This article is about orthophosphoric acid. For other acids commonly called "phosphoric acid", see Phosphoric acids and phosphates.
| Phosphoric acid |
|
|
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IUPAC name
trihydroxidooxidophosphorus
phosphoric acid
|
Other names
Orthophosphoric acid
trihydroxylphosphine oxide
|
| Identifiers |
| CAS number |
7664-38-2 Y,
16271-20-8 (hemihydrate) |
| PubChem |
1004 |
| ChemSpider |
979 Y |
| UNII |
E4GA8884NN Y |
| EC number |
231-633-2 |
| UN number |
1805 |
| KEGG |
D05467 Y |
| ChEBI |
CHEBI:26078 Y |
| ChEMBL |
CHEMBL1187 Y |
| RTECS number |
TB6300000 |
| Jmol-3D images |
Image 1 |
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|
-
InChI=1S/H3O4P/c1-5(2,3)4/h(H3,1,2,3,4) Y
Key: NBIIXXVUZAFLBC-UHFFFAOYSA-N Y
InChI=1/H3O4P/c1-5(2,3)4/h(H3,1,2,3,4)
Key: NBIIXXVUZAFLBC-UHFFFAOYAI
|
| Properties |
| Molecular formula |
H3O4P |
| Molar mass |
98.00 g mol−1 |
| Appearance |
white solid or colourless, viscous liquid (>42 °C)
deliquescent |
| 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 (108.23 °F; 315.50 K)
(anhydrous)
29.32 °C (84.78 °F; 302.47 K)
(hemihydrate) |
| Boiling point |
158 °C (316 °F; 431 K)
213 °C (415 °F; 486 K)
decomposes |
| Solubility in water |
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)[1] |
| Solubility |
soluble in ethanol |
| Acidity (pKa) |
1 = 2.148
2 = 7.198
3 = 12.319 |
| Refractive index (nD) |
1.34203 |
| Viscosity |
2.4–9.4 cP (85% aq. soln.)
147 cP (100%) |
| Structure |
| Crystal structure |
monoclinic |
| Thermochemistry |
Std molar
entropy So298 |
158 J/mol·K[2] |
Std enthalpy of
formation ΔfHo298 |
-1288 kJ/mol[2] |
| Hazards |
| MSDS |
ICSC 1008 |
| GHS pictograms |
[3] |
| GHS signal word |
Danger |
| GHS hazard statements |
H290, H314[3] |
| GHS precautionary statements |
P280, P305+351+338, P310[3] |
| EU Index |
015-011-00-6 |
| EU classification |
C |
| R-phrases |
R34 |
| S-phrases |
(S1/2), S26, S45 |
| NFPA 704 |
|
| Flash point |
Non-flammable |
| LD50 |
1530 mg/kg (rat, oral) |
| Related compounds |
| Related phosphorus oxoacids |
Hypophosphorous acid
Phosphorous acid
Pyrophosphoric acid
Triphosphoric acid
Perphosphoric acid
Permonophosphoric acid |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa) |
| Y (verify) (what is: Y/N?) |
| Infobox references |
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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. Orthophosphoric acid refers to phosphoric acid, which is the IUPAC name for this compound. The prefix ortho is used to distinguish the acid from related phosphoric acids, called polyphosphoric acids. Orthophosphoric acid is a non-toxic acid, which, when pure, is a solid at room temperature and pressure.
The conjugate base of phosphoric acid is the dihydrogen phosphate ion, H
2PO−
4, which in turn has a conjugate base of hydrogen phosphate, HPO2−
4, which has a conjugate base of phosphate, PO3−
4.
In addition to being a chemical reagent, phosphoric acid has a wide variety of uses, including as a rust inhibitor, food additive, dental and orthop(a)edic etchant, electrolyte, flux, dispersing agent, industrial etchant, fertilizer feedstock, and component of home cleaning products.
The most common source of phosphoric acid is an 85% aqueous solution; such solutions are colourless, odourless, and non-volatile. The 85% solution is a rather viscous, syrupy liquid, but still pourable. 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. For the sake of labeling and simplicity, the 85% represents H3PO4 as if it were all orthophosphoric acid. Dilute aqueous solutions of phosphoric acid exist in the ortho- form.
Contents
- 1 Reactions
- 2 Preparation
- 2.1 Wet
- 2.2 Thermal
- 2.3 Laboratory routes
- 3 Uses
- 3.1 Food additive
- 3.2 Niche uses
- 3.2.1 Rust removal
- 3.2.2 In medicine
- 3.2.3 Other applications
- 4 Biological effects
- 5 See also
- 6 References
- 7 External links
Reactions
Anhydrous phosphoric acid, a white low melting solid, is obtained by dehydration of 85% phosphoric acid by heating under a vacuum.[4]
Orthophosphoric acid is a very polar molecule. It is infinitely 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) H3O+(aq) + H2PO4−(aq) Ka1= 7.25×10−3
- H2PO4−(aq)+ H2O(l) H3O+(aq) + HPO42−(aq) Ka2= 6.31×10−8
- HPO42−(aq)+ H2O(l) H3O+(aq) + PO43−(aq) Ka3= 4.80×10−13
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.[5] 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 (phosphorus pentoxide), 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.
In the presence of superacids (acids stronger than H
2SO
4), H
3PO
4 reacts to form mystery products, perhaps corrosive, acidic salts of the hypothetical[6] tetrahydroxylphosphonium ion, which is isoelectronic with orthosilicic acid. The suspected reaction with HSbF
6, for example, is supposed to go:
- H3PO4 + {HSbF6} → [P(OH)4+] [SbF6]−
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 strong acid concentrations, the solution is mainly composed of H3PO4. For [A] = 10−2, the pH is close 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. Since this analysis does not take into account ion activity coefficients, the pH and molarity of a real phosphoric acid solution may deviate substantially from the above values.
Preparation
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
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 + 10 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 further purified by removing fluorine to produce animal-grade phosphoric acid, or by solvent extraction and arsenic removal to produce food-grade phosphoric acid.
The nitrophosphate process is similar to the wet process except that it uses nitric acid in place of sulfuric acid. The advantage to this route is that the coproduct, calcium nitrate is also a plant fertilizer. This method is rarely employed.
Thermal
Very pure phosphoric acid is obtained by burning elemental phosphorus to produce phosphorus pentoxide, which is subsequently dissolved in dilute phosphoric acid. This route 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.
Elemental phosphorus is produced by an electric furnace. At a high temperature, a mixture of phosphate ore, silica and carbonaceous material (coke, coal etc...) produces calcium silicate, phosphorus gas and carbon monoxide. The P and CO off-gases from this reaction are cooled under water to isolate solid phosphorus. Alternatively, the P and CO off-gases can be burned with air to produce phosphorus pentoxide and carbon dioxide.
Laboratory routes
A demonstrative process consists in the oxidation of red phosphorus by nitric acid.[9]
- P + 5 HNO3 → H2O + H3PO4 + 5 NO2
Uses
The dominant use of phosphoric acid is for fertilizers, consuming approximately 90% of production.[10]
| Application |
Demand (2006) in thousands of tons |
Main phosphate derivatives |
| Soaps and detergents |
1836 |
STPP |
| Food industry |
309 |
STPP (Na5P3O10), SHMP, TSP, SAPP, SAlP (NaA, MCP, DSP (Na2HPO4), H3PO4 |
| Water treatment |
164 |
SHMP, STPP, TSPP, MSP (NaH2PO4), DSP |
| Toothpastes |
68 |
DCP (CaHPO4), IMP, SMFP |
| Other applications |
287 |
STPP (Na3P3O9), TCP, APP, DAP, zinc phosphate (Zn3(PO4)2), aluminium phosphate (AlPO4, H3PO4) |
Food additive
Food-grade phosphoric acid (additive E338[11]) is used to acidify foods and beverages such as various colas. It provides a tangy and sour taste. Various salts of phosphoric acid, such as monocalcium phosphate, are used as leavening agents.
Niche uses
Phosphoric acid and its derivatives are pervasive and find many niche applications.
Rust removal
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, Fe2O3 (rust) to ferric phosphate, FePO4. An empirical formula for this reaction is:
- 2 H3PO4 + Fe2O3 → 2 FePO4 + 3 H2O
Liquid phosphoric acid may be used for dipping, but phosphoric acid for rust removal is more often formulated as a gel. As a thick gel, it may be applied to sloping, vertical, or even overhead surfaces. Different phosphoric acid gel formulations are sold as "rust removers" or "rust killers". Multiple applications of phosphoric acid may be required to remove all rust. Rust may also be removed via phosphate conversion coating. This process can leave a black phosphate coating that provides moderate corrosion resistance (such protection is also provided by the superficially similar Parkerizing and blued electrochemical conversion coating processes).
In medicine
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.
Other applications
Among other applications, phosphoric acid is used:
- As an external standard for phosphorus-31 Nuclear magnetic resonance (NMR).
- As a buffer agent in biology and chemistry; For example, a buffer for high-performance liquid chromatography.
- As a chemical oxidizing agent for activated carbon production, as used in the Wentworth Process.[12]
- As the electrolyte in phosphoric acid fuel cells.
- With distilled water (2–3 drops per gallon) as an electrolyte in oxyhydrogen generators.
- As a catalyst in the hydration of alkenes to produce alcohols, predominantly ethanol.
- As an electrolyte in copper electropolishing for burr removal and circuit board planarization.
- As a flux by hobbyists (such as model railroaders) as an aid to soldering.
- In compound semiconductor processing, phosphoric acid is a common wet etching agent: for example, in combination with hydrogen peroxide and water it is used to etch InGaAs selective to InP.[13]
- Heated in microfabrication to etch silicon nitride (Si3N4). It is highly selective in etching Si3N4 instead of SiO2, silicon dioxide.[14]
- As a cleaner by construction trades to remove mineral deposits, cementitious smears, and hard water stains.
- As a chelant in some household cleaners aimed at similar cleaning tasks.
- In hydroponics pH solutions to lower the pH of nutrient solutions. While other types of acids can be used, phosphorus is a nutrient used by plants, especially during flowering, making phosphoric acid particularly desirable.
- As a pH adjuster in cosmetics and skin-care products.[15]
- As a dispersing agent in detergents and leather treatment.
- As an additive to stabilize acidic aqueous solutions within a wanted and specified pH range.
Biological effects
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This section requires expansion with: health effects beyond soft drinks. (January 2013) |
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This section needs more medical references for verification or relies too heavily on primary sources. Please review the contents of the section and add the appropriate references if you can. Unsourced or poorly sourced material may be removed. (February 2014) |
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In soft drinks
Phosphoric acid, used in many soft drinks (primarily cola), has been linked in epidemiological studies to (1) chronic kidney disease and (2) lower bone density.
(1) A study performed by the Epidemiology Branch of the US National Institute of Environmental Health Sciences, concludes that drinking 2 or more colas per day was associated with doubling the risk of chronic kidney disease.[16]
(2) A study[17] 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.
On the other hand, another study 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.[18]
A clinical study by Heaney and Rafferty using calcium-balance methods found no impact of carbonated soft drinks containing phosphoric acid on calcium excretion.[19] 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,[20] 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[19] (another possible confounding factor may be an association between high soft drink consumption and sedentary lifestyle)[citation needed].
See also
- Phosphate fertilizers, such as ammonium phosphate fertilizers
References
- ^ Seidell, Atherton; Linke, William F. (1952). [Google Books Solubilities of Inorganic and Organic Compounds]. Van Nostrand. Retrieved 2014-06-02.
- ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A22. ISBN 0-618-94690-X.
- ^ a b c Sigma-Aldrich Co., Phosphoric acid. Retrieved on 2014-05-09.
- ^ R. Klement "Orthophosphoric Acid" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 543.
- ^ phosphoric acid. The Columbia Encyclopedia, Sixth Edition.
- ^ Gevrey, S.; Luna, A.; Haldys, V.; Tortajada, J.; Morizur, J. P. (1998). "Experimental and theoretical studies of the gas-phase protonation of orthophosphoric acid". The Journal of Chemical Physics 108 (6): 2458. Bibcode:1998JChPh.108.2458G. doi:10.1063/1.475628. edit
- ^ Thomas, W P and Lawton, W S "Stable ammonium polyphosphate liquid fertilizer from merchant grade phosphoric acid" U.S. Patent 4,721,519, Issue date: January 26, 1988
- ^ "Super Phosphoric Acid 0-68-0 Material Safety Data Sheet". J.R. Simplot Company. May 2009. Retrieved 4 May 2010.
- ^ Arthur Sutcliffe (1930) Practical Chemistry for Advanced Students (1949 Ed.), John Murray - London.
- ^ Klaus Schrödter, Gerhard Bettermann, Thomas Staffel, Friedrich Wahl, Thomas Klein, Thomas Hofmann "Phosphoric Acid and Phosphates" in Ullmann’s Encyclopedia of Industrial Chemistry 2008, Wiley-VCH, Weinheim. doi:10.1002/14356007.a19_465.pub3
- ^ "Current EU approved additives and their E Numbers". Foods Standards Agency. 14 March 2012. Retrieved 22 July 2012.
- ^ Toles, C.; Rimmer, S.; Hower, J. C. (1996). "Production of activated carbons from a washington lignite using phosphoric acid activation". Carbon 34 (11): 1419. doi:10.1016/S0008-6223(96)00093-0. edit
- ^ Wet chemical etching. umd.edu
- ^ Wolf, S.; R.N. Tauber (1986). Silicon processing for the VLSI era: Volume 1 – Process technology. p. 534. ISBN 0-9616721-6-1.
- ^ "Ingredient dictionary: P". Cosmetic ingredient dictionary. Paula's Choice. Retrieved 16 November 2007.
- ^ http://www.ncbi.nlm.nih.gov/pubmed/17525693?ordinalpos=1&itool=EntrezSystem2.PEntrez.Pubmed.Pubmed_ResultsPanel.Pubmed_DefaultReportPanel.Pubmed_RVDocSum
- ^ Tucker, K. L.; Morita, K.; Qiao, N.; Hannan, M. T.; Cupples, L. A.; Kiel, D. P. (2006). "Colas, but not other carbonated beverages, are associated with low bone mineral density in older women: The Framingham Osteoporosis Study". The American journal of clinical nutrition 84 (4): 936–942. PMID 17023723. edit
- ^ Elmståhl, S.; Gullberg, B.; Janzon, L.; Johnell, O.; Elmståhl, B. (1998). "Increased incidence of fractures in middle-aged and elderly men with low intakes of phosphorus and zinc". Osteoporosis international : a journal established as result of cooperation between the European Foundation for Osteoporosis and the National Osteoporosis Foundation of the USA 8 (4): 333–340. doi:10.1007/s001980050072. PMID 10024903. edit
- ^ a b Heaney, R. P.; Rafferty, K. (2001). "Carbonated beverages and urinary calcium excretion". The American journal of clinical nutrition 74 (3): 343–347. PMID 11522558. edit
- ^ Barger-Lux, M. J.; Heaney, R. P.; Stegman, M. R. (1990). "Effects of moderate caffeine intake on the calcium economy of premenopausal women". The American journal of clinical nutrition 52 (4): 722–725. PMID 2403065. edit
External links
|
Wikimedia Commons has media related to phosphoric acid. |
- International chemical safety card 1008
- National pollutant inventory – Phosphoric acid fact sheet
- NIOSH Pocket guide to chemical hazards
- Excel spreadsheet containing phosphoric acid titration curve, distribution diagram and buffer pH calculation
|
Hydrogen compounds
|
|
- H3AsO3
- H3AsO4
- HAt
- HSO3F
- HBF4
- HBr
- HBrO
- HBrO2
- HBrO3
- HBrO4
- HCl
- HClO
- HClO2
- HClO3
- HClO4
- HCN
- HCNO
- H2CrO4/H2Cr2O7
- H2CO3
- H2CS3
- HF
- HFΟ
- HI
- HIO
- HIO2
- HIO3
- HIO4
- HMnO4
- H2MoO4
- HNC
- HNCO
- HNO
- HNO3
- H2N2O2
- HNO5S
- H3NSO3
- H2O
- H2O2
- H2O3
- H3PO2
- H3PO3
- H3PO4
- H4P2O7
- H5P3O10
- H2PtCl6
- H2S
- H2Se
- H2SeO3
- H2SeO4
- H4SiO4
- H2SiF6
- HSCN
- HNSC
- H2SO3
- H2SO4
- H2SO5
- H2S2O3
- H2S2O6
- H2S2O7
- H2S2O8
- CF3SO3H
- H2Te
- H2TeO3
- H6TeO6
- H4TiO4
- H2Po
- HCo(CO)4
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