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Higher and foundation tiers

The halogens

 halogen reactivity list The group 7 non-metals are called the halogens. The halogens are fluorine, chlorine, bromine and iodine. Astatine at the bottom of group 7 is a very rare and highly radioactive element; its most stable isotope has a half-life of just over 8 hours. The halogens are all very reactive elements and are not found as elements in nature, instead they are found combined in compounds in rocks and minerals. Fluorine, chlorine and bromine are all toxic and corrosive and great care is needed in handling these reactive elements, though iodine being the least reactive halogen and being a solid is the easiest and safest halogen to handle in the lab.

Fluorine at the top of the group is a toxic pale greeny yellow gas, it is perhaps the most reactive elements in the periodic table. Chlorine is a greeny-yellow gas that is also very toxic and reactive. It has a recognisable smell that most people associate with the swimming baths, though the smell at the baths is not chlorine, as even small amounts of chlorine gas are quite toxic.

Bromine is a volatile red-brown liquid at room temperature with an obnoxious bleachy odour; if a small amount is placed in a flask it will quickly fill with red-brown bromine vapour. Bromine like fluorine and chlorine is a very toxic and dangerous element to handle. Iodine is a greyish/purple solid with a metallic like sheen at room temperature. With gentle heating it sublimes (turns straight from a solid to a gas) to produce beautiful violet vapours. The image below shows each of the halogens in gas jars along with a sample of solid iodine.

The halogens fluorine, chlorine, bromine and iodine in gas jars

Diatomic molecules

The halogens "go around in pairs"- that is they form molecules made up of two atoms as shown in the image. These diatomic molecules or two atom molecules are quite common for non-metal elements e.g. oxygen, nitrogen and hydrogen also form these diatomic molecules in the elemental state.

The halogens consist of small diatomic molecules, fluorine, chlorine, bromine and iodine all consist of small diatomic molecules.

Trends and patterns in the physical properties of the halogens

The table below lists the melting and boiling points of the halogens. The trend or pattern is fairly obvious, as we go down the group the molecules get larger and the relative mass increases. Larger molecules will result in stronger intermolecular bonding and this along with the increase in relative mass results in higher melting and boiling points.

Halogen Colour Melting point/0C Boiling point/0C state at room temperature
fluorine pale yellow -220 -188 gas
chlorine greenish-yellow -101 -34 gas
bromine red-brown -7 59 liquid
iodine greyish-purple 114 131 solid

Trends in the chemical properties of the halogens

Reactions with metals

All the halogens have 7 electrons in their outer shell, so only need to gain one to achieve full last shells. This means that the halogens are used as oxidising agents. That is they accept electrons from other elements, they oxidise them and by accepting electrons they are reduced e.g. All the halogens react with iron wool. The trends are what you might expect:

Iron is a fairly unreactive metal but still displays the trends you would expect. The strongly oxidising fluorine immediately accepts an electron from the iron atoms in the iron wool. The products of the reaction are:

iron(s) + fluorine(g)iron(III) fluoride(s)
2Fe(s) + 3F2(g) 2Fe F3(s)
Similar reactions occur with chlorine and bromine whereas the weakly oxidising iodine reacts slowly and some considerable heat is needed to start the reaction. The other halogens all oxidise the iron to produce Fe3+ ions but the weakly oxidising iodine is only able to oxidise the iron to produce Fe2+ ions.
iron(s) + iodine(s)iron(II) iodide(s)
2Fe(s) + I2(s) Fe I2(s)

The iron can be replaced with more reactive metals and some spectacular reactions can be seen. The halogens all react with aluminium to give some very spectacular reactions. This should be predictable as aluminium is a much more reactive metal than iron, so is more easily oxidised (remember oxidation = loss of electrons). The reaction of aluminium with chlorine gas is shown in the image below. Aluminium and chlorine gas reacting to form aluminium chloride. The reaction of aluminium and chlorine gas is very violent. The equation for this reaction can be shown as:

aluminium(s) + chlorine(g)aluminium chloride(s)
2Al(s) + 3Cl2(s) 2Al Cl3(s)
This time a brilliant white flash is seen as the aluminium and chlorine react, dense white fumes of aluminium chloride can be seen leaving the glass tube. A similar reaction can be carried out with bromine instead of chlorine.

Since bromine is a liquid, a small amount can be placed in a boiling tube and some freshly torn aluminium foil can be pushed into the boiling tube. After a few seconds sparking is seen at the surface of the freshly torn aluminium foil as it reacts with the bromine liquid. The reaction quickly gains speed and become very vigorous.

The reaction of iodine and aluminium is much slower. In fact the two substances can be mixed fairly safely on a clean dry tin lid. A few drops of water are added to catalyse and start the reaction. After about 60 seconds or so the reaction starts. A bright vivid glow is seen as the aluminium iodide forms, but perhaps the most vivid part of the reaction is the dense violet coloured fumes of iodine vapour that are produced. The heat produced by the reaction causes some of the iodine to sublime into iodine vapour.
The reaction of aluminium and iodine In fact any reactive metal will react with the halogens to form salts called metal halides. The salts formed from these reactions are all ionic compounds. The metal atoms lose electrons and are oxidised to form positive metal ions. The halogens always gain electrons; that is they are reduced to form negatively charged halide ions.

The structure of compounds formed between halogens and metals

The sodium chloride lattice structre.  Giant ionic lattice structure of sodium chloride.

When the halogens react with a metal from group I or II in the periodic table the compound formed will be an ionic compound with a giant ionic lattice structure. These compounds will consist of negatively charged halide ions and positively charged metal ions. As an example consider the reaction of the alkali metal sodium with the halogen chlorine.

sodium(s) + chlorine(s) → sodium chloride(s)
2Na(s) + Cl2(s) → 2NaCl(s)

This is a very violent reaction and the product formed, sodium chloride has a giant ionic lattice (shown opposite) which consists of positively charged sodium ions (Na+) and negatively charged chloride ions (Cl-).

Reactions of the halogens with non-metals

While the halogens react with metals to form compound which have giant ionic lattice structures; however when the halogens react with other non-metal elements they form compounds with simple molecular structures. The image below shows a range of different halogen compounds which consist of only non-metals. Here all the compounds have simple molecular structures.

Halogen compounds which consist of only non-metals all have a simple molecular structure.

Reactions of the halogens with hydrogen gas

Perhaps one of the best reactions to show the reactivity trends in the halogens is their reaction with hydrogen gas. All the halogens react with hydrogen to form hydrogen halide vapours:

H2(g) + X2(g) → 2HX(g) where x= F,Cl, Br, I

Explaining the trends

Fluorine being the smallest halogen atom will be able to attract a negatively charged electron from a metal atom more strongly towards its positively charged nucleus and so is the most reactive halogen. Iodine being in period 5 of the periodic table has 5 shells of electrons between its nucleus and any electron it tries to attract, these shells shield the positive nucleus from electrons that it is trying to attract. The iodine nucleus may have a much larger positive charge than the small fluorine nucleus, but the effect of shielding and the fact that the nucleus is a long way from any electrons it may try and attract means that the ability to attract electrons decreases as you descend group 7.

Key points

Practice questions

Check your understanding - Questions on the halogens.

Check your understanding - Quick quiz on the halogens