7.2.2 GROUP VII Halogens

(Halogens : restricted to chlorine, bromine and iodine)

Valid deductions may be expected about other elements in the group.

Physical properties of halogens, limited to colour and physical state at room temperature.

The halogens are a group of reactive non-metals, which are essentially similar to each other with only gradual changes as the atomic number increases.

Physical properties of the halogens

Property	Fluorine F	Chlorine Cl	Bromine Br	Iodine I
Number of protons (atomic number)	9	17	35	53
Outer electron configuration	2s^22p^5	3s^23p^5	3d^104s^24p^5	4d^105s^25p^5
Atomic radius/ nm	0.064	0.099	0.111	0.128
Ionic radius/ nm	0.133	0.181	0.196	0.219
Melting point / ^oC	-220	-101	-7	114
Boiling point / ^oC	-188	-34	58	183
Bond energy / kJ mol^-1	158	242	193	151
Electron affinity/ kJ mol^-1	-361	-388	-365	-332
Standard electrode potential/ V	+2.87	+1.36	+1.09	+0.54
Electronegativity	4.00	2.85	2.75	2.20
Oxidation states	- 1	-1,1,3,5,7	-1,1,3,5,7	-1,1,3,5,7
Standard enthalpy of formation of NaX/ kJ mol^-1	-573	-414	-361	-288
Standard lattice enthalpy of NaX/ kJ mol^-1	902	771	733	684

They are all p-block elements with a simple molecular structure consisting of covalently bonded diatomic molecules, X_2.

o o o o

o o o

o X o X o X X

o o o o

There are only weak Van der Waals forces between the molecules. The strength of the forces increases as the number of electrons (M_r) in the molecule increases.

F₂< Cl₂< Br₂< I₂

In the case of iodine, the forces are sufficiently strong to bind the iodine molecules together in a 3-D crystal lattice. The X-X bond strength decreases down the group

Cl₂> Br₂ > I₂ as the atoms get larger and the attraction of the nucleus for the shared electrons decreases (electronegativity decreases).

There is a slight tendency to metallic character with increasing atomic number. The halogens complete their octet by gaining one electron forming a halide ion, X^- (see electron affinity values) or by sharing one electron. Fluorine is restricted to an oxidation state of -1 but the remaining elements have empty d orbitals and can promote electrons to give oxidation states of +1, +3, +5 and +7.

They are all oxidising agents and combine readily with metals and hydrogen.

Chlorine is a greenish-yellow gas.

Bromine is a red-brown volatile liquid.

Iodine is a black shiny solid, which sublimes on heating to produce a purple vapour.

Reactions of chlorine gas with elements (see Section 7.5.1), water, alkalis (under different conditions to form ClO^- and ClO₃^-), other halides in aqueous solution, iron(II) ions in solution, hydrocarbons (see Sections 7.3.2, 7.3.3 and 7.3.4).

Reaction of halides with elements

Metals

The halogens combine readily with most metals forming the metal halides.

The vigour of the reaction decreases from chlorine to iodine.

Group I and II halides are ionic.

2Na (s) + Cl₂ (g) 2Na⁺Cl^- (s)

Mg (s) + Cl₂ (g) Mg²⁺2Cl^- (s)

The halides of Group III are predominantly covalent.

2Al (s) + 3Cl₂ (g) 2AlCl₃ (s)

Non-metals

The elements react directly with many non-metals the oxidising power decreasing from chlorine to iodine.

The elements combine directly with phosphorus, the oxidation state of the product depending on the oxidising power of the halogen.

2P (s) + 5Cl₂ (g) 2PCl₅ (s)

2P (s) + 3Br₂ (l) 2PBr₃ (l)

Solubility of the elements

All three elements are only slightly soluble in water because of the relatively strong hydrogen-bonding between the water molecules, which does not exist between the halogen molecules

i.e. solvent-solvent attractions > solute-solvent attractions > solute-solute

attractions.

Cl₂> Br₂>I₂

solubility decreasing

They are soluble in non-polar organic solvents such as toluene and TCE.

(Why?)

Chlorine reacts slowly with water forming hydrochloric acid and chloric(I) acid. This reaction involves disproportionation:- a change in which one particular molecule, atom or ion is simultaneously both oxidised and reduced.

reduction

Cl₂ (g) + H₂O (l) HCl (aq) + HClO (aq) (chlorine water)

o.n. 0 -1 +1

oxidation

Bromine and iodine disproportionate in a similar way but to a lesser extent.

Reaction of chlorine with aqueous sodium hydroxide.

Chlorine reacts faster with dilute sodium hydroxide than with water.

When chlorine is added to cold dilute alkali it disproportionates to chloride and chlorate(l).

(i)

reduction

Cl₂ (g) + 2NaOH (aq) NaCl (aq) + NaOCl (aq) + H₂O

o.n. 0 -1 +1

oxidation

( 2OH^- + Cl₂ Cl^- + OCl^- + H₂O )

(ii) In hot concentrated alkali, if the solution is warmed to 70^oC, the chlorate(I) disproportionates further to chlorate(V).

reduction

3NaOCl (aq) 2NaCl (aq) + NaClO₃ (aq)

o.n. +1 -1 +5

oxidation

If chlorine is bubbled directly into hot conc. alkali then

(iii) reduction

3Cl₂ (g) + 6NaOH(aq) 5NaCl (aq) + NaClO₃ (aq)

o.n. 0 -1 +5

oxidation

( 6OH^- + 3Cl₂ 5Cl^- + ClO₃^- + 3H₂O )

For bromine, both reactions (i) and (ii) are fast at 15^oC.

For iodine, decomposition of IO^- occurs rapidly at 0^oC so it is difficult to prepare NaIO free from NaIO₃.

NaClO is a mild antiseptic (Milton).

NaClO₃ powerful weed killer.

Relative oxidising ability of the halogens linked to redox potentials.

Displacement reactions of the halogens

Since they are very electronegative, all the halogens are oxidising agents. Their standard electrode potentials, E^q, become less positive on descending the group, showing that their oxidising power decreases.

X₂ + 2e^- 2X^-

s.e.p. E^q /volts

Cl₂ (g) /2Cl^- + 1.36

Br₂ (g) / 2Br^- + 1.09

I₂ (g) / 2I^- + 0.54

Therefore chlorine oxidises bromide ions to bromine and iodide ions to iodine.

These are displacement reactions.

Cl₂ (g) + 2Br^- (aq) Br₂ (l)+ 2Cl^- (aq)

(colourless) (yellow/orange)

Cl₂ (g) + 2I^- (aq) I₂ (s) + 2Cl^- (aq)

(colourless) (red/brown)

Bromine oxidises iodide to iodine

Br₂ (g) + 2I^- (aq) I₂ (s) + 2Br^- (aq)

Iodine does not oxidise any of the others.

Other oxidising reactions of the halogens

Oxidant	Reaction
All halogens	Sulphite, S03^2- Sulphate, S04^2-
All halogens	Hydrogen sulphide, H2S Sulphur, S
Chlorine, Bromine	Thiosulphate, S203^2- Sulphate, SO4^2-
Iodine	Thiosulphate, S203^2- Tetrathionate, S4O6^2-
Chlorine, Bromine	Iron(II) Iron(III)

Exercise 1

Write balanced equations for each of the above reactions.

(Use the symbol X for a general halogen reaction)

The identification of halide ions in solution by the use of silver ions and aqueous ammonia; the [Ag(NH₃)₂]⁺ ion

Reaction of the halide ions in solution, X^-(aq)

Most metal halides are soluble except lead and silver halide. Therefore solutions of lead and silver ions are used to test for the presence of halide ions in solution.

Reagent	F^- (aq)	Cl^- (aq)	Br^- (aq)	I^- (aq)
Pb(NO3)2 (aq) Pb^2+(aq) + 2X^-(aq) PbX2(s)	White precipitate of PbF2	White precipitate of PbCl2	Cream precipitate of PbBr2	Yellow precipitate of PbI2
AgNO3 (aq) Ag^+ (aq) + X^- (aq) AgX (s)	No reaction AgF soluble in water	White precipitate AgCl	Cream precipitate AgBr	Yellow precipitate AgI
Solubility of silver halide in (a) dil. NH3 (aq) (b) conc. NH3 (c) dil.HNO3 (aq)		soluble soluble insoluble	insoluble soluble insoluble	insoluble insoluble insoluble
Effect of sunlight		White ppt. turns purple/grey	Cream ppt. turns green/ yellow	No effect

Exercise 2

Write an equation for the reaction of sodium chloride solution with