Definitions

# Phosphate

[fos-feyt]
A phosphate, an inorganic chemical, is a salt of phosphoric acid. Inorganic phosphates are mined to obtain phosphorus for use in agriculture and industry. In organic chemistry, a phosphate, or organophosphate, is an ester of phosphoric acid. Organic phosphates are important in biochemistry and biogeochemistry.

## Chemical properties

The phosphate ion is a polyatomic ion with the empirical formula PO43− and a molar mass of 94.973 g/mol; it consists of one central phosphorus atom surrounded by four identical oxygen atoms in a tetrahedral arrangement. The phosphate ion carries a negative three formal charge and is the conjugate base of the hydrogen phosphate ion, HPO42−, which is the conjugate base of H2PO4, the dihydrogen phosphate ion, which in turn is the conjugate base of H3PO4, phosphoric acid. It is a hypervalent molecule (the phosphorus atom has 10 electrons in its valence shell). Phosphate is also an organophosphorus compound with the formula OP(OR)3

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 hydrogenphosphates and the dihydrogenphosphates are slightly more soluble than the corresponding phosphates. The pyrophosphates are mostly water soluble.

In dilute aqueous solution, phosphate exists in four forms. In strongly-basic conditions, the phosphate ion (PO43−) predominates, whereas in weakly-basic conditions, the hydrogen phosphate ion (HPO42−) is prevalent. In weakly-acid conditions, the dihydrogen phosphate ion (H2PO4) is most common. In strongly-acid conditions, aqueous phosphoric acid (H3PO4) is the main form.

More precisely, considering the following three equilibrium reactions:

H3PO4 ⇌ H+ + H2PO4

H2PO4 ⇌ H+ + HPO42−

HPO42− ⇌ H+ + PO43−

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

$K_\left\{a1\right\}=frac\left\{\left[mbox\left\{H\right\}^+\right]\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}\left\{\left[mbox\left\{H\right\}_3mbox\left\{PO\right\}_4\right]\right\}simeq 6.92times10^\left\{-3\right\}$

$K_\left\{a2\right\}=frac\left\{\left[mbox\left\{H\right\}^+\right]\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}simeq 6.17times10^\left\{-8\right\}$

$K_\left\{a3\right\}=frac\left\{\left[mbox\left\{H\right\}^+\right]\left[mbox\left\{PO\right\}_4^\left\{3-\right\}\right]\right\}\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}simeq 4.79times10^\left\{-13\right\}$

For a strongly-basic pH (pH=13), we find

$frac\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}\left\{\left[mbox\left\{H\right\}_3mbox\left\{PO\right\}_4\right]\right\}simeq 7.5times10^\left\{10\right\} mbox\left\{ , \right\}frac\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}simeq 6.2times10^5 mbox\left\{ , \right\} frac\left\{\left[mbox\left\{PO\right\}_4^\left\{3-\right\}\right]\right\}\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}simeq 2.14$

showing that only PO43− and HPO42− are in significant amounts.

For a neutral pH (for example the cytosol pH=7.0), we find

$frac\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}\left\{\left[mbox\left\{H\right\}_3mbox\left\{PO\right\}_4\right]\right\}simeq 7.5times10^4 mbox\left\{ , \right\}frac\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}simeq 0.62 mbox\left\{ , \right\} frac\left\{\left[mbox\left\{PO\right\}_4^\left\{3-\right\}\right]\right\}\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}simeq 2.14times10^\left\{-6\right\}$

so that only H2PO4 and HPO42− ions are in significant amounts (62% H2PO4, 38% HPO42−). Note that in the extracellular fluid (pH=7.4), this proportion is inverted (61% HPO42−, 39% H2PO4).

For a strongly-acid pH (pH=1), we find

$frac\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}\left\{\left[mbox\left\{H\right\}_3mbox\left\{PO\right\}_4\right]\right\}simeq 0.075 mbox\left\{ , \right\}frac\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}\left\{\left[mbox\left\{H\right\}_2mbox\left\{PO\right\}_4^-\right]\right\}simeq 6.2times10^\left\{-7\right\} mbox\left\{ , \right\} frac\left\{\left[mbox\left\{PO\right\}_4^\left\{3-\right\}\right]\right\}\left\{\left[mbox\left\{HPO\right\}_4^\left\{2-\right\}\right]\right\}simeq 2.14times10^\left\{-12\right\}$

showing that H3PO4 is dominant with respect to H2PO4. HPO42− and PO43− are practically absent.

Phosphate can form many polymeric ions such as diphosphate (also pyrophosphate), P2O74−, and triphosphate, P3O105−. The various metaphosphate ions have an empirical formula of PO3 and are found in many compounds.

Phosphate deposits can contain significant amounts of naturally occurring uranium. Uptake of these substances by plants can lead to high uranium concentrations in crops.

## Cellular function

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.

## Mining

Phosphate mines are primarily found in:North America