H2 Chem Summary of Group VII

Summary of Periodic Table – Group VII Halogens (X2) Group VII element Bonding Structure Colour & physical state (at r.t.p) Colour in gaseous state Colour in Organic Solvent (e.g. CCl4) Volatility or Boiling point

Oxidising strength

Displacement reaction Reaction with thiosulphate (S2O32-) Reaction with H2 gas Reaction with NaOH

Cl2 Br2 I2 Covalent bonding (X-X) within the molecule and weak van der waals’ between the molecules (X2 ……… X2) Simple molecular structure Greenish-yellow gas Brown liquid Black solid Greenish-yellow gas Pale yellow

Violet fumes Violet

Volatility  (or b.p ) from Cl2 to I2 Strength of VDW’s forces  due to increase in size of electron cloud  greater extent of distortion of the electron cloud.  More energy is required to overcome these stronger van der Waals’ forces. X2 + 2e  2XE  The element X2 becomes less reactive down the group since electron affinity of the halogen atom decreases  Thus, oxidising power  from Cl2 to I2 as shown by decreasing E values (less likely for X2 to reduce to X-). Cl2 displaces Br2 from Br- and displaces I2 from I-. Br2 displaces I2 from I- but I2 cannot displace Cl- and Br-. 4X2 (aq) + S2O32- (aq) + 5H2O(l)  2SO42-(aq) + 10H+(aq) + 8X-(aq) I2 (aq) +2S2O32- (aq)  2I-(aq) + S4O62Oxidation number of S  from +2 to +6 Cl2 reacts vigorously in presence of sunlight or u.v. light; Cl2 (g) + H2 (g)  2 HCl (g)

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Br2 reacts at 300 C in the presence of Pt catalyst Br2 (g) + H2 (g)  2 HBr (g)

Oxidation number of S  from +2 to +3 I2 reacts reversibly at 400 oC in the presence of Pt catalyst. I2 (g) + H2 (g)  2 HI (g)

At low temperature Cl2 (g) + 2NaOH(aq)  NaCl (aq) + NaClO(aq) + H2O(l) Cl2 undergoes disproportionation to form Cl- and ClO-. At high temperature Cl2 (g) + 6NaOH(aq)  5NaCl (aq) + NaClO3(aq) + 3H2O(l) At high T, ClO- undergoes further disproportionation to form Cl- and ClO3-.

©MJC 2011

Reddish brown gas Orange

At low and high temperature

At low and high temperature

Br2 (g) + 6NaOH(aq)  5NaBr (aq) + NaBrO3(aq) + 3H2O(l)

I2 (g) + 6NaOH(aq)  5NaI (aq) + NaIO3(aq) + 3H2O(l)

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Hydrogen halides (HX) Hydrogen halides Structure and bonding Melting point/ boiling point

Thermal stability

HCl

HBr

HI

Simple molecular structure with permanent dipole-permanent dipole interactions between the molecules. M.p. or b.p.  down the group Strength of VDW’s forces  due to increase in size of electron cloud  greater extent of distortion of the electron cloud.  More energy is required to overcome these stronger van der Waals’ forces between HX molecules.  2HX(g)  H2(g) + X2(g) Thermal stability  from HCl to HI since H-X bond energy  (i.e. HX bonds become weaker).

Down the group,  Atomic radius of halogen increases  Bond length of H-X increases  Bond strength of H-X decreases  Bond energy of H–X decreases Acidity of Hydrogen Halides



HX(aq)  H+(aq) + X-(aq)



Acid strength :

HI > HBr > HCl

Down the group,  Atomic radius of halogen increases  bond length of H–X bond increases  bond energy of H–X bond decreases  weaker H–X bond breaks more easily  H+ is more readily released  hydrogen halide becomes a stronger acid.

©MJC 2011

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Halides (X-) Halides ions

Cl-

Br-

I-

Reaction with AgNO3 followed by NH3

white ppt of AgCl soluble

cream ppt of AgBr insoluble in excess NH3 (only soluble in conc. NH3)

yellow ppt of AgI insoluble

in excess NH3 +

in excess NH3

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Ag (aq) + Cl (aq) AgCl (s) ----- (1) + Ag (aq) + 2NH3(aq) [Ag(NH3)2]+ (aq) ----- (2) When excess aqueous NH3 is added, AgCl (s) is soluble due to the formation of a soluble complex, [Ag(NH3)2]+ (aq). [Ag+] decreases in eqm (1) and AgCl ppt dissolves as the ionic product [Ag+] [Cl-] < Ksp.

Halides as reducing agent

Reaction with Concentrated H2SO4

For AgBr and AgI, the ppts remain insoluble in excess NH3 as both have relatively very small Ksp. Hence, their ionic products > Ksp. E X2 + 2e  2X The reducing power of the halides (X-) increases down the group as shown by decreasing E values. (more likely for X- to oxidise to X2) Down the group Cl2 + 2e 2ClE = +1.36 V  Ionic radius of Cl- to I- anion increases, Br2 + 2e 2Br E = + 1.07V  loss of valence electron occurs more readily I2 + 2e 2I E = + 0.54V  halide ions (X ) more easily oxidised  halide ions (X-) becomes a more powerful reducing agent . NaCl(s)+H2SO4(l)HCl(g)+NaHSO4 (aq)

NaBr(s) + H2SO4 (l)  HBr (g) + NaHSO4 (aq) 2HBr(g) + H2SO4(l)  Br2 (g)+ SO2 (g)+ 2H2O (l)

 White fumes HCl(g) formed.  Conc. H2SO4 is not powerful enough as an oxidizing agent to oxidize HCl to Cl2.

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©MJC 2011

White fumes HBr(g) formed. HBr gas can be further oxidised by concentrated H2SO4 to form reddish brown Br2 gas. Conc. H2SO4 is reduced to form pungent SO2.

2NaI (s) + H2SO4 (l) 2HI(g) + NaHSO4 (aq) 8HI (g) + H2SO4 (l)  4I2(s) +H2S (g) + 4 H2O (l)   

White fumes HI(g) formed. HI gas can be further oxidised by concentrated H2SO4 to form violet I2 gas. Conc. H2SO4 is reduced to form pungent H2S.

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