Phosphate

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A phosphate, in inorganic chemistry, is a salt of phosphoric acid. In organic chemistry, a phosphate, or organophosphate, is an ester of phosphoric acid. 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 hydrogenphosphate 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.

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):

$K\_\{a1\}=frac\{[mbox\{H\}^+][mbox\{H\}\_2mbox\{PO\}\_4^-]\}\{[mbox\{H\}\_3mbox\{PO\}\_4]\}simeq\; 7.5times10^\{-3\}$

$K\_\{a2\}=frac\{[mbox\{H\}^+][mbox\{HPO\}\_4^\{2-\}]\}\{[mbox\{H\}\_2mbox\{PO\}\_4^-]\}simeq\; 6.2times10^\{-8\}$

$K\_\{a3\}=frac\{[mbox\{H\}^+][mbox\{PO\}\_4^\{3-\}]\}\{[mbox\{HPO\}\_4^\{2-\}]\}simeq\; 2.14times10^\{-13\}$

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

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

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

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

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

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

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

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

Phosphate can form many polymeric ions, diphosphate (also pyrophosphate), P₂O₇⁴⁻, triphosphate, P₃O₁₀⁵⁻, et cetera. 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. Subsequent uptake of such soil amendments can lead to crops containing uranium concentrations.

See also

References

3. "Figuring Out Phosphates," Food Product Design, June 2006, Lynn A. Kuntz

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.

Phosphate Fertilizer Industry Environmental Hazards

see the link: http://www.fluoridealert.org/phosphate/overview.htm#9 Contributions/71.114.181.145 (71.114.181.145)

External links

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

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