The elements in Group 17* (the halogens) all exist as diatomic molecules, containing a single covalent bond. They are all non-metals.
* the convention is to number all groups across the periodic table, including the groups of d-block elements. We did not do this at GCSE as it was convenient to have the group number equalling the number of outer-shell electrons, which is why we referred to them as Group 7.
Looking at chlorine, bromine and iodine, we can see that down the group:
- the melting and boiling points increase (gas → liquid → solid at room temp.)
- the colour becomes more intense
The trend in boiling points (Fluorine: -188°C, Chlorine: -34.6°C, Bromine: 58.8°C, Iodine: 184°C) and melting points is explained in terms of the increasing strength of the intermolecular forces which hold the halogen molecules to one another. There are no polar bonds (same atom, so same electronegativity at each end of the bond) and no prospect of hydrogen bonding, so instantaneous dipole-induced dipole (London forces) are the only intermolecular forces present. These increase with the number of electrons present in the molecule.
Note: Iodine is unusual because it SUBLIMES when it is heated (goes straight from solid to vapour without going through a liquid state).
On the basis of the observed properties of chlorine, bromine and iodine we can predict the physical properties of fluorine and astatine (radioactive and decays very quickly, so very low abundance):
- Fluorine is a very pale coloured gas at room temperature
- Astatine is a black solid at room temperature
The halogens can all form halide ions (F–, Cl–, Br– and I–) by gaining an electron to become isoelectronic with noble gases.
e.g. Cl2 + 2e– → 2Cl–
The gain of an electron is reduction, so the halogens are oxidising agents, themselves being reduced and causing the substances they react with to be oxidised.
Halide ions are colourless. This is a common source of mistakes, as the halogens themselves are coloured.
Trend in reactivity
The ease with which the halogens can attract and remove an electron from another atom or ion to become a halide determines their reactivity: The reactivity decreases down the group with fluorine being the most and iodine the least reactive halogen.
Down the group reactivity decreases because:
- The atomic radius increases as more shells are filled – the smaller the halogen atom is, the closer its (positively charged) nucleus can get to other atoms/ions to remove an electron from them. Fluorine will be able to get closer than any other halogen.
- Shielding increases – the more filled shells there are between a halogen’s nucleus and the other atom/ion, the less easily it will be able to attract an electron from them. Fluorine has fewest filled shells surrounding the nucleus.
- Nuclear charge increases due to the increasing number of protons in the nucleus, but the two factors above outweigh this, so the electron to be gained is subjected to a weaker attraction.
Since the halogens get less reactive down the group, they are less effective at oxidising other substances, and hence their oxidising power decreases down the group.
Other compounds of halogens
Halogens will also react by sharing an electron to form (single) covalent bonds. The compounds of halogens include the interhalogens, simple molecules such as ICl, IBr, ICl3 as well as the more familiar HBr, HCl etc. While they normally have a valency of 1, we do see halogens in Period 3 or below forming more than one single covalent bond. This happens in interhalogens such as ICl3 and in ions such as chlorate, bromate and iodate ions: e.g. chlorate(V), ClO3–.
Reactions of Group 7
i) Displacement reaction between halogens and halide ions
These displacement reactions demonstrate the decreasing reactivity down Group 17. A more reactive halogen atom can remove an electron from a halide ion of a less reactive halogen. In doing this, the more reactive halogen is reduced to a halide ion, and the halide ion is oxidised to a halogen, so these are also redox reactions.
We can represent the set of such reactions using ionic equations:
|chloride ions||No reaction||No reaction||No reaction|
|bromide ions||Cl2 + 2Br – → 2Cl– + Br2||No reaction||No reaction|
|iodide ions||Cl2 + 2I– → 2Cl– + I2||Br2 + 2I– → 2Br– + I2||No reaction|
In practice we can do these reactions using iodine solution, bromine water (bromine solution) and chlorine water or chlorine gas, reacting these halogens with sodium or potassium halide solutions. In each case we need to observe the colour of the halogen that is initially present, then note the colour of the halogen present after the halide ions have been mixed, to determine if a reaction has taken place. There is no contribution to the colour from the halide ions, as these are colourless.
It can be difficult to tell the colour of aqueous bromine from the colour of dilute aqueous iodine, so the solution can be shaken with a layer of an inert organic liquid such as cyclohexane, in which the halogens are very soluble.
|Halogen||In aqueous solution||In organic layer|
|Bromine, Br2||Yellow – orange||Orange|
|Iodine, I2||Orange – red/brown||Purple|
Reactions with chlorine
Chlorine undergoes displacements reactions with bromide ions, forming bromine in the solution, along with chloride ions:
Observations: yellow – orange colour appears in the colourless solution. If shaken with an organic layer, this turns orange.
Chlorine also undergoes displacements reactions with iodide ions, forming iodine in the solution, along with chloride ions:
Observations: orange/red/brown colour appears in the colourless solution (colour depends on the amount of iodine released into solution). If shaken with an organic layer, this turns purple.
Reactions with bromine
Bromine undergoes a displacement reactions with iodide ions forming iodine in the solution along with bromide ions, but does not react with choride ions.
Observations: The bromine water is initially orange. After mixing the bromine water and colourless iodide ions, the colour deepens becoming orange/red/brown. The presence of iodine can be seen by shaking with an organic layer, which will turn purple.
On the basis of these reactions we can predict that fluorine will displace chlorine from chloride ion solutions, bromine from bromide ions, and iodine from iodide ions. None of the other halogens will be able to displace fluorine from a solution containing fluoride ions.
ii) Reaction of halide ions with silver ions
Aqueous halide ions react with aqueous silver ions to form precipitates of insoluble silver halides, which have characteristic colours. These are not redox reactions.
Typically silver nitrate solution is used as a source of aqueous silver ions.
The ionic equations for the reactions taking place are:
Ag+(aq) + Cl–(aq) → AgCl(s) white precipitate
Ag+(aq) + Br–(aq) → AgBr(s) cream precipitate*
Ag+(aq) + I–(aq) → AgI(s) yellow precipitate
* the silver bromide precipitate is light-sensitive, and soon darkens to a grey colour. This is the basis of photographic film and papers.
Further test with ammonia
These colours can be difficult to tell apart as they are not very different, but a further test can be used to tell the precipitates apart. This is based on differences in their solubilty in ammonia:
|Silver chloride precipitate||Soluble in dilute ammonia solution, NH3(aq)|
|Silver bromide precipitate||Soluble in concentrated ammonia solution NH3(aq)|
|Silver iodide precipitate||Insoluble in ammonia solution|
iii) Disproportionation reactions
An interesting effect is seen when halogens reacts with water, or with cold dilute sodium hydroxide, known as disproportionation.
Definition: Disproportionation is a reaction in which the same element is simultaneously oxidised and reduced.
Reaction of chlorine with water
Chlorine reacts with the water to form two acids in solution:
Both HCl and HClO split up in aqueous solution to produce H+ ions, turning indicators such as blue litmus red, however the ClO–(aq) ion is also a powerful oxidizing agent, so shortly after the indicator paper becomes bleached to a white colour.
Reaction of chlorine with cold dilute sodium hydroxide
This is the reaction used commercially to produce bleach. Bleach is an aqueous solution of sodium chloride and sodium chlorate(I). The reaction is:
The bleaching action occurs because the ClO– ion is a powerful oxidizing agent, reacting with dyes to oxidize them, and itself becoming reduced to Cl– ions.
Similar disproportionation reactions happen with the other halogens, and with other alkalis, such as potassium hydroxide.
Water treatment with chlorine
Chlorine began to be used as a disinfectant for dinking water treatment over 100 years ago. Harmful bacteria are killed by the chloric acid and chlorate(I) ions rather than by chlorine itself, making the water safe to drink or swim in. The use of chlorine has both risks and benefits, and there are also ethical considerations: not everyone agrees with not having a choice over having our water chlorinated.
- Chlorine gas is toxic in large concentrations, and a respiratory irritant in lower concentrations.
- Chlorine in drinking water can react with organic hydrocarbons such as methane from decaying vegetation, forming chlorinated hydrocarbons which may be carcinogenic.
- Risk to health from chlorinated hydrocarbons is far less than the risk to health from diseases such as typhoid and cholera.
- Some chlorine persists in the water preventing re-infection after treatment.
Alternatives to chlorine
- Ozone, O3, is also a powerful oxidizing agent and can be added to water to kill microorganisms. It is expensive to produce, however, and does not last for long in the water.
- Ultraviolet light can also be used to kill microorganisms by damaging their DNA, but it is ineffective in cloudy water and like O3 does not stop the water becoming infected subsequently to treatment.