Oxide

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An oxide is a chemical compound of oxygen with other chemical elements. In the 18th century, oxides were named calxes or calces after the calcination process used to produce oxides. Calx was later replaced by oxyd.

See Category:Oxides for a list of oxides.

Contents 1 Current naming 2 Chemical properties 3 Types of oxides 4 Common oxides sorted by oxidation state 5 See also 6 References

Current naming

Oxides can be named after the amount of oxygen atoms in the oxide. Oxides containing only one oxygen are called oxide or monoxide, those containing two oxygen atoms dioxide, three trioxide, four tetroxide, and so on following the Greek numerical prefixes.

There are two other types of oxide—peroxide and superoxide. Both count as oxides but have different oxidation states and react in different ways compared to oxides.

Chemical properties

Oxides are formed in redox reactions through oxidation in which a reducing agent is allowed to react with molecular oxygen (O₂) or oxidizing agents which contain oxygen, such as hydrogen peroxide (H₂O₂) and permanganate (MnO₄⁻). Oxides are characterized by a redistribution of electrons, in which the oxygen atoms have a net surplus of electrons and the other atoms a net lack. In oxides of hydrogen, carbon, nitrogen, sulfur, phosphorus and halogens, covalent bonds occur between oxygen and the other elements. Generally, these are gases or fluids at room temperature. Oxides of metals occur as ionic compounds, or salts, which are solid at room temperature. Oxide salts are generally insoluble in water, though some react with it.

Generally, oxides are not conductive to electricity. This property is most commonly taken advantage of with silicon dioxide, as silicon can easily be oxidized and the resulting part can be made into a transistor. This is the basis for much of modern computer technology.

Types of oxides

Oxides of more electropositive elements tend to be basic. They are called basic anhydrides; adding water, they may form basic hydroxides. For example, sodium oxide is basic; when hydrated, it forms sodium hydroxide.

Oxides of more electronegative elements tend to be acids. They are called acid anhydrides; adding water, they form oxoacids. For example, dichlorine heptoxide is acid; perchloric acid is a more hydrated form.

Some oxides can act as both acid and base, at different times. They are amphoteric. An example is aluminium oxide. Some oxides do not show behavior as either acid or base.

The oxides of the chemical elements in their highest oxidation state are predictable and the chemical formula can be derived from the number of valence electrons for that element. Even the chemical formula of ozone 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 that does not exist as expected as F₂O₇ but as OF₂ with the least electronegative element given priority. Template:Ref. Phosphorus pentoxide, the third exception is not properly represented by the chemical formula P₂O₅ but by P₄O₁₀

Common oxides sorted by oxidation state

• Covalent oxides
  • Water (H₂O)
  • Carbon dioxide (CO₂)
  • Nitrogen oxides (NO_x)
  • Sulfur dioxide (SO₂)

• Element in (I) state
  • Sodium oxide (Na₂O)
  • Potassium oxide (K₂O)
  • Copper(I) oxide (Cu₂O)
  • Dichlorine monoxide (Cl₂O)
  • Nitrous oxide (N₂O)

• Element in (II) state
  • Beryllium oxide (BeO)
  • Calcium oxide (CaO)
  • Carbon monoxide (CO)
  • Copper(II) oxide (CuO)
  • Iron(II) oxide (FeO)
  • Lead(II) oxide (PbO₂)
  • Magnesium oxide (MgO)
  • Mercury oxide (Hg{{#if:{{{1|}}}|_{{{1}}}|}}O)
  • Nitrogen oxide (NO)
  • Zinc oxide (ZnO)

• Element in (III) state
  • Aluminium oxide (Al₂O₃)
  • Antimony trioxide (Sb₂O₃)
  • Arsenic trioxide (As₂O₃)
  • Boron oxide (B₂O₃)
  • Dinitrogen trioxide (N₂O₃)
  • Iron(III) oxide (Fe₂O₃)
  • Yttrium(III) oxide (Y₂O₃)

• Element in (IV) state
  • Carbon dioxide (CO₂)
  • Cerium(IV) oxide (CeO₂)
  • Chlorine dioxide (ClO₂)
  • Manganese dioxide (MnO₂)
  • Nitrogen dioxide (NO₂)
  • Ozone (O₃)
  • Plutonium dioxide (PuO₂)
  • Silicon dioxide (SiO₂)
  • Sulfur dioxide (SO₂)
  • Tellurium dioxide (TeO₂)
  • Thorium dioxide (Th{{#if:{{{1|}}}|_{{{1}}}|}}O₂)
  • Titanium dioxide (TiO₂)
  • Uranium dioxide (UO₂)
  • Zirconia (ZrO₂)

• Element in (V) state
  • Dinitrogen pentoxide (N₂O₅)
  • Phosphorus pentoxide (P₂O₅)

• Element in (VI) state
  • Molybdenum trioxide (MoO₃)
  • Sulfur trioxide (SO₃)

• Element in (VII) state
  • Dichlorine heptoxide (Cl₂O₇)

• Element in (VIII) state
  • Osmium tetroxide (OsO₄)

See also

• Other oxygen ions ozonide, O₃⁻, superoxide, O₂⁻, peroxide, O₂²⁻ and dioxygenyl, O₂⁺.

References

1. Template:Note Fully Exploiting the Potential of the Periodic Table through Pattern Recognition Schultz, Emeric. J. Chem. Educ. 2005 82 1649.bg:Оксид

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