Chemical Properties

Oxidation states and trends in chemical reactivity

The elements of Group 15 most commonly display the oxidation states $-3$, $+3$ and $+5$. Descending the group, size and metallic character both increase, and this steadily weakens the tendency to adopt the $-3$ state — by the time bismuth is reached, compounds in the $-3$ oxidation state are essentially unknown. The $+5$ state shows the opposite trend in reverse: it becomes progressively less stable down the group, so that bismuth forms only one well-characterised $+5$ compound, $BiF_5$. This falling stability of $+5$ alongside a rising stability of $+3$ (favoured by the inert-pair effect) is one of the defining chemical trends of the group.

Nitrogen is unusual in also showing the $+1$, $+2$ and $+4$ oxidation states when it combines with oxygen, over and above $-3$, $+3$ and $+5$. Curiously, though, it never reaches a $+5$ state with the halogens, because nitrogen has no available $d$ orbitals to house electrons donated by another atom while forming extra bonds. Phosphorus, similarly, adopts $+1$ and $+4$ states in a handful of its oxoacids.

For nitrogen, every intermediate oxidation state between $+1$ and $+4$ tends to disproportionate in acidic solution — for example, nitrous acid disproportionates to nitric acid and nitric oxide:

Phosphorus behaves similarly: almost all of its intermediate oxidation states disproportionate into the $+5$ and $-3$ states, whether in acid or in alkali. Moving further down the group, however, the $+3$ state of arsenic, antimony and bismuth becomes progressively more resistant to disproportionation.

Because only four orbitals — one $s$ and three $p$ — are available for bonding, nitrogen's maximum covalency is limited to four. The heavier elements of the group, in contrast, have empty $d$ orbitals in their outer shell that can be pressed into service for extra bonds, letting them expand their covalency beyond four, as seen in the octahedral ion $PF_6^-$.

Anomalous properties of nitrogen

Nitrogen stands apart from the rest of the group because of its small size, high electronegativity, high ionisation enthalpy, and the complete absence of $d$ orbitals in its valence shell. It has a distinctive capacity to form strong $p\pi$–$p\pi$ multiple bonds, both with itself and with other small, highly electronegative atoms such as carbon and oxygen. The heavier elements cannot do this: their atomic orbitals are too large and diffuse to overlap effectively side-on. Consequently nitrogen exists as the diatomic molecule $N_2$, held together by a genuine triple bond (one $\sigma$ and two $\pi$), which gives it an exceptionally high bond enthalpy of 941.4 kJ mol$^{-1}$. Phosphorus, arsenic and antimony, unable to form such multiple bonds, instead settle for single $P$–$P$, $As$–$As$ and $Sb$–$Sb$ bonds in their elemental forms, while bismuth exists as a metal held together by metallic bonding.

Despite nitrogen's fondness for multiple bonding, its single $N$–$N$ bond is actually weaker than a single $P$–$P$ bond — the short $N$–$N$ bond length forces the non-bonding electrons on the two nitrogen atoms unusually close together, and the resulting interelectronic repulsion destabilises the bond. This is why nitrogen shows a comparatively weak tendency towards catenation.

The missing $d$ orbitals also mean that nitrogen cannot form $d\pi$–$p\pi$ bonds of the kind seen in the heavier elements — for instance $R_3P=O$ or $R_3P=CH_2$ (where $R$ is an alkyl group) have no nitrogen analogue. Phosphorus and arsenic can go a step further and form $d\pi$–$d\pi$ bonds with transition metals, which is why compounds such as $P(C_2H_5)_3$ and $As(C_6H_5)_3$ are able to act as ligands.

(i) Reactivity towards hydrogen. Every element of Group 15 forms a hydride of the general type $EH_3$ (where $E$ = N, P, As, Sb or Bi). Some properties of these hydrides are set out in Table 7.2 below, and they show a smooth gradation down the group.

The bond-dissociation enthalpy falls steadily from $NH_3$ to $BiH_3$, which is the same trend as the falling stability of the hydrides down the group — and, as a direct consequence, their reducing character rises. Ammonia is only a mild reducing agent, while $BiH_3$ is the strongest reductant of the whole series. Basicity follows the order $NH_3 > PH_3 > AsH_3 > SbH_3 \geq BiH_3$. Because nitrogen is both small and highly electronegative, $NH_3$ engages in hydrogen bonding in both the solid and liquid states — exactly as water does — which is why its melting and boiling points are noticeably higher than its molecular mass alone would predict, and higher than those of $PH_3$. …