Oxides

Oxides contain at least one other element and one oxygen atom. Metal oxides commonly include an oxygen anion in state 2. The majority of the Earth’s crust comprises solid oxides, which form when oxygen oxidises elements found in the air or water. The combustion of hydrocarbon (CO2) produces CO and CO2. Even pure elements get oxide coatings. For example, aluminium foil creates a thin Al2O3 coating that protects it from further corrosion.

Oxides classify on their valency or ability to combine with another element. Based on valency, oxides classify into simple and mixed oxides.

Simple Oxides

Simple oxides are metal-oxygen or semimetal-oxygen compounds. Simple oxides contain atoms allowed by the element’s or metal’s normal valency.

Mixed Oxides

Combining simple oxides produces mixed oxides. You may use the same metal or a different metal to create this combination of simple oxides.

Metal Oxides

Metal oxides are crystalline solids containing oxide anion and metal cation, and they usually react with water or acids to generate bases or salts. The alkali metals and alkaline earth metals produce three types of binary oxygen compounds: oxides (O2−), peroxides (O22−), comprising oxygen-oxygen covalent single bonds, and superoxides (O2-), that have oxygen-oxygen covalent single bonds but a lower negative charge than peroxide ions. 

Alkali metals produce oxides (M2O), peroxides (M2O2), and superoxides (MO2). M refers to the metal atom. The only earth metal oxides to exist are Mo and MO2. Heat the metal nitrate with the elemental metal to form all alkali metal oxides. 

2MNO3 + 10M + heat → 6M2O + N2. The preparation of alkaline earth oxides requires heating metal carbonates. 

MCO3 + heat → MO + CO2. 

Ionic alkali metal oxides as well as alkaline earth metal oxides react with water to generate metal hydroxide solutions.

M2O + H2O → 2MOH (M = group 1 metal)

MO + H2O → M(OH)2 (M = group 2 metal)

They are basic oxides. In typical acid-base reactions, they react with acids to form salts and water, such as M2O + 2HCl → 2MCl + H2O (M = group 1 metal). 

These are also known as neutralising reactions. Magnesium oxide (MgO) used in firebrick and thermal insulation and calcium oxide (CaO) used in steel manufacturing and water purification are the two most important basic oxides.

Scientists thoroughly studied the periodic trends of oxides. In terms of their acid-base nature, oxides can vary from highly basic to weakly basic, amphoteric, weakly acidic, and strongly acidic. Acidity rises with element oxidation number. For example, MnO (with an oxidation state of +2) is the least acidic of the five manganese oxides, while Mn2O7 (Mn7+) is most acidic. Ionic compounds have oxidation numbers +1, +2, and +3 are transition metal oxides.  

Transition metal oxides having oxidation numbers +4, +5, +6, and +7 contain covalent metal-oxygen bonds and behave as covalent compounds. The ionic transition metal oxides are basic in general. 

They will react with aqueous acids to generate salt and water solutions, like 

CoO + 2H3O+ → Co2+ + 3H2O. Acidic +5, +6, and +7 oxides react with hydroxide solutions to create salts and water, as in CrO3 + 2OH– → CrO42− + H2O. Oxides with +4 oxidation values are often amphoteric (from the Greek amphoteros, “in both directions”’), which may act as acids or bases. Amphoteric oxides dissolve in both acidic and basic solutions. For example, vanadium oxide (VO2) dissolves in acid to form [VO]2+ and in base to produce [V4O9]2-. Amphoterism in the main group of oxides occurs mainly with metalloid elements or neighbouring elements. 

Nonmetal Oxides

Nonmetals combine with oxygen to produce covalent oxides that react with water to make acids or bases. Most nonmetal oxides are acidic, forming oxyacids that generate H3O+ ions in water. The behaviour of acidic oxides is summarised in two general statements. First, oxides like SO3 and N2O5 are acid anhydrides. These oxides react with water to form oxyacids, retaining the nonmetal’s oxidation number. 2HNO3 + N2O5 Second, metal oxides with low oxidation numbers, such as NO2 and ClO2, react with water. The nonmetal gets oxidised and reduced in these reactions (i.e., its oxidation number increases and decreases, respectively). A disproportionation reaction happens when oxidises and reduces the same element. Reduce N4+ to N2+ (in NO) and oxidise to N5+ in the following disproportionation process (HNO3).

3NO2 + H2O → 2HNO3 + NO 

Oxides of Nitrogen

Nitrogen (N) generates oxides with positive oxidation numbers from +1 to +5. When we heat ammonium nitrate (NH4NO3), it creates nitrous oxide (N2O). This colourless gas is an anaesthetic for minor procedures in dentistry. It is commonly known as laughing gas for its intoxicating impact and is also widely used in whipped cream aerosol cans. Several processes generate nitric oxide (NO). During thunderstorms, the direct mixing of nitrogen and oxygen forms nitric oxide, which heats the two elements together. 

Commercially burning ammonia (NH3) produces nitric oxide, but reducing dilute nitric acid (HNO3) with copper in a laboratory also creates it. 

For example, 

(Cu). 3Cu + 8HNO3 → 2NO + 3Cu (NO3)2+4H2O. 

Gaseous nitric oxide is the most thermally stable nitrogen oxide and simplest known thermally stable paramagnetic (unpaired electron) molecule. It is found in pollution generated by internal combustion engines when nitrogen and oxygen in the air react during combustion. Nitric oxide is a colourless diatomic gas at room temperature, and due to the unpaired electron, two molecules may combine to form a dimer. 2NO ⇌ N2O2 as a result, liquid nitric oxide is partly dimerised, and the solid only contains dimers.

Upon cooling the combination of nitric oxide and nitrogen dioxide (NO2) to -21°C (-6°F), the gases combine to form dinitrogen trioxide, a blue liquid composed of N2O3 molecules. This molecule is solely liquid or solid. When heated, it forms NO2 and NO. Commercially oxidising NO with the air produces nitrogen dioxide, also made in the lab by heating a heavy metal nitrate, as in 2Pb(NO3)2 + heat → 2PbO + 4NO2 + O2, or by the addition of copper to concentrated nitric acid. Nitrogen dioxide is paramagnetic like nitric oxide, and its unpaired electron gives it colour and dimerisation. NO2 is dark brown at low pressures and high temperatures but dimerisings generate dinitrogen tetroxide, N2O4. At standard temperature, both molecules are in equilibrium. 2NO2 ⇌ N2O4. 

Conclusion

Oxides are a broad and significant family of chemical compounds composed of oxygen and another element. Metal oxides are metal cations and an oxide anion (O2) that react with water to produce bases or acids used to make salts. Nonmetallic oxides are volatile compounds that form a covalent link between oxygen and the nonmetal; they react with water to make acids or with bases to form salts. Aluminium and zinc, for example, produce amphoteric oxides, which combine with acids and bases to create salts. Certain organic compounds generate oxides when the oxygen atom in the organic molecule forms a covalent bond with an atom of phosphorus (phosphine oxides), nitrogen (amine oxides), or sulphur (sulfoxides).