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GCSE Chemistry - Group VII, The Halogens

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Halogens - "Periodic Table" Videos

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Halogens - Properties

The elements of Group 7 of the Periodic table are known as the halogens, from the classical Greek for salt forming. They are all non-metals, have different states at room temperature, different colours and a gradation of physical properties (see the table below).

Element State @ 25 °C Colour Melting pt. (°C) Boiling pt. (°C)
Fluorine gas yellow -220 -188
Chlorine gas green -101 -34
Bromine liquid brown -6 59
Iodine solid black 114 185

The halogens in general are toxic. Fluorine is particularly dangerous; you would already be dead by inhalation of fluorine before you could register its smell. Chlorine was used in the First World War as a poisonous gas, as it is heavier than air and could sink into trenches.

However, iodine is not toxic and can be used as a very effective disinfectant for wounds. Also, whilst solid iodine is black, gaseous iodine has a distinctive purple colour.

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Halogens - Reactions

(1) With metals :

All the halogens react with metals, though the ease of reaction decreases as the group is descended.

Just passing fluorine gas over reactive metals, such as the alkali metals, results in an exothermic reaction and the formation of a metal fluoride salt. However, with iodine the metal must be heated strongly before a reaction will occur.

Exemplar equation -

aluminium + fluorine → aluminium fluoride

2Al(s) + 3F2(g) → 2AlF3(s)

Try balancing this equation,

K(s) + Cl2(g)KCl(s)

(2) Production of chlorine :

diagram of apparatus for generating chlorine gas

When concentrated hydrochloric acid is added to solid potassium manganate(VII), KMnO4, chlorine gas is generated which can be collected using the apparatus shown.

N.B.: this must be done in a fume cupboard because of the toxic nature of chlorine.

The chemical equation for this reaction is not required at GCSE but is shown below,

2KMnO4(s) + 16HCl(aq) → 5Cl2(g) + 2MnCl2(aq) + 2KCl(aq) + 8H2O(l)

The chlorine produced can be tested for by using damp litmus paper. Chlorine is a powerful bleaching agent and the litmus paper will be bleached, i.e. turn white.

N.B.: this test will work for blue or red litmus paper, though blue litmus paper should turn red just before it is bleached, as an aqueous solution of chlorine is slightly acidic.

(3) With other halides :

A halogen high in the group will displace a halogen lower in the group from that halogen's halide salt.

This can be demonstrated in the laboratory by adding some aqueous bromide and iodide ions to test tubes containing chlorine gas generated above.

When a solution of bromide ions are added the green gas turns orange-brown as does the solution, showing that elemental bromine is being formed.

Exemplar equation -

magnesium bromide + chlorine → magnesium chloride + bromine

MgBr2(aq) + Cl2(g) → MgCl2(aq) + Br2(aq)

With a solution of iodide ions, the green gas disappears and the solution turns a yellow colour with a black precipitate, i.e. insoluble solid, in it. The reason for these unusual observations is that iodine is slightly soluble in water and it forms a yellow solution when it does dissolve. Since most of the iodine formed does not dissolve the rest forms the black solid.

Exemplar equation -

magnesium iodide + chlorine → magnesium chloride + iodine

MgI2(aq) + Cl2(g) → MgCl2(aq) + I2(s)

These reactions can also be represented by what is known as an ionic equation. This type of equation does not show any elements that do not change during the reaction, so called spectator ions. So for the above reactions of chlorine with bromide and iodide ions, the ionic equations would look like this -

2Br-(aq) + Cl2(g) → 2Cl-(aq) + Br2(aq)

2I-(aq) + Cl2(g) → 2Cl-(aq) + I2(s)

As can be seen this representation can dramatically simplify a chemical equation by removing the metal ions.

Try balancing these equations using fluorine and some halide salts,

LiBr(aq) + F2(g)LiF(aq) + Br2(aq)

AlCl3(aq) + F2(g)AlF3(aq) + Cl2(g)

(4) Reactions of halide ions:

Experimental sheet for the reactions of halide ions.

Halide ions undergo a series of unique reactions that allow an unknown solid or aqueous sample to be tested for the presence of chloride, bromide or iodide ions.

Aqueous silver ions react with halide ions to produce individually coloured precipitates. These precipitates have different solubilities in ammonia solution and so further differentiation can be achieved, see the table below,

Halide ion Observation with AgNO3(aq) Solubility in NH3(aq)
Chloride white precipitate soluble in dilute NH3(aq)
Bromide cream precipitate soluble in conc. NH3(aq)
Iodide pale yellow precipitate not soluble in NH3(aq)

Aqueous lead(II) ions react with halide ions to produce a different set of halide salts, see the table below,

Halide ion Observation with Pb(NO3)2(aq)
Chloride no precipitate
Bromide white precipitate
Iodide bright yellow precipitate

Concentrated acids react with solid halide salts to produce a variety of results, depending on the acid.  Strong oxidising acids, such as concentrated sulphuric acid, form a variety of products depending on the halide used. With chloride ions, only hydrogen chloride is formed; with bromide ions, bromine and sulphur dioxide are formed; and with iodide ions, iodine, sulphur dioxide and hydrogen sulphide are formed.

A normal strong acid, such as phosphoric(V) acid only form the hydrogen halide and a phosphate(V) salt.

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Halogens - Industrial Resources

Sodium chloride is a great natural resource, present in great abundance in sea water and underground lakes, as a concentration aqueous solution called brine. This brine also contains a myriad of other compounds, including bromide and iodide salts as well as more gold than is present on land.

The sodium chloride can be obtained from sea water by simply evaporating the water from the solution and collecting the solid left behind. This generally requires a lot of energy to evaporate the water so is not always economically viable.

The sodium chloride solution can be used to make pure sodium metal or sodium hydroxide (a powerful caustic base) and chlorine gas itself.

Solid sodium chloride, along with other salts, is spread on icy roads, as it lowers the freezing (or melting) point of water, thus causing any ice already present to melt. This can only work at temperatures higher than -10 oC or so, otherwise other salt mixtures may be needed.

Chlorine by itself is used as a bleach and in the manufacture of sodium chlorate, which can be used as a bleach and a herbicide.

Water purification also relies on chlorine to kill bacteria in the water, after the impure water has passed through various filtration stages.

Chlorine is also used in the production of chlorofluorocarbons, commonly called CFC's, used in the past as refrigerant gases and propellants for aerosol cans. Both these uses have now been banned by international law, in the developed world at least.

The problem these chemicals cause is that when they reach the high atmosphere the molecules break apart to release chlorine atoms. These chlorine atoms then react with molecules of ozone, O3, turning them into molecules of oxygen. One chlorine atom can destroys thousands of molecules of ozone.

Ozone traps UV light from the sun, preventing it hitting the surface of the Earth, and with a depletion of ozone more UV light gets through, which increases occurrences of skin cancer in humans.

Unfortunately, in place of CFC's, alkanes are used and these lead to an increase in the amount of green house gases and an increase in global temperatures.

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Written by Dr Richard Clarkson : © Saturday, 1 November 1997

Updated : Saturday, 14th July, 2012

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