Characteristic Properties of the Halogens 1 Introduction • Group VIIA elements include fluorine chlorine bromine halogens iodine (Salt producers) astatine 2 Introduction • 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. 3 Halogens are p-block elements • outermost shell electronic configuration of ns2np5 4 Halogens are p-block elements • one electron short of the octet structure 5 Introduction • In the free elemental state they complete their octets by sharing their single unpaired p-electrons 6 When halogens react with other elements they either gain an additional electron to form halide ions or share their single unpaired p-electrons to form single covalent bonds 7 High Electronegativity / Electron Affinity highest among the elements in the same period have a high tendency to attract electrons strong oxidizing agents 8 High Electronegativity / Electron Affinity -1 is the most common oxidation state of halogens in their compounds 9 Ionic : NaF, NaCl, NaBr, NaI Covalent : HF, HCl, HBr, HI Variable Oxidation State All halogens (except fluorine) can expand their octet of electrons by utilizing the vacant, energetically low-lying d-orbitals. 10 11 “Electrons-in-boxes” diagrams of the electronic configuration of a halogen atom of the ground state and various excited states +3 +1 O H H Cl Cl O O The half-filled orbital(s) overlap(s) with those of more electronegative atoms (e.g. O) positive oxidation state (+1, +3, +5, +7) O O H O 12 Cl +5 +7 Cl H O O O O Various oxidation states of halogens in their ions or compounds Oxidation state of halogen –1 Ion / Compound F– Cl– Br– I– HF HCl HBr HI Cl2 Br2 I2 OF2 0 +1 +3 13 F2 Cl2O Br2O HOCl HOBr OCl– OBr– HClO2 ClO2– Various oxidation states of halogens in their ions or compounds Oxidation state of halogen +4 +5 Ion / Compound ClO2 BrO2 HClO3 HBrO3 I 2O 5 ClO3– BrO3– HIO3 IO3– +6 +7 14 Cl2O6 BrO3 Cl2O7 H5IO6 HClO4 HIO4 ClO4– IO4– • Fluorine (1) the most electronegative element only one unpaired p electron available for bonding oxidation state is limited to –1 15 • 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, HFO2, HFO3, HFO4 16 Variation in Physical Properties 1. Melting point / boiling point down the group 17 Halogen Melting point (C) Boiling point (C) Fluorine –220 –188 Chlorine –101 –34.7 Bromine –7.2 58.8 Iodine 114 184 Astatine 302 380 Variations in melting point and boiling point of the halogens 18 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 19 2. Colour becomes darker down the group Halogen Colour 20 F2 Pale yellow Cl2 Br2 Greenish Reddish yellow brown I2 Violet black chlorine Appearances of halogens at room temperature and pressure: chlorine 21 bromine Appearances of halogens at room temperature and pressure: bromine 22 iodine Appearances of halogens at room temperature and pressure: iodine 23 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 24 Colour • Fluorine atom has the smallest size absorbs the radiation of relatively high frequency (i.e. blue light) appears yellow (the unabsorbed radiation) 25 Colour • Atoms of other halogens larger sizes absorb radiation of lower frequency 26 Colour • Iodine absorbs the radiation of relatively low frequency (i.e. yellow light) appears violet 27 Q.1 The colour of astatine is black. 28 Colour • Halogens different colours when dissolved in different solvents 29 Colours of halogens in pure form and in solutions Colour Halogen in pure form in water in 1,1,1-trichloroethane F2 Pale yellow Pale yellow Pale yellow Cl2 Greenish yellow Pale yellow Yellow Br2 Reddish brown Yellow Orange Violet black Yellow (only slightly soluble) Violet I2 Brown in KI(aq) 30 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) 31 (a) (b) (c) Colours of halogens in water: (a) chlorine; (b) bromine; (c) iodine 32 (a) (b) (c) Colours of halogens in 1,1,1-trichloroethane: (a) chlorine; (b) bromine; (c) iodine 33 3. Electron Affinity down the group Halogen E.A. kJ/mol1 34 F Cl Br I At -322 -349 -335 -295 -270 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 35 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. 36 4. Electronegativity down the group 37 Halogen F Cl Br I At Electronegativity 4.0 3.0 2.8 2.5 2.2 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 38 Fluorine has the highest electronegativity because it is the most reactive elements. The electronegativity of fluorine is arbitrarily assigned as 4.0. 39 Variation in Chemical Properties Reactivity : F2 > Cl2 > Br2 > I2 React by gaining electrons Oxidizing power : F2 > Cl2 > Br2 > I2 40 1. Reactions with Sodium • All halogens combine directly with sodium to form sodium halides the reactivity decreases down the group from fluorine to iodine 41 1. Reactions with Sodium • Fluorine react explosively to form sodium fluoride 2Na(s) + F2(g) 2NaF(s) 42 1. Reactions with Sodium • Chlorine reacts violently to form sodium chloride 2Na(s) + Cl2(g) 2NaCl(s) 43 1. Reactions with Sodium • Bromine burns steadily in bromine vapour to form sodium bromide 2Na(s) + Br2(g) 2NaBr(s) 44 1. Reactions with Sodium • Iodine burns steadily in iodine vapour to form sodium iodide 2Na(s) + I2(g) 2NaI(s) 45 Na+(g) + X(g) Vigor of reaction depends on 1 +(g) + X(g) 1. The activation energy (endothermic) B.E. Na 2 Na+(g) + 1 2 X2(g) 2. The lattice energy (exothermic) H o atm Activation energy Hovap Na(s) + 1 2 X2 Hof 46 Holattice I.E. NaX(s) E.A. Na+(g) + X(g) 1 2 B.E. F is the most reactive Na+(g) + 1 2 X2(g) Hoatm Na(s) + Hof 47 Na+(g) + X(g) F has an exceptionally low B.E. & zero Hovap Hovap 1 2 X2(g) E.A. Holattice I.E. NaX(s) Na+(g) + X(g) 1 2 B.E. E.A. Na+(g) + Na+(g) + X(g) 1 2 X2(g) H o atm The lattice enthalpy of NaF is most negative Hovap Na(s) + 1 2 X2 Hof 48 Holattice I.E. NaX(s) Na+(g) + X(g) 1 2 B.E. Cl is more reactive than Br & I E.A. Na+(g) + X(g) 1 Na+(g) + 2 X2(g) Cl has zero Hovap Hoatm Hovap Na(s) + 1 2 X2(g) Hof 49 Holattice I.E. NaX(s) Na+(g) + X(g) 1 2 B.E. E.A. Na+(g) + Na+(g) + X(g) 1 2 X2(g) Lattice enthalpy : Hoatm NaCl > NaBr > NaI Hovap Ho lattice Na(s) + 1 2 X2 Hof 50 I.E. NaX(s) Na+(g) + X(g) 1 2 B.E. Br is more reactive than I Na+(g) + Na+(g) + X(g) 1 2 X2(g) Hoatm Hovap : Br2(l) < I2(s) Hovap Na(s) + 1 2 X2(s)/(l) Hof 51 E.A. Holattice I.E. NaX(s) Na+(g) + X(g) 1 2 B.E. E.A. Na+(g) + Na+(g) + X(g) 1 2 X2(g) Lattice enthalpy : Hoatm NaBr > NaI Hovap Na(s) + 1 2 X2 Hof 52 Holattice I.E. NaX(s) Q.2(a) Variation: bond enthalpy decreases from Cl2 to I2 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. F2 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. 53 Q.2(b) The lattice enthalpy becomes less negative down the group. It is because the anionic radius, r- , increases down the group. H 54 o lattice 1 r r 2.1 Reactions with hydrogen X2 + H2(g) 2HX(g) F2 reacts explosively even in the dark at 200C Cl2 reacts explosively in sunlight Br2 reacts moderately on heating with a catalyst I2 reacts slowly and reversibly even on heating 55 Q.3 Explain the extreme reactivity of fluorine in terms of the bond enthalpies of F–F and H–F bonds. F2 + H2(g) 2HF(g) 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 H2 and F2 is the most exothermic. The energy released from the reaction further speeds up the reaction. 56 Chlorine removes hydrogen completely from turpentine(C10H16) C10H16(l) + 8Cl2(g) 10C(s) + 16HCl(g) 57 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. 58 2.2 Reactions with phosphorus F2 + P PF5 Cl2 + P PCl3 + PCl5 Br2 + P PBr3 I2 + P PI3 F2 is the strongest oxidizing agent, it always oxidizes other elements to their highest possible oxidation states. 59 2.2 Reactions with phosphorus F2 + P PF5 Cl2 + P PCl3 + PCl5 Br2 + P PBr3 I2 + P PI3 Br2 and I2 are NOT strong enough to oxidize P to its highest possible oxidation state. 60 2.3 Reactions 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) + 2F2 CF4 Xe + F2 XeF2 Xe + 2F2 XeF4 Xe + 3F2 XeF6 61 2.3 Reactions with xenon It is because (a) Xenon can expand its octet by utilizing vacant, low-lying d-orbitals. 62 By VB Theory, 5s Xe 5p 2s F 2p To form two Xe-F bonds in XeF2, a 5p electron in Xe has to be promoted to a 5d orbital. 5s Xe* 63 5d 5p By VB Theory, Xe 5d 5s 5p To form four Xe-F bonds in XeF4, two 5p electrons in Xe have to be promoted to two 5d orbitals. 5d 5s 5p Xe** 64 By VB Theory, Xe 5d 5s 5p To form six Xe-F bonds in XeF6, three 5p electrons in Xe have to be promoted to three 5d orbitals. 5d 5s 5p Xe*** 65 2.3 Reactions with xenon The gap between np and nd sub-shells down the group, thus, the promotion of electrons from np sub-shell to nd sub-shell becomes easier down the group. Tendency to form bonds down the group : Xe > Kr > Ar > Ne > He 66 Xe 5d 5s 5p 5s Xe*** 5d 5p Also, the energy released by forming more single bonds outweighs the energy required for promoting 5p electrons to 5d orbitals. 67 3 Reactions with other reducing agents I2 is the weakest oxidizing agents among the halogens. 68 3.1 All halogens(except I2) oxidize Fe2+ to Fe3+ Half reaction Standard electrode potential (V) Cl2(aq) + 2e– 2Cl–(aq) +1.36 Br2(aq) + 2e – 2Br–(aq) +1.07 Fe3+(aq) + e– Fe2+(aq) +0.77 I2(aq) + 2e– 2I–(aq) +0.54 X2(aq) + 2Fe2+(aq) 2X(aq) + 2Fe3+(aq) o 0 ( X = F, Cl, Br) Ecell 69 3.1 All halogens(except I2) oxidize Fe2+ to Fe3+ Half reaction Standard electrode potential (V) Cl2(aq) + 2e– 2Cl–(aq) +1.36 Br2(aq) + 2e – 2Br–(aq) +1.07 Fe3+(aq) + e– Fe2+(aq) +0.77 I2(aq) + 2e– 2I–(aq) +0.54 o I2(aq) + 2Fe2+(aq) No reaction Ecell 0 70 3.2 All halogens(except I2) oxidize S2O32 to SO42 4X2(aq) + S2O32(aq) + 5H2O(l) 8X(aq) + 10H+(aq) + 2SO42(aq) (X = F, Cl, Br) S0 I2(aq) + 2S2O32(aq) 2I(aq) + S4O62(aq) Used in iodometric titration O O- O 0 +5 S O 0 S S +5 O +6 S O - O OO- O S O 71 S +4 OO- Determination of [Fe3+(aq)] by iodometric titration (i) (ii) 2I(aq) + 2Fe3+(aq) I2(aq) + 2Fe2+(aq) (excess) (unknown) I2(aq) + 2S2O32(aq) 2I(aq) + S4O62(aq) (standard solution) nFe3 : nS O 2 1 : 1 2 3 Using starch as indicator 72 4 Displacement reactions Cl2(aq) + 2Br(aq) 2Cl(aq) + Br2(aq) Cl2(aq) + 2I(aq) 2Cl(aq) + I2(aq) Br2(aq) + 2I(aq) 2Br(aq) + I2(aq) More reactive Less reactive I2(aq) + I(aq) I3(aq) (yellow) 73 (brown) 4 Displacement reactions Cl2(aq) + 2I(aq) 2Cl(aq) + I2(aq) Br2(aq) + 2I(aq) 2Br(aq) + I2(aq) I2(aq) + I(aq) I3(aq) (yellow) (brown) What would be observed if an excess of Cl2(aq) or Br2(aq) is added to I(aq)? 74 The solution turns cloudy and a black solid settles at the bottom Reactions of halide ions with halogens Aqueous solution 75 Halogen added F2 F– No reaction Cl– A pale yellow solution is formed (Cl2 is formed) Cl2 Br2 I2 No reaction No reaction No reaction No reaction No reaction No reaction Reactions of halide ions with halogens Aqueous solution Halogen added F2 A yellow solution Br– is formed (Br2 is formed) I– 76 Cl2 A yellow solution is formed (Br2 is formed) Br2 I2 No reaction No reaction A yellowish brown A yellowish brown A yellowish brown solution is formed solution is formed solution is formed (I3 is formed) (I3 is formed) (I3 is formed) No reaction Q.5 Shake hexane or 1,1,1-trichloroethane with the two solutions respectively. The one that turns the organic layer violet is I3(aq). The one that turns the organic layer orange or brown is Br2(aq). 77 If hexane is used, the upper layer will be the organic layer Br2(aq) I3(aq) 1,1,1-trichloroethane Br2 78 I2 5. Disproportionation 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. 79 A. Reactions with Water HOCl : chloric(I) acid or hypochlorous acid Chlorine water 80 a mixture of hydrochloric acid and chloric(I) acid A. Reactions with Water • 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) + O2(g) 81 A. Reactions with Water • Chlorate(I) ion bleaches by oxidation Cl2(aq) + H2O(l) 2H+(aq) + Cl–(aq) + OCl–(aq) OCl–(aq) + dye Cl–(aq) + (dye + O) coloured 82 colourless A. Reactions with Water • Bromine only slightly soluble in water mainly exists as molecules in saturated bromine water 83 A. Reactions with Water • When the solution is diluted hydrolysis takes place hydrobromic acid and bromic(I) acid (hydrobromous acid) are formed Br2(l) + H2O(l) 84 HBr(aq) + HOBr(aq) A. Reactions with Water • Bromate(I) ion, OBr also unstable bleaches dyes by oxidation OBr–(aq) + dye coloured Br–(aq) + (dye + O) colourless 85 A. Reactions with Water • Iodine does not react with water only slightly soluble in water 86 A. Reactions with Water • Fluorine reacts vigorously with water to form hydrogen fluoride and oxygen 0 1 2F2(g) + 2H2O(l) 4HF(aq) + O2(g) Being the strongest oxidizing agent, F2 undergoes reduction rather than disproportionation with water. 87 A. Reactions with Water Chlorine reacts similarly at high temperature or when exposed to light 2Cl2(aq) + 2H2O(l) 2HCl(aq) + 2HOCl(aq) 2HOCl(aq) Heat or light 2HCl(aq) + O2(g) Overall : 2Cl2(aq) + 2H2O(l) 88 Heat or light 4HCl(aq) + O2(g) B. Reactions with Alkalis • All halogens react with aqueous alkalis • All halogens (except F2) undergoes disproportionation with alkalis • In general, Reactivity decreases down the group 89 B. Reactions with Alkalis The products formed depend on 1. Temperature 2. The type of halogen reacted 3. The concentration of alkali used 90 B. Reactions with Alkalis Effect of temperature (a) At lower temperatures, T1 0 X2(aq) + 2OH(aq) 91 1 +1 XO(aq) + X(aq) + H2O(l) X2 Cl2 Br2 I2 T1 / C 20 0 <0 B. Reactions with Alkalis Effect of temperature (a) At higher temperatures, +1 3XO(aq) 92 T2 +5 1 XO3(aq) + 2X(aq) XO ClO BrO IO T2 / C 70 20 0 B. Reactions with Alkalis (1) X2(aq) + (2) 3XO(aq) T1 2OH(aq) T2 XO(aq) + X(aq) + H2O(l) XO3(aq) + 2X(aq) Overall reaction : 3(1) + (2) 3X2(aq) + 93 6OH(aq) T2 XO3(aq) + 5X(aq) + 3H2O(l) B. Reactions with Alkalis 3XO(aq) T2 XO3(aq) + 2X(aq) XO ClO BrO IO T2 / C 70 20 <0 On moving down the group, 94 1. stability of XO decreases ClO > BrO > IO 2. stability of XO3 increases ClO3 < BrO3 < IO3 B. Reactions with Alkalis 3X2(aq) + 6OH(aq) XO3(aq) + 5X(aq) + 3H2O(l) At lower pH (when acid is added), the equilibrium position shifts to the left and the reversed process predominates. XO3(aq) + 5X(aq) + 6H+(aq) 3X2(aq) + 3H2O(l) This reaction (when X=I) is often used to prepare standard iodine solution for iodometric titrations 95 B. Reactions with Alkalis • Dissolving a known quantity of KIO3(s) in excess KI(aq) and dilute H2SO4 generates a known amount of I2(aq) KIO3(aq) + 5KI(aq) + 6H+(aq) 3I2(aq) + 3H2O(l) + 6K+(aq) • The iodine produced can be used to standardize thiosulphate solution 3I2(aq) + 6S2O32(aq) 6I(aq) + 3S4O62(aq) 96 B. Reactions with Alkalis • This known amount of iodine generated can also be used to oxidize reducing agents (of unknown concentrations) such as SO32(aq) and ascorbic acid (vitamin C) • The excess iodine can be determined by back titration with sodium thiosulphate solution I2(aq) + 2S2O32–(aq) 2I–(aq) + S4O62–(aq) 97 B. Reactions with Alkalis Effect of concentration of alkali (a) At higher concentrations, XO3(aq) is the major product. (b)At lower concentrations, XO(aq) is the major product. 98 B. Reactions with Alkalis In general, Halogens react with cold, dilute alkali to give halate(I) ions, halide ions and water X2(aq) + 2OH XO(aq) + X(aq) + H2O(l) Halogens react with hot, concentrated alkali to give halate(V) ions, halide ions and water. 3X2(aq) + 6OH XO3(aq) + 5X(aq) + 3H2O(l) 99 B. Reactions with Alkalis 0 2 2F2 + 2OH(aq) 20C +2 1 1 OF2(aq) + 2F(aq) + H2O(l) very dilute 0 2 2F2 + 4OH(aq) 70C 0 1 O2(aq) + 4F(aq) + 2H2O(l) concentrated Being the strongest oxidizing agent, F2 undergoes reduction rather than disproportionation with alkalis. 100 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 101 Reactions with Concentrated Sulphuric(VI) Acid • Concentrated sulphuric acid non-volatile (b.p. ~330C) oxidizing 102 Fluoride and chloride : warm KF(s) + H2SO4(l) KHSO4(s) + HF(g) warm KCl(s) + H2SO4(l) KHSO4(s) + HCl(g) non-volatile volatile Warming is required to speed up the reaction and to drive out the volatile acids 103 warm KF(s) + H2SO4(l) KHSO4(s) + HF(g) warm KCl(s) + H2SO4(l) KHSO4(s) + HCl(g) acid salt Acid salt rather than normal salt is formed because HSO4 is a relatively weak acid A convenient way to prepare HCl in the laboratory 104 warm KF(s) + H2SO4(l) KHSO4(s) + HF(g) warm KCl(s) + H2SO4(l) KHSO4(s) + HCl(g) Observation : White fumes are produced Confirmatory test : Dense white fumes appear with NH3(aq) 105 Bromide: warm KBr(s) + H2SO4(l) KHSO4(s) + HBr(g) reduction -1 +6 warm +4 0 2HBr(g) + H2SO4(l) SO2(g) + Br2(g) + 2H2O(l) oxidation 106 Bromide: warm (1) KBr(s) + H2SO4(l) KHSO4(s) + HBr(g) warm (2) 2HBr(g) + H2SO4(l) SO2(g) + Br2(g) + 2H2O(l) Overall reaction : 2(1) + (2) 2KBr(s) + 3H2SO4(l) warm 2KHSO4(s) + SO2(g) + Br2(g) + 2H2O(l) Not suitable for preparing HBr 107 2KBr(s) + 3H2SO4(l) 2KHSO4(s) + SO2(g) + Br2(g) + 2H2O(l) Halide Br– Observation Product • White fumes are formed HBr • Dense white fumes are formed with aqueous ammonia • A pungent smell is SO2 • It turns orange dichromate solution detected • A brown gas is evolved on warming 108 Confirmatory Test green Br2 • A brown colour is observed when adding hexane iodide: warm KI(s) + H2SO4(l) KHSO4(s) + HI(g) -1 +6 warm -1 +6 warm +4 0 2HI(g) + H2SO4(l) SO2(g) + I2(g) + 2H2O(l) -2 0 8HI(g) + H2SO4(l) H2S(g) + 4I2(g) + 2H2O(l) HI is strong enough to reduce sulphur to its lowest possible oxidation state 109 warm KI(s) + H2SO4(l) KHSO4(s) + HI(g) (1) warm 2HI(g) + H2SO4(l) SO2(g) + I2(s) + 2H2O(l) (2) warm 8HI(g) + H2SO4(l) H2S(g) + 4I2(s) + 2H2O(l) (3) Overall reaction = 10(1) + (2) + (3) 10KI(s) + 12H2SO4(l) 10KHSO4(s) + SO2(g) + H2S(g) + 5I2(s) + 4H2O(l) No suitable for preparing HI 110 10KI(s) + 12H2SO4(l) 10KHSO4(s) + SO2(g) + H2S(s) + 5I2(s) + 4H2O(l) Observation : A bad egg smell is detected Confirmatory test : It turns lead(II) ethanoate paper black (CH3COO)2Pb + H2S PbS(s) + 2CH3COOH 111 10KI(s) + 12H2SO4(l) 10KHSO4(s) + SO2(g) + H2S(s) + 5I2(s) + 4H2O(l) Observation : Violet fumes are formed and condense when cooled to give a black solid Confirmatory test : A violet colour is observed when added to hexane 112 Conclusion : - Increases down the group Reducing power : HI > HBr > HCl > HF 113 Interpretation:Consider the reaction, 2H–X + H2SO4 X–X + SO2 + 2H2O 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 114 2H–X + H2SO4 X–X + SO2 + 2H2O 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. 115 H-X B.E.(kJ mol1) X-X B.E. (kJ mol1 H-Cl 432 Cl-Cl 244 H-Br 366 Br-Br 192 H-I 298 I-I 152 On moving down the group, both H–X bonds and X–X bonds become weaker 116 2H–X + H2SO4 X–X + SO2 + 2H2O 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 117 Reactions with Phosphoric Acid warm NaCl(s) + H3PO4(l) NaH2PO4(s) + HCl(g) warm NaBr(s) + H3PO4(l) NaH2PO4(s) + HBr(g) warm NaI(s) + H3PO4(l) NaH2PO4(s) + HI(g) non-volatile volatile H3PO4(l) + HX(g) no reaction less oxidizing Suitable for preparing HX from solid halids 118 warm NaCl(s) + H3PO4(l) NaH2PO4(s) + HCl(g) warm NaBr(s) + H3PO4(l) NaH2PO4(s) + HBr(g) warm NaI(s) + H3PO4(l) NaH2PO4(s) + HI(g) Halide ion Cl– Br– I– 119 Observation White fumes are formed on warming Confirmatory test of Product the product HCl HBr HI Dense white fumes are formed with aqueous ammonia Reactions with Silver Ions • Aqueous solutions of chlorides, bromides and iodides give precipitates when reacting with acidified silver nitrate solution 120 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) 121 AgI(s) yellow ppt AgCl(s) AgBr(s) AgI(s) Colour intensity down the group 122 Reactions with Silver Ions Silver nitrate solution should be acidified with nitric acid (a) to remove interfering ions like SO32 or CO32 They may form white ppt with Ag+ 123 Reactions with Silver Ions 2H+(aq) + SO32–(aq) SO2(g) + H2O(l) 2H+(aq) + CO32–(aq) CO2(g) + H2O(l) 124 Silver nitrate solution should be acidified with nitric acid (b) to avoid the formation of black ppt of Ag2O in alkaline solution. 2Ag+(aq) + 2OH(aq) Ag2O(s) + H2O(l) 125 The solubility(in water) of AgX down the group soluble AgF insoluble >> Ksp/mol2 dm6 126 AgCl 1.61010 > AgBr 7.71013 > AgI 1.51016 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 127 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. 128 Two confirmatory tests for halides 1.Adding NH3(aq) to the AgX ppt 2.Exposing AgX ppt to sunlight 129 AgX(s) dissolve in NH3(aq) due to the formation of soluble complex ions. AgCl(s) + 2NH3(aq) [Ag(NH3)2]+(aq) + Cl(aq) AgBr(s) + 2NH3(aq) [Ag(NH3)2]+(aq) + Br(aq) AgI(s) + 2NH3(aq) No reaction Solubility in NH3(aq) down the group 130 • When exposed to sunlight silver chloride turns grey light 2AgCl(s) 2Ag(s) + Cl2(g) silver bromide turns yellowish grey light 2AgBr(s) 2Ag(s) + Br2(l) silver iodide remains yellow light 2AgI(s) No reaction 131 Action of acidified silver nitrate solution on halides Confirmatory test of the product Action of acidified Effect of Ion Effect of adding AgNO3 solution exposure aqueous ammonia on halides to sunlight Cl– Br– I– 132 A white ppt is formed The white ppt dissolves The solution turns grey A pale yellow ppt is formed The pale yellow ppt slightly dissolves The solution turns yellowish grey A yellow ppt is formed The yellow ppt does not dissolve The solution remains yellow Anomalous Behaviour of Hydrogen Fluoride 1. Hydrogen fluoride has abnormally high boiling point and melting point among the hydrogen halides 133 HX HF HCl HBr HI b.p./C 19.5 85 66.4 35 • Molecules of all other hydrogen halides held together by weak van der Waal’s forces only Formation of the extensive intermolecular hydrogen bonds among hydrogen fluoride molecules 134 2. Acidic Properties of Hydrogen Halides • The acid strength of hydrogen halides decreases in the order: HI > HBr > HCl >> HF 135 Acid dissociation constants of hydrogen halides and their degrees of dissociation in 0.1 M solutions Hydrogen halide Acid dissociation constant, Ka (mol dm–3) Degree of dissociation in 0.1 M solution (%) Acid strength HF 5.6 × 10–4 8.5 Low HCl 1 × 107 92 Strong HBr 1 × 109 93 Strong HI 1 × 1011 95 Very strong 136 HF(l) + H2O(l) H3O+(aq) + F–(aq) Ka = 5.6 × 10–4 mol dm–3 In dilute (e.g. 0.1M) solution, HF is the weakest acid among all the hydrohalic acids 137 H-bond H H O H + F O H F H H or H3O + F H3OF Very stable ion pair Freedom of F & H3O+ greatly (a drop in entropy of the system) due to H-bond formation Effective concentration of F & H3O+ greatly Thus, Ka & pH 138 In concentrated solution, HF is the strongest acid among all the hydrohalic acids 139 Strength of H-bond:H F H F > O H F H 2. HF is in excess in concentrated solution F ions combine with excess HF rather than with H3O+ free H3O+ & pH H3OF(aq) + HF(aq) H3O+(aq) + HF2(aq) excess 140 For other HX acids, acidity as concentration It is due to the significant interaction between X and H3O+ at high concentrations the effective concentration of H3O+ For HF, interaction between F and H3O+ is significant even at low concentrations due to the smaller size of F. 141 3. Pure, anhydrous liquid HF is ionic due to the formation of HF2 and H2F+ ions Self ionization : 2HF(l) HF(l) + F(l) H2F+(l) + F(l) HF2(l) Overall : - 3HF(l) 142 [H2F]+[HF2](l) F H F F H F Stabilized by resonance Two identical H – F bonds 143 KF(s) + HF(l) • heat KHF2(s) Heating the solid potassium hydrogen difluoride reverses the reaction a convenient way to obtain anhydrous hydrogen fluoride 144 Uses of fluorine and its compounds Sodium hexafluorosilicate, Na2SiF6, 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. 145 Uses of fluorine and its compounds Molten cryolite, Na3AlF6 Lowers the temperature (2517C 1000 C) needed for extracting Al from Al2O3 by electrolysis. 146 Uses of fluorine and its compounds Convert U to UF6 Separate 235UF6 from 238UF6 by diffusion for use in nuclear reactors. The heavier 238UF6 diffuses a bit slower, making the separation possible. 147 Uses of fluorine and its compounds Conc. HF(aq) is used in etching glass (e.g. making scales/graduation marks on glassware) CaSiO3(s) + 6HF(aq) (Glass) 148 CaF2(aq) + SiF4(aq) + 3H2O(l) Uses of fluorine and its compounds • The glass object to be etched coated with wax or a similar acidproof material cutting through the wax layer to expose the glass apply hydrofluoric acid 149 Uses of fluorine and its compounds A glass is etched by hydrofluoric acid 150 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. 151 Uses of fluorine and its compounds Hydrazine/fluorine mixtures are excellent rocket fuels N2H4(g) + 2F2(g) N2(g) + 4HF(g) H = -1166 kJ mol1 (extremely exothermic) Due to the strong NN and H-F bonds 152 Uses of fluorine and its compounds Extraction of fluorine Electrolyte : KF(s) dissolved in pure HF(l) Anode : graphite Cathode : steel 153 Q.8(a) Anode : 2HF2 2HF + F2 + 2e Cathode : 2H2F+ + 2e 2HF + H2 Overall : 2HF2 + 2H2F+ 4HF + F2 + H2 154 8.(b) Overall : 2HF2 + 2H2F+ 4HF + F2 + H2 6HF 4HF + F2 + H2 2HF F2 + H2 KF is added to increase the conductivity of the electrolyte. KHF2 > HF or [H2F][HF2] 155 Q.8(c) OH- (from H2O) rather than HF2 is oxidized at the anode Also, F2 reacts vigorously with water. 2F2(g) + 2H2O(l) 4HF(aq) + O2(g) vigorous reaction 156 8.(d) At high temperatures, fluorine produced can react vigorously with the electrodes, air, etc. 157 Uses of Chlorine and its compounds Polyvinyl chloride, PVC making electrical insulation, bottles, floor tiles, table cloth, shower curtain, etc. 158 CH2=CH2 + Cl2 CH2Cl – CH2Cl CH2Cl – CH2Cl heat CH2=CHCl + HCl Cl n(CH2=CHCl) 159 C H2 C H n Making chlorine bleach Cl2(g) + 2NaOH(aq) NaCl(aq) + NaOCl + H2O(l) Disinfectant in sterilizing water and sewage treatment. Extraction of bromine from sea water Cl2(g) + 2Br(aq) 2Cl(aq) + Br2(aq) 160 Uses of Bromine and its compounds Manufacture of 1,2-dibromoethane to remove Pb from petrol engine Pb(C2H5)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 Air pollutant CH2Br-CH2Br + Pb(C2H5)4 PbBr2 161 volatile and emitted to air easily AgBr is used in black-and-white photography exposure to light coated on film 2AgBr(s) 2Ag(s) + Br2(l) black The excess AgBr(s) is removed as soluble complex ion. AgBr(s) + 2S2O32(aq) [Ag(S2O3)2]3(aq) + Br(aq) hypo 162 Uses of Iodine and its compounds Making iodine tincture (antiseptic) I2 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. 163 Laboratory preparation of halogens(except F2) conc. H2SO4 MnO2 + NaCl 164 -1 +4 2NaCl + MnO2 + 2H2SO4 +2 0 Na2SO4 + MnSO4 + 2H2O + Cl2 NaCl + H2SO4 HCl + NaHSO4 conc. H2SO4 Free from HCl and H2O MnO2 + NaCl 165 To remove HCl To dry Cl2 Laboratory preparation of halogens(except F2) conc. HCl MnO2 166 Laboratory preparation of halogens(except F2) conc. HCl MnO4 167 The END 168