The p-Block Elements: Nitrogen and Oxygen Families

Nitrogen and phosphorus sit in the same group and share the same ns2 np3 valence configuration, yet their elemental forms behave almost like different elements entirely. Nitrogen exists as N2, two nitrogen atoms joined by a genuine triple bond (one sigma and two pi bonds) with a bond enthalpy of roughly 945 kJ/mol, one of the strongest bonds known in the whole of chemistry, which is exactly why N2 is so famously unreactive at ordinary temperatures, resisting breakdown even though breaking it apart is precisely what industrial nitrogen fixation (the Haber process) has to overcome. Phosphorus, by contrast, cannot form an equivalent P#P triple bond, since its larger 3p orbitals overlap far too poorly for strong pi bonding between two phosphorus atoms (the very same orbital-overlap limitation that stops silicon forming Si=Si bonds), so elemental phosphorus instead exists as P4, a tetrahedral cluster of four phosphorus atoms joined only by single bonds, each carrying considerable angle strain (bond angles of just 60 degrees, far from phosphorus's preferred angle), which is exactly why white phosphorus is so much more reactive, even dangerously so, than nitrogen gas ever is.

Ammonia, NH3, is sp3 hybridised with a lone pair occupying the fourth position, giving it a trigonal pyramidal shape (bond angle 107 degrees, slightly compressed from the ideal 109.5 degrees by the lone pair's stronger repulsion against the three bonding pairs), and this same lone pair is what makes ammonia both a Lewis base (donating its electron pair to form coordinate bonds, as in the many metal-ammine complexes covered elsewhere in this syllabus) and a Bronsted base (accepting a proton to form NH4+). Industrially, ammonia is manufactured by the Haber process, N2 + 3H2 to (equilibrium, Fe catalyst, high pressure, moderate temperature) 2NH3, a direct application of Le Chatelier's principle to overcome nitrogen's own famous unreactivity discussed already, and it remains one of the most consequential industrial reactions in human history, since the great majority of synthetic nitrogen fertiliser manufactured worldwide begins with ammonia made this way.

Nitric acid, HNO3, is manufactured industrially by the Ostwald process, a three-stage catalytic oxidation of ammonia: ammonia is first oxidised over a platinum-rhodium gauze catalyst to nitric oxide (4NH3 + 5O2 to 4NO + 6H2O), the nitric oxide is then further oxidised by atmospheric oxygen to nitrogen dioxide (2NO + O2 to 2NO2), and the nitrogen dioxide is finally absorbed in water to give dilute nitric acid (3NO2 + H2O to 2HNO3 + NO, with the NO recycled back into the process). Structurally, HNO3 has one N-OH bond and two equivalent N-O bonds sharing a single delocalised double-bond character, exactly analogous to the resonance seen in the nitrate ion itself. What makes nitric acid genuinely distinctive as an acid is that, unlike hydrochloric or dilute sulphuric acid, it acts as a strong oxidising agent even when reacting with metals, so it typically does not simply liberate hydrogen gas; instead it produces a mixture of reduced nitrogen products (NO2 with concentrated acid, NO with dilute acid, and even NH4NO3 with very dilute acid and a sufficiently reactive metal) depending on the acid's concentration and the metal's own reactivity.

Nitrogen shows a remarkable spread of oxidation states in its oxides, more than almost any other element manages, spanning from +1 in nitrous oxide (N2O, 'laughing gas', used as a mild anaesthetic) through +2 in nitric oxide (NO, a genuine free radical with an odd number of electrons, and, notably, a real biological signalling molecule in the human body), +3 in dinitrogen trioxide (N2O3, the acid anhydride of nitrous acid), +4 in nitrogen dioxide (NO2, a reddish-brown paramagnetic gas that dimerises to colourless N2O4 at low temperature), and finally +5 in dinitrogen pentoxide (N2O5, the acid anhydride of nitric acid). This range exists because nitrogen's own ns2 np3 valence electrons can combine with oxygen in so many different ratios, and it is a genuinely favourite examiner's trick to test recognition of each oxide's formula, oxidation state, and corresponding acid or base character.

White phosphorus (P4, tetrahedral, as introduced already) is soft, waxy, and highly reactive, catching fire spontaneously in air (hence being stored underwater) and glowing faintly green in the dark, a phenomenon called chemiluminescence caused by its slow oxidation, and it is also genuinely toxic. Red phosphorus is formed by heating white phosphorus in the absence of air, and structurally it is polymeric, chains of linked PP4 tetrahedral units, which is exactly why it is far less reactive, does not catch fire spontaneously, is non-toxic, and does not glow, all direct structural consequences of no longer being made of small, strained, discrete P4 units. Black phosphorus, the thermodynamically most stable form, has an even more extended, layered structure (somewhat resembling graphite) and is correspondingly the least reactive of the three, though it is the least industrially significant of the allotropes covered here.

Phosphorus trichloride, PCl3, is sp3 hybridised with one lone pair, giving it the same trigonal pyramidal shape as ammonia, and it hydrolyses readily in moist air, releasing choking fumes of HCl. Phosphorus pentachloride, PCl5, is sp3d hybridised in the gas and liquid phases, giving it a genuine trigonal bipyramidal shape with two distinct bond lengths (the three equatorial P-Cl bonds are shorter and stronger than the two axial ones, since equatorial positions experience less repulsion from neighbouring bonding pairs), but in the solid state PCl5 exists instead as an ionic compound, [PCl4]+[PCl6]-, a tetrahedral cation paired with an octahedral anion, a genuinely striking structural change purely from a change of physical state that examiners love to test directly. Both halides hydrolyse in water, PCl3 giving phosphorous acid (H3PO3) and HCl, and PCl5 giving phosphoric acid (H3PO4) and HCl, in each case the halide behaving as an acid chloride analogue reacting with the nucleophilic water.

Phosphorous acid, H3PO3, and phosphoric acid, H3PO4, look deceptively similar from their formulas alone, both containing exactly three hydrogen atoms, yet H3PO3 is dibasic (only two of its three protons are ionisable) while H3PO4 is genuinely tribasic (all three are ionisable), and the reason is entirely structural, not a matter of counting hydrogen atoms. In H3PO3, one of the three hydrogen atoms is bonded directly to phosphorus (a P-H bond) rather than to an oxygen atom, and only hydrogens bonded through an oxygen (as P-O-H) are acidic enough to ionise as protons in water, since only the electronegative oxygen can adequately stabilise the resulting negative charge; the P-H hydrogen simply isn't acidic at all. H3PO4, by contrast, has all three of its hydrogens bonded through oxygen (three genuine P-OH groups), making all three ionisable and the acid tribasic. This distinction, basicity tracks the number of P-OH groups specifically, not the total hydrogen count in the molecular formula, is one of the most commonly tested, and most commonly misunderstood, ideas in this entire chapter.

Oxygen's other allotrope, ozone (O3), has a bent, angular structure (bond angle about 117 degrees) best described by resonance between two equivalent structures, each with one O=O double bond and one O-O single bond, so the two O-O bonds in real ozone are actually identical, intermediate in length between a true single and true double bond. Ozone is a considerably stronger oxidising agent than ordinary O2, decomposing to release a highly reactive oxygen atom (O3 to O2 + [O]), a property put to use in water purification and air treatment. Ozone's role depends entirely on where it is found: in the stratosphere it forms a protective layer that absorbs the sun's harmful ultraviolet radiation, a genuinely essential shield for life on Earth's surface, while at ground level, produced as a secondary pollutant from vehicle exhaust reacting with sunlight, it is a harmful component of photochemical smog, the same molecule acting as protector in one location and pollutant in another purely as a matter of altitude.

Sulphur's two common crystalline allotropes, rhombic (alpha) sulphur and monoclinic (beta) sulphur, are both built from the same puckered, crown-shaped S8 ring, differing only in how those rings pack together in the solid crystal; rhombic sulphur is stable below 369 K, while heating it converts it to monoclinic sulphur, which is stable above that temperature, and the two forms are genuine, textbook examples of allotropy caused by different crystal packing of an identical molecular unit, directly comparable to diamond and graphite's own difference despite both being built from the same S8 ring rather than differently-bonded atoms. Sulphur dioxide, SO2, itself has a bent, resonance-stabilised structure much like ozone's, and it plays a central role as the direct precursor to sulphuric acid manufacture, formed either by burning elemental sulphur in air or by roasting metal sulphide ores.

Sulphuric acid is manufactured by the Contact process, a three-stage sequence: elemental sulphur is burned in air to sulphur dioxide (S + O2 to SO2), the sulphur dioxide is catalytically oxidised over vanadium(V) oxide to sulphur trioxide (2SO2 + O2 to 2SO3), and the sulphur trioxide is absorbed, not directly into water (which would create an uncontrollable, corrosive mist), but into concentrated sulphuric acid itself, forming oleum (H2S2O7, disulphuric acid), which is only afterward diluted carefully with water to give sulphuric acid at the desired concentration. Concentrated sulphuric acid displays three genuinely distinct chemical roles worth telling apart: as a powerful dehydrating agent, it can pull the elements of water directly out of other compounds (charring sugar to carbon is the classic demonstration); as an oxidising agent, hot concentrated H2SO4 can oxidise both metals and non-metals, being itself reduced to SO2; and as an acid, it is a strong dibasic acid, forming both normal sulphate and acid hydrogensulphate salts, all three roles operating through genuinely different mechanisms despite coming from the very same reagent bottle.