[Chemistry Class Notes] Nonmetal Oxides Pdf for Exam

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 colorless gas has a nice and mild odor and a sweet taste and is used as a local anesthetic 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 odor and a poisonous vapor. It oxidizes slowly in the air and flames when heated to 70°C (158°F) by forming P4O10. It is phosphoric acid (H3PO3) acid.

anhydride, produced as P4O6, that 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 gasses 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

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 metal by the carbon monoxide compound.

 

Carbon Dioxide

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

 

How to Prepare Notes on Non-Metal Oxides?

• Go through Nonmetal Oxides – Reactions, Oxides of Nitrogen, Carbon, and FAQs on .

• Read from this page and try to understand the explanations provided.

• Re-read those portions that seem a bit unclear towards the beginning.

• Write down everything in an organized manner.

• Keep your sentences brief and to the point.

• Highlight all the key parts using a coloured pen.

• Use drawings if possible to retain the concepts.

• Revise from here before appearing for a test on Non-metal oxides.

How Does Prepare Students for a Chemistry Test on Metal Oxides ?

has appropriate study material on Metal oxides that students can read from. They can check out Nonmetal Oxides on and then understand the topic better. The definition and reactions with the other oxides have been explained here. Studying from this section will help students secure higher marks in their exams as the material present here is relevant and in keeping with the Chemistry syllabus.