Oxoacids & Allotropes of Phosphorus

Allotropes of phosphorus: white, red and black

Phosphorus is found in several allotropic forms, the important ones being white, red and black. They share the same element but differ sharply in structure, and that difference dictates everything from reactivity to toxicity. White phosphorus is the reactive, strained form; red and black are progressively more polymerised and far more stable.

White phosphorus is a translucent white waxy solid. It is poisonous, insoluble in water but soluble in carbon disulphide, and glows in the dark (chemiluminescence). It dissolves in boiling NaOH solution in an inert atmosphere, giving phosphine:

$$\ce{P4 + 3NaOH + 3H2O -> PH3 + 3NaH2PO2}$$

Structurally, white phosphorus consists of discrete tetrahedral P₄ molecules in which the P–P–P bond angles are only 60°. This angular strain makes it less stable and therefore more reactive than the other forms — it readily catches fire in air to give dense white fumes of P₄O₁₀:

$$\ce{P4 + 5O2 -> P4O10}$$

Red phosphorus is obtained by heating white phosphorus at 573 K in an inert atmosphere for several days. It possesses an iron-grey lustre, is odourless, non-poisonous and insoluble in water as well as in carbon disulphide, and does not glow in the dark. It is much less reactive than white phosphorus because it is polymeric — consisting of chains of P₄ tetrahedra linked together — so the strain of the isolated tetrahedron is relieved.

Black phosphorus, the most stable form, has two varieties. α-Black phosphorus forms when red phosphorus is heated in a sealed tube at 803 K; it can be sublimed in air, has opaque monoclinic or rhombohedral crystals, and does not oxidise in air. β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure, and does not burn in air up to 673 K.

Phosphine (PH₃): preparation, structure and properties

Phosphine is the hydride of phosphorus, the Group 15 analogue of ammonia. It is prepared by the action of water or dilute HCl on calcium phosphide:

$$\ce{Ca3P2 + 6H2O -> 3Ca(OH)2 + 2PH3}$$

$$\ce{Ca3P2 + 6HCl -> 3CaCl2 + 2PH3}$$

In the laboratory it is prepared by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO₂, giving phosphine and sodium hypophosphite:

$$\ce{P4 + 3NaOH + 3H2O -> PH3 + 3NaH2PO2}$$

When pure, phosphine is non-inflammable, but it becomes spontaneously inflammable owing to the presence of traces of P₂H₄ (diphosphine) or P₄ vapour. To purify it, the gas is absorbed in HI to form phosphonium iodide (PH₄I), which on treatment with KOH gives off pure phosphine:

$$\ce{PH4I + KOH -> KI + H2O + PH3}$$

Phosphine is a colourless gas with a rotten-fish smell and is highly poisonous. It explodes in contact with traces of oxidising agents such as HNO₃, Cl₂ and Br₂ vapours. It is only slightly soluble in water, and its aqueous solution decomposes in the presence of light to give red phosphorus and H₂. When PH₃ is absorbed in copper sulphate or mercuric chloride solution, the corresponding phosphides are obtained:

$$\ce{3CuSO4 + 2PH3 -> Cu3P2 + 3H2SO4}$$

$$\ce{3HgCl2 + 2PH3 -> Hg3P2 + 6HCl}$$

Basic character of phosphine

Phosphine is weakly basic and, like ammonia, forms phosphonium compounds with acids. The lone pair on phosphorus lets it act as a Lewis base:

$$\ce{PH3 + HBr -> PH4Br}$$

$$\ce{PH3 + HI -> PH4I}$$

The formation of PH₄I from PH₃ and HI is the standard proof that PH₃ is basic. Note, however, that PH₃ is a far weaker base than NH₃, and the bond angle in PH₄⁺ is greater than in PH₃, because in PH₄⁺ the lone pair is now used in bonding and the four bond pairs spread to the regular tetrahedral angle.

Phosphorus halides: PCl₃ and PCl₅

Phosphorus forms two series of halides, the trihalides PX₃ (X = F, Cl, Br, I) and the pentahalides PX₅ (X = F, Cl, Br). The two best-studied — and the two NEET asks about — are phosphorus trichloride and phosphorus pentachloride. Both fume in moist air because they hydrolyse, and both are workhorses for converting –OH groups to –Cl.

Phosphorus trichloride, PCl₃

PCl₃ is obtained by passing dry chlorine over heated white phosphorus, or by the action of thionyl chloride on white phosphorus:

$$\ce{P4 + 6Cl2 -> 4PCl3}$$

$$\ce{P4 + 8SOCl2 -> 4PCl3 + 4SO2 + 2S2Cl2}$$

It is a colourless oily liquid that hydrolyses in moisture, which is why it fumes in damp air (the fumes are HCl):

$$\ce{PCl3 + 3H2O -> H3PO3 + 3HCl}$$

It also reacts with organic compounds containing the –OH group, such as acetic acid and ethanol:

$$\ce{3CH3COOH + PCl3 -> 3CH3COCl + H3PO3}$$

$$\ce{3C2H5OH + PCl3 -> 3C2H5Cl + H3PO3}$$

Its shape is pyramidal with phosphorus sp³ hybridised — three P–Cl bonds and one lone pair, the same geometry family as PH₃ and NH₃.

Phosphorus pentachloride, PCl₅

PCl₅ is prepared by the reaction of white phosphorus with an excess of dry chlorine, or by the action of SO₂Cl₂:

$$\ce{P4 + 10Cl2 -> 4PCl5}$$

$$\ce{P4 + 10SO2Cl2 -> 4PCl5 + 10SO2}$$

It is a yellowish-white powder. In moist air it hydrolyses first to phosphoryl chloride (POCl₃) and finally to phosphoric acid:

$$\ce{PCl5 + H2O -> POCl3 + 2HCl}$$

$$\ce{POCl3 + 3H2O -> H3PO4 + 3HCl}$$

On heating it sublimes, but on stronger heating it decomposes back to PCl₃ and chlorine:

$$\ce{PCl5 ->[\Delta] PCl3 + Cl2}$$

It converts –OH-bearing organics to chloro-derivatives, and finely divided metals on heating give the corresponding chlorides:

$$\ce{C2H5OH + PCl5 -> C2H5Cl + POCl3 + HCl}$$

$$\ce{2Ag + PCl5 -> 2AgCl + PCl3}$$

Oxoacids of phosphorus: the master pattern

Phosphorus forms a large family of oxoacids, and their compositions are interrelated by the loss or gain of a water molecule or an oxygen atom. There is one structural rule that ties them all together: in every oxoacid, phosphorus is tetrahedrally surrounded, and each acid contains at least one P=O bond and at least one P–OH bond.

For acids in which phosphorus has a lower oxidation state (less than +5), there is, in addition to P=O and P–OH, either a P–P bond (e.g. H₄P₂O₆) or a P–H bond (e.g. H₃PO₂) — but not both. These +3 acids tend to disproportionate to higher and lower oxidation states; orthophosphorous acid, for instance, disproportionates on heating into phosphoric acid and phosphine:

$$\ce{4H3PO3 -> 3H3PO4 + PH3}$$

*Metaphosphoric acid exists in polymeric forms only; the bonds quoted are for the (HPO₃)₃ ring unit.

Basicity from the number of P–OH groups

This is the single most examined idea in the entire topic. The hydrogen atoms in an oxoacid are not all equivalent. Only the hydrogen attached to oxygen — in the P–OH (P–O–H) form — is ionisable and releases H⁺. The hydrogen attached directly to phosphorus in a P–H bond is not ionisable and plays no part in basicity.

That is why the basicity (number of replaceable hydrogens) equals the number of P–OH groups, not the number of hydrogen atoms in the formula:

All three have the same three hydrogen atoms, yet their basicities are 1, 2 and 3 — entirely set by how many of those hydrogens sit on oxygen rather than on phosphorus. Thus H₃PO₃ is dibasic and H₃PO₄ tribasic, because their structures contain two and three P–OH bonds respectively.

Reducing power comes from the P–H bond

The same P–H bond that is silent for basicity is decisive for redox behaviour. Acids that contain a P–H bond have strong reducing properties. Hypophosphorous acid (H₃PO₂) carries two P–H bonds and is a particularly good reducing agent — it reduces silver nitrate to metallic silver:

$$\ce{4AgNO3 + 2H2O + H3PO2 -> 4Ag + 4HNO3 + H3PO4}$$

H₃PO₃, with one P–H bond, is also a reducing agent, though weaker than H₃PO₂. By contrast, H₃PO₄ has no P–H bond and is not a reducing agent — phosphorus is already at its maximum +5 oxidation state. The pattern is clean: more P–H bonds (and lower oxidation state) means stronger reducing power.

Key oxoacid structures to memorise

Three structures unlock the whole family. In each, phosphorus sits at the centre of a tetrahedron, double-bonded to one oxygen (P=O), and the remaining three positions are filled by P–OH groups and/or P–H bonds.

Read the figure as a single diagnostic: count the teal P–OH arms for basicity, count the coral P–H arms for reducing power. Hypophosphorous acid is the extreme reducing, weakly acidic end; phosphoric acid is the strongly tribasic, non-reducing end.