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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]

Description

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Under Lavoisier's original theory, all acids contained oxygen, which was named from Ancient Greek: ὀξύς + -γενής, romanized: oxys + -genes, lit. 'acid, sharp + 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 and the number of oxygen atoms determine oxyacid acidity. For oxyacids with the same central atom, acid strength increases with the number of oxygen atoms attached to it. With the same number of oxygen atoms attached to it, acid strength increases with increasing electronegativity of the central atom.

Compared to the salts of their deprotonated forms (a class of compounds known as the oxyanions), oxyacids are generally less stable, and many of them only exist formally as hypothetical species, or only exist 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]

Properties

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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. The acidity of oxoacids is also affected by the resonance stabilization of their conjugate bases. Double-bonded oxygen is electron withdrawing by resonance, so the negative charge of a deprotonated hydroxyl group can be distributed to other oxygen atoms. Both acetic acid and methanol contain C-O-H groups that can act as acids, but acetic acid is a far stronger acid because its conjugate base, acetate, can distribute its negative charge over two oxygen atoms. In contrast, the conjugate acid of methanol has the negative charge localized on oxygen, so it is a far stronger base than acetate, making acetic acid the stronger acid.

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 14.0, while that of sodium amide, ammonia, is closer to 40, making sodium hydroxide a much weaker base than sodium amide.[4][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 water, aliphatic alcohols, such as ethanol, and aluminum hydroxide.[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.[5]

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.[6] 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.[7] Such acids which do not contain oxygen are nowadays known as hydroacids.

Names of inorganic oxyacids

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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.[8]

IUPAC, however, recommends against calling future compounds not yet discovered with a name ending with the word acid.[8] 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).[9] 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−),[10] 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.[9] For example, chlorine has the four following oxyacids:

Some elemental atoms can exist in a high enough oxidation state that they can hold one more double-bonded oxygen atom than the perhalic acids do. In that case, any acids regarding such element are given the prefix hyper-. Currently, the only known acid with this prefix is hyperruthenic acid, H2RuO5.

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.[9]

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.[9] 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.[11]

Examples

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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.

Oxyacids and their corresponding anions
Element group Element (central atom) Oxidation state Acid formula Acid name[9][10] Anion formula Anion name
6 Chromium +6 H
2CrO
4		 Chromic acid CrO2−
4		 Chromate
H
2Cr
2O
7		 Dichromic acid Cr
2O2−
7		 Dichromate
7 Manganese +7 HMnO
4		 Permanganic acid MnO−
4		 Permanganate
+6 H
2MnO
4		 Manganic acid MnO2−
4		 Manganate
Technetium +7 HTcO
4		 Pertechnetic acid TcO−
4		 Pertechnetate
+6 H
2TcO
4		 Technetic acid TcO2−
4		 Technetate
Rhenium +7 HReO
4		 Perrhenic acid ReO−
4		 Perrhenate
+6 H
2ReO
4		 Tetraoxorhenic(VI) acid ReO2−
4		 Rhenate(VI)
+5 HReO
3		 Trioxorhenic(V) acid ReO−
3		 Trioxorhenate(V)
H
3ReO
4		 Tetraoxorhenic(V) acid ReO3−
4		 Tetraoxorhenate(V)
H
4Re
2O
7		 Heptaoxodirhenic(V) acid Re
2O4−
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[12]
Osmium +6 H6OsO6 Osmic acid H4OsO62– Osmate
+8 H4OsO6 Perosmic acid H2OsO62– Perosmate (note difference in usage compared to ruthenium)
13 Boron +3 H
3BO
3		 Boric acid
(formerly orthoboric acid)[11]		 BO3−
3		 Borate
(formerly orthoborate)
(HBO
2)
n		 Metaboric acid BO−
2		 Metaborate
14 Carbon +4 H
2CO
3		 Carbonic acid CO2−
3		 Carbonate
Silicon +4 H
4SiO
4		 Silicic acid
(formerly orthosilicic acid)[11]		 SiO4−
4		 Silicate (formerly orthosilicate)
H
2SiO
3		 Metasilicic acid SiO2−
3		 Metasilicate
14, 15 Carbon, nitrogen +4, −3 HOCN Cyanic acid OCN−
 Cyanate
15 Nitrogen +5 HNO
3		 Nitric acid NO−
3		 Nitrate
HNO
4		 Peroxynitric acid NO−
4		 Peroxynitrate
H
3NO
4		 Orthonitric acid NO3−
4		 Orthonitrate
+3 HNO
2		 Nitrous acid NO−
2		 Nitrite
HOONO Peroxynitrous acid OONO−
 Peroxynitrite
+2 H
2NO
2		 Nitroxylic acid NO2−
2		 Nitroxylate
+1 H
2N
2O
2		 Hyponitrous acid N
2O2−
2		 Hyponitrite
Phosphorus +5 H
3PO
4		 Phosphoric acid
(formerly orthophosphoric acid)[11]		 PO3−
4		 Phosphate
(orthophosphate)
HPO
3		 Metaphosphoric acid PO−
3		 Metaphosphate
H
4P
2O
7		 Pyrophosphoric acid
(diphosphoric acid)		 P
2O4−
7		 Pyrophosphate
(diphosphate)
H
3PO
5		 Peroxomonophosphoric acid PO3−
3		 Peroxomonophosphate
+5, +3 (HO)
2POPO(OH)
2		 Diphosphoric(III,V) acid O
2POPOO2−
2		 Diphosphate(III,V)
+4 (HO)
2OPPO(OH)
2		 Hypophosphoric acid
(diphosphoric(IV) acid)		 O
2OPPOO4−
2		 Hypophosphate
(diphosphate(IV))
+3 H
2PHO
3		 Phosphonic acid PHO2−
3		 Phosphonate
H
2P
2H
2O
5		 Diphosphonic acid P
2H
2O5−
3		 Diphosphonate
+1 HPH
2O
2		 Phosphinic acid (hypophosphorous acid) PH
2O−
2		 Phosphinate (hypophosphite)
Arsenic +5 H
3AsO
4		 Arsenic acid AsO3−
4		 Arsenate
+3 H
3AsO
3		 Arsenous acid AsO3−
3		 Arsenite
16 Sulfur +6 H
2SO
4		 Sulfuric acid SO2−
4		 Sulfate
H
2S
2O
7		 Disulfuric acid S
2O2−
7		 Disulfate
H
2SO
5		 Peroxomonosulfuric acid SO2−
5		 Peroxomonosulfate
H
2S
2O
8		 Peroxodisulfuric acid S
2O2−
8		 Peroxodisulfate
+5 H
2S
2O
6		 Dithionic acid S
2O2−
6		 Dithionate
+5, 0 H
2S
xO
6		 Polythionic acids
(x = 3, 4...)		 S
xO2−
6		 Polythionates
+4 H
2SO
3		 Sulfurous acid SO2−
3		 Sulfite
H
2S
2O
5		 Disulfurous acid S
2O2−
5		 Disulfite
+4, 0 H
2S
2O
3		 Thiosulfuric acid S
2O2−
3		 Thiosulfate
+3 H
2S
2O
4		 Dithionous acid S
2O2−
4		 Dithionite
+3, −1 HOSOSH Thiosulfurous acid OSOS2−
 Thiosulfite
+2 H
2SO
2		 Sulfoxylic acid (hyposulfurous acid) SO2−
2		 Sulfoxylate (hyposulfite)
+1 HOSSOH Dihydroxydisulfane OSSO2−
 Disulfanediolate[13]
0 HSOH Sulfenic acid HSO−
 Sulfinite
Selenium +6 H
2SeO
4		 Selenic acid SeO2−
4		 Selenate
+4 H
2SeO
3		 Selenous acid SeO2−
3		 Selenite
Tellurium +6 H
2TeO
4		 Telluric acid TeO2−
4		 Tellurate
H
6TeO
6		 Orthotelluric acid TeO6−
6		 Orthotellurate
+4 H
2TeO
3		 Tellurous acid TeO2−
3		 Tellurite
17 Chlorine +7 HClO
4		 Perchloric acid ClO−
4		 Perchlorate
+5 HClO
3		 Chloric acid ClO−
3		 Chlorate
+3 HClO
2		 Chlorous acid ClO−
2		 Chlorite
+1 HClO Hypochlorous acid ClO−
 Hypochlorite
Bromine +7 HBrO
4		 Perbromic acid BrO−
4		 Perbromate
+5 HBrO
3		 Bromic acid BrO−
3		 Bromate
+3 HBrO
2		 Bromous acid BrO−
2		 Bromite
+1 HBrO Hypobromous acid BrO−
 Hypobromite
Iodine +7 HIO
4		 Periodic acid IO−
4		 Periodate
H
5IO
6		 Orthoperiodic acid IO5−
6		 Orthoperiodate
+5 HIO
3		 Iodic acid IO−
3		 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

Sources

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  • Kivinen, Antti; Mäkitie, Osmo (1988). Kemia (in Finnish). Helsinki, Finland: Otava. ISBN 951-1-10136-6.
  • Nomenclature of Inorganic Compounds, IUPAC Recommendations 2005 (Red Book 2005). International Union of Pure and Applied Chemistry. 2005. ISBN 0-85404-438-8.[dead link]
  • Otavan suuri ensyklopedia, volume 2 (Cid-Harvey) (in Finnish). Helsinki, Finland: Otava. 1977. ISBN 951-1-04170-3.

See also

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References

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[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Oxyacid

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An oxyacid, also known as an oxoacid, is a that contains oxygen, at least one other element, and atoms bound to oxygen, from which it can lose hydrons to form a conjugate base; this distinguishes it from hydracids like HCl, which lack oxygen in the acidic group. These acids typically follow the general formula \ceHmXOn\ce{H_mXO_n}, where \ceX\ce{X} is a central atom—usually a or early —and mm and nn indicate the number of hydrogen and oxygen atoms, respectively, with the acidic hydrogens attached to oxygen atoms. Common examples of oxyacids include (\ceH2SO4\ce{H2SO4}), (\ceHNO3\ce{HNO3}), (\ceH3PO4\ce{H3PO4}), and (\ceHClO4\ce{HClO4}), many of which are derived from the reaction of nonmetallic oxides with and play crucial roles in , biological systems, and laboratory chemistry. Oxyacids are often polyprotic, meaning they can donate multiple protons, with successive dissociation constants decreasing due to the increasing stability of the conjugate bases. Naming of oxyacids follows systematic conventions based on the central atom and the number of oxygen atoms relative to the highest ; for instance, acids with fewer oxygen atoms end in -ous (e.g., \ceH2SO3\ce{H2SO3} as ), while those with more end in -ic (e.g., \ceH2SO4\ce{H2SO4} as ), with prefixes like hypo- for the lowest and per- for the highest oxygen content. Their acidity trends are influenced by the of the central atom—higher electronegativity increases acidity (e.g., \ceHOCl>\ceHOBr>\ceHOI\ce{HOCl} > \ce{HOBr} > \ce{HOI})—and the number of oxygen atoms to it, as additional oxygens stabilize the conjugate base through inductive effects (e.g., \ceHClO4>\ceHClO3>\ceHClO2>\ceHClO\ce{HClO4} > \ce{HClO3} > \ce{HClO2} > \ce{HClO}).

Definition and Nomenclature

Definition

Oxyacids, also known as oxoacids or oxygen acids, are acids containing oxygen in the acidic group, specifically compounds with at least one bound to oxygen, which is further connected to a central atom—typically a , , or early . These acids produce their conjugate base, an oxoanion, upon dissociation by losing one or more hydron ions (H⁺). The general structural formula for oxyacids is often expressed as HmXOnH_m XO_n, where XX represents the central atom, and mm and nn are positive integers denoting the number of and oxygen atoms, respectively. In these structures, the acidic hydrogens are directly attached to oxygen atoms, enabling ionization in aqueous solutions. The term "oxyacid" is primarily used for inorganic acids and differs from binary acids (hydracids), such as (HCl), which consist solely of and a single element without oxygen in the acidic group. Although the structural definition could apply to some organic acids like carboxylic acids (which have oxygen in the acidic ), such compounds are conventionally classified as organic acids rather than oxyacids. The term "oxyacid" originated in the early (first recorded 1830–1840), during a period of advancing chemical understanding that included the formulation of acid-base theory by in the 1880s. Common examples include (H₂SO₄) and (HNO₃).

Nomenclature

The nomenclature of oxyacids follows the recommendations of the International Union of Pure and Applied Chemistry (IUPAC), which provide both systematic and retained traditional names to reflect the of the central atom and the number of oxygen atoms present. In the traditional system, preferred for common use, the name is derived from the root of the central atom, with suffixes indicating the : the "-ous" suffix denotes a lower (fewer oxygen atoms), while the "-ic" suffix denotes a higher (more oxygen atoms). Prefixes modify these when multiple s exist: "hypo-" indicates the lowest state, and "per-" the highest. Systematic names, less commonly used, employ additive based on coordination entities, such as "tetraoxidosulfate(2−) with 2H" for . Common naming patterns illustrate these rules across element families. For sulfur oxyacids, the lower oxidation state compound H2SO3H_2SO_3 is named (+4 oxidation state), while the higher state H2SO4H_2SO_4 is (+6 oxidation state). Similarly, chlorine oxyacids progress with increasing oxygen and oxidation state: HClOHClO as (+1), HClO2HClO_2 as (+3), HClO3HClO_3 as (+5), and HClO4HClO_4 as (+7). These patterns ensure names convey the relative oxygen content and reactivity trends associated with oxidation states. The corresponding oxyanions are named by replacing the acid suffixes with "-ate" or "-ite": for example, the anion from , SO42−SO_4^{2-}, is , while from , SO32−SO_3^{2-}, it is . Prefixes carry over similarly, yielding (ClO−ClO^-) from and (ClO4−ClO_4^-) from . This anion nomenclature extends to salts and other derivatives, maintaining consistency with the parent acid. Certain oxyacids retain traditional or trivial names despite available systematic alternatives, as approved by IUPAC for historical and practical reasons. For instance, HNO3HNO_3 is universally called , a retained name, rather than the additive form "trioxonitrate(1−) with H." Other retained examples include (H3PO4H_3PO_4) and (H2CO3H_2CO_3), which prioritize familiarity in scientific and industrial contexts. These exceptions are listed in IUPAC tables to guide consistent usage.

Properties

Physical Properties

Oxyacids exhibit a range of physical states at room temperature, primarily as liquids or solids, depending on their molecular structure and intermolecular forces. Common examples include nitric acid (HNO₃), which appears as a fuming, pale yellow to reddish-brown liquid with a suffocating odor, and sulfuric acid (H₂SO₄), a colorless, viscous, oily liquid. Pure phosphoric acid (H₃PO₄) is a transparent crystalline solid, though it is typically handled as a concentrated aqueous solution that remains liquid at room temperature. Perchloric acid (HClO₄) is also a clear, colorless liquid in its concentrated form. Most oxyacids are highly soluble in water, owing to extensive hydrogen bonding between their hydroxyl groups and water molecules, often resulting in miscibility. For instance, sulfuric acid is completely miscible with water, releasing significant heat upon dilution, while nitric acid is similarly fully miscible. Many oxyacids form azeotropic mixtures with water, which complicates their purification by distillation; sulfuric acid forms a maximum-boiling azeotrope at approximately 98.3 wt% H₂SO₄, and nitric acid at 68 wt% HNO₃. The melting and s of oxyacids show trends influenced by molecular weight, the number of bonds, and overall polarity, with higher values generally observed for those capable of stronger intermolecular interactions. has a relatively low point of 83 °C and of -42 °C, whereas boils at 337 °C with a of 10 °C, reflecting its greater and -bonding capacity. melts at 42 °C, and at -18 °C, with the latter at 203 °C. These properties establish the scale of thermal stability for handling and processing oxyacids. Densities and viscosities among oxyacids vary significantly, often higher than those of simple binary acids due to their polar nature and molecular size. , for example, has a of 1.84 g/cm³ at 20 °C and a of 21 mPa·s at 25 °C, contributing to its syrupy texture. In contrast, has a lower of 1.51 g/cm³ at 20 °C and of 0.75 mPa·s at 25 °C, making it more fluid. The following table summarizes key physical properties for representative oxyacids:
OxyacidState at 25 °CMelting Point (°C)Boiling Point (°C)Density (g/cm³ at 20–25 °C)Viscosity (mPa·s at 25 °C)
HNO₃831.510.75
H₂SO₄103371.8421
H₃PO₄ (85% aq.)~21~1581.68~40
HClO₄ (70%)-182031.67~3.5

Chemical Properties

Oxyacids are characterized by their ability to donate protons from hydroxyl groups attached to a central atom, leading to ionization in according to the general equilibrium: \ceHmXOn⇌H++Hm−1XOn−\ce{H_m XO_n ⇌ H+ + H_{m-1} XO_n^-} The acidity strength is measured by the pKa value, defined as pKa=−log⁡Ka\mathrm{p}K_a = -\log K_a, where KaK_a is the acid dissociation constant. Strong oxyacids, such as perchloric acid (\ceHClO4\ce{HClO4}), have very low pKa values (approximately -10), indicating nearly complete dissociation, while weak oxyacids like carbonic acid (\ceH2CO3\ce{H2CO3}) have higher pKa values, with the first dissociation constant at 6.35. Several factors govern the acidity of oxyacids. The electronegativity of the central atom plays a key role: higher electronegativity enhances the polarity of the O-H bond, weakening it and promoting proton release; for instance, acids with the same structure but more electronegative central atoms are stronger. Bond strength also influences acidity, as shorter, stronger bonds to oxygen stabilize the conjugate base less effectively. Furthermore, the oxidation state of the central atom affects strength—higher oxidation states increase acidity by drawing electron density away from the O-H bond through inductive effects, as seen in series like \ceH2SO3\ce{H2SO3} (weaker) versus \ceH2SO4\ce{H2SO4} (stronger). In addition to acidity, many oxyacids exhibit oxidation-reduction properties due to the variable oxidation states of their central atoms. These compounds often function as oxidizing agents, particularly when the central atom is in a high oxidation state, allowing reduction to lower states. For example, nitric acid (\ceHNO3\ce{HNO3}), with nitrogen at +5 oxidation state, oxidizes metals such as copper to form nitrates and nitrogen oxides. Sulfuric acid (\ceH2SO4\ce{H2SO4}), featuring sulfur at +6, acts as an oxidant in its concentrated form, dehydrating or oxidizing organic materials and metals. Certain oxyacids also display hydrolysis tendencies and can form polymeric structures via reactions, where is eliminated to link units. (\ceH3PO4\ce{H3PO4}), for instance, undergoes polymerization to yield linear chains of phosphate units, resulting in polyphosphoric acids used in various applications. Some oxyacids or their derived anions further exhibit amphoteric behavior, capable of acting as either acids or bases; hydrogen carbonate ion (\ceHCO3−\ce{HCO3^-}), from , exemplifies this by donating or accepting protons depending on solution .

Classification and Examples

Inorganic Oxyacids

Inorganic oxyacids encompass a diverse group of compounds where a central atom from non-carbon elements, such as , chalcogens, or pnictogens, is bonded to hydroxyl groups and oxygen, exhibiting acidic behavior upon . These acids play crucial roles in , , and , with their properties varying based on the central atom's and .

Halogen Oxyacids

The oxyacids, primarily derived from , illustrate a trend in stability and acidity that increases with the oxidation state of the . (HClO), with chlorine in the +1 oxidation state, is a weak acid (pKa ≈ 7.5) and highly unstable, readily decomposing to release ions used as disinfectants. (HClO₂, +3 state) is stronger (pKa ≈ 2.0) but still prone to , while (HClO₃, +5 state) exhibits greater stability and oxidizing power, often employed in explosives and bleaching agents. (HClO₄, +7 state) is the strongest and most stable, with a pKa < -10, serving as a powerful oxidant and in due to its non-coordinating anion. This progression in stability arises from enhanced delocalization of electrons in higher-oxidation-state species, reducing reactivity toward decomposition.

Sulfur Oxyacids

Sulfur oxyacids form a key family, with (H₂SO₄) being the most prominent due to its industrial significance. (H₂SO₃), existing mainly in aqueous solutions from SO₂ dissolution, is unstable and decomposes readily into and , acting as a weak diprotic acid (pKa₁ ≈ 1.9, pKa₂ ≈ 7.2) with reducing properties. In contrast, H₂SO₄ is a strong diprotic acid (pKa₁ < 0, pKa₂ ≈ 1.9), highly stable, and exhibits dehydrating action on carbohydrates and concentrated oxidizing behavior, essential for producing fertilizers, batteries, and dyes. (H₂S₂O₃), analogous to sulfuric acid but with one atom replaced by , is unstable and decomposes to and H₂SO₃, yet its salts (thiosulfates) are stable reducing agents used in and as antidotes for .

Nitrogen Oxyacids

Nitrogen oxyacids are vital in and explosives, with (HNO₂) and (HNO₃) as primary examples. HNO₂, a weak acid (pKa ≈ 3.3), is unstable in acidic conditions and decomposes to and , serving as a mild oxidant in organic nitrosation reactions. HNO₃, conversely, is a strong monoprotic acid (pKa ≈ -1.4) and potent oxidant capable of dissolving metals and nitrates, widely used in fertilizer production (e.g., ) and as a nitrating agent in explosives like TNT; its stability stems from the high of (+5).

Phosphorus Oxyacids

Phosphorus oxyacids are polyprotic and feature P-H bonds in lower-oxidation forms, influencing their reducing capabilities. (H₃PO₂), with phosphorus in the +1 state, is a monoprotic (pKa ≈ 1.2) and strong reductant due to its P-H bond, used in electroless plating and as an . (H₃PO₃, +3 state) is diprotic (pKa₁ ≈ 2.1, pKa₂ ≈ 7.2), with one ionizable P-OH group, exhibiting reducing properties and applications in inhibition. (H₃PO₄, +5 state) is a triprotic (pKa₁ ≈ 2.1, pKa₂ ≈ 7.2, pKa₃ ≈ 12.7), stable and non-reducing, forming buffer solutions and essential in fertilizers, detergents, and food additives like soft drinks.

Other Families

Carbonic acid (H₂CO₃), formed by CO₂ hydration, is a weak diprotic acid (pKa₁ ≈ 6.4, pKa₂ ≈ 10.3) central to biological regulation and in beverages, though it decomposes readily in solution. (H₄SiO₄) is a very weak acid (pKa ≈ 9.8) that polymerizes to form silica gels and contributes to diatom shells in aquatic environments. (H₃BO₃), a monoprotic weak acid (pKa ≈ 9.2), acts as a Lewis acid through boron-oxygen interactions, used in antiseptics, , and nuclear reactors as a absorber.

Organic Oxyacids

Organic oxyacids are a class of acids featuring a central carbon atom or carbon-based framework bonded to oxygen-containing functional groups that confer acidity, distinguishing them from inorganic oxyacids by their incorporation of organic substituents. These compounds play key roles in , , and biological processes due to their tunable properties and reactivity. The primary examples of organic oxyacids are carboxylic acids, which possess the general R−COOHR-COOH, where RR represents a or an organic group such as an alkyl or aryl chain. In this structure, the carboxyl group (−COOH-COOH) consists of a carbonyl (C=OC=O) bonded to a hydroxyl (−OH-OH) group, enabling proton donation from the acidic . A representative example is acetic acid (CH3COOHCH_3COOH), with a pKapK_a value of 4.76, indicating moderate acidity suitable for applications in buffers and esterifications. Sulfonic acids represent another important family, characterized by the formula R−SO3HR-SO_3H, where the sulfonyl group (−SO3H-SO_3H) imparts significantly greater acidity than carboxylic acids due to the electron-withdrawing effect of the sulfur-oxygen bonds. Methanesulfonic acid (CH3SO3HCH_3SO_3H) exemplifies this class, with a pKapK_a of approximately -1.9, making it a strong acid comparable to mineral acids and useful in catalysis and as a non-oxidizing alternative to sulfuric acid. Other notable organic oxyacids include phosphonic acids, with the general structure R−PO(OH)2R-PO(OH)_2, where a phosphorus atom is bonded to one organic RR group, a double-bonded oxygen, and two hydroxyl groups, facilitating applications in and flame retardants. An example is aminomethylphosphonic acid, which demonstrates the versatility of this motif in coordination chemistry. Sulfinic acids, denoted as R−SO2HR-SO_2H, feature a sulfur atom in the +4 bonded to RR, an oxygen, and a hydroxyl group; they are isoelectronic with carboxylic acids but less stable, often serving as intermediates in sulfur oxidation pathways. In contrast to many inorganic oxyacids, which can exhibit very strong acidity (e.g., pKa<0pK_a < 0), organic oxyacids are generally weaker, with acidity modulated by substituents on the RR group—electron-withdrawing groups enhance dissociation while electron-donating ones reduce it. This tunability is biologically significant, as groups in contribute to the zwitterionic nature of proteins, influencing folding, , and pH-dependent interactions in .

Preparation and Stability

Synthetic Methods

Oxyacids are commonly synthesized through oxidation reactions that increase the of the central atom in lower oxyanions. For instance, (H₂SO₄) is produced industrially via the , where (SO₂), derived from the combustion of or ores, is oxidized to (SO₃) using a vanadium pentoxide (V₂O₅) catalyst at elevated temperatures (400–500°C) and pressures. The SO₃ is then hydrated to form H₂SO₄: SO3+H2O→H2SO4\text{SO}_3 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_4 This method yields high-purity acid on a large scale, with global production of approximately 261 million tonnes as of 2024. Hydrolysis of acid halides or anhydrides provides a laboratory-scale route to several oxyacids by replacing halogen or anhydride linkages with hydroxyl groups. Phosphoric acid (H₃PO₄), an inorganic oxyacid, is prepared by the controlled hydrolysis of phosphorus pentachloride (PCl₅) with excess water: PCl5+4H2O→H3PO4+5HCl\text{PCl}_5 + 4\text{H}_2\text{O} \rightarrow \text{H}_3\text{PO}_4 + 5\text{HCl} This exothermic reaction requires careful temperature control to avoid side products like phosphorous acid. Similarly, the hydration of SO₃, an anhydride, directly yields H₂SO₄, as noted above, and is integral to the final absorption step in the contact process. Electrochemical oxidation enables the synthesis of highly oxidized oxyacids from precursors. (HClO₄), the strongest of the chlorine oxyacids, is produced by anodic oxidation of (HCl) or solutions in electrolytic cells, stepwise forming , , and ions before acidification: Cl−→ClO−→ClO3−→ClO4−\text{Cl}^- \rightarrow \text{ClO}^- \rightarrow \text{ClO}_3^- \rightarrow \text{ClO}_4^- or anodes are typically used, with current efficiencies up to 90% at 50–70°C, though the process is energy-intensive and suited for high-purity needs. On an industrial scale, nitric acid (HNO₃), a key inorganic oxyacid, is synthesized via the Ostwald process, which couples the Haber-Bosch ammonia synthesis with catalytic oxidation. Ammonia (NH₃) is oxidized over a platinum-rhodium gauze catalyst at 800–900°C to nitric oxide (NO), followed by air oxidation to nitrogen dioxide (NO₂) and absorption in water: 4NH3+5O2→4NO+6H2O,2NO+O2→2NO2,3NO2+H2O→2HNO3+NO4\text{NH}_3 + 5\text{O}_2 \rightarrow 4\text{NO} + 6\text{H}_2\text{O}, \quad 2\text{NO} + \text{O}_2 \rightarrow 2\text{NO}_2, \quad 3\text{NO}_2 + \text{H}_2\text{O} \rightarrow 2\text{HNO}_3 + \text{NO} This process accounts for the majority of global HNO₃ production, approximately 58 million tonnes as of 2024, primarily for fertilizers.

Stability and Decomposition

The thermal stability of oxyacids tends to increase with the of the central atom, as more electronegative atoms form stronger bonds with oxygen, reducing the tendency for bond cleavage. For of oxyacids sharing the same central atom, stability also rises with the of that atom, owing to higher bond orders that resist decomposition. For instance, among the oxyacids of , (HClO, Cl in +1 ) is highly unstable and decomposes readily via the reaction 2HClO→2HCl+O22 \text{HClO} \rightarrow 2 \text{HCl} + \text{O}_2, often catalyzed by or metallic impurities. In contrast, (HClO₄, Cl in +7 ) exhibits greater thermal stability than lower oxidation state analogs; aqueous solutions up to 70% concentration are stable at , though heating concentrated solutions requires caution due to potential decomposition. Decomposition reactions of oxyacids typically involve the release of , s, or lower-oxidation-state species, driven by thermodynamic favorability. (H₂SO₃) decomposes upon heating to yield and : H2SO3→H2O+SO2\text{H}_2\text{SO}_3 \rightarrow \text{H}_2\text{O} + \text{SO}_2. Similarly, (HNO₂) undergoes : 3HNO2→HNO3+2NO+H2O3 \text{HNO}_2 \rightarrow \text{HNO}_3 + 2 \text{NO} + \text{H}_2\text{O}, producing and . These processes highlight how lower-oxidation-state oxyacids are more prone to of oxygen or ligands. Several factors influence oxyacid stability, including , temperature, and the presence of catalysts. Elevated temperatures accelerate rates following Arrhenius kinetics, while acidic conditions can either stabilize or hasten breakdown depending on the specific acid— for example, HNO₂ is faster at low due to effects. Catalysts such as transition metals or light further promote instability by lowering activation energies for bond breaking. offers a means to enhance stability; in the case of phosphoric acids, formation of polyphosphoric acids through increases thermal resistance, enabling applications at higher temperatures without rapid degradation. Safety considerations are paramount for certain oxyacids, particularly peroxoacids like peroxomonosulfuric acid (H₂SO₅), which exhibit explosive decomposition under , shock, or , releasing oxygen and generating rapid buildup. These compounds demand strict handling protocols, including storage below critical temperatures and avoidance of initiators, to mitigate risks of .

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