Oxide

Formation

Oxides are associated with all elements except a few noble gases. The pathways for the formation of this diverse family of compounds are correspondingly numerous.

Metal oxides

Many metal oxides arise by decomposition of other metal compounds, e.g. carbonates, hydroxides, and nitrates. In the making of calcium oxide, calcium carbonate (limestone) breaks down upon heating, releasing carbon dioxide:

CaCO₃ → CaO + CO₂

The reaction of elements with oxygen in air is a key step in corrosion relevant to the commercial use of iron especially. Almost all elements form oxides upon heating with oxygen atmosphere. For example, zinc powder will burn in air to give zinc oxide:

2 Zn + O₂ → 2 ZnO

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:

2 MoS₂ + 7 O₂ → 2 MoO₃ + 4 SO₂

Noble metals (such as gold and platinum) are prized because they resist direct chemical combination with oxygen.

Non-metal oxides

Important and prevalent nonmetal oxides are carbon dioxide and carbon monoxide. These species form upon full or partial oxidation of carbon or hydrocarbons. With a deficiency of oxygen, the monoxide is produced:

2 CH₄ + 3 O₂ → 2 CO + 4 H₂O

2 C + O₂ → 2 CO

With excess oxygen, the dioxide is the product, the pathway proceeds by the intermediacy of carbon monoxide:

CH₄ + 2 O₂ → CO₂ + 2 H₂O

C + O₂ → CO₂

Elemental nitrogen (N₂) is difficult to convert to oxides, but the combustion of ammonia gives nitric oxide, which further reacts with oxygen:

4 NH₃ + 5 O₂ → 4 NO + 6 H₂O

2 NO + O₂ → 2 NO₂

These reactions are practiced in the production of nitric acid, a commodity chemical.

The chemical produced on the largest scale industrially is sulfuric acid. It is produced by the oxidation of sulfur to sulfur dioxide, which is separately oxidized to sulfur trioxide:

S + O₂ → SO₂

2 SO₂ + O₂ → 2 SO₃

Finally the trioxide is converted to sulfuric acid by a hydration reaction:

SO₃ + H₂O → H₂SO₄

Structure

Oxides have a range of structures, from individual molecules to polymeric and crystalline structures. At standard conditions, oxides may range from solids to gases. Solid oxides of metals usually have polymeric structures at ambient conditions.

Molecular oxides

Although most metal oxides are crystalline solids, many non-metal oxides are molecules. Examples of molecular oxides are carbon dioxide and carbon monoxide. All simple oxides of nitrogen are molecular, e.g., NO, N₂O, NO₂ and N₂O₄. Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the real formula being P₄O₁₀. Tetroxides are rare, with a few more common examples being ruthenium tetroxide, osmium tetroxide, and xenon tetroxide.

Reactions

Reduction

Reduction of metal oxide to the metal is practiced on a large scale in the production of some metals. Many metal oxides convert to metals simply by heating (thermal decomposition). For example, silver oxide decomposes at 200 °C:

2 Ag₂O → 4 Ag + O₂

Most often, however, metal oxides are reduced by a chemical reagent. A common and cheap reducing agent is carbon in the form of coke. The most prominent example is that of iron ore smelting. Many reactions are involved, but the simplified equation is usually shown as:

2 Fe₂O₃ + 3 C → 4 Fe + 3 CO₂

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.

Hydrolysis and dissolution

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.

Dissolution of oxides often gives oxyanions. Adding aqueous base to P₄O₁₀ gives various phosphates. Adding aqueous base to MoO₃ gives polyoxometalates. Similarly, Metal peroxide compounds, e.g. [K₄Ti(O₂)₄], arise by reaction of the metal oxides with a alkaline solution of hydrogen peroxide.

Oxycations are somewhat rare, one examples being nitrosonium (NO⁺). Some oxycations, e.g. vanadyl, are occasionally represented with abbreviated formula such as VO²⁺. VO²⁺ is in fact an aquo complex [VO(H₂O)₅]²⁺. Similarly uranyl (UO2+2) refers to hydrated cations. Related to the oxycations are the oxyhalides, e.g. vanadium oxytrichloride (VOCl₃).

Nomenclature and formulas

The chemical formulas of the oxides of the chemical elements in their highest oxidation state are predictable and are derived from the number of valence electrons for that element. Even the chemical formula of O₄, tetraoxygen, is predictable as a group 16 element. One exception is copper, for which the highest oxidation state oxide is copper(II) oxide and not copper(I) oxide. Another exception is fluoride, which does not exist as one might expect—as F₂O₇—but as OF₂.