Oxide - Formation

Formation

Due to its electronegativity, oxygen forms stable chemical bonds with almost all elements to give the corresponding oxides. Noble metals (such as gold or platinum) are prized because they resist direct chemical combination with oxygen, and substances like gold(III) oxide must be generated by indirect routes.

Two independent pathways for corrosion of elements are hydrolysis and oxidation by oxygen. The combination of water and oxygen is even more corrosive. Virtually all elements burn in an atmosphere of oxygen, or an oxygen rich environment. In the presence of water and oxygen (or simply air), some elements— sodium—react rapidly, even dangerously, to give the hydroxides. In part for this reason, alkali and alkaline earth metals are not found in nature in their metallic, i.e., native, form. Caesium is so reactive with oxygen that it is used as a getter in vacuum tubes, and solutions of potassium and sodium, so called NaK are used to deoxygenate and dehydrate some organic solvents. The surface of most metals consists of oxides and hydroxides in the presence of air. A well known example is aluminium foil, which is coated with a thin film of aluminium oxide that passivates the metal, slowing further corrosion. The aluminium oxide layer can be built to greater thickness by the process of electrolytic anodising. Though solid magnesium and aluminium react slowly with oxygen at STP—they, like most metals, burn in air, generating very high temperatures. Finely grained powders of most metals can be dangerously explosive in air. Consequently, they are often used in Solid-fuel rockets.

In dry oxygen, iron readily forms iron(II) oxide, but the formation of the hydrated ferric oxides, Fe₂O_3−x(OH)_2x, that mainly comprise rust, typically requires oxygen and water. Free oxygen production by photosynthetic bacteria some 3.5 billion years ago precipitated iron out of solution in the oceans as Fe₂O₃ in the economically important iron ore hematite.