phosphate

phosphate, salt or ester of phosphoric acid, H₃PO₄. Because phosphoric acid is tribasic (having three replaceable hydrogen atoms), it forms monophosphate, diphosphate, and triphosphate salts in which one, two, or three of the hydrogens of the acid are replaced, respectively. Because replaceable hydrogens remain in monophosphates and diphosphates, they are sometimes called acid phosphates. The most important inorganic phosphate is calcium phosphate, Ca₃(PO₄)₂. It makes up the larger part of phosphate rock, a mineral that is abundantly distributed throughout the world. Since calcium phosphate is only slightly soluble in water, it is not very suitable as a source of the phosphorus necessary for plant life; however, by treating it with sulfuric acid the soluble calcium acid phosphate known as superphosphate of lime is formed. Other important inorganic phosphates include ammonium phosphate, important as a fertilizer; trisodium phosphate, used in detergents and for softening water; and disodium phosphate, used to some extent in medicine and in preparing baking powders. Various acid phosphates, e.g., those of calcium, magnesium, and sodium, are sometimes present in carbonated beverages. Microcosmic salt, used in certain bead tests in chemical analysis, is sodium ammonium phosphate. Organic phosphates play an important role in metabolism. For example, in the metabolism of sugars (which have hydroxyl groups, -OH, in their molecules), phosphate esters are often formed as an intermediate compound. Formation of these esters is called phosphorylation. Nucleotides are phosphate esters that play an important role in the conservation and use of the energy released in the metabolism of foods in the body; adenosine triphosphate is an important nucleotide. DNA and RNA (see nucleic acid) are complex polymeric organic phosphates.

or phosphate rock

Rock with a high concentration of phosphates in nodular or compact masses. The phosphates may be derived from a variety of sources, including marine invertebrates that secrete shells of calcium phosphate and the bones and excrement of vertebrates. Typical phosphorite beds contain about 30percnt phosphorus pentoxide (P₂O₅) and constitute the primary source of raw materials for phosphate fertilizers. Significant deposits in the U.S. include the Phosphoria Formation in Idaho and the Monterey Formation in California. Major deposits also occur in the Sechura Desert in Peru.

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Any of numerous chemical compounds related to phosphoric acid (H₃PO₄). Phosphate salts are inorganic compounds containing the phosphate ion (PO₄³⁻), the hydrogen phosphate ion (HPO₄²⁻), or the dihydrogen phosphate ion (H₂PO₄⁻), along with any cation. Phosphate esters are organic compounds in which the hydrogens of phosphoric acid are replaced by organic groups (e.g., methyl, ethyl, phenyl), with one of their carbon atoms bonding to an oxygen atom in the phosphate group. Nucleic acids and ATP both contain phosphate; bones and teeth contain calcium phosphate. Phosphate rock (mainly calcium phosphate) is one of the four most important basic chemical commodities. Phosphates were formerly used in detergents, which washed into rivers and lakes, causing water blooms of algae and bacteria (see eutrophication); such use is now generally outlawed or regulated. Phosphates are still used in fertilizers, baking powder, and toothpaste.

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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):

$K\_\{a1\}=frac\{[mbox\{H\}^+][mbox\{H\}\_2mbox\{PO\}\_4^-]\}\{[mbox\{H\}\_3mbox\{PO\}\_4]\}simeq\; 6.92times10^\{-3\}$

$K\_\{a2\}=frac\{[mbox\{H\}^+][mbox\{HPO\}\_4^\{2-\}]\}\{[mbox\{H\}\_2mbox\{PO\}\_4^-]\}simeq\; 6.17times10^\{-8\}$

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

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 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

• United States of America, especially FloridaAfrica
• Morocco mainly near Khouribga and Youssoufia and Western Sahara
• Senegal
• TunisiaOceania
• 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