Phosphate

A phosphate, an inorganic chemical, is a salt of phosphoric acid. Inorganic phosphates are mined to obtain phosphorus for use in agriculture and industry.^[1]^[2] In organic chemistry, a phosphate, or organophosphate, is an ester of phosphoric acid. Organic phosphates are important in biochemistry and biogeochemistry.

Chemical properties

File:Phosphate.png

The general chemical structure of a phosphate

File:Phosphate Group.PNG

This is the structural formula of the phosphoric acid functional group as found in a weakly acidic aqueous solution. In more basic aqueous solutions, the group donates the two hydrogen atoms and ionizes as a phosphate group with a negative charge of 2. ^[3]

The phosphate ion is a polyatomic ion with the empirical formula P O₄³⁻ 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.

H₃PO₄
H₂PO₄⁻
HPO₄²⁻
PO₄³⁻

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}}_{2}{\mbox{PO}}_{4}^{-}]}{[{\mbox{H}}_{3}{\mbox{PO}}_{4}]}}\simeq 6.92\times 10^{-3}

(pK_a1 2.16)

K_{a2}={\frac {[{\mbox{H}}^{+}][{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 6.17\times 10^{-8}

(pK_a2 7.21)

K_{a3}={\frac {[{\mbox{H}}^{+}][{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 4.79\times 10^{-13}

(pK_a3 12.32)

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\times 10^{10}{\mbox{ , }}{\frac {[{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 6.2\times 10^{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}}_{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}

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}}_{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\times 10^{-7}{\mbox{ , }}{\frac {[{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 2.14\times 10^{-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), 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

Rock phosphate or phosphorite mines are primarily found in:

North America

United States of America, especially Florida, with lesser deposits in North Carolina, Idaho and Tennessee.

Africa

Morocco mainly near Khouribga and Youssoufia and Western Sahara
Senegal
Tunisia

Oceania

References

^ Phosphate Primer, website of the Florida Institute of Phosphate Research
^ "Figuring Out Phosphates," Food Product Design, June 2006, Lynn A. Kuntz
^ Campbell, Neil A. (2005). Biology (Seventh Edition ed.). San Francisco, California: Benjamin Cummings. p. 65. ISBN 0-8053-7171-0. {{cite book}}: |edition= has extra text (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

External links

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

[1] Phosphate Primer, website of the Florida Institute of Phosphate Research

[2] "Figuring Out Phosphates," Food Product Design, June 2006, Lynn A. Kuntz

[3] Campbell, Neil A. (2005). Biology (Seventh Edition ed.). San Francisco, California: Benjamin Cummings. p. 65. ISBN 0-8053-7171-0. {{cite book}}: |edition= has extra text (help); Unknown parameter |coauthors= ignored (|author= suggested) (help)

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