Oxides

You have probably heard of the element Oxygen. Oxygen is a fascinating element with a wide range of uses in our daily lives. Oxygen can react with metals and nonmetals to generate oxides, which are chemical compounds. Many elements are naturally oxidised in air or water. This is why numerous Oxides can be found as minerals in nature. The majority of the Earth’s crust is made up of solid Oxides, which are formed when elements are oxidised by oxygen in the air or water.

A chemical molecule with at least one oxygen atom and one additional element in its chemical formula is known as an oxide. Oxide is an O2– (molecular) ion, which is the dianion of oxygen. We will study the substance – Oxide – in this article.

What are Oxides?

An element that merges with the binary compound of oxygen is called oxide. Oxygen reacts with most of the metals of the periodic table to form oxides. In many cases, one element forms two or more oxides. The oxides vary widely in their nature and properties. For example, when the elements combine with oxygen it forms dioxides \(\left( {{\rm{M}}{{\rm{O}}_{\rm{2}}}} \right)\) and tri oxides \(\left( {{\rm{M}}{{\rm{O}}_3}} \right)\) (where \({\rm{M}}\)=Sulphur, Selenium, Tellurium, etc).

Examples of Oxides

Carbon dioxide, Sulphur dioxide, Calcium oxide, Carbon monoxide, Zinc oxide, Barium peroxide, Water, etc. These are termed oxides because oxygen is in combination with only one element.

Dioxides

An oxide that contains two atoms of oxygen, each bonded directly to an atom of a second element, is called dioxide. These can be produced by burning the element in the air.
\({\rm{S}} + {{\rm{O}}_2} \to {\rm{S}}{{\rm{O}}_2}\)

Trioxides

An oxide that contains three atoms of oxygen, each bonded directly to an atom of a second element is called trioxide. \({\rm{S}}{{\rm{O}}_3}\) is the most important trioxide.
\(2{\rm{S}}{{\rm{O}}_2} + {{\rm{O}}_2} \to 2{\rm{S}}{{\rm{O}}_3}\)
It is the anhydride of sulphuric acid.
\({\rm{S}}{{\rm{O}}_3} + {{\rm{H}}_2}{\rm{O}} \to {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}\)
Selenium and Tellurium also give trioxides.

The Formula of Oxide

An oxide ion is a negatively charged oxygen atom. This means it has gained two electrons from another atom. It is represented as \({{\rm{O}}^{2 – }}.\)
\(\frac{1}{2}{{\rm{O}}_2}\left( {\rm{g}} \right) + 2{{\rm{e}}^ – } \to {{\rm{O}}^{2 – }}\)

Classification of Oxides

They are two types of oxides:

Simple Oxides

Simple oxides are oxides that carry only that number of oxygen atoms as is allowed by the normal valency of its metal. Simple oxides can be classified based on their acidic, basic, neutral, and amphoteric oxides. Magnesium oxide, Aluminium oxide are classic examples of simple oxides.

Mixed Oxides

The oxides made up of two similar oxides are called mixed oxides. The two simple oxides may have the same element or different elements. Let’s discuss a bit more to understand mixed oxides.

1. When you combine Lead dioxide \(\left( {{\rm{Pb}}{{\rm{O}}_2}} \right)\) and Lead monoxide \(\left( {{\rm{PbO}}} \right)\) it produces Red Lead \(\left( {{\rm{P}}{{\rm{b}}_3}{{\rm{O}}_4}} \right)\)
2. Ferric oxide \(\left( {{\rm{F}}{{\rm{e}}_2}{{\rm{O}}_3}} \right)\) and Ferrous Oxide \(\left( {{\rm{FeO}}} \right)\) form the mixed oxide of Ferro-Ferric Oxide \(\left( {{\rm{F}}{{\rm{e}}_3}{{\rm{O}}_4}} \right).\) This is also called Magnetic Oxide.
3. Ilmenite is a mixed oxide of Ferrous Oxide \(\left( {{\rm{FeO}}} \right)\) and Titanium Oxide \(\left( {{\rm{Ti}}{{\rm{O}}_2}} \right).\)
4. A Strontium Titanate \(\left( {{\rm{SrTi}}{{\rm{O}}_3}} \right)\)  is a mixture of \({\rm{SrO}}\) and \({\rm{Ti}}{{\rm{O}}_2}.\)

Nature of Oxides

Different properties can help us distinguish between the three types of oxides.

Acidic Oxide

An oxide that combines with water to give an acid termed an acidic oxide. They neutralize bases like sodium hydroxide. They dissolve in water to give acids and are called acidic anhydrides. Examples: Sulphur dioxide, \({{\rm{P}}_{\rm{4}}}{{\rm{O}}_{{\rm{10}}}}{\rm{,}}\,{\rm{C}}{{\rm{l}}_{\rm{2}}}{{\rm{O}}_{\rm{7}}}{\rm{,}}\) etc.

For example, sulphur trioxide \(\left( {{\rm{S}}{{\rm{O}}_3}} \right)\) dissolves in water to give sulphuric acid \(({{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}).\)
\({\rm{S}}{{\rm{O}}_3}\; + \;{{\rm{H}}_2}{\rm{O}}\; \to \;{{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}\;\)

Similarly, carbon dioxide \(\left( {{\rm{C}}{{\rm{O}}_2}} \right)\) dissolves in water to produce carbonic acid \(({{\rm{H}}_2}{\rm{C}}{{\rm{O}}_3}).\)
\({\rm{C}}{{\rm{O}}_2}\; + \;{{\rm{H}}_2}{\rm{O}}\; \to \;{{\rm{H}}_2}{\rm{C}}{{\rm{O}}_3}\;\)

Non- metallic oxides are acidic. The acidic oxides of non-metals dissolve in water to form acids. The acidic oxides of non-metals turn the blue litmus solution to red.
As a general rule, only non-metal oxides are acidic, but some metals in the high oxidation state also have an acidic character. For example \({\rm{M}}{{\rm{n}}_2}{{\rm{O}}_7},{\rm{Cr}}{{\rm{O}}_3},{{\rm{V}}_2}{{\rm{O}}_5}.\)

Basic Oxide

The oxides which give a base with water are known as basic oxides. Example: \({\rm{N}}{{\rm{a}}_2}{\rm{O}},{\rm{CaO}},{\rm{BaO,}}\) etc. They neutralize acids like hydrochloric acid. They dissolve in water to give bases. They are called basic anhydrides.
\({\rm{N}}{{\rm{a}}_{\rm{2}}}{\rm{O + }}{{\rm{H}}_{\rm{2}}}{\rm{O}} \to 2{\rm{NaOH}},\)
\({\rm{CaO + }}{{\rm{H}}_2}{\rm{O}} \to {\rm{Ca}}{\left( {{\rm{OH}}} \right)_2}\)
Generally, metallic oxides are basic. When they react with water, the resultant solution turns red litmus blue. Such oxides are called basic oxides.

Amphoteric Oxide

Amphoteric oxides are metallic oxides that show both basic as well as acidic properties. When they react with an acid, they produce salt and water, showing basic properties. However, while responding with alkalis, they form salt and water showing acidic properties. For example, Zinc oxide reacts with concentrated sodium hydroxide and behaves as an acidic oxide but reacts with hydrochloric acid as a basic oxide.

\({\rm{ZnO}} + 2{{\rm{H}}_2}{\rm{O}} + 2{\rm{NaOH}} \to {\rm{N}}{{\rm{a}}_2}{\rm{Zn}}{\left( {{\rm{OH}}} \right)_4} + {{\rm{H}}_2}\)

\({\rm{ZnO}} + 2{\rm{HCl}} \to {\rm{ZnC}}{{\rm{l}}_2}{\rm{ + }}{{\rm{H}}_2}{\rm{O}}\)

Aluminium oxide is another example that reacts with alkalis as well as acids.
Aluminium oxide acts as a basic oxide on reaction with hydrochloric acid.
\(\mathop {{\rm{A}}{{\rm{l}}_2}{{\rm{O}}_3}\left( {\rm{S}} \right)}\limits_{{\rm{Aluminium oxide}}}  + \mathop {6{\rm{HCl}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Hydrochloric acid}}}  \to \mathop {2{\rm{AlC}}{{\rm{l}}_3}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Aluminium chloride}}}  + \mathop {{\rm{3}}{{\rm{H}}_2}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \)

Aluminium oxide acts as an acidic oxide on reaction with sodium hydroxide
\(\mathop {{\rm{A}}{{\rm{l}}_2}{{\rm{O}}_3}\left( {\rm{S}} \right)}\limits_{{\rm{Aluminium}}{\kern 1pt} {\kern 1pt} {\rm{oxide}}}  + \mathop {2{\rm{NaOH}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Sodium}}{\kern 1pt} {\rm{hydroxide}}}  \to \mathop {2{\rm{NaAl}}{{\rm{O}}_2}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Sodium}}{\kern 1pt} {\kern 1pt} {\rm{aliminate}}}  + \mathop {{{\rm{H}}_2}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \)

Neutral Oxides

Some oxides are neither acidic nor basic. Such oxides are known as neutral oxides. Examples of neutral oxides are nitrous oxide and carbon monoxide.

Acidic – Basic Nature of Oxides in a Period and Group

1. The basic nature of the oxides generally increases with an increase in the electropositivity of metal forming oxide. Oxides of all elements of group 1 are basic, and group \(17\) are acidic.
2. Down the group the basic nature of oxides increases and acidic nature decreases. Across a period, the basic nature of oxides gradually decreases, and acidic nature gradually increases.
3. If the same element forms many oxides, acidic nature increases with an increase in the number of oxygen atoms forming the oxide.
4. The most basic oxide is caesium oxide. Caesium hydroxide is the strongest base.
5. The most acidic oxide is chlorine heptoxide. Perchloric acid is the strongest acid.

Metal Oxides

Metal oxides are crystalline solids that contain an oxide anion and a metal cation. They typically react with acids to form salts or with water to form bases.

Preparation of Metal Oxides: There are two ways for preparing metal oxides:

1. Direct method
2. Indirect method

Direct method: It involves the direct combination of the metal with oxygen.
Example:
\(2{\rm{Ca\;}} + {{\rm{O}}_2} \to 2{\rm{CaO}}\)
\(2{\rm{Mg}}\; + {{\rm{O}}_2} \to 2{\rm{MgO}}\)

Indirect method: It involves thermal decomposition of salt of carbonate, hydroxides and nitrates.
Example:
\({\rm{CaC}}{{\rm{O}}_3}\; \to \;{\rm{CaO}} + {\rm{C}}{{\rm{O}}_2}\)
\(2{\rm{Pb}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}\; \to 2{\rm{PbO}} + 4{\rm{N}}{{\rm{O}}_2} + {{\rm{O}}_2}\)

Reaction of Metal Oxides with Acids:

An oxide of the metal, when it reacts with acids, forms salts. Examples are Magnesium Oxide, Calcium Oxide, Sodium Oxide, Potassium Oxide.
\({\rm{MgO}} + 2{\rm{HCl}} \to {\rm{MgC}}{{\rm{l}}_2} + {{\rm{H}}_2}{\rm{O}}\)
\({\rm{CaO}} + {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{CaS}}{{\rm{O}}_4} + {{\rm{H}}_2}{\rm{O}}\)

Reaction of Metals with Metal Oxide:

This type of reaction is an apparent ‘competition’ between two metal atoms for oxygen atoms. The more reactive metal atoms will remove the oxygen from the less reactive metal oxide. This results in a transfer of oxygen from the less reactive metal oxide to the more reactive metal. The rule is that the more reactive metal always takes the oxygen away from the less reactive metal oxide.
\({\rm{Mg + CuO}} \to {\rm{Cu + MgO}}\)

Uses of Oxides

1. Oxides are used in the preparation of salts in the laboratory i.e.,
   \({\rm{CaO + }}{{\rm{H}}_{\rm{2}}}{\rm{S}}{{\rm{O}}_{\rm{4}}} \to {\rm{CaS}}{{\rm{O}}_{\rm{4}}}{\rm{ + }}{{\rm{H}}_{\rm{2}}}{\rm{O}}\)
2. It used in the formation of slag.
3. Oxides used as a drying agent.
4. Oxides used in the manufacture of motors.

Summary

From this article, we can conclude that most elements form oxides. To be an oxide, the oxidation state of oxygen must be \(-2,\) and the oxygen must act as an anion. We have learned the classification, nature of oxides, and their uses.

Frequently Asked Questions

Q.1. What is oxide used for?
Ans: Oxides used in the preparation of salts in the laboratory, in the formation of slag. Some oxides used as a drying agent.

Q.2. What are the basic oxides? Give an example?
Ans: The oxides which give a base with water are known as basic oxides.
Example: \({\rm{N}}{{\rm{a}}_{\rm{2}}}{\rm{O,\;CaO,\;BaO}}\) etc.

Q.3. What are two examples of oxides?
Ans: The two examples of oxides are Sulphur dioxide \(\left( {{\rm{S}}{{\rm{O}}_2}} \right),\) Sodium oxide \(\left( {{\rm{N}}{{\rm{a}}_2}{\rm{O}}} \right).\)

Q.4. How is an oxide formed?
Ans: A binary compound of oxygen with another element is called oxide. Oxygen is highly reactive. They react with metals and non-metals and form oxides. The amine forms amine oxides, Sulphides forms sulfoxides, respectively, in which the oxygen atom covalently bonded to the Nitrogen or Sulphur atom.

Q.5. With the help of examples, describe how metal oxide differs from non-metal oxides?
Ans: Metal oxides are basic and turn red litmus blue. For example, magnesium oxide.
Non-metal oxides are acidic or neutral. The acidic oxides turn blue litmus red. For example, sulphur dioxide.

Q.6. How metal oxides react with acids. Give an example?
Ans: An oxide of the metal, when it reacts with acids, forms salt and water. Examples are Magnesium oxide, Calcium oxide, Sodium oxide and Potassium oxide.
\({\rm{MgO}} + 2{\rm{HCl}} \to {\rm{MgC}}{{\rm{l}}_2} + {{\rm{H}}_2}{\rm{O}}\)
\({\rm{CaO}} + {{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to {\rm{CaS}}{{\rm{O}}_4} + {{\rm{H}}_2}{\rm{O}}\)

Q.7. What are the neutral oxides, and give some examples?
Ans: Some oxides are neither acidic nor basic. Such oxides are known as neutral oxides. Examples of neutral oxides are nitrous oxide, water, nitric oxide, and carbon monoxide.

Q.8. What type of oxide is carbon dioxide?
Ans: Carbon dioxide is an acidic oxide. The acidic oxides of non-metals dissolve in water to form acids.

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