The phosphate ion is a polyatomic ion with the empirical formula PO3−
4
and a molar mass of 94.97 g/mol. It consists of one central phosphorus atom surrounded by four oxygen atoms in a tetrahedral arrangement. The phosphate ion carries a −3 formal charge and is the conjugate base of the hydrogen phosphate ion, HPO2−
4
, which is the conjugate base of H
2PO−
4
, the dihydrogen phosphate ion, which in turn is the conjugate base of H
3PO
4
, phosphoric acid. A phosphate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound. Many phosphates are not soluble in water at standard temperature and pressure. The sodium, potassium, rubidium, caesium, and ammonium phosphates are all water-soluble. Most other phosphates are only slightly soluble or are insoluble in water. As a rule, the hydrogen and dihydrogen phosphates are slightly more soluble than the corresponding phosphates. The pyrophosphates are mostly water-soluble.

Aqueous phosphate exists in four forms. In strongly basic conditions, the phosphate ion (PO3−
4
) predominates, whereas in weakly basic conditions, the hydrogen phosphate ion (HPO2−
4
) is prevalent. In weakly acidic conditions, the dihydrogen phosphate ion (H
2PO−
4
) is most common. In strongly acidic conditions, trihydrogen phosphate (H
3PO
4
) is the main form.

More precisely, considering these three equilibrium reactions:

H
3PO
4
⇌ H+ + H
2PO−
4
H
2PO−
4
⇌ H+ + HPO2−
4
HPO2−
4
⇌ H+ + PO3−
4

the corresponding constants at 25 °C (in mol/L) are (see phosphoric acid):

K a 1 = [ H + ] [ H 2 PO 4 − ] [ H 3 PO 4 ] ≃ 7.5 × 10 − 3 {\displaystyle K_{\mathrm {a1} }={\frac {[{\mbox{H}}^{+}][{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}{[{\mbox{H}}_{3}{\mbox{PO}}_{4}]}}\simeq 7.5\times 10^{-3}} {\displaystyle K_{\mathrm {a1} }={\frac {[{\mbox{H}}^{+}][{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}{[{\mbox{H}}_{3}{\mbox{PO}}_{4}]}}\simeq 7.5\times 10^{-3}}      (pKa1 ≈ 2.12)
K a 2 = [ H + ] [ HPO 4 2 − ] [ H 2 PO 4 − ] ≃ 6.2 × 10 − 8 {\displaystyle K_{\mathrm {a2} }={\frac {[{\mbox{H}}^{+}][{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 6.2\times 10^{-8}} {\displaystyle K_{\mathrm {a2} }={\frac {[{\mbox{H}}^{+}][{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 6.2\times 10^{-8}}      (pKa2 ≈ 7.21)
K a 3 = [ H + ] [ PO 4 3 − ] [ HPO 4 2 − ] ≃ 2.14 × 10 − 13 {\displaystyle K_{\mathrm {a3} }={\frac {[{\mbox{H}}^{+}][{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 2.14\times 10^{-13}} {\displaystyle K_{\mathrm {a3} }={\frac {[{\mbox{H}}^{+}][{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 2.14\times 10^{-13}}      (pKa3 ≈ 12.67)
Phosphoric acid speciation.png

The speciation diagram obtained using these pK values shows three distinct regions. In effect, H
3PO
4
, H
2PO−
4
and HPO2−
4
behave as separate weak acids. This is because the successive pK values differ by more than 4. For each acid, the pH at half-neutralization is equal to the pK value of the acid. The region in which the acid is in equilibrium with its conjugate base is defined by pH ≈ pK ± 2. Thus, the three pH regions are approximately 0–4, 5–9 and 10–14. This is idealized, as it assumes constant ionic strength, which will not hold in reality at very low and very high pH values.

For a neutral pH as in the cytosol, pH = 7.0

[ H 2 PO 4 − ] [ H 3 PO 4 ] ≃ 7.5 × 10 4  ,  [ HPO 4 2 − ] [ H 2 PO 4 − ] ≃ 0.62  ,  [ PO 4 3 − ] [ HPO 4 2 − ] ≃ 2.14 × 10 − 6 {\displaystyle {\frac {[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}{[{\mbox{H}}_{3}{\mbox{PO}}_{4}]}}\simeq 7.5\times 10^{4}{\mbox{ , }}{\frac {[{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 0.62{\mbox{ , }}{\frac {[{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 2.14\times 10^{-6}}  \frac{[\mbox{H}_2\mbox{PO}_4^-]}{[\mbox{H}_3\mbox{PO}_4]}\simeq 7.5\times10^4 \mbox{ , }\frac{[\mbox{HPO}_4^{2-}]}{[\mbox{H}_2\mbox{PO}_4^-]}\simeq 0.62 \mbox{ , } \frac{[\mbox{PO}_4^{3-}]}{[\mbox{HPO}_4^{2-}]}\simeq 2.14\times10^{-6}

so that only H
2PO−
4
and HPO2−
4
ions are present in significant amounts (62% H
2PO−
4
, 38% HPO2−
4
. Note that in the extracellular fluid (pH = 7.4), this proportion is inverted (61% HPO2−
4
, 39% H
2PO−
4
).

Phosphate can form many polymeric ions such as pyrophosphate), P
2O4−
7
, and triphosphate, P
3O5−
10
. The various metaphosphate ions (which are usually long linear polymers) have an empirical formula of PO−
3
and are found in many compounds.

Biochemistry of phosphates

In biological systems, phosphorus is found as a free phosphate ion in solution and is called inorganic phosphate, to distinguish it from phosphates bound in various phosphate esters. Inorganic phosphate is generally denoted Pi and at physiological (homeostatic) pH primarily consists of a mixture of HPO2−
4
and H
2PO−
4
ions.

Inorganic phosphate can be created by the hydrolysis of pyrophosphate, which is denoted PPi:

P
2O4−
7
+ H2O ⇌ 2 HPO2−
4

However, phosphates are most commonly found in the form of adenosine phosphates (AMP, ADP, and ATP) and in DNA and RNA, and can be released by the hydrolysis of ATP or ADP. Similar reactions exist for the other nucleoside diphosphates and triphosphates. Phosphoanhydride bonds in ADP and ATP, or other nucleoside diphosphates and triphosphates, contain high amounts of energy which give them their vital role in all living organisms. They are generally referred to as high-energy phosphate, as are the phosphagens in muscle tissue. Compounds such as substituted phosphines have uses in organic chemistry, but do not seem to have any natural counterparts.

The addition and removal of phosphate from proteins in all cells is a pivotal strategy in the regulation of metabolic processes. Phosphorylation and dephosphorylation are important ways that energy is stored and released in living systems. Cells use ATP in this manner.

Reference ranges for blood tests, showing ‘inorganic phosphorus’ in purple at right, being almost identical to the molar concentration of phosphate

Phosphate is useful in animal cells as a buffering agent. Phosphate salts that are commonly used for preparing buffer solutions at cell pHs include Na2HPO4, NaH2PO4, and the corresponding potassium salts.

An important occurrence of phosphates in biological systems is as the structural material of bone and teeth. These structures are made of crystalline calcium phosphate in the form of hydroxyapatite. The hard dense enamel of mammalian teeth consists of fluoroapatite, a hydroxy calcium phosphate where some of the hydroxyl groups have been replaced by fluoride ions.

Plants take up phosphorus through several pathways: the arbuscular mycorrhizal pathway and the direct uptake pathway.