperature (a)At higher temperatures, 3XO  (aq) XO 3  (aq) + 2X  (aq) T2T2 XO  ClO  BrO  IO  T 2 /  C  70  20 0 11 B.Reactions with Alkalis

94 93 (2) 3XO  (aq) XO 3  (aq) + 2X  (aq) T2T2 (1) X 2 (aq) + 2OH  (aq) XO  (aq) + X  (aq) + H 2 O(l) T1T1 Overall reaction : 3  (1) + (2) 3X 2 (aq) + 6OH  (aq) XO 3  (aq) + 5X  (aq) + 3H 2 O(l) T2T2 B.Reactions with Alkalis

95 94 3XO  (aq) XO 3  (aq) + 2X  (aq) T2T2 XO  ClO  BrO  IO  T 2 /  C7020<0 On moving down the group, 1.stability of XO  decreasesClO  > BrO  > IO  2.stability of XO 3  increasesClO 3  < BrO 3  < IO 3  B.Reactions with Alkalis

96 95 3X 2 (aq) + 6OH  (aq) XO 3  (aq) + 5X  (aq) + 3H 2 O(l) At lower pH (when acid is added), the equilibrium position shifts to the left and the reversed process predominates. XO 3  (aq) + 5X  (aq) + 6H + (aq) 3X 2 (aq) + 3H 2 O(l) This reaction (when X=I) is often used to prepare standard iodine solution for iodometric titrations B.Reactions with Alkalis

97 96 Dissolving a known quantity of KIO 3 (s) in excess KI(aq) and dilute H 2 SO 4 generates a known amount of I 2 (aq) KIO 3 (aq) + 5KI(aq) + 6H + (aq)  3I 2 (aq) + 3H 2 O(l) + 6K + (aq) The iodine produced can be used to standardize thiosulphate solution 3I 2 (aq) + 6S 2 O 3 2  (aq)  6I  (aq) + 3S 4 O 6 2  (aq) B.Reactions with Alkalis

98 97 This known amount of iodine generated can also be used to oxidize reducing agents (of unknown concentrations ) such as SO 3 2  (aq) and ascorbic acid (vitamin C) The excess iodine can be determined by back titration with sodium thiosulphate solution I 2 (aq) + 2S 2 O 3 2– (aq)  2I – (aq) + S 4 O 6 2– (aq) B.Reactions with Alkalis

99 98 Effect of concentration of alkali (a)At higher concentrations, XO 3  (aq) is the major product. (b)At lower concentrations, XO  (aq) is the major product. B.Reactions with Alkalis

100 99 B.Reactions with Alkalis In general, Halogens react with cold, dilute alkali to give halate(I) ions, halide ions and water X 2 (aq) + 2OH   XO  (aq) + X  (aq) + H 2 O(l) Halogens react with hot, concentrated alkali to give halate(V) ions, halide ions and water. 3X 2 (aq) + 6OH   XO 3  (aq) + 5X  (aq) + 3H 2 O(l)

101 100 2F 2 + 2OH  (aq) OF 2 (aq) + 2F  (aq) + H 2 O(l) very dilute 20  C 2F 2 + 4OH  (aq) O 2 (aq) + 4F  (aq) + 2H 2 O(l) concentrated 70  C 0 0 11 11 11 22 2 Being the strongest oxidizing agent, F 2 undergoes reduction rather than disproportionation with alkalis. B.Reactions with Alkalis

102 101 Variation in chemical properties of halides A Comparative study 1.Reactions with conc. sulphuric acid 2.Reactions with conc. phosphoric acid 3.Reactions with silver ion

103 102 Concentrated sulphuric acid  non-volatile (b.p. ~330  C)  oxidizing Reactions with Concentrated Sulphuric(VI) Acid

104 103 KF(s) + H 2 SO 4 (l)  KHSO 4 (s) + HF(g) KCl(s) + H 2 SO 4 (l)  KHSO 4 (s) + HCl(g) warm non-volatile volatile Warming is required to speed up the reaction and to drive out the volatile acids Fluoride and chloride : -

105 104 KF(s) + H 2 SO 4 (l)  KHSO 4 (s) + HF(g) KCl(s) + H 2 SO 4 (l)  KHSO 4 (s) + HCl(g) warm acid salt Acid salt rather than normal salt is formed because HSO 4  is a relatively weak acid A convenient way to prepare HCl in the laboratory

106 105 KF(s) + H 2 SO 4 (l)  KHSO 4 (s) + HF(g) KCl(s) + H 2 SO 4 (l)  KHSO 4 (s) + HCl(g) warm Observation : - White fumes are produced Confirmatory test : - Dense white fumes appear with NH 3 (aq)

107 106 Bromide: - 0 oxidation +6+4 reduction KBr(s) + H 2 SO 4 (l)  KHSO 4 (s) + HBr(g) warm 2HBr(g) + H 2 SO 4 (l)  SO 2 (g) + Br 2 (g) + 2H 2 O(l) warm

108 107 (1)KBr(s) + H 2 SO 4 (l)  KHSO 4 (s) + HBr(g) (2) 2HBr(g) + H 2 SO 4 (l)  SO 2 (g) + Br 2 (g) + 2H 2 O(l) Bromide: - Overall reaction : 2  (1) + (2) Not suitable for preparing HBr warm 2KBr(s) + 3H 2 SO 4 (l)  2KHSO 4 (s) + SO 2 (g) + Br 2 (g) + 2H 2 O(l) warm

109 108 A brown gas is evolved on warming A pungent smell is detected A brown colour is observed when adding hexane Br 2 It turns orange dichromate solution green SO 2 Confirmatory Test Dense white fumes are formed with aqueous ammonia HBrWhite fumes are formed Br – ProductObservationHalide 2KBr(s) + 3H 2 SO 4 (l)  2KHSO 4 (s) + SO 2 (g) + Br 2 (g) + 2H 2 O(l)

110 109 iodide: - KI(s) + H 2 SO 4 (l)  KHSO 4 (s) + HI(g) warm 2HI(g) + H 2 SO 4 (l)  SO 2 (g) + I 2 (g) + 2H 2 O(l) warm 8HI(g) + H 2 SO 4 (l)  H 2 S(g) + 4I 2 (g) + 2H 2 O(l) warm HI is strong enough to reduce sulphur to its lowest possible oxidation state

111 110 KI(s) + H 2 SO 4 (l)  KHSO 4 (s) + HI(g) (1) warm 2HI(g) + H 2 SO 4 (l)  SO 2 (g) + I 2 (s) + 2H 2 O(l) (2) warm 8HI(g) + H 2 SO 4 (l)  H 2 S(g) + 4I 2 (s) + 2H 2 O(l) (3) warm Overall reaction = 10  (1) + (2) + (3) 10KI(s) + 12H 2 SO 4 (l)  10KHSO 4 (s) + SO 2 (g) + H 2 S(g) + 5I 2 (s) + 4H 2 O(l) No suitable for preparing HI

112 111 10KI(s) + 12H 2 SO 4 (l)  10KHSO 4 (s) + SO 2 (g) + H 2 S(s) + 5I 2 (s) + 4H 2 O(l) Observation : - A bad egg smell is detected Confirmatory test : - It turns lead(II) ethanoate paper black (CH 3 COO) 2 Pb + H 2 S  PbS(s) + 2CH 3 COOH

113 112 10KI(s) + 12H 2 SO 4 (l)  10KHSO 4 (s) + SO 2 (g) + H 2 S(s) + 5I 2 (s) + 4H 2 O(l) Observation : - Confirmatory test : - A violet colour is observed when added to hexane Violet fumes are formed and condense when cooled to give a black solid

114 113 Conclusion : - Reducing power : HI > HBr > HCl > HF Increases down the group

115 114 Interpretation:- Consider the reaction, 2H–X + H 2 SO 4  X–X + SO 2 + 2H 2 O The feasibility of the reaction depends on 1.the strength of H–X bond to be broken the stronger the bond, the less feasible is the rx 2.the strength of X–X bond to be formed the stronger the bond, the more feasible is the rx

116 115 2H–X + H 2 SO 4  X–X + SO 2 + 2H 2 O The feasibility of the reaction depends on 1.the strength of H–X bond the stronger the bond, the less feasible is the rx 2.the strength of X–X bond the stronger the bond, the more feasible is the rx The reaction with HF is least feasible because 1.H-F bond is the strongest 2.F-F bond is exceptionally weak due to repulsion between lone pairs of bonding atoms.

117 116 H-XB.E.(kJ mol  1 )X-XB.E. (kJ mol  1 H-Cl432Cl-Cl244 H-Br366Br-Br192 H-I298I-I152 On moving down the group, both H–X bonds and X–X bonds become weaker

118 117 2H–X + H 2 SO 4  X–X + SO 2 + 2H 2 O The strength of H-X bond is more important Since two H-X bonds have to be broken for each X-X bond formed. Reactivity : H-Cl < H-Br < H-I

119 118 Reactions with Phosphoric Acid non-volatile volatile H 3 PO 4 (l) + HX(g)  no reaction less oxidizing Suitable for preparing HX from solid halids NaCl(s) + H 3 PO 4 (l)  NaH 2 PO 4 (s) + HCl(g) NaBr(s) + H 3 PO 4 (l)  NaH 2 PO 4 (s) + HBr(g) NaI(s) + H 3 PO 4 (l)  NaH 2 PO 4 (s) + HI(g) warm

120 119 Halide ion ObservationProduct Confirmatory test of the product Cl – White fumes are formed on warming HCl Dense white fumes are formed with aqueous ammonia Br – HBr I–I– HI NaCl(s) + H 3 PO 4 (l)  NaH 2 PO 4 (s) + HCl(g) NaBr(s) + H 3 PO 4 (l)  NaH 2 PO 4 (s) + HBr(g) NaI(s) + H 3 PO 4 (l)  NaH 2 PO 4 (s) + HI(g) warm

121 120 Reactions with Silver Ions Aqueous solutions of chlorides, bromides and iodides  give precipitates when reacting with acidified silver nitrate solution

122 121 Reactions with Silver Ions Ag + (aq) + Cl – (aq)  AgCl(s) white ppt Ag + (aq) + Br – (aq)  AgBr(s) pale yellow ppt Ag + (aq) + I – (aq)  AgI(s) yellow ppt

123 122 AgI(s)AgBr(s)AgCl(s) Colour intensity  down the group

124 123 Reactions with Silver Ions Silver nitrate solution should be acidified with nitric acid (a)to remove interfering ions like SO 3 2  or CO 3 2  They may form white ppt with Ag +

125 124 Reactions with Silver Ions 2H + (aq) + SO 3 2– (aq)  SO 2 (g) + H 2 O(l) 2H + (aq) + CO 3 2– (aq)  CO 2 (g) + H 2 O(l)

126 125 Silver nitrate solution should be acidified with nitric acid (b)to avoid the formation of black ppt of Ag 2 O in alkaline solution. 2Ag + (aq) + 2OH  (aq)  Ag 2 O(s) + H 2 O(l)

127 126 The solubility(in water) of AgX  down the group AgF >> AgCl > AgBr > AgI Ksp/mol 2 dm   10   10   10  16 soluble insoluble

128 127 Q.7 On moving down the group, the size of the halide anions   The electron cloud of the anions becomes more easily polarized by Ag+  The halides become more covalent and less ionic  The halides become less soluble in polar solvents like water

129 128 Reactions with Silver Ions The reaction can be used as a test to show the presence of halide ions. Different halides give ppt with different colours. Sometimes ambiguous. Confirmatory tests are needed.

130 129 Two confirmatory tests for halides 1.Adding NH 3 (aq) to the AgX ppt 2.Exposing AgX ppt to sunlight

131 130 AgX(s) dissolve in NH 3 (aq) due to the formation of soluble complex ions. AgCl(s) + 2NH 3 (aq)  [Ag(NH 3 ) 2 ] + (aq) + Cl  (aq) AgBr(s) + 2NH 3 (aq)  [Ag(NH 3 ) 2 ] + (aq) + Br  (aq) AgI(s) + 2NH 3 (aq)  No reaction Solubility in NH 3 (aq)  down the group

132 131 When exposed to sunlight light 2AgCl(s)  2Ag(s) + Cl 2 (g) 2AgBr(s)  2Ag(s) + Br 2 (l) light 2AgI(s)  No reaction light  silver bromide turns yellowish grey  silver iodide remains yellow  silver chloride turns grey

133 132 Ion Action of acidified AgNO 3 solution on halides Confirmatory test of the product Effect of adding aqueous ammonia Effect of exposure to sunlight Cl – A white ppt is formed The white ppt dissolves The solution turns grey Br – A pale yellow ppt is formed The pale yellow ppt slightly dissolves The solution turns yellowish grey I–I– A yellow ppt is formed The yellow ppt does not dissolve The solution remains yellow Action of acidified silver nitrate solution on halides

134 133 Anomalous Behaviour of Hydrogen Fluoride 1.Hydrogen fluoride has abnormally high boiling point and melting point among the hydrogen halides HXHFHClHBrHI b.p./  C19.5  85  66.4  35

135 134 Formation of the extensive intermolecular hydrogen bonds among hydrogen fluoride molecules Molecules of all other hydrogen halides  held together by weak van der Waal’s forces only

136 135 The acid strength of hydrogen halides decreases in the order: HI > HBr > HCl >> HF 2.Acidic Properties of Hydrogen Halides

137 136 Hydrogen halide Acid dissociation constant, K a (mol dm –3 ) Degree of dissociation in 0.1 M solution (%) Acid strength HF HCl HBr HI 5.6 × 10 –4 1 × × × Low Strong Very strong Acid dissociation constants of hydrogen halides and their degrees of dissociation in 0.1 M solutions

138 137 In dilute (e.g. 0.1M) solution, HF is the weakest acid among all the hydrohalic acids HF(l) + H 2 O(l) H 3 O + (aq) + F – (aq) Ka = 5.6 × 10 –4 mol dm –3

139 138 Very stable ion pair Freedom of F  & H 3 O + greatly  (a drop in entropy of the system) due to H-bond formation Effective concentration of F  & H 3 O + greatly  Thus, Ka  & pH 

140 139 In concentrated solution, HF is the strongest acid among all the hydrohalic acids

141 140 Strength of H-bond:- 2.HF is in excess in concentrated solution F  ions combine with excess HF rather than with H 3 O +  free H 3 O+  & pH  excess H 3 OF(aq) + HF(aq)  H 3 O + (aq) + HF 2  (aq)

142 141 For other HX acids, acidity  as concentration  It is due to the significant interaction between X  and H 3 O + at high concentrations  the effective concentration of H 3 O +  For HF, interaction between F  and H 3 O + is significant even at low concentrations due to the smaller size of F .

143 142 3.Pure, anhydrous liquid HF is ionic due to the formation of HF 2  and H 2 F + ions Self ionization : - 2HF(l) H 2 F + (l) + F  (l) HF(l) + F  (l) HF 2  (l) Overall : - 3HF(l) [H 2 F] + [HF 2 ]  (l)

144 143  Stabilized by resonance  Two identical H – F bonds

145 144 Heating the solid potassium hydrogen difluoride  reverses the reaction  a convenient way to obtain anhydrous hydrogen fluoride KF(s) + HF(l) KHF 2 (s) heat

146 145 Uses of fluorine and its compounds Sodium hexafluorosilicate, Na 2 SiF 6, is used in water fluoridation. F , being isoelectronic to OH , can replace the OH  in the tooth enamel, making it less soluble in acidic solutions.

147 146 Uses of fluorine and its compounds Molten cryolite, Na 3 AlF 6 Lowers the temperature (2517  C  1000  C) needed for extracting Al from Al 2 O 3 by electrolysis.

148 147 Uses of fluorine and its compounds Convert U to UF 6 Separate 235 UF 6 from 238 UF 6 by diffusion for use in nuclear reactors. The heavier 238 UF 6 diffuses a bit slower, making the separation possible.

149 148 Uses of fluorine and its compounds Conc. HF(aq) is used in etching glass (e.g. making scales/graduation marks on glassware) CaSiO 3 (s) + 6HF(aq)  CaF 2 (aq) + SiF 4 (aq) + 3H 2 O(l) (Glass)

150 149 Uses of fluorine and its compounds The glass object to be etched  coated with wax or a similar acid- proof material  cutting through the wax layer to expose the glass  apply hydrofluoric acid

151 150 Uses of fluorine and its compounds A glass is etched by hydrofluoric acid

152 151 Uses of fluorine and its compounds Making fluorocarbon compounds Used as refrigerants, aerosol propellants, anaesthetics and fire-fighting agents(BTM, BCF) PTFE (teflon) used in electrical insulation, coating on surface of non-stick saucepans, etc.

153 152 Uses of fluorine and its compounds Hydrazine/fluorine mixtures are excellent rocket fuels N 2 H 4 (g) + 2F 2 (g)  N 2 (g) + 4HF(g)  H = kJ mol  1 (extremely exothermic) Due to the strong N  N and H-F bonds

154 153 Uses of fluorine and its compounds Extraction of fluorine Electrolyte : KF(s) dissolved in pure HF(l) Anode : graphite Cathode : steel

155 154 Q.8(a) Anode : 2HF 2   2HF + F 2 + 2e  Cathode : 2H 2 F + + 2e   2HF + H 2 Overall : 2HF 2  + 2H 2 F +  4HF + F 2 + H 2

156 155 8.(b)Overall : 2HF 2  + 2H 2 F +  4HF + F 2 + H 2 6HF  4HF + F 2 + H 2 2HF  F 2 + H 2 KF is added to increase the conductivity of the electrolyte. KHF 2 > HF or [H 2 F][HF 2 ]

157 156 Q.8(c) OH - (from H 2 O) rather than HF 2  is oxidized at the anode 2F 2 (g) + 2H 2 O(l)  4HF(aq) + O 2 (g) vigorous reaction Also, F 2 reacts vigorously with water.

158 157 8.(d) At high temperatures, fluorine produced can react vigorously with the electrodes, air, etc.

159 158 Uses of Chlorine and its compounds Polyvinyl chloride, PVC making electrical insulation, bottles, floor tiles, table cloth, shower curtain, etc.

160 159 CH 2 =CH 2 + Cl 2  CH 2 Cl – CH 2 Cl CH 2 Cl – CH 2 Cl CH 2 =CHCl + HCl heat n(CH 2 =CHCl) 

161 160 Making chlorine bleach Cl 2 (g) + 2NaOH(aq)  NaCl(aq) + NaOCl + H 2 O(l) Disinfectant in sterilizing water and sewage treatment. Extraction of bromine from sea water Cl 2 (g) + 2Br  (aq)  2Cl  (aq) + Br 2 (aq)

162 161 Uses of Bromine and its compounds Manufacture of 1,2-dibromoethane to remove Pb from petrol engine Pb(C 2 H 5 ) 4, TEL : anti-knock agent added to petrol engine to prevent premature ignition. TEL decomposes to give Pb that may cause damage to the engine CH 2 Br-CH 2 Br + Pb(C 2 H 5 ) 4  PbBr 2 volatile and emitted to air easily Air pollutant

163 162 AgBr is used in black-and-white photography exposure to light 2AgBr(s) 2Ag(s) + Br 2 (l) coated on filmblack The excess AgBr(s) is removed as soluble complex ion. AgBr(s) + 2S 2 O 3 2  (aq)  [Ag(S 2 O 3 ) 2 ] 3  (aq) + Br  (aq) hypo

164 163 Uses of Iodine and its compounds Making iodine tincture (antiseptic) I 2 in alcohol or KI(aq) Radioactive iodine-131 as tracer in medical diagnosis Iodide is used to make iodized table salt for preventing development of goitre.

165 164 Laboratory preparation of halogens(except F 2 ) conc. H 2 SO 4 MnO 2 + NaCl

166 165 2NaCl + MnO 2 + 2H 2 SO 4  Na 2 SO 4 + MnSO 4 + 2H 2 O + Cl 2 conc. H 2 SO 4 MnO 2 + NaCl Free from HCl and H 2 O NaCl + H 2 SO 4  HCl + NaHSO 4 To remove HClTo dry Cl 2

167 166 Laboratory preparation of halogens(except F 2 ) conc. HCl MnO 2

168 167 Laboratory preparation of halogens(except F 2 ) conc. HCl MnO 4 

169 168 The END


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