X−O−H ⇄ (X−O)− + H+

s

• An oxyacid molecule contains the structure X−O−H, where other atoms or atom groups can be connected to the central atom X. In a solution, such a molecule can be dissociated into ions in two distinct ways:
• X−O−H ⇄ (X−O)− + H+
• X−O−H ⇄ X+ + OH−[3]
• If the central atom X is strongly electronegative, then it strongly attracts the electrons of the oxygen atom. In that case, the bond between the oxygen and hydrogen atom is weak, and the compound ionizes easily in the way of the former of the two chemical equations above. In this case, the compound XOH is an acid, because it releases a proton, that is, a hydrogen ion. For example, nitrogen, sulfur and chlorine are strongly electronegative elements, and therefore nitric acid, sulfuric acid, and perchloric acid, are strong acids.
• If, however, the electronegativity of X is low, then the compound is dissociated to ions according to the latter chemical equation, and XOH is an alkaline hydroxide. Examples of such compounds are sodium hydroxide NaOH and calcium hydroxide Ca(OH)2.[3] Owing to the high electronegativity of oxygen, however, most of the common oxobases, such as sodium hydroxide, while strongly basic in water, are only moderately basic in comparison to other bases. For example, the pKa of the conjugate acid of sodium hydroxide, water, is 15.7, while that of sodium amide, ammonia, is closer to 40, making sodium hydroxide a much weaker base than sodium amide.[3]
• If the electronegativity of X is somewhere in between, the compound can be amphoteric, and in that case it can dissociate to ions in both ways, in the former case when reacting with bases, and in the latter case when reacting with acids. Examples of this include aliphatic alcohols, such as ethanol.[3]
• Inorganic oxyacids typically have a chemical formula of type HmXOn, where X is an atom functioning as a central atom, whereas parameters m and n depend on the oxidation state of the element X. In most cases, the element X is a nonmetal, but some metals, for example chromium and manganese, can form oxyacids when occurring at their highest oxidation states.[3]

• Names of inorganic oxyacids
• Many inorganic oxyacids are traditionally called with names ending with the word acid and which also contain, in a somewhat modified form, the name of the element they contain in addition to hydrogen and oxygen. Well-known examples of such acids are sulfuric acid, nitric acid and phosphoric acid.
• This practice is fully well-established, and IUPAC has accepted such names. In light of the current chemical nomenclature, this practice is an exception, because systematic names of compounds are formed according to the elements they contain and their molecular structure, not according to other properties (for example, acidity) they have.[7]
• IUPAC, however, recommends against calling future compounds not yet discovered with a name ending with the word acid.[7] Indeed, acids can be called with names formed by adding the word hydrogen in front of the corresponding anion; for example, sulfuric acid could just as well be called hydrogen sulfate (or dihydrogen sulfate).[8] In fact, the fully systematic name of sulfuric acid, according to IUPAC's rules, would be dihydroxidodioxidosulfur and that of the sulfate ion, tetraoxidosulfate(2−),[9] Such names, however, are almost never used.
• However, the same element can form more than one acid when compounded with hydrogen and oxygen. In such cases, the English practice to distinguish such acids is to use the suffix -ic in the name of the element in the name of the acid containing more oxygen atoms, and the suffix -ous in the name of the element in the name of the acid containing fewer oxygen atoms. Thus, for example, sulfuric acid is H2SO4, and sulfurous acid, H2SO3. Analogously, nitric acid is HNO3, and nitrous acid, HNO2. If there are more than two oxyacids having the same element as the central atom, then, in some cases, acids are distinguished by adding the prefix per- or hypo- to their names. The prefix per-, however, is used only when the central atom is a halogen or a group 7 element.[8] For example, chlorine has the four following oxyacids:
• hypochlorous acid HClO
• chlorous acid HClO2
• chloric acid HClO3
• perchloric acid HClO4

• The suffix -ite occurs in names of anions and salts derived from acids whose names end to the suffix -ous. On the other hand, the suffix -ate occurs in names of anions and salts derived from acids whose names end to the suffix -ic. Prefixes hypo- and per- occur in the name of anions and salts; for example the ion ClO−
• 4 is called perchlorate.[8]
• In a few cases, the prefixes ortho- and para- occur in names of some oxyacids and their derivative anions. In such cases, the para- acid is what can be thought as remaining of the ortho- acid if a water molecule is separated from the ortho- acid molecule. For example, phosphoric acid, H3PO4, has sometimes been called orthophosphoric acid, in order to distinguish it from metaphosphoric acid, HPO3.[8] However, according to IUPAC's current rules, the prefix ortho- should only be used in names of orthotelluric acid and orthoperiodic acid, and their corresponding anions and salts.[10]
• Examples
• In the following table, the formula and the name of the anion refer to what remains of the acid when it loses all its hydrogen atoms as protons. Many of these acids, however, are polyprotic, and in such cases, there also exists one or more intermediate anions. In name of such anions, the prefix hydrogen- (in older nomenclature bi-) is added, with numeral prefixes if needed. For example, SO2−
• 4 is the sulfate anion, and HSO−
• 4, the hydrogensulfate (or bisulfate) anion. Similarly, PO3−
• 4 is phosphate, HPO2−
• 4 is hydrogenphosphate, and H
• 2PO−
• 4 is dihydrogenphosphate.

• oxyacid, any oxygen-containing acid. Most covalent nonmetallic oxides react with water to form acidic oxides; that is, they react with water to form oxyacids that yield hydronium ions (H3O+) in solution. There are some exceptions, such as carbon monoxide, CO, nitrous oxide, N2O, and nitric oxide, NO.
• The strength of an oxyacid is defined by the extent to which it dissociates in water (i.e., its ability to form H+ ions). In general, the relative strength of oxyacids can be predicted on the basis of the electronegativity and oxidation number of the central nonmetal atom.

• The acid strength increases as the electronegativity of the central atom increases. For example, because the electronegativity of chlorine (Cl) is greater than that of sulfur (S), which is in turn greater than that of phosphorus (P), it can be predicted that perchloric acid, HClO4, is a stronger acid than sulfuric acid, H2SO4, which should be a stronger acid than phosphoric acid, H3PO4. For a given nonmetal central atom, the acid strength increases as the oxidation number of the central atom increases. For example, nitric acid, HNO3, in which the nitrogen (N) atom has an oxidation number of +5, is a stronger acid than nitrous acid, HNO2, where the nitrogen oxidation state is +3. In the same manner, sulfuric acid, H2SO4, with sulfur in its +6 oxidation state, is a stronger acid than sulfurous acid, H2SO3, where a +4 oxidation number of sulfur exists.
• The salt of an oxyacid is a compound formed when the acid reacts with a base: acid + base → salt + water. This type of reaction is called neutralization, because the solution is made neutral.

• Oxyacids of nitrogen
• Nitric acid and nitrate salts
• Nitric acid, HNO3, was known to the alchemists of the 8th century as “aqua fortis” (strong water). It is formed by the reaction of both dinitrogen pentoxide (N2O5) and nitrogen dioxide (NO2) with water. Small amounts of nitric acid are found in the atmosphere after thunderstorms, and its salts, called nitrates, occur widely in nature. Enormous deposits of sodium nitrate, NaNO3, also known as Chile saltpetre, are found in the desert region near the boundary of Chile and Peru. These deposits can be 3 km (2 miles) wide, 300 km (200 miles) long, and up to 2 metres (7 feet) thick. Potassium nitrate, KNO3, sometimes called Bengal saltpetre, is found in India and other countries in East Asia. Nitric acid can be prepared in the laboratory by heating a nitrate salt, such as those mentioned above, with concentrated sulfuric acid; for example,
• NaNO3 + H2SO4 + heat → NaHSO4 + HNO3.
• Since HNO3 boils at 86 °C (187 °F) and H2SO4 boils at 338 °C (640 °F) and NaNO3 and NaHSO4 are nonvolatile salts, nitric acid is easily removed by distillation.

• Oxyacids of phosphorus
• Orthophosphoric acid and phosphate salts
• Orthophosphoric acid, H3PO4, is usually called simply phosphoric acid. When pure, it is a colourless crystalline solid that melts at 42 °C (108 °F). It rapidly absorbs moisture from the air and liquefies. It is typically available commercially as syrupy phosphoric acid, which is an 85 percent solution in water. Pure H3PO4 is produced by dissolving phosphorus pentoxide (P4O10) in water, although it is more commonly prepared by treating calcium phosphate, Ca3(PO4)2, with concentrated sulfuric acid, H2SO4.
• Ca3(PO4)2 + 3H2SO4 → 2H3PO4 + 3CaSO4
• The products are diluted with water, and the insoluble CaSO4 is removed by filtration. The dilute acid produced is contaminated with calcium dihydrogen phosphate, Ca(H2PO4)2, and other compounds found with naturally occurring Ca3(PO4)2.

• Orthophosphoric acid is a triprotic acid—i.e., it can donate all three of its hydrogen atoms as protons in aqueous solution. Thus, it can form three series of salts: dihydrogen phosphates, containing the H2PO4− ion; hydrogen phosphates, containing the HPO42− ion; and orthophosphates, containing the PO43− ion. When dissolved in water, soluble dihydrogen phosphate salts form solutions that are weakly acidic, because H2PO4− is a weak acid. Aqueous solutions of hydrogen phosphates are basic, because the HPO42− ion is stronger as a base (i.e., a proton acceptor) than as an acid. The PO43− ion is a moderately strong base, so orthophosphate salts form strongly basic solutions. The hydrogen phosphate salts, as well as H3PO4, decompose with loss of water when heated to form compounds containing P―O―P bonds. The ease of polymerization via P―O―P linkages and the possibility of the formation of P―P and P―H bonds allow innumerable oxyacids and their salts to be formed. These acids are termed the lower oxyacids of phosphorus.

• Phosphorous acid and phosphite salts
• Pure phosphorous acid, H3PO3, is best prepared by hydrolysis of phosphorus trichloride, PCl3.
• PCl3 + 3H2O → H3PO3 + 3HCl
• The resulting solution is heated to drive off the HCl, and the remaining water is evaporated until colourless crystalline H3PO3 appears on cooling. The acid can also be obtained by the action of water on P4O6, PBr3, or PI3. Colourless crystalline H3PO3 melts at 70.1 °C (158.2 °F), is very soluble in water, and has an odour similar to that of garlic. Heating phosphorous acid to about 200 °C (392 °F) causes it to disproportionate into phosphine, PH3, and orthophosphoric acid.
• 4H3PO3 + heat → PH3 + 3H3PO4
• Phosphorous acid and its salts are active reducing agents, because of their easy oxidation to phosphoric acid and phosphate salts, respectively. For example, phosphorous acid reduces the silver ion (Ag+) to elemental silver (Ag), mercury(II) salts to mercury(I) salts, and sulfurous acid, H2SO3, to elemental sulfur.