Oxide

Oxides (from Greek ὀξύς, oxys = sharp, pointed, sour) are oxygen compounds in which this has the oxidation number II. Most oxides are formed when flammable substances react with oxygen ( oxidation origin of the word ): At its oxidation they give electrons to the oxidant is oxygen, so that oxides are formed.

Depending on the binding partner, a distinction in the two groups of materials chemistry of oxides:

• Metal oxides (these are ionic ( salt-like ) or covalent oxides, oxides of base metals react with water to form bases and alkalis),
• Non-metal oxides (these are molecular, ia volatile and react with water to form acids) and

According to its stoichiometric composition, a distinction Monoxide, Dioxide, Trioxide, tetroxides, Pentoxide, as for carbon monoxide, chlorine dioxide and sulfur trioxide. The majority of the earth's crust and mantle consisting of oxides (mainly of silica (quartz) and salts derived therefrom, the silicates and aluminum oxide ). Water also belongs to the group of oxides. Ethylene oxide is an example of an organic oxide.

Production

Oxides are prepared by:

• Heat of hydroxides and oxide hydrates (for example: Copper ( II) hydroxide to copper ( II) oxide and water vapor, will rust iron oxides and water vapor),
• Heat of salts with volatile anhydrides (for example, the burning of limestone / calcium carbonate to calcium oxide, by heating the copper ( II) nitrate to copper ( II) oxide and nitrous gases )
• Reaction of elements with oxygen ( oxidation in the strict sense, formerly referred to as oxygenation ).

That shown above black copper ( II) oxide can thus be synthesized by the following reactions for example:

Furthermore, the red copper (I ) oxide by oxygen could be in black copper convert (II ) oxide. Also in the roasting of sulfidic copper ores of copper (II ) oxide is prepared by adding copper (II ) sulphide in the air or in an oxygen flow glows (by-product sulfur dioxide).

Such a metal easily forms an oxide depends on the oxygen affinity and electronegativity of the element. Each a metal less noble, the more violently, it can generally react with oxygen to form oxides. In addition, the reactivity also depends on the passivation of the element, as with many elements of a tightly adherent oxide layer prevents further reaction. Only if it is permeable to oxygen, or is removed, the metal may react further.

Properties of the oxides

• Amphoteric oxides and hydroxides have the property of being able to react acidic and basic, depending on the reactants (see under acid -base reaction ). They react with acids and with bases to form salts.
• Metal oxides are salt-like ( ionic), oxides of base metals react with water to form bases and alkalis.
• Non-metal oxides have molecular and react with water to form acids,

Oxides of more noble metals are therefore often transformed via a detour as salts into hydroxides in order reaction with water: copper (II ) oxide, for example, in conc. Hydrochloric acid to be dissolved into copper (II ) chloride. This together with sodium hydroxide solution of copper ( II) hydroxide, which can be converted by heating with copper (II ) oxide, as described above.

Hydroxides are flocculent precipitates which often have characteristic staining (copper (II ) hydroxide light blue, nickel (II ) hydroxide apple green, chromium ( III ) hydroxide gray-green, manganese (II ) hydroxide pink and in air due to oxidation becoming brown, cobalt (II ) hydroxide blue or pink, iron (III ) hydroxide rust brown, iron (II ) hydroxide gray green).

Oxide ion and hydroxide ion

The metal oxide underlying O2 - ion is formed in the oxidation-reduction reaction of the oxidant of oxygen with a metal. It is existent only in melts and in combination with cations (in the form of salts ), but not as a free ion, since it is an extremely strong base and is thus quantitatively protonated in aqueous solution to the hydroxide ion ( acid -base reaction ). Metal hydroxides containing the OH - ion and are mostly derived from salt solutions and bases.

In non-metal oxides is usually in the no oxide anion, as non-metals with each other form a covalent bond atom. The peroxide ion, the oxide ion similar has instead II to an oxidation number of -I, since two oxygen atoms are connected to each other. Nonmetal oxides react with water to form acids ( with oxo - anions such as sulfate, carbonate, etc.). They are therefore to be regarded as hydroxides of acidic character.

The binding capacity of the oxygen

Oxygen is a strong oxidizing agent and, together with almost all elements isolable oxides, with the exception of the noble gases helium, neon, argon, krypton and halogen fluorine ( fluorine here occupies a special position because, although the oxygen compounds OF2, O2F2 and O4F2 are displayed, these but substances not referred to as fluorine oxides, but as oxygen fluorides because of the higher electronegativity of fluorine ).

Oxygen forms oxides in addition also oxo anions: Here, have more oxygen atoms bound to an atom, which usually has the highest possible oxidation number (examples: phosphate, sulfate, chromate, permanganate, nitrate, carbonate ). They usually arise when non-metal group or transition metal oxides react with very high oxidation with water to form acids.

In addition, oxygen-oxygen compounds exist such as in the bleaching agent is hydrogen peroxide (see above). Inorganic peroxides are highly corrosive and oxidizing, organic peroxides usually explosive.

Use

Natural metal oxides are used as ores for metal recovery. You will be smelting - for example, by carbon ( blast furnace process ) - deprived of oxygen and thus obtained the pure metal.

Metal oxides have been used as pigments in the Stone Age and also called earth pigments.

Another application in modern times is the use as an insulator in Information Technology.

Individual oxides and other oxygen compounds

Known oxides

• Aluminum oxide (a white, slightly basic solids )
• Lead oxide (there is a yellow and a black-brown oxide and a mixed oxide, red lead called )
• Calcium oxide ( Quicklime, corrosive, strongly basic )
• Dihydrogen (mon ) oxide, hydrogen oxide (water)
• Iron oxides such as iron ( III) oxide or rust
• Carbon dioxide ( carbonic acid forms with water )
• Carbon monoxide ( toxic, colorless and odorless, flammable)
• Copper (there is red copper (I ) oxide and black copper (II ) oxide )
• Magnesium oxide (magnesia, a white, basic powder)
• Phosphorus pentoxide ( phosphoric acid forms with water )
• Sulfur ( sour odor, with sour reaction in water soluble)
• Sulfur trioxide ( sulfuric acid forms with water )
• Silica ( quartz, forms with water may silica)
• Nitrogen oxides ( mostly brown colored by nitrogen dioxide, can may react further to form nitric acid)
• Zinc oxide ( an earthy- white, pale yellow in heat powder)

Oxygen compounds with oxygen in other oxidation states are:

• Hyperoxide ( - ½),
• Ozonides (-1 / 3)
• Peroxides (-1)
• Salts with the cation dioxygenyl O2 ( ½).