3.1.3 - The Halogens


The halogens are the first four elements found in group 7 of the periodic table. As a result, their atoms all have electronic configurations ending in p$\small^5$, meaning they are only one electron away from filling the highest occupied energy level.

The halogens are non-metals, which means they make covalent bonds with other non-metals, including themselves. In fact, since halogen atoms are one electron away from stability, they are commonly encountered as diatomic molecules (molecules with two of the same atom), making one covalent bond with another.

The fact that their molecules are two of the same atom make them non-polar, as the two atoms both have the same electronegativity. As a result, certain interactions do not exit between the molecules, such as the permanent dipole-permanent dipole and permanent dipole-induced dipole interactions. The lack of these relatively strong intermolecular forces means that the melting and boiling points of the non-metal halogens are low. This is why fluorine and chlorine are gases at room temperature, and why bromine is one of only two elements (the other is mercury) that is liquid at room temperature.

That being said however, as the period increases down the halogens, the atomic radius also increases. The bigger the atoms, the greater the polarisation that cause LDFs, and so the boiling increases down the group.

Redox Reactions of the Halogens

As halogens are non-metals, they form anions and negative oxidative states in redox reactions. Because they are only one electron away from having a complete electron energy level, they typically form ions or atoms with charge or oxidative state -1.

However, fluorine is the only halogen that forms only one oxidative state. The others can be found with different oxidation numbers. For example, chlorine in bleach (NaClO) has oxidation number +1.

The reaction with which bleach is formed is a special case of redox reaction. The equation for the reaction is shown below (note that the elemental chlorine is a diatomic molecule):

$$ \text{Cl$_2$} _\text{(g)} + \text{2NaOH} _\text{(aq)} \rightarrow \text{NaClO} _\text{(aq)} + \text{NaCl} _\text{(aq)} + \text{H$_2$O} _\text{(l)} $$

As shown in 2.1.5 - Redox, the chlorine, which starts off with oxidation number 0, is both oxidised and reduced to form oxidative states +1 and -1. This is a disproportionation reaction.

Another example of a disproportionation reaction is the one between chlorine and water, in the absence of sunlight. In the presence of sunlight, the chloric(I) acid decomposes into hydrochloric acid, making it not a disproportionation reaction overall. The reaction for the first step is shown below:

$$ \text{Cl$_2$} _\text{(aq)} + \text{H$_2$O} _\text{(l)} \leftrightharpoons \text{HCl} _\text{(aq)} + \text{HClO} _\text{(aq)} $$

where chlorine goes from 0 to -1 in HCl and +1 HClO.

Reactivity of the Halogens

The ionisation energies of elements in any group decrease as the period increases. It is the same for halogens. However, in most reactions, halogens do not lose electrons, but gain one. The reactivity of the halogens is therefore not down to ionisation energy. Instead, the measure used to indicate the ease of gaining an electron to form a 1- ion is the first electron affinity of an element, also in $kJ mol^{-1}$.

As is perhaps expected, given that ionisation energy and electron affinity are opposites, the energy required to gain one electron instead increases as the period increases down the group. This is because the atomic radius increases, and the highest occupied energy level is increasingly further from the positively charged nucleus, and therefore it’s inward pulling attraction is less. In addition to atomic radius, there is more electron shielding from the previously occupied energy levels, acting as an electrostatic barrier between the highest energy level and the nucleus. Thus the atoms have less affinity for the extra electron, and so more energy is required to force them to accept it.

This is why the reactivity of the halogens decreases as the period increases down the group.

Characteristic Reactions of Halide Compounds

Those who did triple award science GCSEs will have already completed work on qualitative tests for ions in solution. Qualitative test are test-tube reactions conducted on a small sample of a solution to identify the ions in the solution. The qualitative aspect is a non-numerical positive result, usually a colour change or the production of an insoluble precipitate.

For halide ions, the qualitative test involves the addition of nitric acid to react with any non-halide ions, and then silver nitrate (AgNO$\small_3$) solution. The soluble silver (1+) ions combine with the soluble halide to form an insoluble precipitate of silver halide (AgX). This is the positive test result for the presence of halides.

The colour of the precipitate depends of the halide present. Chloride ions, when combined with silver ions, will produce a white precipitate; bromide ions, a cream precipitate; and iodide, a yellow precipitate.

For further clarification, if the colour is not discernible, the precipitates will dissolve in aqueous ammonia under certain conditions. Silver chloride will dissolve in dilute ammonia. Silver bromide will dissolve in concentrated ammonia, but not dilute ammonia. Silver iodide will not dissolve in dilute or concentrated ammonia.