The unit cell of rutile, an important oxide of titanium. Ti(IV) centers are grey; oxygen centers are red. Notice that oxygen forms three bonds to titanium and titanium forms six bonds to oxygen.
The unit cell of rutile, an important oxide of titanium. Ti(IV) centers are grey; oxygen centers are red. Notice that oxygen forms three bonds to titanium and titanium forms six bonds to oxygen.

An oxide (/ˈɒksd/) is a chemical compound that contains at least one oxygen atom and one other element[1] in its chemical formula. "Oxide" itself is the dianion of oxygen, an O2– (molecular) ion. with oxygen in the oxidation state of −2. Most of the Earth's crust consists of oxides. Even materials considered pure elements often develop an oxide coating. For example, aluminium foil develops a thin skin of Al2O3 (called a passivation layer) that protects the foil from further corrosion.[2]

Stoichiometry

The above fact 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 or other anions. 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.[3] 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.[4]

For simplicity 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:[2]

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:[5]

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 molybdenum trioxide, the precursor to virtually all molybdenum compounds:[6]

2 MoS2 + 7 O2 → 2MoO3 + 4 SO2

Noble metals (such as gold and platinum) are prized because they resist direct chemical combination with oxygen.[2]

NiS + 1.5 O2 → NiO + SO2

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 + 1.5 O2 → CO + 2 H2O
C + 0.5 O2 → 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
NO + 0.5 O2 → NO2

These reactions are practiced in the production of nitric acid, a commodity chemical.[7]

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:[8]

S + O2 → SO2
SO2 + 0.5 O2 → 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.[9]

Molecular oxides

Although most metal oxides are crystalline solids, some 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.[2]

Reactions

Reduction

See also: Carbothermic 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, (see Thermal decomposition). For example, silver oxide decomposes at 200 °C:[10]

2 Ag2O → 4 Ag + O2

Most often, however, metals 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]

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.[11]

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.[2]

Dissolution of oxides often gives oxyanions. Adding aqueous base to P4O10 gives various phosphates. Adding aqueous base to MoO3 gives polyoxometalates. Oxycations are rarer, some examples being nitrosonium (NO+), vanadyl (VO2+), and uranyl (UO2+
2
). Of course many compounds are known with both oxides and other groups. In organic chemistry, these include ketones and many related carbonyl compounds. For the transition metals, many oxo complexes are known as well as oxyhalides.[2]

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.[12]

See also

References

  1. ^ Hein, Morris; Arena, Susan (2006). Foundations of College Chemistry (12th ed.). Wiley. ISBN 9780471741534.
  2. ^ a b c d e f g h Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
  3. ^ C. N. R. Rao, B. Raveau (1995). Transition Metal Oxides. New York: VCH. ISBN 1-56081-647-3.
  4. ^ Herbert W. Roesky, Ionel Haiduc, and Narayan S. Hosmane (2003). "Organometallic Oxides of Main Group and Transition Elements Downsizing Inorganic Solids to Small Molecular Fragments". Chem. Rev. 103: 2579–2596. doi:10.1021/cr020376q.((cite journal)): CS1 maint: multiple names: authors list (link)
  5. ^ Graf, Günter G. (2000). "Zinc". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a28_509. ISBN 3527306730.
  6. ^ Roger F. Sebenik et al. (2005). "Molybdenum and Molybdenum Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a16_655.((cite encyclopedia)): CS1 maint: uses authors parameter (link)
  7. ^ Thiemann, Michael; Scheibler, Erich; Wiegand, Karl Wilhelm. "Nitric Acid, Nitrous Acid, and Nitrogen Oxides". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a17_293.
  8. ^ Müller, Hermann (2000). "Sulfuric Acid and Sulfur Trioxide". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_635.
  9. ^ P.A. Cox (2010). Transition Metal Oxides. An Introduction to Their Electronic Structure and Properties. Oxford University Press. ISBN 9780199588947.((cite book)): CS1 maint: uses authors parameter (link)
  10. ^ "Silver oxide".
  11. ^ Cornell, R. M.; Schwertmann, U. (2003). The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses, Second Edition. p. 323. doi:10.1002/3527602097. ISBN 9783527302741.
  12. ^ Schultz, Emeric (2005). "Fully Exploiting the Potential of the Periodic Table through Pattern Recognition". J. Chem. Educ. 82 (11): 1649. Bibcode:2005JChEd..82.1649S. doi:10.1021/ed082p1649.