Oxide

Oxides are extraordinarily diverse in terms of stoichiometries (the measurable relationship between reactants and chemical equations of an equation or reaction) and in terms of the structures of each stoichiometry. Most elements form oxides of more than one stoichiometry. A well known example is carbon monoxide and carbon dioxide. This applies to binary oxides, that is, compounds containing only oxide and another element. Far more common than binary oxides are oxides of more complex stoichiometries. Such complexity can arise by the introduction of other cations (a positively charged ion, i.e. one that would be attracted to the cathode in electrolysis) or other anions (a negatively charged ion). Iron silicate, Fe2SiO4, the mineral fayalite, is one of many examples of a ternary oxide. For many metal oxides, the possibilities of polymorphism and nonstoichiometry exist as well. The commercially important dioxides of titanium exists in three distinct structures, for example. Many metal oxides exist in various nonstoichiometric states. Many molecular oxides exist with diverse ligands as well.{{cite journal |doi=10.1021/cr020376q|title=Organometallic Oxides of Main Group and Transition Elements Downsizing Inorganic Solids to Small Molecular Fragments|first1=Herbert W. |last1=Roesky|first2=Ionel |last2=Haiduc|first3=Narayan S. |last3=Hosmane For simplicity's sake, most of this article focuses on binary oxides. ==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: : The production of metals from ores often involves the production of oxides by roasting (heating) metal sulfide minerals in air. In this way, (molybdenite) is converted to molybdenum trioxide, the precursor to virtually all molybdenum compounds: : Noble metals (such as gold and platinum) are prized because they resist direct chemical combination with oxygen. 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: : : Finally the trioxide is converted to sulfuric acid by a hydration reaction: : ==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 File:Carbon-dioxide-2D-dimensions.svg|Carbon dioxide is the main product of fossil fuel combustion. File:Carbon monoxide 2D.svg|Carbon monoxide is the product of the incomplete combustion of carbon-based fuels and a precursor to many useful chemicals. File:Nitrogen-dioxide-2D-dimensions-vector.svg|Nitrogen dioxide is a problematic pollutant from internal combustion engines. File:Sulfur-dioxide-2D.svg|Sulfur dioxide, the principal oxide of sulfur, is emitted from volcanoes. File:Nitrous-oxide-2D-dimensions.png|Nitrous oxide ("laughing gas") is a potent greenhouse gas produced by soil bacteria. 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: : 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: 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. Oxycations are somewhat rare, one examples being nitrosonium (). Some oxycations, e.g. vanadyl, are occasionally represented with abbreviated formula such as . is in fact an aquo complex . Similarly uranyl () refers to hydrated cations. Related to the oxycations are the oxyhalides, e.g. vanadium oxytrichloride (). ==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. == See also ==