Oxide

An oxide (

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]

${\displaystyle {\ce {CaCO3 -> CaO + CO2}}}$

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]

${\displaystyle {\ce {2 Zn + O2 -> 2 ZnO}}}$

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:^[6]

${\displaystyle {\ce {2 MoS2 + 7 O2 -> 2MoO3 + 4 SO2}}}$

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

${\displaystyle {\ce {NiS + 3/2 O2 -> NiO + SO2}}}$

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]

${\displaystyle {\ce {CH4 + 3/2 O2 -> CO + 2 H2O}}}$

${\displaystyle {\ce {C + 1/2 O2 -> CO}}}$

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

${\displaystyle {\ce {CH4 + 2 O2 -> CO2 + 2 H2O}}}$

${\displaystyle {\ce {C + O2 -> CO2}}}$

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

${\displaystyle {\ce {4 NH3 + 5 O2 -> 4 NO + 6 H2O}}}$

${\displaystyle {\ce {NO + 1/2 O2 -> NO2}}}$

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]

${\displaystyle {\ce {S + O2 -> SO2}}}$

${\displaystyle {\ce {SO2 + 1/2 O2 -> SO3}}}$

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

${\displaystyle {\ce {SO3 + H2O -> H2SO4}}}$

Structure

Oxides have a range of structures, from individual molecules to

• Some important gaseous oxides
• 
  Carbon dioxide is the main product of fossil fuel combustion.
• 
  Carbon monoxide is the product of the incomplete combustion of carbon-based fuels and a precursor to many useful chemicals.
• 
  Nitrogen dioxide is a problematic pollutant from internal combustion engines.
• 
  Sulfur dioxide, the principal oxide of sulfur, is emitted from volcanoes.
• 
  Nitrous oxide ("laughing gas") is a potent greenhouse gas produced by soil bacteria.

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 iridium tetroxide,^[10] ruthenium tetroxide, osmium tetroxide, and xenon tetroxide.^[2]

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, (see Thermal decomposition). For example, silver oxide decomposes at 200 °C:^[11]

${\displaystyle {\ce {2 Ag2O -> 4 Ag + O2}}}$

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]

${\displaystyle {\ce {2 Fe2O3 + 3 C -> 4 Fe + 3 CO2}}}$

Some

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

Look up oxide in Wiktionary, the free dictionary.

• Other oxygen ions ozonide, O−3, superoxide, O−2, peroxide, O2−2 and dioxygenyl, O+2.
• Suboxide
• Oxohalide
• Oxyanion
• Complex oxide
• See Category:Oxides for a list of oxides.
• Salt
• Wet electrons

References