Chemical Properties

Oxidation States and Trends in Chemical Reactivity

Group 16 elements display a wide range of oxidation states (Table 7.6). The stability of the $-2$ state falls steadily down the group, and polonium barely shows it at all. Because oxygen's electronegativity is so high, it is essentially restricted to the $-2$ state, with a rare exception in $OF_2$, where fluorine — being even more electronegative — forces oxygen into a $+2$ state. The heavier members of the group can instead show $+2$, $+4$ and $+6$ states, of which $+4$ and $+6$ are the most commonly encountered: sulphur, selenium and tellurium typically show $+4$ in their oxygen compounds and $+6$ with fluorine. Down the group, the $+6$ state becomes progressively less stable while the $+4$ state becomes more stable — the familiar inert-pair effect seen elsewhere in the p-block. Bonding in both the $+4$ and $+6$ states is predominantly covalent.

Anomalous Behaviour of Oxygen

As with the first member of every p-block group, oxygen behaves anomalously compared with its heavier congeners, owing to its small size and high electronegativity. One direct consequence is the strong hydrogen bonding present in $H_2O$ — an effect essentially absent in $H_2S$, despite the structural similarity of the two molecules.

A second consequence is oxygen's limited covalency: lacking accessible d orbitals, oxygen's covalency is capped at four and in practice rarely goes beyond two. The heavier elements of the group, which do have accessible d orbitals, can expand their valence shells and readily exceed a covalency of four.

(i) Reactivity with Hydrogen

All Group 16 elements form hydrides of the general type $H_2E$ (E = O, S, Se, Te, Po), whose properties are compared in Table 7.7:

Reading across the table, the acidic character of these hydrides increases from $H_2O$ to $H_2Te$, a trend traceable to the falling $H-E$ bond dissociation enthalpy down the group — a weaker bond ionises more readily. That same weakening bond also means thermal stability falls from $H_2O$ to $H_2Po$. Every hydride except water is a reducing agent, and this reducing power strengthens further from $H_2S$ to $H_2Te$.

(ii) Reactivity with Oxygen

The elements form oxides of two general types, $EO_2$ and $EO_3$ (E = S, Se, Te or Po). Ozone ($O_3$) and sulphur dioxide ($SO_2$) are gases, while selenium dioxide ($SeO_2$) is a solid. The reducing power of the dioxides falls on descending the group: $SO_2$ is a reducing agent, while $TeO_2$ has flipped to being an oxidising agent. Sulphur, selenium and tellurium additionally form the $EO_3$-type trioxides ($SO_3$, $SeO_3$, $TeO_3$). Both the dioxide and trioxide series are acidic in nature.

(iii) Reactivity towards the Halogens

Group 16 elements form an extensive series of halides of the general types $EX_6$, $EX_4$ and $EX_2$. Across all of these, halide stability falls in the order $F^- > Cl^- > Br^- > I^-$.

Among the hexahalides, only the hexafluorides are stable, and all of these are gases with an octahedral structure; sulphur hexafluoride, $SF_6$, is exceptionally stable for purely steric reasons — its six fluorine atoms pack tightly and shield the central sulphur from attack. …