Group 7 consists of the elements fluorine, chorine, bromine, iodine and astatine, and they are known as the halogens. All the isotopes of astatine are radioactive and short-lived so we can mostly ignore it!
Group 7 elements are covalently bonded diatomic molecules, with instantaneous dipole – induced dipole intermolecular bonding (London forces). As we descend the group the strength of these intermolecular bonds increases (an iodine molecule has a lot more electron density than a fluorine molecule) which takes more energy to break, hence the trend in melting and boiling points. Iodine crystals will sublime if heated gently forming a purple vapour which condenses back to the solid on contact with a cool surface. It’s important to remember that we are not breaking the covalent bond in the halogen molecules themselves when we melt or boil these substances, we only need enough energy to break / disrupt the intermolecular bonding.
The strength of the X-X covalent bond decreases from chlorine to astatine which is to be expected – the overlap between the valence orbitals is less with larger atoms (alternatively we can think of it in terms of the distance between the shared pair of electrons and the two nuclei is far greater in iodine and so the attraction is weaker). We can see this if we look at the trend in bond dissociation enthalpy …
| bond dissociation enthalpy / kJ mol-1 | |
|---|---|
| F-F | +159 |
| Cl-Cl | +242 |
| Br-Br | +194 |
| I-I | +152 |
| At-At | +80 |
But how do we explain the weakness of the F-F bond?
Chlorine, bromine and iodine are all soluble in water to some extent, although iodine is far more soluble in a solution of potassium iodide which is how it is commonly made up in the lab. However, they are all considerably more soluble in non-polar solvents such as cyclohexane.
The halogens become less reactive (and less electronegative) as you go down the group. The vast majority of their reactions are redox reactions with fluorine being the strongest oxidising agent and iodine the weakest.
Displacement reactions – halogens behaving as oxidising agents
A more reactive halogen will displace a less reactive halide from solution causing a colour change (which is made more obvious if cyclohexane is added as the displaced halogen will dissolve in the upper non-polar layer preferentially).
Cl2(aq) + 2Br–(aq) ⇾ 2Cl–(aq) + Br2(aq)
Essentially, chlorine has oxidised the bromide from an oxidation state of -1 to 0 and the chlorine has itself been reduced from an oxidation state of 0 to -1. The solution changes colour from colourless to pale orange – addition of cyclohexane results in the free bromine moving into the solvent layer which turns deeper orange in colour.
Both chlorine and bromine will oxidise iodide to iodine as they are stronger oxidising agents than iodine.
Br2(aq) + 2I–(aq) ⇾ 2Br–(aq) + I2(aq)
The solution changes from colourless to brown as the iodine is displaced – addition of cyclohexane results in the free iodine moving into the solvent layer which turns violet in colour. You need to know the colours of the halogens in aqueous solution as well as in a non-polar solvent.
Halide ions behaving as reducing agents
If the halogens become stronger oxidising agents as we move up the group, then we can predict that the halide ions become stronger reducing agents as we move down the group.
A reducing agent causes another species in a chemical reaction to be reduced by providing electrons for the reduction reaction. The reducing agent is itself oxidised.
We can illustrate this trend by looking at the reaction of the sodium halides with concentrated sulphuric acid …
- sodium iodide
If we add concentrated sulphuric acid to solid sodium iodide drop by drop, the following reaction occurs:
H2SO4(l) + NaI(s) ⇾ NaHSO4(s) + HI(g)
The iodide ions then reduce the sulphur in the sulphuric acid from an oxidation state of +6 to -2 in hydrogen sulphide. Iodine is oxidised from an oxidation state of -1 in the iodide ion to 0 in iodine.
8H+ + 8I– + H2SO4(l) ⇾ H2S(g) + 4H2O(l) + 4I2(s)
There is a distinctive bad egg smell of hydrogen sulphide and grey-black iodine solid can be seen.
- sodium bromide
If we add concentrated sulphuric acid to solid sodium bromide drop by drop, a similar reaction occurs:
H2SO4(l) + NaBr(s) ⇾ NaHSO4(s) + HBr(g)
Bromide ions are not as strong oxidising agents as iodide ions so although a further reaction happens, the sulphur in the sulphuric acid is reduced from an oxidation state of +6 to +4 in sulphur dioxide.
2H+ + 2Br– + H2SO4(l) ⇾ SO2(g) + 2H2O(l) + Br2(l)
Orange-brown fumes of bromine are seen as the reaction is exothermic and some of the bromine is vaporised.
As a result of the reducing power of iodide and bromide ions, hydrogen iodide and hydrogen bromide must be prepared in the laboratory by adding concentrated phosphoric acid to the sodium halide (as conc. phosphoric acid will not be reduced and a pure sample of the hydrogen halide is obtained).
- sodium chloride
And finally, if we add concentrated sulphuric acid to solid sodium chloride drop by drop, the only products are sodium hydrogen sulphate and hydrogen chloride because chloride ions are not strong enough reducing agents to reduce the sulphur in sulphuric acid. This is how hydrogen chloride is prepared in the laboratory.
H2SO4(l) + NaCl(s) ⇾ NaHSO4(s) + HCl(g)
Identifying halide ions in solution
Addition of silver nitrate solution to an acidified solution of a metal chloride, bromide or iodide causes the silver halide to precipitate out because these salts are become increasingly more insoluble as we go down the group (silver fluoride is soluble).
Ag+(aq) + Br–(aq) ⇾ AgBr(s)
The acid of choice for acidification is nitric acid which is needed to react with any carbonate or hydroxide ions in the solution before the addition of Ag+ (silver carbonate and silver hydroxide are both white precipitates).
CO32-(aq) + 2H+(aq) ⇾ CO2(g) + H2O(l)
OH–(aq) + H+(aq) ⇾ H2O(l)
It can be difficult in practice to distinguish between the colours of the silver halide precipitates. On addition of dilute ammonia solution, silver chloride will dissolve completely to give a colourless solution, silver bromide is less soluble giving a slightly cloudy solution and silver iodide is insoluble.
Reactions of chlorine – oxoacids and oxoanions
Chlorine forms a number of acids (oxoacids) and anions (oxoanions) with oxygen in which the chlorine displays multiple oxidation states. The acids are mostly used in aqueous solution and cannot be isolated as pure substances.
- chlorine dissolves in water to form a solution of chloric(I) acid / hypochlorous acid, HOCl.
This is an example of a disproportionation reaction in which the oxidation state of chlorine both increases and decreases. Chloric(I) acid is an oxidising agent which can be used to purify drinking water.
- when we dissolve chlorine in aqueous sodium hydroxide, chloric(I) acid loses its proton forming a chlorate(I) anion, ClO–. This solution of sodium chlorate(I) and sodium chloride is more commonly known as bleach.
- the chlorate(I) ion is unstable and very slowly decomposes in another disproportionation reaction forming chlorate(V) and chloride ions.
Extraction and uses of the halogens
Fluorine
Fluorine is extracted from mineral ores such as fluorspar, CaF2. The ore is reacted with sulphuric acid to produce hydrogen fluoride which is dissolved in molten potassium fluoride before it is electroysed.
CaF2 + H2SO4 → 2HF + CaSO4
2F– ⇾ F2 + 2e– at the carbon anode
2H+ + 2e– ⇾ H2 at the steel cathode
Fluorine is mainly used to produce uranium hexafluoride for the enrichment of uranium in the nuclear industry.
Chlorine
Chlorine is a product of the electrolysis of brine (aqueous sodium chloride) in the chlor-alkali process.
Chlorine is used in many organic compounds such as chloroethene which is used to make polymers, in inorganic chlorides, in bleach and water treatment.
Bromine
Bromide ions are found in sea water at a concentration of around 65ppm but in inland seas such as the Dead Sea in Israel the concentration can be as high as 10000ppm. Chlorine is added to the sea water oxidising bromide ions to bromine in a displacement reaction and then steam is blown through the solution. The volatile bromine vapour produced is condensed, distilled and dried.
Bromine is used in the manufacture of flame retardants and in a wide range of pharmaceuticals.
Iodine
Iodine is extracted from iodate minerals such as lauterite, Ca(IO3)2, kelp seaweed and some naturally occurring brine solutions. It is used in a wide range of applications, from animal feed supplements to dyes, pharmaceuticals and in photography.
Practice questions
- The following reaction is an example of disproportionation
2HClO3(aq) + 2HCl(aq) ⇾ 2ClO2(aq) + Cl2(g) + 2H2O(l)
(a) What is a disproportionation reaction?
(b) Use oxidation states to illustrate your answer to part (a)
- In the reaction between chlorine and cold, aqueous hydroxide ions, the chlorine is both reduced to form chloride ions and oxidised to form chlorate(I) ions. Write an ionic equation for this reaction.
- A hydrogen halide is oxidised by concentrated sulphuric acid to form the foul smelling gas, hydrogen sulphide amongst other products.
(a) Identify the hydrogen halide
(b) Write an equation for the reaction
- A student plans to carry out a series of experiments to determine the order of reactivity of the halogens. They are given solutions of chlorine, bromine and iodine in cyclohexane as well as aqueous solutions of sodium chloride, sodium bromide and sodium iodide.
(a) Describe a method the student could use to achieve this
(b) Describe all expected observations
(c) Write ionic equations for the reactions that happen.
Answers
- (a) a disproportionation reaction is a redox reaction in which an element is both oxidised and reduced.
(b) chlorine has an oxidation state of +5 in HClO3 which is reduced to +4 in ClO2 and an oxidation state of -1 in HCl which is oxidised to 0 in Cl2.
- Cl2 + 2OH– ⇾ Cl– + ClO– + H2O
- (a) hydrogen iodide (b) 8HI + H2SO4 ⇾ H2S + 4H2O + 4I2
- The student should carry out the following experiments:
- mix 2cm3 of sodium chloride solution with 2cm3 of bromine in cyclohexane, stopper the test tube and shake – no reaction, top layer remains orange
- mix 2cm3 of sodium chloride solution with 2cm3 of iodine in cyclohexane, stopper the test tube and shake – no reaction, top layer remains violet
- mix 2cm3 of sodium bromide solution with 2cm3 of chlorine in cyclohexane, stopper the test tube and shake – chlorine displaces bromine, top layer changes from pale yellow-green to orange; Cl2(aq) + 2Br–(aq) ⇾ 2Cl–(aq) + Br2(aq)
- mix 2cm3 of sodium bromide solution with 2cm3 of iodine in cyclohexane, stopper the test tube and shake – no reaction, top layer remains violet
- mix 2cm3 of sodium iodide solution with 2cm3 of bromine in cyclohexane, stopper the test tube and shake – bromine displaces iodine, top layer changes from orange to violet; Br2(aq) + 2I–(aq) ⇾ 2Br–(aq) + I2(aq)
- mix 2cm3 of sodium iodide solution with 2cm3 of chlorine in cyclohexane, stopper the test tube and shake – chlorine displaces iodine, top layer changes from yellow-green to violet; Cl2(aq) + 2I–(aq) ⇾ 2Cl–(aq) + I2(aq)
Chlorine displaces both bromide ions and iodide ions from solution so is the most reactive halogen. Iodine cannot displace chloride ions or bromide ions from solution so is the least reactive halogen. Reactivity decreases as we go down the group.