General Introduction to p-Block Elements (Groups 15-18)

Groups 15 through 18 close out the p-block of the periodic table. Group 15, the nitrogen family (N, P, As, Sb, Bi), has the general valence configuration $ns^2np^3$; Group 16, the

oxygen family (O, S, Se, Te, Po), has $ns^2np^4$; Group 17, the halogens (F, Cl, Br, I,

At), has $ns^2np^5$; and Group 18, the noble gases (He, Ne, Ar, Kr, Xe, Rn), has a complete

outer shell, $ns^2np^6$ (or $1s^2$ for helium). Moving across these four groups, the p sub-shell

fills from three electrons to a full six, and this single fact drives almost everything that

follows: increasing electronegativity and oxidising power left to right (Group 15 elements are

comparatively mild, halogens are ferociously reactive, noble gases are essentially unreactive),

a widening range of accessible oxidation states as more electrons become available for bonding

or lone pairs, and a systematic tightening of atomic radius across each period as the increasing

nuclear charge pulls the (poorly-shielding) added p-electrons closer in.

As in the earlier p-block groups, the first member of each group behaves anomalously compared

with its heavier congeners: nitrogen, oxygen and fluorine are all markedly smaller, more

electronegative, and structurally different from phosphorus, sulphur and chlorine respectively,

largely because they lack accessible $d$ orbitals in their valence shell ($n=2$) and so cannot

expand their coordination number the way period-3-and-below elements can. This is why phosphorus

forms $\text{PCl}_5$ but nitrogen cannot form an analogous $\text{NCl}_5$, why sulphur reaches

$+6$ in $\text{SO}_3/\text{H}_2\text{SO}_4$ but oxygen itself is essentially capped at $-2$

(showing positive states only with fluorine), and why chlorine, bromine and iodine build up a

full family of oxoacids from $+1$ to $+7$ while fluorine, the most electronegative element of

all, shows only the $-1$ state.

The inert pair effect, met already in Groups 13 and 14, recurs here too: down Group 15,

bismuth increasingly favours $+3$ over the group state $+5$ ($\text{Bi}^{5+}$ compounds are

strong oxidisers), and an analogous trend runs down Group 16. Down Group 17, oxidising power

falls steadily from fluorine to iodine even though ionisation enthalpy and electronegativity both

fall too -- a reminder that halogen reactivity is governed less by electronegativity alone than

by the overall thermodynamics of the reduction step, which favours the smaller halogens.

Finally, Group 18 stands apart from the rest of the p-block precisely because its members already

possess a complete octet (or, for helium, a full $1s^2$ shell): with no vacant low-energy orbital

readily available for ordinary covalent bonding, the noble gases were long believed entirely

inert, until xenon's chemistry (discussed in detail later in this chapter) showed that even a

"closed shell" atom can be made to bond under the right conditions. Across these four groups,

then, the chapter traces one continuous story: the same $ns^2np^x$ framework, applied four times

over, producing four completely different chemical personalities -- covalent and life-essential

(nitrogen), oxidatively central (oxygen and sulphur), aggressively reactive (the halogens), and,

almost, entirely unreactive (the noble gases).