any oxygen-containing acid. Most covalent nonmetallic oxides (oxide) 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.

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 (nitrate), 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,

Commercially, nitric acid is produced by the Ostwald process. This process involves oxidation of ammonia, NH₃, to nitric oxide, NO, further oxidation of the NO to nitrogen dioxide, NO₂, and then conversion of the NO₂ to nitric acid (HNO₃). This is a flow process in which a mixture of ammonia and excess air is heated to 600 to 700 °C (1,100 to 1,300 °F) and passed through a platinum-rhodium catalyst. (A catalyst increases the rate of a reaction without itself being consumed in the reaction.) As the oxidation to NO occurs, this gaseous mixture literally burns with a flame. Additional air is added to oxidize the NO to NO₂. The NO₂, excess oxygen, and the unreactive nitrogen from the air are passed through a water spray, where HNO₃ and NO form as the NO₂ disproportionates. The gaseous NO is recycled through the process with more air, and the liquid HNO₃ is drawn off and concentrated. About 7 billion kg (16 billion pounds) of HNO₃ are produced commercially in the United States each year, with the bulk of it made by the Ostwald process.

When pure, nitric acid is a colourless liquid that boils at 86 °C (187 °F) and freezes at −42 °C (−44 °F). Upon being exposed to light or heat, it decomposes to produce oxygen, water, and a mixture of nitrogen oxides (primarily NO₂).

The action of nitric acid on a metal usually results in reduction of the acid (i.e., a decrease in the oxidation state of the nitrogen). The products of the reaction are determined by the concentration of HNO₃, the metal involved (i.e., its reactivity), and the temperature. In most cases, a mixture of nitrogen oxides, nitrates, and other reduction products is formed. Relatively unreactive metals such as copper (Cu), silver (Ag), and lead (Pb) reduce concentrated HNO₃ primarily to NO₂. The reaction of dilute HNO₃ with copper produces NO, whereas more reactive metals, such as zinc (Zn) and iron (Fe), react with dilute HNO₃ to yield N₂O. When extremely dilute HNO₃ is used, either nitrogen gas (N₂) or the ammonium ion (NH₄⁺) may be formed. Nitric acid reacts with proteins (protein), such as those in human skin, to produce a yellow material called xanthoprotein.

Nitrates, which are salts of nitric acid, are produced when metals or their oxides, hydroxides (hydroxide), or carbonates (carbonate) react with nitric acid. Most nitrates are soluble in water, and a major use of nitric acid is to produce soluble metal nitrates. All nitrates decompose when heated and may do so explosively. For example, when potassium nitrate (KNO₃) is heated, a nitrite (a compound containing NO₂⁻) is formed, and oxygen gas is evolved.

Nitric acid is heavily used in the laboratory and in chemical industries as a strong acid and as an oxidizing agent. The manufacture of explosives (explosive), dyes (dye), plastics (plastic), and drugs (drug) makes extensive use of the acid. Nitrates are valuable as fertilizers (fertilizer). gunpowder is a mixture of potassium nitrate, sulfur, and charcoal. Ammonal, an explosive, is a mixture of ammonium nitrate and aluminum powder.

In general, the salts of all oxyacids are more stable than the acids themselves; such is the case with nitrites. They are much more stable than nitrous acid. Most nitrites are soluble in water and in concentrated forms, like nitrates, can explode upon heating or detonation.

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.

Orthophosphoric acid is a triprotic acid—i.e., it can donate all three of its hydrogen atoms as protons (proton) 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.

Pure phosphorous acid, H₃PO₃, is best prepared by hydrolysis of phosphorus trichloride, PCl₃.

Sulfuric acid is sometimes referred to as the “king of chemicals” because it is produced worldwide in such large quantities. In fact, per capita use of sulfuric acid has been taken as one index of the technical development of a country. Annual production in the United States, which is the world's leading producer, is well over 39 billion kg (86 billion pounds). It is the cheapest bulk acid.

Pure sulfuric acid is a colourless, oily, dense (1.83 grams per cc) liquid that freezes at 10.5 °C (50.9 °F). It fumes when heated because of its decomposition to water and sulfur trioxide. Because SO₃ has a lower boiling point than water, more SO₃ is lost during heating. When a concentration of 98.33 percent acid is reached, the solution boils at 338 °C without any further change in concentration. This is called a constant boiling solution, and it is this concentration that is sold as concentrated sulfuric acid. Anhydrous sulfuric acid mixes with water in all proportions in a very exothermic reaction. Adding water to concentrated acid can cause explosive spattering. Because it reacts with organic compounds in the skin, concentrated sulfuric acid can cause severe burns (burn). Thus, to decrease the risk of injury in the laboratory, sulfuric acid should always be added to water slowly and with stirring to distribute the heat.

Sulfuric acid has its two hydrogen atoms bonded to oxygen, ionizes in two stages, and is a strong diprotic acid. In aqueous solution, loss of the first hydrogen (as a hydrogen ion, H⁺) is essentially 100 percent. The second ionization takes place to an extent of about 25 percent, but HSO₄⁻ is nonetheless considered a moderately strong acid. Because it is a diprotic acid, H₂SO₄ forms two series of salts: hydrogen sulfates, HSO₄⁻, and sulfates, SO₄²⁻. The sulfates of the alkaline-earth metals (alkaline-earth metal)— calcium (Ca), strontium (Sr), and barium (Ba)—as well as that of lead (Pb) are virtually insoluble, and these salts are found as naturally occurring minerals. These important minerals include gypsum (CaSO₄ · 2H₂O), celestine (SrSO₄), barite (BaSO₄), and anglesite (PbSO₄). These insoluble salts can be prepared in the laboratory by metathesis reactions. A metathesis reaction is one in which compounds exchange anion- cation partners. For example, if a solution of barium nitrate, Ba(NO₃)₂, is added to a solution of sodium sulfate, Na₂SO₄, a precipitation of barium sulfate, BaSO₄, occurs. This is an important reaction because it can be used as both a qualitative and quantitative test for the sulfate ion and the barium ion. (Qualitative tests are used to determine the presence or absence of a substance, whereas quantitative tests are used to measure the amount of a constituent.) In addition to metathesis reactions, sulfate salts can generally be prepared by dissolution of metals in aqueous H₂SO₄, neutralization of aqueous H₂SO₄ with metal oxides (oxide) or hydroxides (hydroxide), oxidation of metal sulfides (a sulfide contains S²⁻) or sulfites (SO₃²⁻), or decomposition of salts of volatile acids, such as carbonates, with aqueous H₂SO₄. Some important soluble sulfate salts are Glauber's salt, Na₂SO₄ · 10H₂O; Epsom salt, MgSO₄ · 7H₂O; blue vitriol, CuSO₄ · 5H₂O; green vitriol, FeSO₄ · 7H₂O; and white vitriol, ZnSO₄ · 7H₂O.

Pure H₂SO₄ undergoes extensive self- ionization (sometimes called autoprotolysis).

Concentrated sulfuric acid is not a very strong oxidizing agent unless it is hot. When it acts as an oxidizing agent, however, it can be reduced to several different sulfur species, including SO₂, HSO₃⁻, SO₃²⁻, elemental sulfur (S₈), hydrogen sulfide (H₂S), and the sulfide anion (S²⁻). Concentrated sulfuric acid is a good dehydrating agent, as it reacts with many organic materials to remove the elements of water.

The amount of sulfuric acid used in industry exceeds that of any other manufactured compound. In the United States approximately 67 percent of the acid is utilized to convert phosphate rock to phosphoric acid. The phosphoric acid is then converted to phosphate fertilizers (fertilizer). Other major uses include the refining of petroleum, the removal of impurities from gasoline and kerosene, the pickling of steel (the cleaning of its surface), and the manufacture of other chemicals, such as nitric and hydrochloric acids. It also is utilized in lead storage batteries (battery) and in the production of paints (paint), plastics (plastic), explosives (explosive), and textiles (textile).

When sulfur dioxide is dissolved in water, an acidic solution results. This has long been loosely called a sulfurous acid, H₂SO₃, solution. However, pure anhydrous sulfurous acid has never been isolated or detected, and an aqueous solution of SO₂ contains little, if any, H₂SO₃. Studies of these solutions indicate that the predominant species are hydrated SO₂ molecules, SO₂ · nH₂O. The ions present in these solutions are dependent on concentration, temperature, and pH and include H₃O⁺, HSO₃⁻, S₂O₅²⁻, and perhaps SO₃²⁻. However, “sulfurous acid” has two acid dissociation constants. It acts as a moderately strong acid with an apparent ionization of about 25 percent in the first stage and much less in the second stage. These ionizations produce two series of salts—sulfites, containing SO₃²⁻, and hydrogen sulfites, containing HSO₃⁻. Only with large cations, such as Rb⁺ ( rubidium) or Cs⁺ ( cesium), have solid HSO₃⁻ salts been isolated. Attempts to isolate these salts with smaller cations tend to yield disulfites as a product of dehydration.

With the exception of the alkali metal sulfites, these salts are relatively insoluble. The HSO₃⁻ ion has an interesting structure in that the hydrogen atom is bonded to the sulfur atom and not to the oxygen atom, as might be expected. There is some suggestion that in solution both the sulfur-hydrogen and oxygen-hydrogen structures may exist in equilibrium with one another, but there is no concrete evidence for this phenomenon. Heating solid hydrogen sulfite salts (shown by the equation above) or passing gaseous sulfur dioxide into their aqueous solutions produces disulfites.

Carbonic acid (H₂CO₃) is formed in small amounts when its anhydride, carbon dioxide (CO₂), dissolves in water.

These salts can be prepared by the reaction of carbon dioxide with metal oxides and metal hydroxides, respectively.

Carbonates are moderately strong bases. Aqueous solutions are basic because the carbonate anion can accept a hydrogen ion from water.

An interesting use of lithium carbonate, Li₂CO₃, stems from the discovery that small doses of the salt, orally administered, are an effective treatment for manic-depressive psychoses (psychosis). It is not entirely understood how this treatment works, but it is almost certainly related to the effect of the Li⁺ ion on the Na⁺:K⁺ or the Mg²⁺:Ca²⁺ balance in the brain.

When water containing carbon dioxide comes in contact with limestone rocks, the rocks dissolve because Ca(HCO₃)₂, a water-soluble compound that forms aqueous Ca²⁺ and HCO₃⁻ ions, is formed. Thus, this reaction is responsible for the formation of the caves (cave) that are often found in limestone rock beds. On the other hand, if water containing Ca(HCO₃)₂ liberates carbon dioxide, calcium carbonate may again be deposited.

Two other carbon-containing acids are sometimes referred to as carbonic acids. formic acid (HCOOH) is the acid that formally has carbon monoxide (CO) as its acid anhydride. This acid has a low solubility in water. As noted previously, carbon suboxide, C₃O₂, is the acid anhydride of malonic acid, CH₂(COOH)₂, which is considered by some to be a carbonic acid.