Revision:Chemistry Unit 1.7 - Group 7 (The Halogens)

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Revision:Chemistry Unit 1.7 - Group 7 (The Halogens)

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TSR Wiki > Study Help > Subjects and Revision > Revision Notes > Chemistry > Chemistry Unit 1.7 - Group 7 (The Halogens)

A Level Chemistry Unit 1.7 - Group 7 (The Halogens).

1 Physical properties
2 Chemical tests
3 Hydrogen halides
4 Identification of the halides
5 Reactions of halide salts with H₂SO₄
6 The chlorate(I) and chlorate(V) ions
7 Disproportionation
8 The halogens as oxidising agents
9 Comments

Physical properties

The halogens are diatomic molecules $\mathsf{(X_2)}$ that exist in group 7 (VII) of the periodic table. They have the following physical properties:

Symbol	Name	Formula	State at r.t.p.	Boiling point (°C)	Atomic radius (pm)	Ionic radius (pm)
$\mathsf{F}$	Fluorine	$\mathsf{F_2}$	Pale yellow gas	-188	72	136
$\mathsf{Cl}$	Chlorine	$\mathsf{Cl_2}$	Pale Green gas	-34.7	99	181
$\mathsf{Br}$	Bromine	$\mathsf{Br_2}$	Brown liquid	58.8	114	195
$\mathsf{I}$	Iodine	$\mathsf{I_2}$	Dark grey solid	184.0	133	216

They also have many common properties:

They are coloured.
They exist as diatomic molecules.
Melting points increase as you go down group 7.
They are toxic.
The reactivity decreases as you go down group 7.

Chemical tests

Chlorine – Turns damp litmus red, and then bleaches it.
Bromine – Bromine also turns damp litmus red and bleaches it but slower. When sodium hydroxide is added, bromine loses its colour.
Iodine – When starch solution is added, a blue/black colour forms.

Hydrogen halides

Hydrogen halides are very soluble in water, for example hydrogen chloride gas dissolves readily in water.

$\mathsf{HCl_{(g)} \longrightarrow H^+_{(aq)} + Cl^-_{(aq)}}$ (in the presence of water)

When the hydrogen halides dissolve in water they form acidic solution due the high levels of hydrogen ions.

HCl dissolves well in water because HCl is a polar molecule as is water and so the intermolecular forces are favourable. Hence the HCl and H₂O molecules will have permanent dipole interactions to such an extent the HCl dissociates.

Identification of the halides

	Chloride	Bromide	Iodide
Addition of $\mathsf{AgNO_3}$	White precipitate	Off-white precipitate	Yellow precipitate
Addition of dilute ammonia solution	Precipitate dissolves	No change	No change
Addition of conc. ammonia solution	Precipitate dissolves	Precipitate dissolves	No change

Reactions of halide salts with H₂SO₄

Chloride salts

Oxidation state	-1	+6		+6	-1
Reaction	$\mathsf{NaCl}$	$\mathsf{+ H_2SO_4}$	$\mathsf{\longrightarrow}$	$\mathsf{NaHSO_4}$	$\mathsf{+ HCl}$

Bromide salts

Oxidation state	-1	+6		+6	-1
Reaction	$\mathsf{NaBr}$	$\mathsf{+ H_2SO_4}$	$\mathsf{\longrightarrow}$	$\mathsf{NaHSO_4}$	$\mathsf{+ HBr}$

Oxidation state	-1	+6			+4	0
Reaction	$\mathsf{2HBr}$	$\mathsf{+ H_2SO_4}$	$\mathsf{\longrightarrow}$	$\mathsf{2H_2O}$	$\mathsf{+ SO_2}$	$\mathsf{+ Br_2}$

Iodide salts

Oxidation state	-1	+6		+6	-1
Reaction	$\mathsf{NaI}$	$\mathsf{+ H_2SO_4}$	$\mathsf{\longrightarrow}$	$\mathsf{NaHSO_4}$	$\mathsf{+ HI}$

Oxidation state	-1	+6			+4	0
Reaction	$\mathsf{2HI}$	$\mathsf{+ H_2SO_4}$	$\mathsf{\longrightarrow}$	$\mathsf{2H_2O}$	$\mathsf{+ SO_2}$	$\mathsf{+ I_2}$

Oxidation state	-1	+6			-2	0
Reaction	$\mathsf{8HI}$	$\mathsf{+ H_2SO_4}$	$\mathsf{\longrightarrow}$	$\mathsf{4H_2O}$	$\mathsf{+ H_2S}$	$\mathsf{+ 4I_2}$

As it can be seen, as you descend the group the halide ions become more easily oxidised. In fact $\mathsf{Cl^-}$ ions are not oxidised at all with $\mathsf{H_2SO_4}$ . Hence from the equations we can conclude that the reducing power of the halide ions increase going down the group.

The chlorate(I) and chlorate(V) ions

There are two oxo-anions of chlorine each with different oxidation numbers:

The chlorate(I) ion – $\mathsf{OCl^-}$ (where chlorine is in the +1 state)

The chlorate(v) ion – $\mathsf{ClO_3^-}$ (where chlorine is in the +5 state)

Disproportionation

This occurs when a species is simultaneously oxidised and reduced. Chlorine undergoes this in an alkaline environment:

$\mathsf{Cl_2 + 2OH^- \longrightarrow Cl^- + OCl^- + H_2O}$

The chlorate(I) ion also undergoes disproportionation but on heating:

$\mathsf{3OCl^- \longrightarrow 2Cl^- + ClO_3^-}$

The halogens as oxidising agents

As you descend group 7, the halogens become less effective as an oxidising agent. This property can be used in the extraction of bromine from seawater. If Chlorine gas is passed through seawater then the bromide ions are oxidised to bromine. Therefore Cl₂ is a better oxidising agent than bromine.