4_5 Oxides and Oxyacids of Phosphorus: P₄O₆, P₄O₁₀, H₃PO₃ and H₃PO₄

When phosphorus burns, the product depends on how much oxygen is available. Each of its two oxides dissolves in water to give a different acid. The oxides have elegant cage structures built on the

\mathrm{P_4}

tetrahedron from the previous lesson, and the acids show how molecular structure, not just formula, controls acidity.

Phosphorus(III) Oxide, P₄O₆

When phosphorus burns with too little oxygen to oxidize it fully, the product is phosphorus(III) oxide. Its simplest formula is

\mathrm{P_2O_3}

, but the real molecule is

\mathrm{P_4O_6}

. Picture the

\mathrm{P_4}

tetrahedron with an oxygen atom inserted into each of its six P–P bonds. The result is a cage in which every phosphorus atom bonds to three bridging oxygens. It is a white, waxy, poisonous solid with a garlic smell that melts at about 24 °C. Heated in air it oxidizes to

\mathrm{P_4O_{10}}

, and with cold water it gives phosphorous acid:

O1P2OP3OP1OP(O2)O3.O.O.O.O.O.O>>[H]P(=O)(O)O.[H]P(=O)(O)O.[H]P(=O)(O)O.[H]P(=O)(O)O

With hot water,

\mathrm{P_4O_6}

instead disproportionates, giving phosphine and phosphoric acid.

Phosphorus(V) Oxide, P₄O₁₀

Burning phosphorus in excess dry air gives phosphorus(V) oxide. Its simplest formula is

\mathrm{P_2O_5}

, and it is often called phosphorus pentoxide, but the molecule is

\mathrm{P_4O_{10}}

. It has the same cage as

\mathrm{P_4O_6}

, plus one extra oxygen double-bonded to the outside of each phosphorus. It is a white, fluffy powder that sublimes at about 360 °C.

Its outstanding property is a fierce attraction for water. It reacts violently with water, first forming metaphosphoric acid and then, with more water and heat, orthophosphoric acid:

O1P2(=O)OP3(=O)OP1(=O)OP(=O)(O2)O3.O.O.O.O.O.O>>OP(O)(O)=O.OP(O)(O)=O.OP(O)(O)=O.OP(O)(O)=O

This makes

\mathrm{P_4O_{10}}

one of the most powerful dehydrating agents known. It dries gases and removes the elements of water even from other acids. It turns nitric acid into

\mathrm{N_2O_5}

, sulphuric acid into

\mathrm{SO_3}

, ethanoic acid into ethanoic anhydride, and ethanol into ethene.

Phosphorous Acid, H₃PO₃

Phosphorous acid, also called phosphonic acid, is made by dissolving

\mathrm{P_4O_6}

in cold water or by hydrolysing

\mathrm{PCl_3}

. It is a white crystalline solid that melts at about 74 °C. Its formula suggests three acidic hydrogens, but only two are acidic. Its structure is

\mathrm{HP(O)(OH)_2}

: one hydrogen is bonded directly to phosphorus, and a P–H bond does not ionize. Only hydrogens on oxygen are released as

\mathrm{H^+}

, so the acid is dibasic and forms only two series of salts.

That P–H bond also makes phosphorous acid a strong reducing agent. It reduces silver and copper(II) salts to the metals and is itself oxidized to phosphoric acid. When heated, it disproportionates into phosphoric acid and phosphine:

\mathrm{4H_3PO_3 \rightarrow 3H_3PO_4 + PH_3}

.

Orthophosphoric Acid, H₃PO₄

Orthophosphoric acid, usually just called phosphoric acid, has a tetrahedral phosphorus bonded to one oxygen and three OH groups. It is made industrially in two ways. The "wet process" treats phosphate rock with sulphuric acid. The "thermal process" burns phosphorus to

\mathrm{P_4O_{10}}

and dissolves it in water, which gives the pure acid for food use. In the laboratory it can be made by oxidizing red phosphorus with concentrated nitric acid.

The pure acid is a colourless, deliquescent solid that melts at 42 °C, and it is usually sold as a syrupy 85% solution. Its three OH groups make it a weak tribasic acid. It neutralizes in three steps, forming three series of salts:

\mathrm{H_3PO_4 \xrightarrow{\ NaOH\ } NaH_2PO_4 \xrightarrow{\ NaOH\ } Na_2HPO_4 \xrightarrow{\ NaOH\ } Na_3PO_4}

Strong heating removes water between molecules. At about 250 °C two molecules join to form pyrophosphoric acid,

\mathrm{H_4P_2O_7}

, and above about 320 °C metaphosphoric acid,

\mathrm{HPO_3}

, forms. Phosphoric acid is used to make phosphate fertilizers, to flavour cola drinks, in rust removers and in detergents and food additives.

Tips for remembering

• Count the OH groups, not the H atoms: H₃PO₄ has three P–OH groups and is tribasic, while H₃PO₃ has one H bonded directly to P and is only dibasic.
• P₄O₁₀ is thirsty: it pulls water out of almost anything, including H₂SO₄ (giving SO₃) and HNO₃ (giving N₂O₅).
• Limited air gives the III oxide, excess air gives the V oxide: P₄O₆ burns on to P₄O₁₀ when more oxygen is available.