Chemical Properties

Chemical Properties of the Halogens

Oxidation States and the Origin of +1, +3, +5, +7

Every halogen shows the –1 oxidation state, simply by picking up the one electron it needs to complete its octet. But chlorine, bromine and iodine go further: each of them can also display the +1, +3, +5 and +7 oxidation states.

The orbital-promotion diagram accompanying this section explains how. It lays out a halogen atom's valence-shell boxes for the $ns$, $np$ (three boxes) and $nd$ (five boxes) sub-shells across four successive electronic states:

• In its ground state (for any halogen other than fluorine), the atom has configuration $ns^2np^5$ — the $ns$ pair and two of the three $np$ pairs are already paired up, leaving just one unpaired electron. This single unpaired electron accounts for the –1 oxidation state (by gaining an electron) or the +1 state (by losing/sharing it).
• In the first excited state, one of the paired $np$ electrons is promoted into an empty $nd$ orbital, freeing up a third unpaired electron and giving the +3 oxidation state.
• In the second excited state, a second electron moves from $np$ into $nd$, producing five unpaired electrons in all and the +5 oxidation state.
• In the third excited state, even the previously untouched $ns$ pair is broken open and promoted, giving a full seven unpaired electrons — the maximum possible — and the +7 oxidation state.

Each step of this promotion sequence needs a spare $nd$ orbital to receive the promoted electron, which is exactly why the trick fails for fluorine: fluorine has no d orbitals in its valence shell at all, so it can never expand its octet. Being also the most electronegative element there is, fluorine is confined strictly to the –1 oxidation state.

The higher oxidation states (+3, +5, +7) of chlorine, bromine and iodine only really appear when these halogens are bonded to small, strongly electronegative partners — chiefly fluorine and oxygen — as seen in interhalogen compounds, halogen oxides and oxoacids. Chlorine and bromine also form oxides and oxoacids showing the +4 and +6 oxidation states.

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Reactivity Trends and Oxidising Power

All the halogens are highly reactive, combining readily with both metals and non-metals to form halides — though this reactivity steadily decreases down the group.

The reason the halogens are such strong oxidising agents is the flip side of their electron-gain-enthalpy story: they accept an extra electron very readily. $F_2$ is the most powerful oxidiser of the whole group — strong enough to oxidise the halide ions Cl⁻, Br⁻ or I⁻, whether in solution or even in the solid state. As a general rule, a lighter halogen will oxidise the halide ion of a heavier halogen:

This falling oxidising power down the group is captured quantitatively in the standard electrode potentials ($E^\circ$ for $X_2(g) + 2e^- \rightarrow 2X^-(aq)$: F 2.87 V, Cl 1.36 V, Br 1.09 V, I 0.54 V). These potentials are really the net outcome of three successive energy steps — bond dissociation of $X_2$, electron gain by the atom, and hydration of the resulting ion:

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Reaction with Water

The halogens' relative oxidising strength also shows up clearly in how each reacts with water. Fluorine actually oxidises water itself, liberating oxygen gas:

Chlorine and bromine, being milder oxidisers, instead undergo a simple addition-type reaction with water to give a hydrohalic acid together with the corresponding hypohalous acid:

Iodine's reaction with water is non-spontaneous — in fact the reverse process happens instead, with iodide ion being oxidised by dissolved oxygen under acidic conditions, exactly the mirror image of fluorine's behaviour:

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The Anomalous Behaviour of Fluorine

As the first member of its group and a period-2 p-block element, fluorine breaks rank with the rest of Group 17 in several ways. Its ionisation enthalpy, electronegativity and electrode potential are all higher than the group trend would predict, while its ionic and covalent radii, melting and boiling points, bond dissociation enthalpy, and electron gain enthalpy are all lower than expected. The root causes are fluorine's very small size, its unmatched electronegativity, the surprisingly weak F–F bond, and the total absence of d orbitals in its valence shell.

Because the bonds fluorine forms with other elements are small and strong, most of its reactions are strongly exothermic. It is also restricted to forming just one oxoacid, unlike the other halogens, which form several. And whereas the other hydrogen halides are gases at ordinary temperatures, hydrogen fluoride is a liquid (boiling point 293 K) — a direct consequence of the strong hydrogen bonding that fluorine's small size and extreme electronegativity make possible.

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(i) Reactivity Towards Hydrogen

Every halogen reacts with hydrogen to form the corresponding hydrogen halide, though the ease of this reaction — the halogen's affinity for hydrogen — falls off from fluorine to iodine. Dissolved in water, these hydrogen halides give the hydrohalic acids.

Two opposite trends jump out of this table. The acid strength of the hydrohalic acids increases down the group, $HF < HCl < HBr < HI$ (notice how the pKa values fall sharply and go negative), while the thermal stability of the H–X molecule itself decreases down the group, tracking the steady fall in the H–X bond dissociation enthalpy: $H\text{–}F > H\text{–}Cl > H\text{–}Br > H\text{–}I$.

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(ii) Reactivity Towards Oxygen

The halogens form a wide range of oxides with oxygen, though most of these are thermally unstable.

Fluorine forms only two: $OF_2$ and $O_2F_2$, and of these only $OF_2$ is thermally stable at 298 K. Because fluorine is more electronegative than oxygen, these compounds are more accurately described as oxygen fluorides rather than fluorine oxides, and both are powerful fluorinating agents. $O_2F_2$, notably, can oxidise plutonium metal all the way to $PuF_6$ — a reaction actually used to help recover plutonium from spent nuclear fuel.

Chlorine, bromine and iodine form a much broader range of oxides, spanning oxidation states from +1 to +7. Their thermal stability follows the order $I > Cl > Br$, arising from a combination of kinetic and thermodynamic factors:

• Iodine's oxides are the most stable, thanks to the high polarisability of the I–O bond.
• Chlorine oxides gain extra stability from multiple-bond character between Cl and O, made possible by chlorine's accessible d orbitals.
• Bromine, lacking both of these stabilising features, forms the least stable oxides of the three — sometimes called the "middle-row anomaly."

As a general rule, the higher oxides of a given halogen tend to be more stable than its lower oxides.

• Chlorine oxides — $Cl_2O$, $ClO_2$, $Cl_2O_6$ and $Cl_2O_7$ — are all highly reactive, powerful oxidising agents, and several tend to explode. $ClO_2$ in particular is used as a bleaching agent for paper pulp and textiles and in water treatment. …