Oxide

Even materials considered pure elements often develop an oxide coating.

[3] The commercially important dioxides of titanium exists in three distinct structures, for example.

In the making of calcium oxide, calcium carbonate (limestone) breaks down upon heating, releasing carbon dioxide:[2] The reaction of elements with oxygen in air is a key step in corrosion relevant to the commercial use of iron especially.

For example, zinc powder will burn in air to give zinc oxide:[5] The production of metals from ores often involves the production of oxides by roasting (heating) metal sulfide minerals in air.

In this way, MoS2 (molybdenite) is converted to molybdenum trioxide, the precursor to virtually all molybdenum compounds:[6] Noble metals (such as gold and platinum) are prized because they resist direct chemical combination with oxygen.

With a deficiency of oxygen, the monoxide is produced:[2] With excess oxygen, the dioxide is the product, the pathway proceeds by the intermediacy of carbon monoxide: Elemental nitrogen (N2) is difficult to convert to oxides, but the combustion of ammonia gives nitric oxide, which further reacts with oxygen: These reactions are practiced in the production of nitric acid, a commodity chemical.

Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the real formula being P4O10.

Many reactions are involved, but the simplified equation is usually shown as:[2] Some metal oxides dissolve in the presence of reducing agents, which can include organic compounds.

Reductive dissolution of ferric oxides is integral to geochemical phenomena such as the iron cycle.

[11] Because the M-O bonds are typically strong, metal oxides tend to be insoluble in solvents, though they may be attacked by aqueous acids and bases.

In organic chemistry, these include ketones and many related carbonyl compounds.