Oxide

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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. Rutile-unit-cell-3D-balls.png
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.

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– atom. Metal oxides thus 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. 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] Certain elements can form multiple oxides, differing in the amounts of the element combining with the oxygen. Examples are carbon, iron, nitrogen (see nitrogen oxide), silicon, titanium, and aluminium. In such cases the oxides are distinguished by specifying the numbers of atoms involved, as in carbon monoxide and carbon dioxide, or by specifying the element's oxidation number, as in iron(II) oxide and iron(III) oxide.

Contents

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, 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. Cesium 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 aluminum oxide layer can be built to greater thickness by the process of electrolytic anodizing. Though solid magnesium and aluminum 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.

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 Rust screw.jpg
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

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.

Structure

Oxides have a range of different structures, from individual molecules to polymeric and crystalline structures. At standard conditions, oxides may range from solids to gases.

Oxides of metals

Oxides of most metals adopt polymeric structures. [3] The oxide typically links three metal atoms (e.g., rutile structure) or six metal atoms (carborundum or rock salt structures). Because the M-O bonds are typically strong and these compounds are crosslinked polymers, the solids tend to be insoluble in solvents, though they are attacked by acids and bases. The formulas are often deceptively simple where many are nonstoichiometric compounds. [2]

Molecular oxides

Although most metal oxides are polymeric, 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. Some polymeric oxides depolymerize when heated to give molecules, examples being selenium dioxide and sulfur trioxide. Tetroxides are rare. The more common examples: ruthenium tetroxide, osmium tetroxide, and xenon tetroxide.

Many oxyanions are known, such as polyphosphates and 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.

Reduction

Conversion of a metal oxide to the metal is called reduction. The reduction can be induced by many reagents. Many metal oxides convert to metals simply by heating.

Reduction by carbon

Metals are "won" from their oxides by chemical reduction, i.e. by the addition of 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

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 center transfers an oxide ligand to the organic compound followed by regeneration of the metal oxide, often by oxygen in the air.

Reduction by heating

Metals that are lower in the reactivity series can be reduced by heating alone. For example, silver oxide decomposes at 200 °C: [4]

2 Ag2O → 4 Ag + O2

Reduction by displacement

Metals that are more reactive displace the oxide of the metals that are less reactive. For example, zinc is more reactive than copper, so it displaces copper (II) oxide to form zinc oxide:

Zn + CuO → ZnO + Cu

Reduction by hydrogen

Apart from metals, hydrogen can also displace metal oxides to form hydrogen oxide, also known as water:

H2 + CuO → Cu + H2O

Reduction by electrolysis

Since metals that are reactive form oxides that are stable, some metal oxides must be electrolyzed to be reduced. This includes sodium oxide, potassium oxide, calcium oxide, magnesium oxide, and aluminium oxide. The oxides must be molten before immersing graphite electrodes in them:

2Al2O3 → 4Al + 3O2


Hydrolysis and dissolution

Oxides typically react with acids or bases, sometimes both. Those reacting only with acids are labeled basic oxides. Those reacting only by bases are called "acidic oxides" (Anhydrous)[]. Oxides that react with both 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.

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 an acid anhydride; perchloric acid is its fully 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 solids 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.

Reductive dissolution

The reductive dissolution of a transition metal oxide occurs when dissolution is coupled to a redox event. [5] For example, ferric oxides dissolve in the presence of reductants, which can include organic compounds. [6] or bacteria [7] Reductive dissolution is integral to geochemical phenomena such as the iron cycle. [8]

Reductive dissolution does not necessarily occur at the site where the reductant adsorbs. Instead, the added electron travel through the particle, causing reductive dissolution elsewhere on the particle. [9] [10]

Nomenclature and formulas

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 adding the suffix -a to the element's name. Hence alumina, magnesia and chromia, are, respectively, Al2O3, MgO and 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. [11]

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

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.

NameFormulaFound/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 (O2) 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 ZnORequired 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

Related Research Articles

Barium Chemical element with atomic number 56

Barium is a chemical element with the symbol Ba and atomic number 56. It is the fifth element in group 2 and is a soft, silvery alkaline earth metal. Because of its high chemical reactivity, barium is never found in nature as a free element. Its hydroxide, known in pre-modern times as baryta, does not occur as a mineral, but can be prepared by heating barium carbonate.

Carboxylic acid oxoacid having the structure RC(=O)OH, used as a suffix in systematic name formation to denote the –C(=O)OH group including its carbon atom

A carboxylic acid is an organic compound that contains a carboxyl group. The general formula of a carboxylic acid is R–COOH, with R referring to the alkyl group. Carboxylic acids occur widely. Important examples include the amino acids and acetic acid. Deprotonation of a carboxyl group gives a carboxylate anion.

Hydroxide family of the hydroxide salts

Hydroxide is a diatomic anion with chemical formula OH. It consists of an oxygen and hydrogen atom held together by a covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile, and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. A hydroxide attached to a strongly electropositive center may itself ionize, liberating a hydrogen cation (H+), making the parent compound an acid.

Hydrolysis is any chemical reaction in which a molecule of water ruptures one or more chemical bonds. The term is used broadly for substitution, elimination, and fragmentation reactions in which water is the nucleophile.

Inorganic chemistry deals with the synthesis and behavior of inorganic and organometallic compounds. This field covers all chemical compounds except the myriad organic compounds, which are the subjects of organic chemistry. The distinction between the two disciplines is far from absolute, as there is much overlap in the subdiscipline of organometallic chemistry. It has applications in every aspect of the chemical industry, including catalysis, materials science, pigments, surfactants, coatings, medications, fuels, and agriculture.

Rust type of iron oxide

Rust is an iron oxide, a usually red oxide formed by the redox reaction of iron and oxygen in the presence of water or air moisture. Several forms of rust are distinguishable both visually and by spectroscopy, and form under different circumstances. Rust consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3).

Silicon Chemical element with atomic number 14

Silicon is a chemical element with the symbol Si and atomic number 14. It is a hard and brittle crystalline solid with a blue-grey metallic lustre; and it is a tetravalent metalloid and semiconductor. It is a member of group 14 in the periodic table: carbon is above it; and germanium, tin, and lead are below it. It is relatively unreactive. Because of its high chemical affinity for oxygen, it was not until 1823 that Jöns Jakob Berzelius was first able to prepare it and characterize it in pure form. Its melting and boiling points of 1414 °C and 3265 °C respectively are the second-highest among all the metalloids and nonmetals, being only surpassed by boron. Silicon is the eighth most common element in the universe by mass, but very rarely occurs as the pure element in the Earth's crust. It is most widely distributed in dusts, sands, planetoids, and planets as various forms of silicon dioxide (silica) or silicates. More than 90% of the Earth's crust is composed of silicate minerals, making silicon the second most abundant element in the Earth's crust after oxygen.

Sulfuric acid chemical compound

Sulfuric acid (alternative spelling sulphuric acid), also known as vitriol, is a mineral acid composed of the elements sulfur, oxygen and hydrogen, with molecular formula H2SO4. It is a colorless, odorless, and viscous liquid that is soluble in water and is synthesized in reactions that are highly exothermic.

Electrolysis technique that uses a direct electric current to drive an otherwise non-spontaneous chemical reaction

In chemistry and manufacturing, electrolysis is a technique that uses a direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential.

Sodium hydroxide Chemical compound with formula NaOH

Sodium hydroxide, also known as lye and caustic soda, is an inorganic compound with the formula NaOH. It is a white solid ionic compound consisting of sodium cations Na+
and hydroxide anions OH
.

Redox Chemical reaction

Redox is a type of chemical reaction in which the oxidation states of atoms are changed. Redox reactions are characterized by the transfer of electrons between chemical species, most often with one species undergoing oxidation while another species undergoes reduction. The chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. In other words:

Weathering Breaking down of rocks, soil and minerals as well as artificial materials through contact with the Earths atmosphere, biota and waters

Weathering is the breaking down of rocks, soil, and minerals as well as wood and artificial materials through contact with the Earth's atmosphere, water, and biological organisms. Weathering occurs in situ, that is, in the same place, with little or no movement, and thus should not be confused with erosion, which involves the movement of rocks and minerals by agents such as water, ice, snow, wind, waves and gravity and then being transported and deposited in other locations.

Base (chemistry) substance that can react with an acid, accepting hydrogen ions (protons) or more generally, donating a pair of valence electrons

In chemistry, bases are substances that, in aqueous solution, release hydroxide (OH) ions, are slippery to the touch, can taste bitter if an alkali, change the color of indicators (e.g., turn red litmus paper blue), react with acids to form salts, promote certain chemical reactions (base catalysis), accept protons from any proton donor or contain completely or partially displaceable OH ions. Examples of bases are the hydroxides of the alkali metals and the alkaline earth metals (NaOH, Ca(OH)2, etc.—see alkali hydroxide and alkaline earth hydroxide).

Reducing agent is an element or compound that loses an electron to an electron recipient in a redox chemical reaction.

A period 3 element is one of the chemical elements in the third row of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when the periodic table skips a row and a chemical behaviour begins to repeat, meaning that elements with similar behavior fall into the same vertical columns. The third period contains eight elements: sodium, magnesium, aluminium, silicon, phosphorus, sulfur, chlorine, and argon. The first two, sodium and magnesium, are members of the s-block of the periodic table, while the others are members of the p-block. All of the period 3 elements occur in nature and have at least one stable isotope.

In chemistry, a reactivity series (or activity series) is an empirical, calculated, and structurally analytical progression of a series of metals, arranged by their "reactivity" from highest to lowest. It is used to summarize information about the reactions of metals with acids and water, double displacement reactions and the extraction of metals from their ores.

Acidic oxides, or acid anhydride, are oxides that react with water to form an acid, or with a base to form a salt. They are oxides of either nonmetals or of metals in high oxidation states. Their chemistry can be systematically understood by taking an oxoacid and removing water from it, until only an oxide remains. The resulting oxide belongs to this group of substances. For example, sulfurous acid (SO2), sulfuric acid (SO3), and carbonic acid (CO2) are acidic oxides. An inorganic anhydride (a somewhat archaic term) is an acid anhydride without an organic moiety.

Pyrometallurgy is a branch of extractive metallurgy. It consists of the thermal treatment of minerals and metallurgical ores and concentrates to bring about physical and chemical transformations in the materials to enable recovery of valuable metals. Pyrometallurgical treatment may produce products able to be sold such as pure metals, or intermediate compounds or alloys, suitable as feed for further processing. Examples of elements extracted by pyrometallurgical processes include the oxides of less reactive elements like iron, copper, zinc, chromium, tin, and manganese.

A combination reaction (also known as a synthesis reaction) is a reaction where two or more elements or compounds (reactants) combine to form a single compound (product). Such reactions may be represented by equations of the following form: X + Y → XY. The combination of two or more elements and form one compound is called combination reaction. | a) Between elements | C + O2 → CO2 | Carbon completely burnt in oxygen yields carbon dioxide |- | b) Between compounds | CaO + H2O → Ca(OH)2 | Calcium oxide (lime) combined with water gives calcium hydroxide (slaked lime) |- | c) Between elements and compounds | 2CO + O2 → 2CO2 | Oxygen combines with carbon monoxide, and carbon dioxide is formed. |}

Compounds of oxygen any chemical compound having at least one oxygen atom

The oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as peroxides. Compounds containing oxygen in other oxidation states are very uncommon: −​12 (superoxides), −​13 (ozonides), 0, +​12 (dioxygenyl), +1, and +2.

References

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