| IUPAC name
| Other names
|Molar mass||97.99 g/mol|
|Appearance|| Colorless solid|
Clear viscous liquid (>42 °C)
|Density|| 1.885 g/ml (liquid)|
1.685 g/ml (85% solution)
2.030 g/ml (crystal at 25 °C)
|Melting point|| anhydrous|
42.35 °C (108.23 °F; 315.50 K)
hemihydrate</br>29.32 °C (84.78 °F; 302.47 K)
|Boiling point||158–213 °C (316–415 °F; 431–486 K) (decomposes)|
| 392.2 g/100 g (−16.3 °C)|
369.4 g/100 ml (0.5 °C)
446 g/100 ml (14.95 °C)
miscible (>42.3 °C)
|Solubility||Soluble in alcohols|
|Vapor pressure||0.03 mmHg (at 25 °C)|
|Acidity (pKa)|| pKa1 = 2.148|
pKa2 = 7.198
pKa3 = 12.319
Std enthalpy of
|Safety data sheet|| Sigma-Aldrich (crystalline)|
Sigma-Aldrich (85% solution)
|Lethal dose or concentration (LD, LC):|
LD50 (Median dose)
|1,530 mg/kg (rat, oral)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Phosphoric acid is a mid-strength mineral acid having the chemical formula H3PO4 that is more correctly called orthophosphoric acid, to distinguish it from other phosphate-containing acids. Orthophosphoric acid is a relatively non-toxic acid, which, when pure, is a solid at room temperature and pressure. It is, however, extremely deliquescent and hard to isolate as a solid, so a viscous 80-85% solution, known as concentrated phosphoric acid, is usually used.
The oxidation state of phosphorus in ortho- and other phosphoric acids is +5. The acid is triprotic, which means that an orthophosphoric acid molecule can dissociate up to three times, giving up an H+ each time, which typically combines with a water molecule, H2O.
Phosphoric acid may be used to remove rust by direct application to rusted iron, steel tools, or other surfaces. The phosphoric acid changes the reddish-brown iron(III) oxide to ferric phosphate, FePO4:
- 2 H3PO4 + Fe2O3 → 2 FePO4 + 3 H2O
A mixture of concentrated phosphoric acid (sometimes the residue from the combination of a phosphate and sulfuric acid can be used alone) and carbon can be heated to redness(at least 1000°C) in a retort within a furnace to generate elemental phosphorus vapor, which can be condensed underwater to form usable white phosphorus.
Since this acid is particularly non-volatile, it can displace even stronger acids from their salts if heated enough. Such strong acids as hydroiodic acid can be displaced from their salts by phosphoric acid. It is also exceptionally gentle redox-wise, it does not tend to oxidize even very strong reducing agents. However, it is reported to attack glass if heated over 80-90 Celsius, especially at high concentrations, so only those acids can be obtained by using phosphoric acid without damaging glassware that can exit the reaction flask in gaseous form below this temperature.
Orthophosphoric acid is a very polar molecule, and is miscible with water. The more concentrated the solution, the more viscous and syrupy it becomes, causing it to form large bubbles when reacted with gas-evolving substances, such as carbonates. Anhydrous phosphoric acid, a white low melting solid, is obtained by dehydration of 85% phosphoric acid by gentle heating under a vacuum. Heating alone will decompose the phosphoric acid. Evaporation alone will dry phosphoric acid to the solid dihydrate.
Different concentrations of phosphoric acid are available as rust removers, floor cleaners, or etching solutions in hardware stores or body shops. Concentrated food-grade phosphoric acid can rarely be found as a drink additive. In some countries, 80-85% phosphoric acid is sold OTC in small bottles as soldering flux; this acid does not require purification and can be used for experiments as is.
Phosphoric acid is produced industrially by two general routes – the thermal process and the wet process, which includes two sub-methods. The wet process dominates in the commercial sector. The more expensive thermal process produces a purer product that is used for applications in the food industry.
Wet process phosphoric acid is prepared by adding sulfuric acid (or another strong acid) to tricalcium phosphate rock, typically found in nature as apatite, which is also the major component of bone. The reaction is:
- Ca5(PO4)3X + 5 H2SO4 + 10 H2O → 3 H3PO4 + 5 CaSO4·2 H2O + HX
- where X may include OH, F, Cl, and Br
Production of phosphoric acid by this method should preferably be done outside, as deadly hydrogen fluoride and other dangerous byproducts may be produced.
Very pure phosphoric acid is obtained by burning phosphorus to produce phosphorus pentoxide, which is subsequently dissolved in dilute phosphoric acid, or even plain water if you don't have any acid. However this should be done very carefully as the reaction is highly exothermic. Alternatively, water steam can be used.
- Make elemental phosphorus (white)
- Make metal phosphates
- Convert the rusty coating on old equipment to a more stable and appealing form
- Hydroiodic acid synthesis
- It can also be used to make any hydrohalic acid (any volatile acid, in fact)
Like many stronger mineral acids, phosphoric acid can cause chemical burns at high concentrations and irritation or dehydration of the skin at low concentrations. Very highly concentrated samples may contain elemental phosphorus, which has its own set of hazards; this should be taken into account before reacting phosphoric acid with things like metals, which can cause phosphine to form.
Best stored in closed bottles, thick plastic or glass, away from moisture and ammonia.
Phosphoric acid should be neutralized with calcium hydroxide or carbonate before disposal. It does not pose a threat for the environment, so it can be disposed almost everywhere.