Phosphoric acid

Phosphoric acid
IUPAC name
trihydroxidooxidophosphorus
phosphoric acid
Other names
Orthophosphoric acid
Identifiers
CAS number 7664-38-2 YesY, 
16271-20-8 (hemihydrate)
ChemSpider 979
EC number 231-633-2
UN number 1805
RTECS number TB6300000
Properties
Molecular formula H3PO4
Molar mass 98.00 g/mol
Appearance white solid or colourless, 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

42.35 °C (anhydrous)
29.32 °C (hemihydrate)
Boiling point

158 °C (decomp)
Solubility in water 548 g/100 mL
Acidity (pKa) 2.148, 7.198, 12.375
Viscosity 2.4–9.4 cP (85% aq. soln.)
147 cP (100 %)
Hazards
MSDS ICSC 1008
EU Index 015-011-00-6
EU classification Corrosive (C)
R-phrases R34
S-phrases (S1/2) S26 S45
NFPA 704
NFPA 704.svg
0
2
0
COR
Flash point Non-flammable
Related compounds
Related phosphorus oxoacids Hypophosphorous acid
Phosphorous acid
Pyrophosphoric acid
Triphosphoric acid
Perphosphoric acid
Permonophosphoric acid
 YesY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Phosphoric acid, also known as orthophosphoric acid or phosphoric(V) acid, is a mineral (inorganic) acid having the chemical formula H3PO4. Orthophosphoric acid molecules can combine with themselves to form a variety of compounds which are also referred to as phosphoric acids, but in a more general way. The term phosphoric acid can also refer to a chemical or reagent consisting of phosphoric acids, usually orthophosphoric acid.

Contents

Orthophosphoric acid chemistry

Pure anhydrous phosphoric acid is a white solid that melts at 42.35 °C to form a colorless, viscous liquid.

Most people and even chemists refer to orthophosphoric acid as phosphoric acid, which is the IUPAC name for this compound. The prefix ortho is used to distinguish the acid from other phosphoric acids, called polyphosphoric acids two(ii) . Orthophosphoric acid is a non-toxic, inorganic, rather weak triprotic acid, which, when pure, is a solid at room temperature and pressure. The chemical structure of orthophosphoric acid is shown above in the data table. Orthophosphoric acid is a very polar molecule; therefore it is highly soluble in water. The oxidation state of phosphorus (P) in ortho- and other phosphoric acids is +5; the oxidation state of all the oxygen atoms (O) is −2 and all the hydrogen atoms (H) is +1. Triprotic 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, as shown in these reactions:

H3PO4(s)   + H2O(l) is in equilibrium with H3O+(aq) + H2PO4−(aq)       Ka1= 7.25×10−3
H2PO4−(aq)+ H2O(l) is in equilibrium with H3O+(aq) + HPO42−(aq)       Ka2= 6.31×10−8
HPO42−(aq)+ H2O(l) is in equilibrium with H3O+(aq) +  PO43−(aq)        Ka3= 3.98×10−13

The anion after the first dissociation, H2PO4−, is the dihydrogen phosphate anion. The anion after the second dissociation, HPO42−, is the hydrogen phosphate anion. The anion after the third dissociation, PO43−, is the phosphate or orthophosphate anion. For each of the dissociation reactions shown above, there is a separate acid dissociation constant, called Ka1, Ka2, and Ka3 given at 25 °C. Associated with these three dissociation constants are corresponding pKa1=2.12 , pKa2=7.21 , and pKa3=12.67 values at 25 °C. Even though all three hydrogen (H ) atoms are equivalent on an orthophosphoric acid molecule, the successive Ka values differ since it is energetically less favorable to lose another H+ if one (or more) has already been lost and the molecule/ion is more negatively-charged.

Because the triprotic dissociation of orthophosphoric acid, the fact that its conjugate bases (the phosphates mentioned above) cover a wide pH range, and, because phosphoric acid/phosphate solutions are, in general, non-toxic, mixtures of these types of phosphates are often used as buffering agents or to make buffer solutions, where the desired pH depends on the proportions of the phosphates in the mixtures. Similarly, the non-toxic, anion salts of triprotic organic citric acid are also often used to make buffers. Phosphates are found pervasively in biology, especially in the compounds derived from phosphorylated sugars, such as DNA, RNA, and adenosine triphosphate (ATP). There is a separate article on phosphate as an anion or its salts.

Upon heating orthophosphoric acid, condensation of the phosphoric units can be induced by driving off the water formed from condensation. When one molecule of water has been removed for each two molecules of phosphoric acid, the result is pyrophosphoric acid (H4P2O7). When an average of one molecule of water per phosphoric unit has been driven off, the resulting substance is a glassy solid having an empirical formula of HPO3 and is called metaphosphoric acid.[1] Metaphosphoric acid is a singly anhydrous version of orthophosphoic acid and is sometimes used as a water- or moisture-absorbing reagent. Further dehydrating is very difficult, and can be accomplished only by means of an extremely strong desiccant (and not by heating alone). It produces phosphoric anhydride, which has an empirical formula P2O5, although an actual molecule has a chemical formula of P4O10. Phosphoric anhydride is a solid, which is very strongly moisture-absorbing and is used as a desiccant.

pH and composition of a phosphoric acid aqueous solution

For a given total acid concentration [A] = [H3PO4] + [H2PO4−] + [HPO42−] + [PO43−] ([A] is the total number of moles of pure H3PO4 which have been used to prepare 1 liter of solution), the composition of an aqueous solution of phosphoric acid can be calculated using the equilibrium equations associated with the three reactions described above together with the [H+][OH−] = 10−14 relation and the electrical neutrality equation. Possible concentrations of polyphosphoric molecules and ions is neglected. The system may be reduced to a fifth degree equation for [H+] which can be solved numerically, yielding:

[A] (mol/L) pH [H3PO4]/[A] (%) [H2PO4−]/[A] (%) [HPO42−]/[A] (%) [PO43−]/[A] (%)
1 1.08 91.7 8.29 6.20×10−6 1.60×10−17
10−1 1.62 76.1 23.9 6.20×10−5 5.55×10−16
10−2 2.25 43.1 56.9 6.20×10−4 2.33×10−14
10−3 3.05 10.6 89.3 6.20×10−3 1.48×10−12
10−4 4.01 1.30 98.6 6.19×10−2 1.34×10−10
10−5 5.00 0.133 99.3 0.612 1.30×10−8
10−6 5.97 1.34×10−2 94.5 5.50 1.11×10−6
10−7 6.74 1.80×10−3 74.5 25.5 3.02×10−5
10−10 7.00 8.24×10−4 61.7 38.3 8.18×10−5

For large acid concentrations, the solution is mainly composed of H3PO4. For [A] = 10−2, the pH is closed to pKa1, giving an equimolar mixture of H3PO4 and H2PO4−. For [A] below 10−3, the solution is mainly composed of H2PO4− with [HPO42−] becoming non negligible for very dilute solutions. [PO43−] is always negligible.

Chemical reagent

Pure 75–85% aqueous solutions (the most common) are clear, colourless, odourless, non-volatile, rather viscous, syrupy liquids, but still pourable. Phosphoric acid is very commonly used as an aqueous solution of 85% phosphoric acid or H3PO4. Because it is a concentrated acid, an 85% solution can be corrosive, although nontoxic when diluted. Because of the high percentage of phosphoric acid in this reagent, at least some of the orthophosphoric acid is condensed into polyphosphoric acids in a temperature-dependent equilibrium, but, for the sake of labeling and simplicity, the 85% represents H3PO4 as if it were all orthophosphoric acid. Other percentages are possible too, even above 100%, where the phosphoric acids and water would be in an unspecified equilibrium, but the overall elemental mole content would be considered specified. When aqueous solutions of phosphoric acid and/or phosphate are dilute, they are in or will reach an equilibrium after a while where practically all the phosphoric/phosphate units are in the ortho- form.

Preparation of hydrogen halides

Phosphoric acid reacts with halides to form the corresponding hydrogen halide gas (steamy fumes are observed on warming the reaction mixture). This is a common practice for the laboratory preparation of hydrogen halides.

NaCl(s) + H3PO4(l) → NaH2PO4(s) + HCl(g)
NaBr(s) + H3PO4(l) → NaH2PO4(s) + HBr(g)
NaI(s) + H3PO4(l) → NaH2PO4(s) + HI(g)

Rust removal

Phosphoric acid may be used as a "rust converter", by direct application to rusted iron, steel tools, or surfaces. The phosphoric acid converts reddish-brown iron(III) oxide (rust) to black ferric phosphate, FePO4.

"Rust converter" is sometimes a greenish liquid suitable for dipping (in the same sort of acid bath as is used for pickling metal), but it is more often formulated as a gel, commonly called naval jelly. It is sometimes sold under other names, such as "rust remover" or "rust killer". As a thick gel, it may be applied to sloping, vertical, or even overhead surfaces.

After treatment, the black ferric-phosphate coating can be scrubbed off, leaving a fresh metal surface. Multiple applications of phosphoric acid may be required to remove all rust. The black phosphate coating can also be left in place, where it will provide moderate further corrosion resistance. (Such protection is also provided by the superficially similar Parkerizing and blued electrochemical conversion coating processes.)

Processed food use

Food-grade phosphoric acid (additive E338) is used to acidify foods and beverages such as various colas, but not without controversy regarding its health effects. It provides a tangy or sour taste and, being a mass-produced chemical, is available cheaply and in large quantities. The low cost and bulk availability is unlike more expensive seasonings that give comparable flavors, such as citric acid which is obtainable from lemons and limes. (However most citric acid in the food industry is not extracted from citrus fruit, but fermented by Aspergillus niger mold from scrap molasses, waste starch hydrolysates and phosphoric acid.)

Biological effects on bone calcium and kidney health

Phosphoric acid, used in many soft drinks (primarily cola), has been linked to lower bone density in epidemiological studies. For example, a study[2] using dual-energy X-ray absorptiometry rather than a questionnaire about breakage, provides reasonable evidence to support the theory that drinking cola results in lower bone density. This study was published in the American Journal of Clinical Nutrition. A total of 1672 women and 1148 men were studied between 1996 and 2001. Dietary information was collected using a food frequency questionnaire that had specific questions about the number of servings of cola and other carbonated beverages and that also made a differentiation between regular, caffeine-free, and diet drinks. The paper cites significant statistical evidence to show that women who consume cola daily have lower bone density. Total phosphorus intake was not significantly higher in daily cola consumers than in nonconsumers; however, the calcium-to-phosphorus ratios were lower. The study also suggests that further research is needed to confirm the findings.

On the other hand, a study funded by Pepsi suggests that insufficient intake of phosphorus leads to lower bone density. The study does not examine the effect of phosphoric acid, which binds with magnesium and calcium in the digestive tract to form salts that are not absorbed, but rather studies general phosphorus intake.[3]

However, a well-controlled clinical study by Heaney and Rafferty using calcium-balance methods found no impact of carbonated soft drinks containing phosphoric acid on calcium excretion.[4] The study compared the impact of water, milk, and various soft drinks (two with caffeine and two without; two with phosphoric acid and two with citric acid) on the calcium balance of 20- to 40-year-old women who customarily consumed ~3 or more cups (680 ml) of a carbonated soft drink per day. They found that, relative to water, only milk and the two caffeine-containing soft drinks increased urinary calcium, and that the calcium loss associated with the caffeinated soft drink consumption was about equal to that previously found for caffeine alone. Phosphoric acid without caffeine had no impact on urine calcium, nor did it augment the urinary calcium loss related to caffeine. Because studies have shown that the effect of caffeine is compensated for by reduced calcium losses later in the day,[5] Heaney and Rafferty concluded that the net effect of carbonated beverages—including those with caffeine and phosphoric acid—is negligible, and that the skeletal effects of carbonated soft drink consumption are likely due primarily to milk displacement.

Other chemicals such as caffeine (also a significant component of popular common cola drinks) were also suspected as possible contributors to low bone density, due to the known effect of caffeine on calciuria. One other study, involving 30 women over the course of a week, suggests that phosphoric acid in colas has no such effect, and postulates that caffeine has only a temporary effect, which is later reversed. The authors of this study conclude that the skeletal effects of carbonated beverage consumption are likely due primarily to milk displacement.[4] (Another possible confounding factor may be an association between high soft drink consumption and sedentary lifestyle.)

Cola consumption has also been associated with chronic kidney disease and kidney stones through medical research.[6] The preliminary results suggest that cola consumption may increase the risk of chronic kidney disease.

Medical use

Phosphoric acid is used in dentistry and orthodontics as an etching solution, to clean and roughen the surfaces of teeth where dental appliances or fillings will be placed. Phosphoric acid is also an ingredient in over-the-counter anti-nausea medications that also contain high levels of sugar (glucose and fructose). This acid is also used in many teeth whiteners to eliminate plaque that may be on the teeth before application.

Preparation

Phosphoric acid can be prepared by three routes - the Thermal Process, the Wet Process and the dry Kiln Process.

Thermal phosphoric acid

This very pure phosphoric acid is obtained by burning elemental phosphorus to produce phosphorus pentoxide and dissolving the product in dilute phosphoric acid. This produces a very pure phosphoric acid, since most impurities present in the rock have been removed when extracting phosphorus from the rock in a furnace. The end result is food-grade, thermal phosphoric acid; however, for critical applications, additional processing to remove arsenic compounds may be needed.

Wet phosphoric acid

Wet process phosphoric acid is prepared by adding sulfuric acid to tricalcium phosphate rock, typically found in nature as apatite.

The reaction is:

Ca5(PO4)3X + 5 H2SO4 + 2 H2O → 3 H3PO4 + 5 CaSO4 · 2 H2O + HX
where X may include OH, F, Cl, and Br

The initial phosphoric acid solution may contain 23-33% P2O5, but can be concentrated by the evaporation of water to produce commercial- or merchant-grade phosphoric acid, which contains about 54% P2O5. Further evaporation of water yields superphosphoric acid with a P2O5 concentration above 70%.[7][8]

Digestion of the phosphate ore using sulfuric acid yields the insoluble calcium sulfate (gypsum), which is filtered and removed as phosphogypsum. Wet-process acid can be furthered purified by removing fluorine to produce animal-grade phosphoric acid, or by solvent extraction and arsenic removal to produce food-grade phosphoric acid.

Kiln Phosphoric Acid

Kiln phosphoric acid (KPA) process technology is the most recent technology. Called the “Improved Hard Process”,[9] this technology will both make low grade phosphate rock reserves commercially viable and will increase the P2O5 recovery from existing phosphate reserves. This may significantly extend the commercial viability of phosphate reserves.

Other applications

See also

References

  1. ↑ phosphoric acid. The Columbia Encyclopedia, Sixth Edition. 2001-05
  2. ↑ Katherine L Tucker, Kyoko Morita, Ning Qiao, Marian T Hannan, L Adrienne Cupples and Douglas P Kiel (2006). "Colas, but not other carbonated beverages, are associated with low bone mineral density in older women: The Framingham Osteoporosis Study". Am. J Clin. Nut. 84 (4): 936–42. PMID 17023723. http://www.ajcn.org/cgi/content/full/84/4/936. 
  3. ↑ S. Elmståhl, B. Gullberg, L. Janzon, O. Johnell and B. Elmståhl (1998). "Increased Incidence of Fractures in Middle-aged and Elderly Men with Low Intakes of Phosphorus and Zinc". Osteoporosis International 8 (4): 333–340. doi:10.1007/s001980050072. PMID 10024903. 
  4. ↑ 4.0 4.1 R. P. Heaney and K. Rafferty (2001). "Carbonated beverages and urinary calcium excretion". Am J Clin Nutr 74 (3): 343–347. PMID 11522558. 
  5. ↑ M. J. Barger-Lux, R. P. Heaney and M. R. Stegman (1990). "Effects of moderate caffeine intake on the calcium economy of premenopausal women [published erratum appears in Am J Clin Nutr 1991 Jan;53(1):182]". Am J Clin Nutr 52 (4): 722–725. PMID 2403065. 
  6. ↑ Tina M. Saldana, Olga Basso, Rebecca Darden, and Dale P. Sandler (2007). "Carbonated beverages and chronic kidney disease". Epidemiology 18 (4): 501–6. doi:10.1097/EDE.0b013e3180646338. PMID 17525693. 
  7. ↑ US 4721519 
  8. ↑ "Super Phosphoric Acid 0-68-0 Material Safety Data Sheet". J.R. Simplot Company. May 2009. http://msds.simplot.com/datasheets/12002.pdf. Retrieved 4 May 2010. 
  9. ↑ Fertilizer International, May-June 2008, article, page 81, "A credible alternative to the Wet Phosphoric Acid Process"
  10. ↑ Wolf, S.; R.N. Tauber (1986). Silicon Processing for the VLSI Era: Volume 1 - Process Technology. pp. 534. ISBN 0961672161. 
  11. ↑ "Ingredient Dictionary: P". Cosmetic Ingredient Dictionary. Paula's Choice. http://www.cosmeticscop.com/learn/cosmetic_dictionary.asp?id=21&letter=P. Retrieved 2007-11-16. 
  12. ↑ C. Toles, S. Rimmera and J. C. Hower (1996). "Production of activated carbons from a Washington lignite using phosphoric acid activation". Carbon 34 (11): 1419–1426. doi:10.1016/S0008-6223(96)00093-0. 
  13. ↑ http://terpconnect.umd.edu/~browns/wetetch.html

External links