Chemical Properties of Metals and Non-metals: Metals and non-metals are the elements that surround us. As a result, it’s crucial to determine if a given element is a metal or a non-metal. The two sorts of materials are metals and non-metals. Non-metals, such as sulphur and phosphorus, are insulators, whereas metals, such as aluminium and copper, have excellent electrical and thermal conductivity.

The metals are present on the left-hand side, in the middle and at the bottom of the periodic table. On the other hand, non-metals are placed on the extreme right-hand side. In this article, we will discuss about chemical properties of metals and non-metals in detail. Scroll down to learn more!

Metals and Non-metals

At present, one hundred fourteen \((114)\) elements are known to us. These include metals, non-metals and metalloids. The properties and characteristics of metalloids are intermediate to metals and non-metals.

Learn About Chemical Properties Of Metals Here

Metals

Out of \(114\) elements known so far, most of them are metals. The knowledge of metals and their uses has been known since ancient times. Copper, silver and gold coins have been used as currency in a different era. The use of iron metals in forming weapons is reported in history. Copper vessels have been used in domestic cookware for a long.

Metals are the electropositive elements that tend to lose one or more of their valence electrons attaining octet and form cations.

The metals usually have \(1,\,2\,{\rm{\;or\;\, 3}}\) electrons (some have \(4\) electrons) in their valence shell. Metals lose their valency electrons and form cations (positive ions) on gaining energy from an external source. e.g.,

\(\mathop {{\rm{Na}}}\limits_{\left( {2,8,1} \right)} {\mkern 1mu} \, \to \,\mathop {{\mkern 1mu} {\rm{N}}{{\rm{a}}^{\rm{ + }}}}\limits_{\left( {{\rm{2}},{\rm{8}}} \right){\mkern 1mu} } \,{\rm{ + }}\,{\mkern 1mu} {\rm{e}}\)

\(\mathop {{\rm{Mg}}}\limits_{\left( {2,8,2} \right)} \,{\mkern 1mu} \to \,{\mkern 1mu} \mathop {{\rm{M}}{{\rm{g}}^{2 + }}}\limits_{\left( {{\rm{2}},{\rm{8}}} \right)} \,{\mkern 1mu} {\rm{ + }}{\mkern 1mu} \,2{\rm{e}}\)

The number of electrons lost by an atom of a metal to form the positive ion is called the valency of that metal. The ion of the metal carries the same number of unit positive charges. For example, sodium, magnesium and aluminium atoms lose \(1,\,\,2\) and \(3\) electrons respectively to form \({\rm{N}}{{\rm{a}}^ + },{\rm{M}}{{\rm{g}}^{2 + }}\) and \({\rm{A}}{{\rm{l}}^{3 + }}\) and show valency of \(1,\,2\) and \(3,\) respectively.

Non-Metals

Non-metals are the electronegative elements that tend to form anion (negative ions) by gaining one or more electrons.

The atoms of non-metals usually have \(4,\,5\,,\,6\,\,{\rm{or\;\, 7}}\) electrons in their valency shell. To gain a stable electronic configuration to attain octet, they tend to gain electrons and form anions (negative ions). e.g.,

\({\rm{N}} + \,3{\rm{e}} \to {{\rm{N}}^{3 – }}\)

\({\rm{O}} + \,2{\rm{e}} \to {{\rm{O}}^{2 – }}\)

The number of electrons gained by an atom of a non-metal to form the negative ion is called the valency of that non-metal. The ion thus formed carries the same number of unit negative charges. For example, an oxygen atom gains \(2\) electrons to form \({{\rm{O}}^{2 – }}\) ions. Therefore, the valency of oxygen is \(2\) and \({{\rm{O}}^{2 – }}\) ion carries two negative charges.

Position of Metals and Non-Metals in the Periodic Table

In the periodic table, metals are placed on the left-hand side, in the middle and at the bottom of the periodic table, whereas non-metals are placed on the extreme right-hand side of the periodic table. Hydrogen, although a non-metal but also resembles alkali metals, is an exception that is placed on the extreme left of the periodic table. Only a few metalloids exist, e.g., \({\rm{B}},{\mkern 1mu} \,{\rm{Si}},{\mkern 1mu} \,{\rm{Ge}},{\mkern 1mu} \,{\rm{As}},{\mkern 1mu} \,{\rm{Sb}},{\mkern 1mu} \,{\rm{Te}},\,{\rm{Po}}\) and \({\rm{At}}{\rm{.}}\)

Chemical Properties of Metals

The chemical properties of metals include the following reactions:
(a) Reaction of metals with oxygen
(b) Reaction of metals with water
(c) Reaction of metals with dilute acids
(d) Reaction of metals with salt solutions
(e) Reaction of metals with chlorine
(f) Reaction of metals with hydrogen

1. Reaction of Metals with Oxygen

Most of the metals react with oxygen \(({{\rm{O}}_2})\) in the air to form metal oxides.

\({\rm{Metal}}\, + \,{\rm{Oxygen}}\, \to {\rm{Metal}}\,{\rm{oxides}}\)

The metal oxide formed is basic in nature. So, these oxides turn red litmus solution blue. Different metals react with oxygen at different temperatures. Some metals react with oxygen at room temperature, some react with oxygen on heating, and some react with oxygen only on strong heating.

Example 1: Reactive metals like sodium and potassium react with oxygen at room temperature to form oxides.

\(\mathop {{\rm{4Na}}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}} \,{\rm{ + }}\,\mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} \, \to \,\mathop {{\rm{2N}}{{\rm{a}}_{\rm{2}}}{\rm{O}}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}\,\,{\rm{oxide}}}\)

\(\mathop {{\rm{4K}}\left( {\rm{s}} \right)}\limits_{{\rm{Potassium}}} \,{\rm{ + }}\,\mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} \, \to \,\mathop {{\rm{2}}{{\rm{K}}_{\rm{2}}}{\rm{O}}\left( {\rm{s}} \right)}\limits_{{\rm{Potassium}}\,\,{\rm{oxide}}} \)

Example 2: Magnesium metal does not react with oxygen at room temperature. It burns with dazzling white light on heating in air forming magnesium oxide.

\(\mathop {{\rm{2Mg}}\left( {\rm{s}} \right)}\limits_{{\rm{Magnesium}}} \,{\rm{ + }}\,\mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} \, \to \,\mathop {{\rm{2MgO}}\left( {\rm{s}} \right)}\limits_{{\rm{Magnesium}}\,\,{\rm{oxide}}} \)

Similarly, aluminium metal burns on heating in the air to form aluminium oxide.

\(\mathop {{\rm{4Al}}\left( {\rm{s}} \right)}\limits_{{\rm{Aluminium}}} \,{\rm{ + }}\,\mathop {{\rm{3}}{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} \, \to \,\mathop {{\rm{2A}}{{\rm{l}}_2}{{\rm{O}}_3}\left( {\rm{g}} \right)}\limits_{{\rm{Aluminium}}\,\,{\rm{oxide}}} \)

Example 3: Zinc metal burns on strong heating in air to form zinc oxide.

\(\mathop {{\rm{2Zn}}\left( {\rm{s}} \right)}\limits_{{\rm{Zinc}}} \,{\rm{ + }}\,\mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} \, \to \,\mathop {{\rm{2ZnO}}\left( {\rm{s}} \right)}\limits_{{\rm{Zinc}}\,\,{\rm{oxide}}} \)

2. Solubility of Metal Oxide in Water

The oxides formed by most of the metals are basic in nature. For example, \({\rm{N}}{{\rm{a}}_2}{\rm{O}},\,{{\rm{K}}_2}{\rm{O}},\,{\rm{MgO,}}\,{\rm{CuO}}\) are basic in nature. Their reaction to water shows the basic nature of these oxides.

\({\rm{Metal}}\,{\rm{Oxide}}\, + \,{\rm{Water}}\, \to \,{\rm{Metal}}\,{\rm{Hydroxide}}\)

\(\mathop {{\rm{N}}{{\rm{a}}_2}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}\,\,{\rm{Oxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \, \to \,\mathop {{\rm{2NaOH}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Sodium  hydroxide}}} \)

\(\mathop {{{\rm{K}}_2}{\rm{O}}\left( {\rm{s}} \right)}\limits_{{\rm{Potassium}}\,\,{\rm{Oxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \, \to \,\mathop {{\rm{2KOH}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Potassium  hydroxide}}} \)

\(\mathop {{\rm{MgO}}\left( {\rm{s}} \right)}\limits_{{\rm{Magnesium}}\,\,{\rm{Oxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \, \to \,\mathop {{\rm{Mg}}{{\left( {{\rm{OH}}} \right)}_2}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Magnesium  hydroxide}}} \)

Sodium and potassium oxides are completely soluble in water, while magnesium oxide is partially soluble. The basic nature of hydroxides formed by the reaction of these oxides with water can be checked by red litmus solution.

Certain metal oxides like \({\rm{A}}{{\rm{l}}_2}{{\rm{O}}_3}\) and \({\rm{ZnO}}\) are amphoteric and show both acidic as well as basic behaviour. They react with both acids and bases to form salts.

3. Reaction of Metals with Water

Metals react with water to produce metal oxide and hydrogen gas. Water-soluble metal oxides dissolve in water to form metal hydroxide.

(i) \({\rm{Metal}}\, + \,{\rm{Cold}}\,{\rm{Water}}\, \to \,{\rm{Metal}}\,{\rm{Hydroxide}}\, + \,{\rm{Hydrogen}}\,\)

(ii) \({\rm{Metal}}\, + {\rm{Steam}}\,\, \to \,{\rm{Metal}}\,{\rm{oxide}}\, + {\rm{Hydrogen}}\)

The reactivity of each metal with water is different. Some reactive metals like sodium and potassium react with cold water, whereas some metals like magnesium do not react with cold water but react with hot water. Some metals do not react with water at all. They can react with steam. Certain metals like lead, copper, silver, gold and platinum do not show any reaction with water at any condition (neither water nor steam).

If the metal oxide is soluble in water, metal hydroxide and hydrogen gas are formed.

a. Sodium and potassium react violently with cold water. The reaction is highly exothermic, and the evolved hydrogen gas immediately catches fire.

\({\rm{2K}}\,{\rm{ + }}\,{\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O}}\, \to \,{\rm{2 KOH}}\,{\rm{ + }}\,{{\rm{H}}_{\rm{2}}}\)

\({\rm{2Na}}\,{\rm{ + }}\,{\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O}}\, \to \,{\rm{2NaOH}}\,{\rm{ + }}\,{{\rm{H}}_{\rm{2}}}\)

(b) Magnesium reacts with hot water to form magnesium hydroxide and hydrogen.

\(\mathop {{\rm{Mg}}\left( {\rm{s}} \right)}\limits_{{\rm{Magnesium}}} \,{\rm{ + }}\,\mathop {{\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}\,\left( {{\rm{hot}}} \right)} \, \to \,\mathop {{\rm{Mg}}{{\left( {{\rm{OH}}} \right)}_{\rm{2}}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Magnesium}}\,\,{\rm{hydroxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \)

(c) Calcium also reacts with cold water, but the reaction is not so violent. The reaction is also exothermic, but less heat is evolved, and the evolved hydrogen gas does not burn.

\({\rm{Ca}}{\mkern 1mu} + 2{{\rm{H}}_2}{\rm{O}} \to {\rm{Ca}}{\left( {{\rm{OH}}} \right)_2} + {{\rm{H}}_2}\)

Metals like zinc and aluminium react only with steam to form their corresponding oxides and hydrogen.

\(\mathop {{\rm{2Al}}\left( {\rm{s}} \right)}\limits_{{\rm{Aluminium}}} \,{\rm{ + }}\,\mathop {{\rm{3}}{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{g}} \right)}\limits_{{\rm{Steam}}} \, \to \,\mathop {{\rm{A}}{{\rm{l}}_2}{{\rm{O}}_3}\left( {\rm{s}} \right)}\limits_{{\rm{Aluminium}}\,\,{\rm{oxide}}} \,{\rm{ + }}\,\mathop {{\rm{3}}{{\rm{H}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \)

\(\mathop {{\rm{Zn}}\left( {\rm{s}} \right)}\limits_{{\rm{Zinc}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{g}} \right)}\limits_{{\rm{Steam}}} \, \to \,\mathop {{\rm{ZnO}}\left( {\rm{s}} \right)}\limits_{{\rm{Zinc}}\,\,{\rm{oxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \)

Under normal conditions, iron metal does not react with water. The reactions occur when steam is passed over red-hot iron, and the products are iron (II, III) oxides and hydrogen.

\(\mathop {{\rm{3Fe}}\left( {\rm{s}} \right)}\limits_{{\rm{Iron}}} \,{\rm{ + }}\,\mathop {{\rm{4}}{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{g}} \right)}\limits_{{\rm{Steam}}} \, \to \,\mathop {{\rm{F}}{{\rm{e}}_3}{{\rm{O}}_4}\left( {\rm{s}} \right)}\limits_{{\rm{Iron }}\left( {{\rm{II,}}\,{\rm{III}}} \right)\,\,{\rm{oxide}}} \,{\rm{ + }}\,\mathop {{\rm{4}}{{\rm{H}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \)

The order of reactivities of different metals with water decreases in the order:

\({\rm{Na}}\, > {\rm{Mg}}\,{\rm{ > Zn}}\,{\rm{ > Fe}}\,{\rm{ > Cu}}\)

Note: Metals like copper, lead, silver, gold and platinum do not react with water at all.

4. Reaction of Metals with Dilute Acids

Some reactive metals like sodium, magnesium, aluminium, zinc and iron react with dilute acids to form salt and hydrogen gas. This is due to the displacement of hydrogen gas from dilute acids by the reactive metal.

\({\rm{Metal}}\,{\rm{ + }}\,{\rm{Dilute}}\,{\rm{acid}}\, \to \,{\rm{Metal}}\,{\rm{salt}}\,{\rm{ + }}\,{\rm{Hydrogen}}\)

The vigour of the reaction depends on the chemical reactivity of the metal.

Some metals of less reactivity like copper, silver, gold and platinum do not react with dilute acids.

(i) Sodium reacts with dilute hydrochloric acid vigorously to form sodium chloride and hydrogen gas.

\(2{\rm{Na}}\,{\rm{ + }}\,{\rm{2HCl}} \to {\rm{2NaCl}}\,{\rm{ + }}{{\rm{H}}_2}\)

This is because sodium metal is highly reactive.

(ii) Copper does not react with dilute hydrochloric acid. This shows that copper is less reactive than zinc.

\({\rm{Cu}}\,{\rm{ + }}\,{\rm{HCl}} \to {\rm{No\, reaction}}\)

In fact, all weak acids react with metals slowly.

Metals below hydrogen in the activity series of metals do not liberate hydrogen from acids. Still, they react with oxidizing acids (conc. \({{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}\,{\rm{or\, HN}}{{\rm{O}}_3}\) ) to form water, salt of the metal and the reduction products of the oxidizing acids. For example, copper reacts with dilute nitric acid to form water, copper nitrate and nitric oxide.

\(3{\mkern 1mu} {\rm{Cu}}{\mkern 1mu} {\rm{ + }}{\mkern 1mu} {\rm{8HN}}{{\rm{O}}_3}\, \to \,{\mkern 1mu} 3{\mkern 1mu} {\rm{Cu}}\left( {{\rm{N}}{{\rm{O}}_3}} \right){{\mkern 1mu} _2} + {\mkern 1mu} 4{{\rm{H}}_2}{\rm{O}}{\mkern 1mu} {\rm{ + }}{\mkern 1mu} {\rm{2NO}}\)

Copper reacts with hot concentrated sulphuric acid to produce copper sulphate, sulphur dioxide and water.

\({\rm{Cu}}\,{\rm{ + }}\,2{{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4} \to \,\,{\rm{CuS}}{{\rm{O}}_4}{\rm{ + S}}{{\rm{O}}_2}{\rm{ + 2}}{{\rm{H}}_2}{\rm{O}}\)

5. Reaction of Metals with Salt Solutions

It is clear after studying the reaction of metals with oxygen, water and dilute acids that some metals are more reactive, and some are less reactive. When a more reactive metal is put in the salt solution of the less reactive metal, the less reactive metal is displaced from its salt solution by the more reactive metal.

\({\rm{Metal}}\,{\rm{A}}\, + \,{\rm{Salt}}\,{\rm{solution}}\,{\rm{of}}\,{\rm{B}} \to \,{\rm{Salt}}\,{\rm{solution A}}\,{\rm{ + }}\,{\rm{Metal}}\,{\rm{B}}\)

Or,

\({\rm{A}}\, + \,\,{\rm{BC}}\,\, \to \,\,{\rm{B}}\,{\rm{ + }}\,{\rm{AC}}\)

Where, \({\rm{A}}\,\) is a more reactive metal than \({\rm{A}}\,\).

Only a more reactive metal can displace less reactive metal from its salt solution.

Example: When zinc granules are put in copper sulphate solution, its blue colour fades and becomes colourless. Also, a reddish-brown coating of copper is observed on the surface of zinc. This is because zinc is more reactive than copper and displaces copper from copper sulphate solution to form colourless zinc sulphate and copper deposited on the zinc surface.

\(\mathop {{\rm{Zn}}\left( {\rm{s}} \right)}\limits_{{\rm{Zinc}}} \,{\rm{ + }}\,\mathop {{\rm{CuS}}{{\rm{O}}_4}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Copper}}\,\,{\rm{sulphate}}} \, \to \,\mathop {{\rm{ZnS}}{{\rm{O}}_4}\left( {\rm{aq}} \right)}\limits_{{\rm{Zinc}}\,\,{\rm{sulphate}}} \,{\rm{ + }}\,\mathop {{\rm{Cu}}\left( {\rm{s}} \right)}\limits_{{\rm{Copper}}} \)

6. Reaction of Metals with Chlorine

Metals react with chlorine to form metal chlorides.

\({\rm{Metal}}\, + \,{\rm{Chlorine}}\, \to \,{\rm{Metal}}\,{\rm{Chloride}}\)

For example,

\(\mathop {{\rm{2Na}}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}} \,{\rm{ + }}\,\mathop {{\rm{C}}{{\rm{l}}_2}\left( {\rm{g}} \right)}\limits_{{\rm{Chloride}}} \, \to \,\mathop {{\rm{2NaCl}}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}\,\,{\rm{chloride}}} \)

\(\mathop {{\rm{Ca}}\left( {\rm{s}} \right)}\limits_{{\rm{Calcium}}} {\mkern 1mu} {\rm{ + }}{\mkern 1mu} \mathop {{\rm{C}}{{\rm{l}}_2}\left( {\rm{g}} \right)}\limits_{{\rm{Chloride}}} {\mkern 1mu}  \to \,{\mkern 1mu} \mathop {{\rm{CaC}}{{\rm{l}}_2}\left( {\rm{s}} \right)}\limits_{{\rm{Calcium}}\,\,{\rm{chloride}}} \)

7. Reaction with Hydrogen

All metals do not react with hydrogen. Only a few reactive metals like sodium, potassium, magnesium and calcium react with hydrogen to form metal hydrides.

\({\rm{Metal}}\, + \,{\rm{Hydrogen}}\, \to \,{\rm{Metal}}\,{\rm{Hydride}}\)

For example,

\(\mathop {{\rm{2Na}}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}} {\mkern 1mu} \,{\rm{ + }}{\mkern 1mu} \mathop {\,{{\rm{H}}_2}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \,{\mkern 1mu}  \to \,{\mkern 1mu} \mathop {2{\rm{NaH}}\left( {\rm{s}} \right)}\limits_{{\rm{Sodium}}\,\,{\rm{hydride}}} \)

\(\mathop {{\rm{Ca}}\left( {\rm{s}} \right)}\limits_{{\rm{Calcium}}} {\mkern 1mu} \,{\rm{ + }}{\mkern 1mu} \mathop {\,{{\rm{H}}_2}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \,{\mkern 1mu}  \to \,{\mkern 1mu} \mathop {{\rm{Ca}}{{\rm{H}}_2}\left( {\rm{s}} \right)}\limits_{{\rm{Calcium}}\,\,{\rm{hydride}}} \)

Chemical Properties of Non-metals

Non-metals have the following important chemical properties:

1. Reaction with Oxygen

Non-metals react with oxygen to form oxides.

\({\rm{Non\, Metal}}\, + \,{\rm{Oxygen}}\, \to \,{\rm{Non\, Metal}}\,{\rm{Oxide}}\)

Example 1: Carbon burns in a supply of oxygen to form carbon dioxide.

\(\mathop {{\rm{C}}\left( {\rm{s}} \right)}\limits_{{\rm{Carbon}}} \,{\mkern 1mu} {\mkern 1mu} {\rm{ + }}\,{\mkern 1mu} \mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} {\mkern 1mu} {\mkern 1mu} \, \to \,\mathop {{\rm{C}}{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Carbon}}\,\,{\rm{dioxide}}} \)

When the supply of oxygen is limited, carbon burns to form carbon monoxide.

\(\mathop {{\rm{2C}}\left( {\rm{s}} \right)}\limits_{{\rm{Carbon}}} {\mkern 1mu} \, + \,\mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} \, \to {\mkern 1mu} \,\mathop {{\rm{2CO}}\left( {\rm{g}} \right)}\limits_{{\rm{Carbon}}{\kern 1pt} {\kern 1pt} {\rm{monoxide}}} \)

Example 2: Sulphur is a yellow solid that burns in oxygen to form sulphur dioxide.

\(\mathop {{\rm{S}}\left( {\rm{s}} \right)}\limits_{{\rm{Sulphur}}} \,{\mkern 1mu} {\mkern 1mu} {\rm{ + }}\,{\mkern 1mu} \mathop {{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Oxygen}}} {\mkern 1mu} {\mkern 1mu} \, \to \,\mathop {{\rm{S}}{{\rm{O}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Sulphur}}\,\,{\rm{dioxide}}} \)

Nature of Oxide

The oxides formed by most of the non-metals are acidic in nature. For example, CO₂ and SO₂ are acidic in nature. Their reaction with water shows the acidic nature of these oxides.

\({\rm{Non}}\, – {\rm{Metal}}\,{\rm{Oxide}}\, + \,{\rm{Water}}\, \to \,{\rm{Acid}}\)

Carbon dioxide dissolves in water to form carbonic acid. Thus, it is acidic in nature.

\(\mathop {{\rm{C}}{{\rm{O}}_{\rm{2}}}}\limits_{{\rm{Carbon}}\,\,{\rm{dioxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \, \to \,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{C}}{{\rm{O}}_{\rm{3}}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Carbonic}}\,\,{\rm{acid}}} \)

Sulphur dioxide also dissolves in water to form an acid called sulphurous acid. Thus, it is also acidic in nature.

\(\mathop {{\rm{S}}{{\rm{O}}_{\rm{2}}}}\limits_{{\rm{Sulphur}}\,\,{\rm{dioxide}}} \,{\rm{ + }}\,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}\left( {\rm{l}} \right)}\limits_{{\rm{Water}}} \, \to \,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{S}}{{\rm{O}}_{\rm{3}}}\left( {{\rm{aq}}} \right)}\limits_{{\rm{Sulphrous}}\,\,{\rm{acid}}} \)

Certain non-metal oxides like \({\rm{CO}},{\mkern 1mu} {\rm{NO}}\) and \({{\rm{N}}_2}{\rm{O}}\) are neutral in nature. These oxides do not show acidic or basic behaviour.

2. Reaction with Water

Unlike metals, non-metals do not react with water to evolve hydrogen gas. However, carbon reacts with steam at high temperatures to form a mixture of \({\rm{CO \;and\;}}\,{{\rm{H}}_2}\) known as water gas.

\(\mathop {\rm{C}}\limits_{{\rm{Coke}}} \,{\rm{ + }}{\mkern 1mu} \,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{O}}}\limits_{{\rm{Steam}}} \, \to \,{\mkern 1mu} {\mkern 1mu} {\mkern 1mu} {\mkern 1mu} \underbrace {{\rm{CO}} + {{\rm{H}}_2}}_{{\rm{Water}}{\kern 1pt} {\rm{gas}}}\)

3. Reaction with Dilute Acids

Non-metals do not react with dilute acids. They do not displace hydrogen from dilute acids. However, non-metals can be oxidized in the presence of conc. HNO₃ and conc. H₂SO₄ due to the strong oxidizing nature of these acids.

4. Reaction of metals with salt solution

The reaction of non-metals with salt solutions is significant only in the case of halogens \(({{\rm{F}}_{\rm{2}}}{\rm{,}}\,{\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{,}}\,{\rm{B}}{{\rm{r}}_{\rm{2}}}\,{\rm{and}}\,{{\rm{I}}_{\rm{2}}})\) A more reactive halogen can displace a less reactive halogen from its salt solution.

\(\mathop {{\rm{Sodium}}\,{\rm{bromide}}}\limits^{2{\rm{NaBr(aq)}}} \,\, + \,\mathop {{\rm{Chlorine}}}\limits^{{\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{(g}})}  \to \mathop {{\rm{Sodium}}\,{\rm{chloride}}}\limits^{2{\rm{NaCl}}({\rm{aq}})} \,\, + \,\mathop {{\rm{Bromine}}}\limits^{{\rm{B}}{{\rm{r}}_2}({\rm{aq}})} \,\)

5. Reaction with Chlorine

Non-metals also react with chlorine to form chlorides which are usually liquids or gases.

Example: Hydrogen reacts with chlorine to form hydrogen chloride, which is gas in a dry state.

\(\mathop {{\rm{Hydrogen}}}\limits^{{{\rm{H}}_{\rm{2}}}{\rm{(g)}}} \,\, + \,\mathop {{\rm{Chlorine}}}\limits^{{\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{(g}})\,\,\,\,\,\,\,\,\,\, \to } \,\,\mathop {{\rm{Hydrogen}}\,{\rm{chloride}}}\limits^{2{\rm{HCl}}({\rm{aq}})} \,\)

Hydrogen chloride gas dissolves in water to form hydrochloric acid.

6. Reaction with Hydrogen

Non-metals also react with hydrogen to form hydrides. For example, sulphur and nitrogen react with hydrogen to form hydrogen sulphide and ammonia, respectively.

\(\mathop {{\rm{S}}\left( {\rm{l}} \right)}\limits_{{\rm{Sulphur}}} \,{\rm{ + }}{\mkern 1mu} \mathop {{{\rm{H}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} \,\, \to \,\mathop {{{\rm{H}}_{\rm{2}}}{\rm{S}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}{\kern 1pt} {\kern 1pt} {\rm{sulphide}}} \)

\(\mathop {{{\rm{N}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Nitrogen}}} {\mkern 1mu} {\mkern 1mu} \,{\rm{ + }}\,\mathop {{\rm{3}}{{\rm{H}}_{\rm{2}}}\left( {\rm{g}} \right)}\limits_{{\rm{Hydrogen}}} {\mkern 1mu} \, \to \,{\mkern 1mu} \mathop {{\rm{2N}}{{\rm{H}}_3}\left( {\rm{g}} \right)}\limits_{{\rm{Ammonia}}} {\mkern 1mu} \)

Difference between Metals and Non-metals Based on their Chemical Properties

The difference between metals and non-metals according to their chemical properties are as follows:

S.No	Reactionwith	Metals	Non- Metals
1.	Oxygen	Metal oxides are formed, which are generally basic in nature.	Non-metal oxides are formed, which are generally acidic in nature
2.	Water	Metal hydroxides are formed.	Generally, do not react
3.	Dilute acid	Metals more reactive than hydrogen displace hydrogen from dilute acids to form metal salts and hydrogen gas	Do not react with dilute acids
4.	Salt solution	More reactive metal displaces less reactive metal from its salt solution	Non-metal that is more reactive displaces less reactive non-metal from its salt solution
5.	Chlorine	Metal chlorides are formed which are ionic in nature	Non- metal Chlorides are formed, which are covalent in nature.
6.	Hydrogen	Reactive metals from ionic metal hydrides.	Covalent hydrides are formed

Summary

Metals and non-metals are necessary components of our daily existence. We can’t live without some non-metals, such as oxygen, and we’d have a hard time living without metals. Elements are distinguished as either metals or non-metals based on their properties. In this article, we have learned about the reaction of metal and non-metals with oxygen, water, acid, chlorides, and hydrogen.

FAQs on Chemical Properties of Metals and Non Metals

Q.1. What are the five chemical properties of metals?
Ans: The five chemical properties of metals are:
a. Metals form basic oxides.
b. Metals displace hydrogen from water or steam.
c. Metals above hydrogen in the reactivity series displace hydrogen from dilute acids.
d. Metals behave as reducing agents.
e. Metals form ionic hydrides.

Q.2. What are the five chemical properties of non-metals?
Ans: The five chemical properties of non-metals are
a. Non-metals form acidic or neutral oxides.
b. Non-metals do not displace hydrogen from water or steam.
c. Non-metals do not displace hydrogen from dilute acids and hence do not react with dilute acids.
d. Non-metals form covalent chlorides.
e. Non-metals behave as oxidising agents.

Q.3. How do the chemical properties of metals differ from non-metals?
Ans:

Reaction with	Metals	Non-Metal
Oxygen	Metal oxides are formed, which are generally basic in nature.	Non-Metal oxides are formed which are generally acidic in nature.
Water	Metal hydroxides are formed.	Generally, do not react.
Dilute acid	Metals more reactive than hydrogen displace hydrogen from dilute acids to form metal salts and hydrogen gas.	Do not react with dilute acids.
Salt solution	Metal that is more reactive displaces less reactive metal from its salt solution.	Non-metal that is more reactive displaces less reactive-metal from its salt solution.
Chlorine	Metal chlorides are formed, which are ionic in nature.	Non-metal chlorides are formed which are covalent in nature.
Hydrogen	Ionic metal hydrides are formed by reactive metals.	Covalent hydrides are formed.

Q.4. Name a few metals and non-metals, which are in a liquid state at room temperature.
Ans: Bromine is the only liquid non-metal at room temperature. Mercury is a liquid metal at room temperature.

Q.5. Why does iron corrode?
Ans: The chemical reaction known as oxidation causes iron and iron alloys to rust. Oxidation happens when the iron is exposed to moisture or oxygen. Iron is transformed into iron oxide during this chemical process. Rusting is caused by both oxygen and moisture.

Q.6. Is colour a chemical property?
Ans: No, physical property is a characteristic of a substance that can be observed or measured without changing the identity of the substance. Physical properties include colour, density, hardness, and melting and boiling points. A chemical property describes the ability of a substance to undergo a specific chemical change.