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 a highly radioactive element; the most stable isotope of astatine 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 commonly found combined in compounds in rocks and minerals. Fluorine, chlorine and bromine toxic and corrosive elements 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. The image below shows the state and appearance of the first four halogens at room temperature.
Fluorine at the top of group 7 is a pale green yellow toxic gas, it is perhaps one of the most reactive elements in the periodic table. Chlorine is a greeny-yellow gas that is also very toxic and reactive. Chlorine 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, if a small amount if placed in a flask it will quickly evaporate to fill the flask with red-brown bromine vapour. Bromine like fluorine and chlorine is a very toxic and dangerous element with a very unpleasant bleach like smell to it. Iodine is a grey-black solid with a metallic sheen at room temperature. If heated it sublimes; that is it turns directly from a solid to a gas when heated; to form a purple iodine gas
The halogens "go around in pairs"- that is they form molecules made up of two atoms as shown in the image below. 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 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 halogen molecules get larger and their relative mass molecular increases. Larger molecules with more electrons will results in stronger intermolecular Van der Waals 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 | outer electron configuration | atomic radius/nm | electronegativity |
|---|---|---|---|---|---|---|---|
| fluorine | pale yellow | -220 | -188 | gas | 2s22p5 | 0.071 | 4.0 |
| chlorine | greenish-yellow | -101 | -34 | gas | 3s23p5 | 0.099 | 3.0 |
| bromine | red-brown | -7 | 59 | liquid | 4s24p5 | 0.114 | 2.8 |
| iodine | greyish-purple | 114 | 131 | solid | 5s25p5 | 0.133 | 2.5 |
The electronegativity is the power or ability of an atom to attract the electron density in a covalent bond. Looking at the electronegativity values for the halogens in the table above the trend again is obvious:
The halogens are covalent non-polar molecules and as such are not particularly soluble in water but they will readily dissolve in organic solvents such as cyclohexane or tetrachloromethane. Chlorine will dissolve to a certain extent in water to form a pale green solution called chlorine water. This is an example of a disproportionation reaction; here the chlorine is both oxidised and reduced. When chlorine dissolves in water it forms a mixture of the weak acid chloric (I) acid and the strong acid hydrochloric acid.
Chlorine being an element has an oxidation state
of O, but when it forms hydrochloric acid its oxidation state changes to -1
whereas in
chloric (I) acid the oxidation state of the chlorine is +1, so the
chlorine has been both oxidise and
reduced in this reaction. Bromine
dissolves
in a very similar way to chlorine to form hydrobromic
and bromic (I) acids. Bromine is also more
soluble in water than chlorine.
Iodine is practically insoluble in water but it does dissolve in an aqueous solution of potassium iodide. When added to a potassium iodide solution the iodine molecules react with the soluble iodide ions (I-) to form triiodide ions (I3-). The solution formed is often labelled as iodine solution; as shown in the image opposite.
In the lab we may use a bottle of "iodine solution" which has a pale yellow colour when dilute but its colour becomes a dark orange-brown colour when its concentration increases. However despite limited solubility in water the halogens are very soluble in organic solvents such as cyclohexane. In organic solvents the halogens generally dissolve to form solutions with bright clear colours. The rather dull red-orange colour of an aqueous iodine solution is replaced by a vivid purple solution when iodine dissolves in cyclohexane, as shown in the image below. Bromine dissolves freely in cyclohexane to form a red-brown solution and chlorine forms a yellow-green solution when dissolved in cyclohexane; as shown in the image below:
Fluorine does not dissolve in water but instead reacts violently with water to form a mixture of hydrofluoric acid, oxygen and ozone (03) gases.
Or if an excess of fluorine is used:
All the halogens have 7 electrons in their outer valence shell and have a np6 electronic configuration so they only need to gain one electron to achieve a full octet of electrons. 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. The reactions of the halogens with reactive metals such as those in groups I and II in the periodic table follow the trend you might expect, the more reactive the metal and the more reactive the halogen the more violent is the reaction. For example the alkali metal sodium reacts violently with chlorine to form the ionic compound sodium chloride, this is shown in the equations and image below:
and for magnesium:
Even when less reactive metals such as iron (in iron wool for example) are used the results are the same; in each case the metal is oxidised and the halogen is reduced e.g. all the halogens react with iron wool. The trends are what you might expect:
Perhaps some of the most visually spectacular reactions involving the halogens you are likely to see in school is in their reactions with aluminium metal. Here:
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 any 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.
Perhaps one of the best reactions to show the reactivity trends in the halogens is in their reaction with hydrogen gas. All the halogens react with hydrogen to form hydrogen halide vapours, as shown in the equation below:
This trend in reactivity is due mainly is to the decreasing bond strength of hydrogen halides as you go down the group, which makes the formation of hydrogen halides less favorable energetically.