The group 7 non-metals are called halogens. The halogens are fluorine, chlorine, bromine and iodine. Astatine at the bottom of group 7 is a very rare element and 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 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.
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 temperarture, 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.
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 molecules get larger and their relative mass molecular increases. Larger molecules 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|
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 is obvious:
The halogens are covalent non-polar molecules and as such are not particularly soluble in water, the solubility decreases as we descend group 7. Chlorine will dissolve in water to form a 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 hydrochloric aci.
Chlorine being an element has an oxidation state
of O, but when it forms hydrochloric aci its oxidation state changes to -1
chloric (I) acid the oxidation state of the chlorine is +1, so the
chlorine has been both oxidise and
reduced in this reaction. Bromine
in a very similar way to chlorine to forming hydrobromic
and bromic (I) acids but bromine is less
soluble in water than chlorine.
In the lab we may use a bottle of "iodine solution" which has a pale yellow when dilute but its colour becomes a dark orange-brown colour when its concentration increases. However 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 iodine solution is replaced by a vivid purple solution when iodine dissolves in cyclohexane, as shown in the image. Bromine dissolves freely in cyclohexane to form a red-brown solution and chlorine forms a yellow-green solution when dissolved in cyclohexane.
All the halogens have 7 electrons in their outer shell and have a np6 electronic configuration, so only need to gain one electron to achieve a full outer electron shell. 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:
Perhaps one of the best reactions to show the reactivity trends in the halogens is the reaction with hydrogen gas. All the halogens react with hydrogen to form hydrogen halide vapours:
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.