Phosphate

A phosThis page sucks nutphate, 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.

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.

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.

• 
  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):

$K_{a1}={\frac {[{\mbox{H}}^{+}][{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}{[{\mbox{H}}_{3}{\mbox{PO}}_{4}]}}\simeq 7.5\times 10^{-3}$ (pK_a1 2.12)

$K_{a2}={\frac {[{\mbox{H}}^{+}][{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 6.2\times 10^{-8}$ (pK_a2 7.21)

$K_{a3}={\frac {[{\mbox{H}}^{+}][{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 2.14\times 10^{-13}$ (pK_a3 12.67)

The speciation diagram obtained using these pK values shows three distinct regions. In effect H₃PO₄, H₂PO₄^- and HPO₄^2- 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

${\frac {[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}{[{\mbox{H}}_{3}{\mbox{PO}}_{4}]}}\simeq 7.5\times 10^{4}{\mbox{ , }}{\frac {[{\mbox{HPO}}_{4}^{2-}]}{[{\mbox{H}}_{2}{\mbox{PO}}_{4}^{-}]}}\simeq 0.62{\mbox{ , }}{\frac {[{\mbox{PO}}_{4}^{3-}]}{[{\mbox{HPO}}_{4}^{2-}]}}\simeq 2.14\times 10^{-6}$

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

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

See also phosphorylation for more information on the use of phosphate by cells.

The addition and removal of phosphate from proteins in all cells is a pivotal strategy in the regulation of metabolic processes.

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 North Carolina, with lesser deposits in Florida, Idaho and Tennessee.

Africa

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

Middle East

• Israel
• Jordan

Oceania

• Australia
• Makatea
• Nauru
• Ocean Island

See also

• Organophosphorus compounds
• Phosphate conversion coating
• Phosphine - PR₃
• Phosphine oxide - OPR₃
• Phosphinite - P(OR)R₂
• Phosphonite - P(OR)₂R
• Phosphite - P(OR)₃
• Phosphinate - OP(OR)R₂
• Phosphonate - OP(OR)₂R
• Phosphate - OP(OR)₃, such as triphenyl phosphate
• Polyphosphate - P_n

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.

External links

• Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery
• Website on recycling of phosphates from sewage water