All the nonmetals form covalent oxides with oxygen that react with water to produce acids or with bases to produce salts. The majority of nonmetal oxides are acidic, forming oxyacids, which contain hydronium ions (H3O+) in aqueous solutions.
There exist two general statements, which describe the acidic oxide behaviour. The oxides, such as dinitrogen pentoxide (N2O5) and sulphur trioxide (SO3), are called acid anhydrides because the nonmetal exhibits one of its typical oxidation numbers.
Reactions of Nonmetal Oxides
Nonmetal oxides react with water to produce oxyacids, with no change in the nonmetal’s oxidation number. For example:
N2O5 + H2O → 2HNO3
Second, metal oxides that do not have one of the metal's typical oxidation numbers, such as nitrogen dioxide (NO2) and chlorine dioxide (ClO2), react with water as well. The nonmetal, on the other hand, is both oxidised and reduced in these reactions (i.e., its oxidation number is increased and decreased, respectively). A disproportionation reaction occurs when the same element is oxidised and reduced at the same time. N0+ is reduced to N2+ (in NO) and oxidised to N5+ in the next disproportionation reaction (in HNO3).
3NO2 + H2O → 2HNO3 + NO
Oxides of Nitrogen
The oxides are formed by nitrogen (N), which exhibit each of its positive oxidation numbers ranging from +1 to +5. Nitrous oxide (or dinitrogen oxide), N2O, is formed when the ammonium nitrate (NH4NO3) is heated. This colourless gas has a nice and mild odour and a sweet taste and is used as a local anaesthetic for minor procedures, especially in dentistry. It is also called laughing gas due to its intoxicating effect. And, it is used widely as a propellant in aerosol cans of whipped cream.
Nitric oxide (NO) is created in many ways. The lightning that takes place during thunderstorms brings up the direct union of oxygen and nitrogen in the air to form fewer amounts of nitric oxide, as does heating these two elements together. Nitric oxide can be produced commercially by burning ammonia (NH3), but it can also be made in the lab by reducing dilute nitric acid (HNO3) with, for example, copper (Cu).
3Cu + 8HNO3 → 2NO + 3Cu(NO3)2 + 4H2O
Oxides of Phosphorus
Phosphorus (III) oxide (also known as tetraphosphorus hexoxide), P4O6, and phosphorus (V) oxide (also known as tetraphosphorus decaoxide), P4O10, are two popular oxides. Both these oxides contain a structure based on the tetrahedral structure of the elemental white phosphorus. Phosphorus(III) oxide comes in the form of a white crystalline solid with a garlic-like odour and a poisonous vapour. It oxidizes slowly in the air and in flames when heated to 70°C (158°F) by forming P4O10. It is phosphoric acid (H3PO3) acid anhydride, produced as P4O6, dissolves slowly in cold water.
Phosphorus (V) oxide is a white flocculent powder made by heating elemental phosphorus in the presence of excess oxygen. It is a poor oxidizing agent and is very stable. The molecule P4O10 is an acid anhydride of H3PO4, an orthophosphoric acid. When this P4O10 is dropped into water, heat is liberated, the acid is formed and makes a hissing sound. Due to its great affinity for water, P4O10 can be used extensively as a drying agent for the gases and to remove water from several compounds.
P4O10 + 6H2O → 4H3PO4
Oxides of Carbon
Carbon forms two well-known oxides, which are carbon monoxide (CO), and carbon dioxide (CO2). In addition, it also forms C3O2 and carbon suboxide.
Carbon monoxide can be produced when graphite (a naturally occurring form of elemental carbon) is burned or heated in a limited amount of oxygen. Steam with red-hot coke reaction also produces carbon monoxide, including hydrogen gas (H2). Coke is given as an impure carbon residue resulting from coal burning.
This CO and H2 mixture is called water gas and can be used as an industrial fuel. In the laboratory, carbon monoxide can be prepared by heating the oxalic acid (H2C2O4) or formic acid (HCOOH) with the conc. sulfuric acid (H2SO4). The sulfuric acid removes the water elements (it means H2O) from the oxalic or formic acid and absorbs the produced water because the carbon monoxide burns readily in oxygen to form carbon dioxide,
2CO + O2 → 2CO2
It is also useful as a gaseous fuel and as a metallurgical reducing agent because it reduces several metal oxides to the elemental metal at high temperatures. For example, iron (III) oxide (Fe2O3) and copper (II) oxide (CuO) are both reduced to the metal by the carbon monoxide compound.
Carbon dioxide can be produced when almost any carbon compound or any form of carbon is burned in excess oxygen. Several metal carbonates liberate CO2 when heated. Calcium carbonate (CaCO3), for example, contains calcium oxide (CaO) and carbon dioxide.
CaCO3 + heat → CO2 + CaO