Chemistry » Metals, Metalloids, and Nonmetals » Occurrence, Preparation, and Compounds of Oxygen

# Nonmetal Oxyacids and Their Salts

## Nonmetal Oxyacids and Their Salts

Nonmetal oxides form acids when allowed to react with water; these are acid anhydrides. The resulting oxyanions can form salts with various metal ions.

## Nitrogen Oxyacids and Salts

Nitrogen pentaoxide, N2O5, and NO2 react with water to form nitric acid, HNO3. Alchemists, as early as the eighth century, knew nitric acid (shown in the figure below) as aqua fortis (meaning “strong water”). The acid was useful in the separation of gold from silver because it dissolves silver but not gold. Traces of nitric acid occur in the atmosphere after thunderstorms, and its salts are widely distributed in nature. There are tremendous deposits of Chile saltpeter, NaNO3, in the desert region near the boundary of Chile and Peru. Bengal saltpeter, KNO3, occurs in India and in other countries of the Far East.

This image shows the molecular structure (left) of nitric acid, HNO3 and its resonance forms (right).

In the laboratory, it is possible to produce nitric acid by heating a nitrate salt (such as sodium or potassium nitrate) with concentrated sulfuric acid:

$${\text{NaNO}}_{3}(s)+{\text{H}}_{2}{\text{SO}}_{4}(l)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;{\text{NaHSO}}_{4}(s)+{\text{HNO}}_{3}(g)$$

The Ostwald process is the commercial method for producing nitric acid. This process involves the oxidation of ammonia to nitric oxide, NO; oxidation of nitric oxide to nitrogen dioxide, NO2; and further oxidation and hydration of nitrogen dioxide to form nitric acid:

$$4{\text{NH}}_{3}(g)+5{\text{O}}_{2}(g)\;⟶\;\text{4NO}(g)+6{\text{H}}_{2}\text{O}(g)$$

$$\text{2NO}(g)+{\text{O}}_{2}(g)\;⟶\;2{\text{NO}}_{2}(g)$$

$$3{\text{NO}}_{2}(g)+{\text{H}}_{2}\text{O}(l)\;⟶\;2{\text{HNO}}_{3}(aq)+\text{NO}(g)$$

Or

$$4{\text{NO}}_{2}(g)+{\text{O}}_{2}(g)+2{\text{H}}_{2}\text{O}(g)\;⟶\;4{\text{HNO}}_{3}(l)$$

Pure nitric acid is a colorless liquid. However, it is often yellow or brown in color because NO2 forms as the acid decomposes. Nitric acid is stable in aqueous solution; solutions containing 68% of the acid are commercially available concentrated nitric acid. It is both a strong oxidizing agent and a strong acid.

The action of nitric acid on a metal rarely produces H2 (by reduction of H+) in more than small amounts. Instead, the reduction of nitrogen occurs. The products formed depend on the concentration of the acid, the activity of the metal, and the temperature.

Normally, a mixture of nitrates, nitrogen oxides, and various reduction products form. Less active metals such as copper, silver, and lead reduce concentrated nitric acid primarily to nitrogen dioxide. The reaction of dilute nitric acid with copper produces NO. In each case, the nitrate salts of the metals crystallize upon evaporation of the resultant solutions.

Nonmetallic elements, such as sulfur, carbon, iodine, and phosphorus, undergo oxidation by concentrated nitric acid to their oxides or oxyacids, with the formation of NO2:

$$\text{S}(s)+6{\text{HNO}}_{3}(aq)\;⟶\;{\text{H}}_{2}{\text{SO}}_{4}(aq)+6{\text{NO}}_{2}(g)+2{\text{H}}_{2}\text{O}(l)$$

$$\text{C}(s)+4{\text{HNO}}_{3}(aq)\;⟶\;{\text{CO}}_{2}(g)+4{\text{NO}}_{2}(g)+2{\text{H}}_{2}\text{O}(l)$$

Nitric acid oxidizes many compounds; for example, concentrated nitric acid readily oxidizes hydrochloric acid to chlorine and chlorine dioxide. A mixture of one part concentrated nitric acid and three parts concentrated hydrochloric acid (called aqua regia, which means royal water) reacts vigorously with metals. This mixture is particularly useful in dissolving gold, platinum, and other metals that are more difficult to oxidize than hydrogen. A simplified equation to represent the action of aqua regia on gold is:

$$\text{Au}(s)+\text{4HCl}(aq)+3{\text{HNO}}_{3}(aq)\;⟶\;{\text{HAuCl}}_{4}(aq)+3{\text{NO}}_{2}(g)+3{\text{H}}_{2}\text{O}(l)$$

### Optional Video:

Although gold is generally unreactive, you can watch this video of the complex mixture of compounds present in aqua regia dissolving it into solution.

Nitrates, salts of nitric acid, form when metals, oxides, hydroxides, or carbonates react with nitric acid. Most nitrates are soluble in water; indeed, one of the significant uses of nitric acid is to prepare soluble metal nitrates.

Nitric acid finds extensive use in the laboratory and in chemical industries as a strong acid and strong oxidizing agent. It is important in the manufacture of explosives, dyes, plastics, and drugs. Salts of nitric acid (nitrates) are valuable as fertilizers. Gunpowder is a mixture of potassium nitrate, sulfur, and charcoal.

The reaction of N2O3 with water gives a pale blue solution of nitrous acid, HNO2. However, HNO2 (shown in the figure below) is easier to prepare by the addition of an acid to a solution of nitrite; nitrous acid is a weak acid, so the nitrite ion is basic in aqueous solution:

$${\text{NO}}_{2}{}^{\text{−}}(aq)+{\text{H}}_{3}{\text{O}}^{\text{+}}(aq)\;⟶\;{\text{HNO}}_{2}(aq)+{\text{H}}_{2}\text{O}(l)$$

Nitrous acid is very unstable and exists only in solution. It disproportionates slowly at room temperature (rapidly when heated) into nitric acid and nitric oxide. Nitrous acid is an active oxidizing agent with strong reducing agents, and strong oxidizing agents oxidize it to nitric acid.

This image shows the molecular structure of a molecule of nitrous acid, HNO2.

Sodium nitrite, NaNO2, is an additive to meats such as hot dogs and cold cuts. The nitrite ion has two functions. It limits the growth of bacteria that can cause food poisoning, and it prolongs the meat’s retention of its red color. The addition of sodium nitrite to meat products is controversial because nitrous acid reacts with certain organic compounds to form a class of compounds known as nitrosamines. Nitrosamines produce cancer in laboratory animals. This has prompted the FDA to limit the amount of NaNO2 in foods.

The nitrites are much more stable than the acid, but nitrites, like nitrates, can explode. Nitrites, like nitrates, are also soluble in water (AgNO2 is only slightly soluble).