Oxide

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Silicon dioxide (SiO2) is one of the most common oxides on the surface of earth. Like most oxides, it adopts a polymeric structure.

An oxide /ˈɒksd/ is a chemical compound that contains at least one oxygen atom and one other element[1] in its chemical formula. Metal oxides typically contain an anion of oxygen in the oxidation state of −2. Most of the Earth's crust consists of solid oxides, the result of elements being oxidized by the oxygen in air or in water. Hydrocarbon combustion affords the two principal carbon oxides: carbon monoxide and carbon dioxide. 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] Different oxides of the same element are distinguished by Roman numerals denoting their oxidation number, e.g. iron(II) oxide versus iron(III) oxide.

Formation[edit]

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, Fe2O3−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.

Oxides, such as iron(III) oxide or rust, which consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), form when oxygen combines with other elements

Structure[edit]

Oxides of most metals adopt polymeric structures with M-O-M crosslinks. Because these crosslinks are strong, the solids tend to be insoluble in solvents, though they are attacked by acids and bases. The formulas are often deceptively simple. Many are nonstoichiometric compounds. In these oxides, the coordination number of the oxide ligand is two for most electronegative elements and 3–6 for most metals.[2]

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

Molecular oxides[edit]

Although most metal oxides are polymeric, some oxides are molecules. The most famous molecular oxides are carbon dioxide and carbon monoxide. Phosphorus pentoxide is a more complex molecular oxide with a deceptive name, the formula being P4O10. Some polymeric oxides when heated depolymerize to give molecules, examples being selenium dioxide and sulfur trioxide. Tetroxides are rare, and there are only five known examples: ruthenium tetroxide, osmium tetroxide, hassium tetroxide, iridium tetroxide, and xenon tetroxide.

Many oxyanions are known, such as polyphosphates and polyoxometalates. Oxycations are rarer, an example being nitrosonium (NO+). 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.

Reactivity[edit]

Oxides can be attacked by acids and bases. Those attacked only by acids are basic oxides; those attacked only by bases are acidic oxides. Oxides that react with both acids and bases are amphoteric. Metals tend to form basic oxides, non-metals tend to form acidic oxides, and amphoteric oxides are formed by elements near the boundary between metals and non-metals (metalloids).

This reactivity is the basis of many practical processes such, as the extraction of some metals from their ores in the process called hydrometallurgy.

Reduction[edit]

Metals are "won" from their oxides by chemical reduction. 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

Metal oxides can be reduced by organic compounds. This redox process is the basis for many important transformations in chemistry, such as the detoxification of drugs by the P450 enzymes and the production of ethylene oxide, which is converted to antifreeze. In such systems the metal centre transfers an oxide ligand to the organic compound followed by regeneration of the metal oxide, often by oxygen in air.

Hydrolysis[edit]

Oxides of more electropositive elements tend to be basic. They are called basic anhydrides. Exposed to 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 acidic. 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. They are amphoteric. An example is aluminium oxide. Some oxides do not show behavior as either acid or base.

The oxide ion has the formula O2−. It is the conjugate base of the hydroxide ion, OH, and is encountered in ionic solid such as calcium oxide. O2− is unstable in aqueous solution − its affinity for H+ is so great (pKb ~ −38) that it abstracts a proton from a solvent H2O molecule:

O2− + H2O → 2 OH

The equilibrium constant of aforesaid reactions is pKeq ~ −22

In the 18th century, oxides were named calxes or calces after the calcination process used to produce oxides. Calx was later replaced by oxyd.

Nomenclature and formulas[edit]

Sometimes, metal-oxygen ratios are used to name oxides. Thus, NbO would be called niobium monoxide and TiO2 is titanium dioxide. This naming follows the Greek numerical prefixes. In the older literature and continuing in industry, oxides are named by contracting the element name with "a." Hence alumina, magnesia, chromia, are, respectively, Al2O3, MgO, Cr2O3.

Special types of oxides are peroxide, O22−, and superoxide, O2. In such species, oxygen is assigned higher oxidation states than oxide.

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.[3]

Since fluorine is more electronegative than oxygen, oxygen difluoride (OF2) does not represent an oxide of fluorine, but instead represents a fluoride of oxygen.

Examples of oxides[edit]

The following table gives examples of commonly encountered oxides. Only a few representatives are given, as the number of polyatomic ions encountered in practice is very large.

Name Formula Found/Usage
Water (hydrogen oxide) H
2
O
Common solvent, Required by carbon-based life
Nitrous oxide N
2
O
Laughing gas, anesthetic (used in a combination with diatomic oxygen to make Nitrous oxide and oxygen anesthesia), produced by Nitrogen-fixing bacteria, Nitrous, oxidizing agent in rocketry, aerosol propellant, recreational drug, greenhouse gas. Other nitrogen oxides such as NO
2
(Nitrogen dioxide), NO(Nitrogen oxide), N
2
O
3
(Dinitrogen trioxide) and N
2
O
4
(Dinitrogen tetroxide) exist, particularly in areas with notable air pollution. They are also strong oxidisers, can add Nitric acid to Acid rain, and are harmful to health.
Silicon dioxide SiO
2
Sand, quartz
Iron(II,III) oxide Fe
3
O
4
Iron Ore, Rust, along with iron(III) oxide (Fe
2
O
3
)
Aluminium oxide Al
2
O
3
Aluminium Ore, Alumina, Corundum, Ruby (Corundum with impurities of Chromium).
Zinc oxide ZnO Reqiured for vulcanization of rubber, additive to concrete, sunscreen, skin care lotions, antibacterial and antifungal properties, food additive, white pigment.
Carbon dioxide CO
2
Constituent of the atmosphere of Earth, the most abundant and important greenhouse gas, used by plants in photosynthesis to make sugars, product of biological processes such as respiration and chemical reactions such as combustion and chemical decomposition of carbonates. CO or Carbon monoxide exists as a product of incomplete combustion and is a highly toxic gas.
Calcium oxide CaO Quicklime (used in construction to make mortar and concrete), used in Self-heating cans due to exothermic reaction with water to produce Calcium hydroxide, possible ingredient in Greek fire and produces limelight when heated over 2,400 °Celsius.

See also[edit]

References[edit]

  1. ^ Foundations of College Chemistry, 12th Edition
  2. ^ a b c Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
  3. ^ Schultz, Emeric (2005). "Fully Exploiting the Potential of the Periodic Table through Pattern Recognition". J. Chem. Education 82: 1649. Bibcode:2005JChEd..82.1649S. doi:10.1021/ed082p1649.