1. 3.2.3 Group 7, the Halogens

2. The halogens are non-metals
The halogens are non-metals. As elements they exist as diatomic molecules.
Br2
(brown liquid) I2
(black solid) F2
(yellow gas)
Cl2
(green gas)

3. Physical Properties
At room temperature fluorine is a pale yellow gas, chlorine is a greenish gas, bromine is a red-brown liquid and iodine is a black solid. The halogens get darker and denser going down the group. 

4. Electronegativity
Electronegativity is the power of an atomic to attract electron density towards itself in a covalent bond.

5. The electronegativity of the halogens decreases going down the group.

6. Electronegativity depends on the attraction between the nucleus and the shared electrons in the covalent bond.
This, in turn, depends on a balance between the shielding of the nucleus, the distance between the bonding electron and number of protons in the nucleus (nuclear charge).

7. The halogen atoms get larger going down the group
The halogen atoms get larger going down the group. Therefore in covalent compounds the shared electrons are further from the halogen nucleus and there is more shielding of the halogen nucleus due to the larger number of inner shells of electrons.

8. These factors are more important than the increase in nuclear charge and so the electronegativity decreases going down the group.

9. Boiling point
The boiling points of the halogens increase going down the group. This is because the magnitude of Van der Waals’ forces between the diatomic halogen molecules increases with the size of the molecules.

10. Oxidising ability of the halogens
Halogens usually react by gaining electrons to become negative ions with a charge of -1. The halogens are oxidising agents and are themselves reduced (gain electrons) in these redox reactions.
X2 + 2e-  2X-
The oxidising ability of the halogens decreases going down the group.

11. Displacement Reactions
Halogens react with metal halides such that the halogen in the compound will be displaced by a more reactive halogen but not by a less reactive halogen.

12. Cl2(aq) + 2NaBr(aq)  Br2(aq) +2NaCl(aq) Oxidation state: 0 -1 0 -1
For example, chlorine will displace bromine from sodium bromide but iodine will not displace bromine from sodium bromide.
Cl2(aq) + 2NaBr(aq)  Br2(aq) +2NaCl(aq)
Oxidation state:
(Cl) (Br) (Br) (Cl)
Ionic equation:
Cl2(aq) + 2Br-(aq)  Br2(aq) + 2Cl-(aq)
the sodium ions are spectator ions.
The chlorine acts as an oxidising agent, it oxidises the bromide ions.

13. Cl-
Br- I-
Cl2 -
displaces Br2
displaces I2
Br2
no displacement I2
You cannot investigate fluorine in aqueous solution because it reacts with water.

14. Halide ions as reducing agents
Halide ions can act as reducing agents. In these reactions the halide ions lose electrons and become halogen molecules.
The larger the ion the more easily it loses an electron since the electron is lost from the outer shell which is further from the nucleus as the ion gets larger and so the attraction to the outer electron is less.
The reducing power of the halides increases as you go down the group.

15. The reactions of sodium halides with concentrated sulphuric acid
Solid sodium halides react with concentrated sulphuric acid. The products are different and reflect the different reducing powers of the halide ions.

16. Black solid, purple fumes
NaX
Observations
Product
Type of reaction
NaF
Steamy fumes HF
acid-base
(F- accepting H+ to act as a base)
NaCl
HCl
(Cl- acting as base)
NaBr
Colourless gas
Brown fumes
HBr
SO2
Br2
(Br- acting as base)
redox
(reduction product of H2SO4)
(oxidation product of Br-)
NaI
Yellow solid
Smell of bad eggs
Black solid, purple fumes HI S
H2S I2
(I- acting as base)
(oxidation product of I-)

17. Iodide ions can reduce sulfur in H2SO4 from oxidation state +6 to +4 as SO2, then to 0 as the element S, and finally to -2 as H2S.
Bromide ions can reduce sulfur in H2SO4 from oxidation state +6 to +4 as SO2.
Chloride and fluoride ions cannot reduce the sulphur in H2SO4.

18. Make sure that you can work out the half equations for:
Reduction of H2SO4 to SO2
Reduction of H2SO4 to S
Reduction of H2SO4 to H2S
Oxidation of Br- to Br2
Oxidation of I- to I2

19. Combine your half equations to deduce the overall equations for:
The reaction between Br2 and H2SO4 to form SO2
The reaction between I2 and H2SO4 to form SO2
The reaction between I2 and H2SO4 to form S
The reaction between I2 and H2SO4 to form H2S

20. Identifying metal halides with silver ions
Metal halides (except fluorides) react with silver ions in aqueous solution to form a precipitate of the insoluble silver halide. Silver fluoride does not precipitate because it is soluble in water.
Dilute nitric acid is first added to the halide solution to get rid of any soluble carbonate or hydroxide impurities (which can form precipitates, confusing the results):
CO32-(aq) + 2H+(aq)  CO2(aq) + H2O(l)
OH-(aq) + H+(aq)  H2O(l)
These would interfere with the test by forming insoluble silver carbonate or silver hydroxide:
2Ag+(aq) + CO32-(aq)  Ag2CO3(s)
Ag+(aq) + OH- (aq)AgOH(s)
Then a few drops of silver nitrate solution are added and the halide precipitate forms.

21. Identifying metal halides with silver ions
After the addition of nitric acid, silver nitrate solution is added and the silver halide precipitate forms.
White
AgCl
Cream
AgBr
Yellow
AgI

22. Identifying metal halides with silver ions
The colours of silver bromide and silver iodide are similar but if you add a few drops of concentrated ammonia solution, silver bromide dissolves but silver iodide does not.
Halide
Silver fluoride
Silver chloride
Silver bromide
Silver iodide
Colour
No precipitate
White precipitate
Cream precipitate
Pale yellow precipitate
Further test
Dissolves in dilute ammonia
Dissolves in concentrated ammonia (,but not in dilute ammonia)
Does not dissolve in (dilute or) concentrated ammonia

23. Reaction of chlorine with water
Chlorine reacts with water in a reversible reaction to form chloric (I) acid and hydrochloric acid:
Cl2 + H2O → HClO + HCl
Oxidation
state of Cl
The reaction is an example of a disproportionation reaction where one species (chlorine) is simultaneously oxidised and reduced.

24. Reaction of chlorine with water
Cl2 + H2O → HClO + HCl
Or Cl2 + H2O → HClO + H+ + Cl-
If universal indicator is added to solution of chlorine water, initially turns red (hydrochloric acid and chloric (I) acid formed) then red colour disappears and colourless solution left (chloric (I) acid bleaches very effectively).
This reaction takes place when chlorine is used to purify water for drinking and in swimming pools, to kill disease-causing bacteria.
The other halogens react similarly, albeit much more slowly going down the group.

25. Reaction of chlorine with water
If chlorine reacts with water in the presence of sunlight a different reaction occurs. Chlorine oxidises water to oxygen and is itself reduced to chloride ions.
2Cl2 + 2H2O  4H+ + 4Cl- + O2
Chlorine is lost rapidly from pool water in sunlight and so shallow pools need frequent addition of chlorine.

26. NaClO(s) + H2O(l) ⇌ Na+(aq) + OH-(aq) + HClO(aq)
An alternative to the direct chlorination of swimming pools is to add solid sodium or calcium chlorate (I). This dissolves in water to form chloric (I) acid in a reversible reaction.
NaClO(s) + H2O(l) ⇌ Na+(aq) + OH-(aq) + HClO(aq)
In alkaline solution the equilibrium shifts to the left and the HClO is removed as ClO- ions. To prevent this from happening swimming pools must be kept slightly acidic.

27. Cl2 + 2NaOH  NaCl + NaClO + H2O
Reaction with Alkali
Chlorine reacts with cold dilute sodium hydroxide to form sodium chlorate(I) and sodium chloride.
Cl2 + 2NaOH  NaCl + NaClO + H2O
Sodium chlorate(I) is an oxidising agent and the active ingredient in household bleach.
The other halogens behave similarly.