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 (PO₄^3?) predominates, whereas in weakly-basic conditions, the hydrogen phosphate ion (HPO₄^2?) 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.

<gallery> Image:3-phosphoric-acid-3D-balls.png|

More precisely, considering the following three equilibrium reactions:

H₃PO₄ H^{+ + H₂PO₄?}

H₂PO₄^? H^{+ + HPO₄2?}

HPO₄^2? H^{+ + PO₄3?}

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

K_{a1}=\frac{[\mbox{H}^+][\mbox{H}_2\mbox{PO}_4^-]}{[\mbox{H}_3\mbox{PO}_4]}\simeq 6.92\times10^{-3}

K_{a2}=\frac{[\mbox{H}^+][\mbox{HPO}_4^{2-}]}{[\mbox{H}_2\mbox{PO}_4^-]}\simeq 6.17\times10^{-8}

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

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

\frac{[\mbox{H}_2\mbox{PO}_4^-]}{[\mbox{H}_3\mbox{PO}_4]}\simeq 7.5\times10^{10} \mbox{ , }\frac{[\mbox{HPO}_4^{2-}]}{[\mbox{H}_2\mbox{PO}_4^-]}\simeq 6.2\times10^5 \mbox{ , } \frac{[\mbox{PO}_4^{3-}]}{[\mbox{HPO}_4^{2-}]}\simeq 2.14

showing that only PO₄^3? and HPO₄^2? are in significant amounts.

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

\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 H₂PO₄^? and HPO₄^2? ions are in significant amounts (62% H₂PO₄^?, 38% HPO₄^2?). Note that in the extracellular fluid (pH=7.4), this proportion is inverted (61% HPO₄^2?, 39% H₂PO₄^?).

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

\frac{[\mbox{H}_2\mbox{PO}_4^-]}{[\mbox{H}_3\mbox{PO}_4]}\simeq 0.075 \mbox{ , }\frac{[\mbox{HPO}_4^{2-}]}{[\mbox{H}_2\mbox{PO}_4^-]}\simeq 6.2\times10^{-7} \mbox{ , } \frac{[\mbox{PO}_4^{3-}]}{[\mbox{HPO}_4^{2-}]}\simeq 2.14\times10^{-12}

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), P₂O₇^4?, and triphosphate, P₃O₁₀^5?. 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

• United States of America, especially Florida

Africa

• Morocco mainly near Khouribga and Youssoufia and Western Sahara
• Senegal
• Tunisia

Oceania

• Australia
• Makatea
• Nauru
• Ocean Island

See also

Further reading

• Schmittner Karl-Erich and Giresse Pierre, 1999. Micro-environmental controls on biomineralization: superficial processes of apatite and calcite precipitation in Quaternary soils, Roussillon, France. Sedimentology 46/3: 463-476.
• http://www.fluoridealert.org/phosphate/overview.htm#9, discusses environmental hazards of the phosphate fertilizer industry

References

External links

• Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery

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