General Introduction to p-Block Elements

The p-block of the periodic table consists of Groups 13 through 18 — the elements in which the differentiating (last-entering) electron occupies a $p$-orbital of the outermost shell. Helium, though its configuration is $1s^2$, is conventionally placed with the noble gases of Group 18 rather than the $s$-block. The general valence-shell configuration across the p-block runs from $ns^2np^1$ at Group 13 to $ns^2np^6$ at Group 18, with one additional $p$-electron added at each successive group. This chapter concentrates on the first two p-block groups: Group 13, the boron family (B, Al, Ga, In, Tl), with the general configuration $ns^2np^1$, and Group 14, the carbon family (C, Si, Ge, Sn, Pb), with the general configuration $ns^2np^2$.

Several broad trends recur across every p-block group and set the stage for the detailed chemistry of boron, aluminium, carbon and silicon that follows. Atomic and ionic size generally increases on descending a group as successive shells are added, though the increase from the second-period element to the third-period element is often unusually small (boron to aluminium, or carbon to silicon) because the intervening $d$- and $f$-electrons in the heavier members shield the nuclear charge poorly, partially offsetting the extra shell. Ionisation enthalpy generally decreases down a group as the valence electrons lie farther from the nucleus and are more shielded, but this decrease is often irregular — for instance the first ionisation enthalpy of gallium (Group 13) is anomalously close to that of aluminium rather than clearly lower, because gallium's valence electron experiences poor shielding from the ten intervening $3d$ electrons. Electronegativity tends to decrease down a group overall, though again with irregularities in the heavier p-block metals. Metallic character increases down every p-block group — the lightest member of a group (boron, carbon) is typically a nonmetal or metalloid, while the heaviest members (thallium, lead) are distinctly metallic, softer, and form more ionic compounds.

Oxidation states in the p-block are dominated by the group oxidation state — the number of valence electrons available for bonding, equal to the group number minus 10 (so $+3$ for Group 13, $+4$ for Group 14). However, on descending a group, an oxidation state two less than the group state becomes progressively more stable — the well-known inert pair effect, in which the outermost $ns^2$ electron pair becomes increasingly reluctant to participate in bonding for the heavier elements, due to a combination of relativistic contraction of the $ns$ orbital and the poor shielding provided by filled inner $d$/$f$ subshells. This is why thallium favours $+1$ over $+3$, and lead favours $+2$ over $+4$, even though $+3$ and $+4$ respectively are each group's nominal oxidation state.

Finally, the trend in chemical reactivity across a p-block group mirrors its trend in metallic character: compounds of the lighter members tend to be covalent and, where acidic/basic character applies, more acidic (e.g. $\text{B}_2\text{O}_3$, $\text{CO}_2$), while compounds of the heavier members become increasingly ionic and basic or amphoteric (e.g. $\text{Tl}_2\text{O}_3$, $\text{PbO}$). These themes — anomalous first-member behaviour, the inert pair effect, and increasing metallic/basic character down the group — recur throughout the detailed study of boron, aluminium, carbon and silicon chemistry in the sections that follow.