AS LEVEL -The Periodic Table

The organisation of elements in the Periodic Table according to their proton numbers and electronic structures. The terms group and period. The trends in the physical properties across the period sodium to argon limited to melting points, electrical conductivity, first ionisation energies and atomic radii.
Group VII (fluorine, chlorine, bromine and iodine)
Practical work restricted to chlorine, bromine and iodine and their compounds.
Trends within the group limited to colour, physical state, melting and boiling points, atomic and ionic radii, first ionisation energies, bond energies of halogen molecules, hydrogen halides and carbon-halogen bonds; electronegativies.

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
Colour
Pale yellow
Pale green
Red/brown
Black
Physical state
Gas
Gas
Liquid
Solid


They are all p-block elements with a simple molecular structure consisting of covalently bonded diatomic molecules, X2.

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

F2<> 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.

Solubility in water and non-aqueous solvents eg hexane.

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?)





Chemical trends: reactivity with hydrogen, sodium and phosphorus.

All the halogens are oxidising agents and combine readily with metals and non-metals.

Reaction of halogens with elements
Hydrogen
The halogens combine enthusiastically with hydrogen, the vigour of the reaction decreasing from fluorine to iodine.
H2 (g) + X2 (g) 2HX (g)
Fluorine reacts explosively even in the dark at –200 oC.
Chlorine reacts explosively in sunlight, or slowly in the dark below 200 oC.
Bromine reacts above 200 oC and at lower temperatures with a platinum catalyst.
Iodine reacts to form an equilibrium mixture
H2 (g) + I2 (g) ⇌ 2HI (g)

Metals
The halogens combine readily with most metals forming the metal halides.
The vigour of the reaction decreases from fluorine 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 fluorine 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)

Reactions of the elements illustrated by use of chlorine gas (or chlorine water), bromine water and aqueous iodine (in potassium iodide) with water, aqueous alkalis, other halides in solution and iron (II) and iron (III) ions as appropriate. Disproportionation.

Reaction of halogens with water
Fluorine and chlorine can oxidise water. Fluorine oxidises water to oxygen.
2F2 (g) + 2H2O (l) 4HF (aq) + O2 (g)

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

Chlorine water contains chloric (I) acid HClO (aq), (hypochlorous acid). This is a weak acid which ionises to give the chlorate (I) ion ClO-, (hypochlorite ion). The hypochlorite ion is a powerful disinfectant and bleach.

Bromine disproportionates in a similar way but to a lesser extent.
Iodine has a very low solubility in water.



Reaction of halogens 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(I).
(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 70oC, 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 15oC.
For iodine, decomposition of IO- occurs rapidly at 0oC so it is difficult to prepare NaIO free from NaIO3.
NaClO is a mild antiseptic (Milton).
NaClO3 is a powerful weed killer.

Displacement reactions of the halogens
Since they are very electronegative, all the halogens are oxidising agents. As the group is descended their oxidising power decreases.

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
The trend in oxidising power is illustrated by the compounds formed by iron when it combines directly with the halogens.
Fluorine and chlorine form iron(III) fluoride and iron(III) chloride respectively.
2Fe (s) + 3F2 (g) 2FeF3 (s)

Bromine forms both iron(II) bromide and iron(III) bromide.
Iodine is too weak an oxidising agent and only forms iron (II) iodide.
Fe (s) + I2 (g) 2FeI2 (s)

Aqueous solutions of chlorine, bromine and iodine oxidise iron (II) to iron (III).
Cl2 (aq) + 2Fe2+ (aq) 2Cl- (aq) + 2Fe3+ (aq)

Iodine is so weak an oxidising agent that iron (III) ions oxidise iodide ions to iodine.

2Fe3+ (aq) + 2I- (aq) 2Fe2+ (aq) + I2 (s)


















Thermal stability of hydrogen halides related to bond enthalpies. The relative strength of the acids, HF, HCI, HBr and HI.

Thermal stability of hydrogen halides
The thermal stability of the hydrogen halides decreases as the group is descended. This is in keeping with the trend in bond enthalpies

The size of the halogen atom increases from fluorine to iodine; therefore the bond length increases and the bond enthalpy decreases.
Hydrogen fluoride and hydrogen chloride are stable to heating. Hydrogen bromide decomposes on strong heating.
2HBr (g) H2 (g) + Br2 (g)
Hydrogen iodide decomposes on gentle heating. If a hot wire is dipped into hydrogen iodide gas violet clouds of iodine are produced.
2HI (g) H2 (g) + I2 (g)

Hydrogen halides as acids
The dry hydrogen halides are not acidic and do not affect litmus paper. The hydrogen halides dissolve readily in water forming acid solutions.
HX (g) + H2O (l) H3O+ (aq) + X- (aq)

The acid strength increases in the order
HF << HCl < HBr < HI
An aqueous solution of HF is weakly acidic. This is because the H-F bond is very strong and there is strong hydrogen bonding between the HF molecules.
The rest are all strong acids.

Ionic halides. The identification of halide ions in solution by use of silver ions followed by aqueous ammonia. The effect of light on silver halides. Presence of halide ions in sea water. The reaction of solid halides with concentrated sulphuric acid to illustrate the relative reducing ability of halides ions and hydrogen halide. The effects of fluoridation of public water supplies on dental health and an appreciation of the debate between public health policy and practice and the rights of the individual.


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)

Pb2+(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)


Exercise 2
Write an equation for the reaction of sodium chloride solution with
(a) lead nitrate solution and





(b) silver nitrate solution followed by the addition of ammonia.





Halide ions in sea water

Sea water contains over 3% of dissolved chlorides (mainly sodium, potassium, calcium and magnesium) and large solid deposits are found where inland seas have undergone evaporation. Bromine is also found in sea water as the bromides of sodium, potassium and magnesium. Iodine is present in sea water at a concentration of less than one part per million. Some seaweeds and sponges extract this iodine during growth and their ash contains about 0.5% by mass of iodine as iodides.


Reaction of the solid halides with conc. sulphuric acid, H2SO4

When concentrated sulphuric acid is added to a sodium halide the first product is fumes of the hydrogen halide HX, because each of these compounds is more volatile than sulphuric acid.
NaCl (s) + H2SO4 (1) HCl (g) + NaHSO4 (s)

NaBr (s) + H2SO4 (1) HBr(g) + NaHSO4 (s)

However conc. sulphuric acid is also a strong oxidising agent and will oxidise
HBr Br2 and HI I2, but not HF and HCl.



oxidised

2HBr(g) + H2SO4 Br2 + SO2 (g) + 2H2O

reduced


Similarly 2HI(g) + H2SO4 I2 + SO2(g) + 2H2O

Therefore conc. sulphuric acid cannot be used for the preparation of
HBr(g) and HI(g).
However conc. phosphoric(V) acid, H3PO4, can be used for the preparation since it is relatively non volatile and a poor oxidising agent.

NaBr (s) + H3PO4 (1) HBr (g) + NaH2PO4
NaI (s) + H3PO4 (1) HI (g) + NaH2PO4

Reagent
Fluoride
Chloride
Bromide
Iodide
Conc. H2SO4

X- + H2SO4 HX (g) + HSO4-
HF (g) colourless, pungent, corrosive gas
HCl (g)
formed
HBr (g) + a little Br2
A little HI (g) but mainly I2
Conc. H3PO4

HF (g)
HCl (g)
HBr (g)
HI (g)

Exercise 3
Make notes on the uses of the halogens and their compounds, with particular reference to the fluoridation of public water supplies and the effects on dental health.

KEY SKILLS