Chapter – 3

Metals and Non-metals

In this post we have given the detailed notes of class 10 Science Chapter 3 (Metals and Non-metals) in English. These notes are useful for the students who are going to appear in class 10 board exams.

Chapter 3: Metals and Non-metals

Metals 🔩

Substances that are hard, shiny, malleable, ductile, sonorous, and good conductors of heat and electricity are called metals. They are generally elements that are lustrous and produce a sound when struck, for example — iron, tin, copper, gold, zinc, steel, etc. We see different kinds of materials around us and use many of them in our daily life. But have you ever thought about what these materials are made of?

Examples: Sodium (Na), potassium (K), magnesium (Mg), iron (Fe), aluminium (Al), calcium (Ca), and barium (Ba) are metals.

Physical Properties of Metals

• Metals are solid and lustrous.
• They are good conductors of heat and electricity.
• Metals are ductile.
• Metals are malleable.
• Metals are sonorous.

Malleability

• Some metals can be beaten into thin sheets; this property is called malleability. 🔨

Ductility

• The ability of a metal to be drawn into thin wires is called ductility. Silver and copper are the best conductors of heat. In comparison, lead and mercury are poor conductors of heat.

Full Form of PVC

• Materials such as polyvinyl chloride (PVC) and rubber are poor conductors of both heat and electricity.

Sonorous

• Being sonorous is a physical property of metals. Because of this property, metals produce a sound when struck. School bells are made using this property of metals. 🔔

Non-metals

• Substances that are soft, dull, brittle, and poor conductors of heat and electricity, and which are not sonorous, are called non-metals. Being soft, they break into pieces with a light blow of a hammer, are not sonorous, and are poor conductors of heat and electricity. Examples of non-metals are coal, sulphur, oxygen, phosphorus, etc.

Examples: Oxygen (O), hydrogen (H), nitrogen (N), sulphur (S), phosphorus (P), fluorine (F), chlorine (Cl), bromine (Br), and iodine (I) are non-metals.

Physical Properties of Non-metals

• Non-metals are not solid and lustrous.
• They are not good conductors of heat and electricity.
• Non-metals are not ductile.
• Non-metals are not malleable.
• Non-metals are not sonorous, i.e., they do not produce sound when struck.

Non-metals, except for bromine, are either solids or gases; bromine is a liquid.

Some Other Properties of Metals and Non-metals

• All metals, except mercury, are found in the solid state at room temperature.
• Mercury is found in the liquid state at room temperature.
• Gallium and caesium are two metals that have very low melting points; they melt if you just hold them in your palm.
• Iodine is a non-metal, but it is lustrous.
• Alkali metals (lithium, sodium, and potassium) are so soft that they can be cut with a knife. They have a low density and low melting point.

Carbon and Its Allotropes 💎

• Carbon is a non-metal that is found in different forms. Each of these forms is called an allotrope of carbon.

Allotropes of Carbon

• Diamond
• Graphite
• Buckminsterfullerene

(i) Diamond: This is an allotrope of carbon and is the hardest substance known so far. It has a very high boiling point and melting point.

(ii) Graphite: This is another allotrope of carbon which is a very good conductor of electricity.

(iii) Buckminsterfullerene: This is another allotrope of carbon made up of 60 carbon atoms. Its structure is like that of a football.

Note: Most non-metals form acidic oxides when dissolved in water, while metals form basic oxides when dissolved in water.

Chemical Properties of Metals

All metals combine with oxygen to form the corresponding metal oxide.

Metal + Oxygen → Metal oxide

For example, when copper is heated in air, it reacts with oxygen to form copper (II) oxide, which is a black oxide.

2Cu + O₂ → 2CuO

(Copper) (Oxygen) (Copper (II) oxide)

Similarly, aluminium forms aluminium oxide.

4Al + 3O₂ → 2Al₂O₃

(Aluminium) (Aluminium oxide)

Amphoteric Oxides

• Some metal oxides, such as aluminium oxide and zinc oxide, show both acidic and basic behaviour. Metal oxides that react with both acids and bases to form a salt and water are called amphoteric oxides.

Example: Aluminium oxide and zinc oxide.

Reaction of Metal Oxides with Acids

• Aluminium oxide reacts with hydrochloric acid to give aluminium chloride salt and water.
• The equation for this reaction is as follows:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Reaction of Metal Oxides with Bases

• Aluminium oxide reacts with sodium hydroxide to give sodium aluminate and water:
• The equation for this reaction is as follows:

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

(Sodium aluminate)

Solubility of Metal Oxides in Water

Most metal oxides are insoluble in water, but some of them dissolve in water to form a base. Sodium oxide and potassium oxide are two such oxides that dissolve in water to form a base. On dissolving, sodium oxide and potassium oxide give sodium hydroxide base and potassium hydroxide base respectively.

Na₂O(s) + H₂O(l) → 2NaOH(aq)

K₂O(s) + H₂O(l) → 2KOH(aq)

Reactivity of Metals with Oxygen

• Different metals show different reactivity when they react with oxygen. Metals like gold, platinum, and silver do not react with oxygen at all.

Reaction of Sodium and Potassium with Oxygen

• Some metals like sodium and potassium react so vigorously with oxygen that they catch fire quickly if left exposed to air. This is why they are kept immersed in kerosene oil, to protect them from catching fire suddenly.

Some Metal Oxides Form a Protective Layer

• At ordinary temperature, the surfaces of metals such as magnesium, aluminium, zinc, and lead get coated with a thin layer of oxide. This protective layer prevents them from further oxidation. This gives metals a big advantage, as it prevents them from getting corroded because of this oxide layer.

Some Metals Do Not React with Oxygen

• Iron does not burn on heating, but when iron filings are sprinkled into a flame, they burn vigorously. Copper does not burn either, but a black layer of copper (II) oxide is formed on the hot metal. Silver and gold do not react with oxygen even at very high temperatures.

Anodising

• Anodising is the process of forming a thick oxide layer on aluminium. When aluminium comes into contact with air, a thin layer of oxide is formed on it. This layer of aluminium oxide protects it from corrosion. By making this layer thicker, aluminium can be made more resistant to corrosion.

Anodising of Aluminium

• For anodising, a clean aluminium object is made the anode and electrolysed with dilute sulphuric acid. The oxygen gas evolved at the anode reacts with the aluminium to form a thick oxide layer. This oxide layer can easily be dyed to make attractive coloured articles of aluminium.

Reaction of Metals with Water

• Metals react with water to produce hydrogen gas and a metal oxide. Metal oxides that are soluble in water dissolve in it to give the corresponding metal hydroxide.

General Equation

Metal + Water → Metal oxide + Hydrogen

Metal oxide + Water → Metal hydroxide

Reaction of Sodium and Potassium with Cold Water

• Metals like potassium and sodium react vigorously with cold water. The reaction of sodium and potassium is so fast and exothermic that the hydrogen evolved immediately catches fire. 🔥

2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat energy

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat energy

Reaction of Calcium with Water

• The reaction of calcium with water is less violent. The heat generated is not enough for the hydrogen to catch fire.

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal.

Reaction of Metals with Hot Water

Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. The metal here also starts floating due to the bubbles of hydrogen gas sticking to its surface.

Reaction of Metals with Steam

• Metals such as aluminium, iron, and zinc do not react with either cold or hot water. But they react with steam to form the metal oxide and hydrogen.

2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g)

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Some Metals Do Not React with Water

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

Reaction of Metals with Acids

• Metals react with acids to give the corresponding salt and hydrogen gas.

Metal + Dilute acid → Salt + Hydrogen

• When a metal reacts with nitric acid, hydrogen gas is not produced. This is because HNO₃ is a strong oxidising agent. It oxidises the H₂ produced to water,
• and itself gets reduced to one of the nitrogen oxides (N₂O, NO, NO₂). However, magnesium (Mg) and manganese (Mn) react with very dilute HNO₃ to give H₂ gas.

Aqua Regia

• Aqua regia is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1.
• It can dissolve gold, even though neither of these acids can do so on its own. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to dissolve gold and platinum.

Reaction of Metals with Solutions of Other Metal Salts

• More reactive metals can displace less reactive metals from their compounds in solution or molten form. This is called a displacement reaction.

Metal A + Salt solution of B → Salt solution of A + Metal B

Reactivity Series

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au

Reaction of Metals and Non-metals

• Most metals form positively charged ions (cations), while non-metals form negatively charged ions (anions).

Cations and Anions

• To understand these cations and anions, we need to understand the electronic configuration of elements and their valency.

Valency

The number of valence electrons present in the outermost shell of an atom is called its valency. For example, the electronic configuration of sodium (Na) is 2, 8, 1. A sodium atom has three shells, and the outermost shell has 1 electron that can be shared, so the valence electron of sodium is 1.

• If there are 1, 2, 3, or 4 electrons in the outermost shell, these can be given away in sharing electrons. So the valence electrons would be 1, 2, 3, and 4 respectively.
• If there are 5, 6, or 7 electrons in the outermost shell, these cannot be given away in sharing, because such atoms need electrons to complete their octet.

• A sodium atom has one electron in its outermost shell. If it loses the electron from its M shell, its L shell now becomes the outermost shell, and it has a stable octet. This atom’s nucleus still has 11 protons, but the number of electrons has become 10, so there is a net positive charge, which gives us the sodium cation Na⁺. On the other hand, a chlorine atom has seven electrons in its outermost shell and needs one more electron to complete its octet. If sodium and chlorine react, the electron lost by sodium can be gained by chlorine. After gaining an electron, the chlorine atom acquires one unit of negative charge, since its nucleus has 17 protons while its K, L, and M shells have 18 electrons. This gives us the chloride ion Cl^–, so there can be a transfer relationship between these two elements.

For example:

Na → Na⁺ + e^–

2, 8, 1      2, 8

(Sodium cation)

Cl + e^– → Cl^–

2, 8, 7      2, 8, 8

(Chloride ion)

Ionic Compounds

• Compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds or electrovalent compounds.

Properties of Ionic Compounds

• Physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.
• Melting and boiling point: Ionic compounds have high melting and boiling points because a considerable amount of energy is required to break the strong inter-ionic attraction.
• Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
• Conduction of electricity: The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move towards the oppositely charged electrode when electricity is passed through the solution. Ionic compounds in the solid state do not conduct electricity because the movement of ions is not possible in the solid due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state because the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

Key Points to Remember

• Metals are hard, lustrous, malleable, ductile, sonorous, and good conductors of heat and electricity; non-metals are generally soft, dull, brittle, and poor conductors. 🔩
• Mercury is the only metal that is liquid at room temperature; bromine is the only non-metal that is liquid at room temperature.
• Carbon exists in different allotropic forms — diamond, graphite, and buckminsterfullerene. 💎
• Amphoteric oxides, such as aluminium oxide and zinc oxide, react with both acids and bases to form salt and water.
• Metals react with water to form a metal oxide/hydroxide and hydrogen gas; reactivity with water varies widely across metals.
• Metals react with dilute acids to give the corresponding salt and hydrogen gas (except when reacting with nitric acid).
• The reactivity series arranges metals in order of decreasing reactivity: K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au.
• Metals lose electrons to form cations, while non-metals gain electrons to form anions, resulting in the formation of ionic compounds.
• Ionic compounds have high melting and boiling points, are generally soluble in water, and conduct electricity only in the molten or aqueous state.

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