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An oxyacid, oxoacid, or ternary acid is an acid that contains oxygen. Specifically, it is a compound that contains hydrogen, oxygen, and at least one other element, with at least one hydrogen atom bonded to oxygen that can dissociate to produce the H+ cation and the anion of the acid.[1]


Under Lavoisier's original theory, all acids contained oxygen, which was named from the Greek ὀξύς (oxys: acid, sharp) and the root -γενής (-genes: creator). It was later discovered that some acids, notably hydrochloric acid, did not contain oxygen and so acids were divided into oxo-acids and these new hydroacids.

All oxyacids have the acidic hydrogen bound to an oxygen atom, so bond strength (length) is not a factor, as it is with binary nonmetal hydrides. Rather, the electronegativity of the central atom (X) and the number of O atoms determine oxyacid acidity. With the same central atom X, acid strength increases as the number of oxygens attached to X increases. With the same number of oxygens around E, acid strength increases with the electronegativity of X.

Compared to salt of their deprotonated forms, the oxyanions, oxyacids are generally less stable, and many of them only exist formally as hypothetical species, or exist only in solution and cannot be isolated in pure form. There are several general reasons for this: (1) they may condense to form oligomers (e.g., H2CrO4 to H2Cr2O7), or dehydrate all the way to form the anhydride (e.g., H2CO3 to CO2), (2) they may disproportionate to one compound of higher and another of lower oxidation state (e.g., HClO2 to HClO and HClO3), or (3) they might exist almost entirely as another, more stable tautomeric form (e.g., phosphorous acid P(OH)3 exists almost entirely as phosphonic acid HP(=O)(OH)2). Nevertheless, perchloric acid (HClO4), sulfuric acid (H2SO4), and nitric acid (HNO3) are a few common oxyacids that are relatively easily prepared as pure substances.

Imidic acids are created by replacing =O with =NR in an oxyacid.[2]


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]

When oxyacids are heated, many of them dissociate to water and the anhydride of the acid. In most cases, such anhydrides are oxides of nonmetals. For example, carbon dioxide, CO2, is the anhydride of carbonic acid, H2CO3, and sulfur trioxide, SO3, is the anhydride of sulfuric acid, H2SO4. These anhydrides react quickly with water and form those oxyacids again.[4]

Many organic acids, like carboxylic acids and phenols, are oxyacids.[3] Their molecular structure, however, is much more complicated than that of inorganic oxyacids.

Most of the commonly encountered acids are oxyacids.[3] Indeed, in the 18th century, Lavoisier assumed that all acids contain oxygen and that oxygen causes their acidity. Because of this, he gave to this element its name, oxygenium, derived from Greek and meaning acid-maker, which is still, in a more or less modified form, used in most languages.[5] Later, however, Humphry Davy showed that the so-called muriatic acid did not contain oxygen, despite its being a strong acid; instead, it is a solution of hydrogen chloride, HCl.[6] Such acids which do not contain oxygen are nowadays known as hydroacids.

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:

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 ClO4 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]


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 HSO4, the hydrogensulfate (or bisulfate) anion. Similarly, PO3−4 is phosphate, HPO2−4 is hydrogenphosphate, and H2PO4 is dihydrogenphosphate.

Oxyacids and their corresponding anions
Element group Element (central atom) Oxidation state Acid formula Acid name[8][9] Anion formula Anion name
6 Chromium +6 H2CrO4 Chromic acid CrO2−4 Chromate
H2Cr2O7 Dichromic acid Cr2O2−7 Dichromate
7 Manganese +7 HMnO4 Permanganic acid MnO4 Permanganate
+6 H2MnO4 Manganic acid MnO2−4 Manganate
Technetium +7 HTcO4 Pertechnetic acid TcO4 Pertechnetate
+6 H2TcO4 Technetic acid TcO2−4 Technetate
Rhenium +7 HReO4 Perrhenic acid ReO4 Perrhenate
+6 H2ReO4 Tetraoxorhenic(VI) acid ReO2−4 Rhenate(VI)
+5 HReO3 Trioxorhenic(V) acid ReO3 Trioxorhenate(V)
H3ReO4 Tetraoxorhenic(V) acid ReO3−4 Tetraoxorhenate(V)
H4Re2O7 Heptaoxodirhenic(V) acid Re2O4−7 Dirhenate(V)
8 Iron +6 H2FeO4 Ferric acid FeO42– Ferrate
Ruthenium +6 H2RuO4 Ruthenic acid RuO42– Ruthenate
+7 HRuO4 Perruthenic acid RuO4 Perruthenate (note difference in usage compared to osmium)
+8 H2RuO5 Hyperruthenic acid HRuO5 Hyperruthenate[11]
Osmium +6 H6OsO6 Osmic acid H4OsO62– Osmate
+8 H4OsO6 Perosmic acid H2OsO62– Perosmate (note difference in usage compared to ruthenium)
13 Boron +3 H3BO3 Boric acid
(formerly orthoboric acid)[10]
BO3−3 Borate
(formerly orthoborate)
(HBO2)n Metaboric acid BO2 Metaborate
14 Carbon +4 H2CO3 Carbonic acid CO2−3 Carbonate
Silicon +4 H4SiO4 Silicic acid
(formerly orthosilicic acid)[10]
SiO4−4 Silicate (formerly orthosilicate)
H2SiO3 Metasilicic acid SiO2−3 Metasilicate
14, 15 Carbon, nitrogen +4, −3 HOCN Cyanic acid OCN Cyanate
15 Nitrogen +5 HNO3 Nitric acid NO3 Nitrate
HNO4 Peroxynitric acid NO4 Peroxynitrate
H3NO4 Orthonitric acid NO3−4 Orthonitrate
+3 HNO2 Nitrous acid NO2 Nitrite
HOONO Peroxynitrous acid OONO Peroxynitrite
+2 H2NO2 Nitroxylic acid NO2−2 Nitroxylate
+1 H2N2O2 Hyponitrous acid N2O2−2 Hyponitrite
Phosphorus +5 H3PO4 Phosphoric acid
(formerly orthophosphoric acid)[10]
PO3−4 Phosphate
HPO3 Metaphosphoric acid PO3 Metaphosphate
H4P2O7 Pyrophosphoric acid
(diphosphoric acid)
P2O4−7 Pyrophosphate
H3PO5 Peroxomonophosphoric acid PO3−3 Peroxomonophosphate
+5, +3 (HO)2POPO(OH)2 Diphosphoric(III,V) acid O2POPOO2−2 Diphosphate(III,V)
+4 (HO)2OPPO(OH)2 Hypophosphoric acid
(diphosphoric(IV) acid)
O2OPPOO4−2 Hypophosphate
+3 H2PHO3 Phosphonic acid PHO2−3 Phosphonate
H2P2H2O5 Diphosphonic acid P2H2O5−3 Diphosphonate
+1 HPH2O2 Phosphinic acid (hypophosphorous acid) PH2O2 Phosphinate (hypophosphite)
Arsenic +5 H3AsO4 Arsenic acid AsO3−4 Arsenate
+3 H3AsO3 Arsenous acid AsO3−3 Arsenite
16 Sulfur +6 H2SO4 Sulfuric acid SO2−4 Sulfate
H2S2O7 Disulfuric acid S2O2−7 Disulfate
H2SO5 Peroxomonosulfuric acid SO2−5 Peroxomonosulfate
H2S2O8 Peroxodisulfuric acid S2O2−8 Peroxodisulfate
+5 H2S2O6 Dithionic acid S2O2−6 Dithionate
+5, 0 H2SxO6 Polythionic acids
(x = 3, 4...)
SxO2−6 Polythionates
+4 H2SO3 Sulfurous acid SO2−3 Sulfite
H2S2O5 Disulfurous acid S2O2−5 Disulfite
+4, 0 H2S2O3 Thiosulfuric acid S2O2−3 Thiosulfate
+3 H2S2O4 Dithionous acid S2O2−4 Dithionite
+3, −1 H2S2O2 Thiosulfurous acid S2O2−2 Thiosulfite
+2 H2SO2 Sulfoxylic acid (hyposulfurous acid) SO2−2 Sulfoxylate (hyposulfite)
+1 H2S2O2 Dihydroxydisulfane S2O2−2
0 HSOH Sulfenic acid HSO Sulfinite
Selenium +6 H2SeO4 Selenic acid SeO2−4 Selenate
+4 H2SeO3 Selenous acid SeO2−3 Selenite
Tellurium +6 H2TeO4 Telluric acid TeO2−4 Tellurate
H6TeO6 Orthotelluric acid TeO6−6 Orthotellurate
+4 H2TeO3 Tellurous acid TeO2−3 Tellurite
17 Chlorine +7 HClO4 Perchloric acid ClO4 Perchlorate
+5 HClO3 Chloric acid ClO3 Chlorate
+3 HClO2 Chlorous acid ClO2 Chlorite
+1 HClO Hypochlorous acid ClO Hypochlorite
Bromine +7 HBrO4 Perbromic acid BrO4 Perbromate
+5 HBrO3 Bromic acid BrO3 Bromate
+3 HBrO2 Bromous acid BrO2 Bromite
+1 HBrO Hypobromous acid BrO Hypobromite
Iodine +7 HIO4 Periodic acid IO4 Periodate
H5IO6 Orthoperiodic acid IO5−6 Orthoperiodate
+5 HIO3 Iodic acid IO3 Iodate
+1 HIO Hypoiodous acid IO Hypoiodite
18 Xenon +6 H2XeO4 Xenic acid HXeO4 Hydrogenxenate (dibasic xenate is unknown)
+8 H4XeO6 Perxenic acid XeO64– Perxenate


  • Kivinen, Antti; Mäkitie, Osmo (1988) (in fi). Kemia. Helsinki, Finland: Otava. ISBN 951-1-10136-6. 
  • (in fi) Otavan suuri ensyklopedia, volume 2 (Cid-Harvey). Helsinki, Finland: Otava. 1977. ISBN 951-1-04170-3. 

See also


  1. Chemistry, International Union of Pure and Applied. IUPAC Compendium of Chemical Terminology. IUPAC. doi:10.1351/goldbook.O04374. 
  2. Chemistry, International Union of Pure and Applied. IUPAC Compendium of Chemical Terminology. IUPAC. doi:10.1351/goldbook.I02949. 
  3. 3.0 3.1 3.2 3.3 3.4 3.5 3.6 Kivinen, Mäkitie: Kemia, p. 202-203, chapter=Happihapot
  4. "Hapot". Otavan iso Fokus, Part 2 (El-Io). Otava. 1973. pp. 990. ISBN 951-1-00272-4. 
  5. Otavan suuri Ensyklopedia, s. 1606, art. Happi
  6. Otavan suuri Ensyklopedia, s. 1605, art. Hapot ja emäxet
  7. 7.0 7.1 Red Book 2005, s. 124, chapter IR-8: Inorganic Acids and Derivatives
  8. 8.0 8.1 8.2 8.3 8.4 Kivinen, Mäkitie: Kemia, p. 459-461, chapter Kemian nimistö: Hapot
  9. 9.0 9.1 Red Book 2005, p. 129-132, table IR-8-1
  10. 10.0 10.1 10.2 10.3 Red Book 2005, p. 132, note a
  11. Encyclopedia of electrochemical power sources. Garche, Jürgen., Dyer, Chris K.. Amsterdam: Academic Press. 2009. pp. 854. ISBN 978-0444527455. OCLC 656362152. 

External links