Chapter Notes: Metals and Nonmetals - Class 10 Science


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Metals and Non-metals

Elements are generally classified into metals and non-metals based on their properties.

Physical properties of metals and non-metals:

  1. Malleability: Property of a substance due to which it can be beaten into thin sheet. Metals are malleable while non-metals are not.
  2. Ductility: Property of a substance by which it can be drawn into thin wires. Metals are ductile while non-metals are not.
  3. Conduction of heat: Metals are good conductors of heat. Silver and copper are best conductors while lead and mercury are the poor conductors of heat. Non-metals are bad conductors of heat.
  4. Conduction of electricity: Metals are good conductors of electricity while non-metals are bad conductors of electricity with an exception of graphite(an allotrope of carbon)
  5. Appearance of surface: Metals, in their pure state, have a shining surface also called metallic lustre. Non-metals are generally non-lustrous with an exception of iodine. Metals on reacting with gases in atmosphere lose its shiny appearance when kept in air for a long time.
  6. Hardness: Metals are generally hard with the exceptions of sodium and potassium that can be cut by knife.
  7. Density: Metals have high density except sodium and potassium.
  8. Melting and boiling points: Generally metals have melting and boiling points except for sodium, potassium, mercury, cesium, gallium.
  9. State at room temperature: Metals are generally solid at room temperature with an exception of mercury which is liquid. Non-metals are present in all three states, solids, liquids and gaseous, at room temperature.
  10. Sonority: Property of producing sound on striking a hard surface. Metals are sonorous while non-metals are not.

Exceptions in Physical Properties

  1. Graphite, a non-metal, is a good conductor of electricity.
  2. Iodine is a lustrous non-metal.
  3. Diamond, an allotrope of carbon, which is a non-metal is the hardest substance while sodium and potassium, being metals are soft enough to be cut by knife.
  4. Mercury, which is a metal, is liquid at room temperature while rest are solids.
  5. Sodium, potassium, mercury, cesium and gallium are metals with low melting and boiling points.
  6. Diamond is the non-metal with high melting and boiling point.
  7. Sodium, potassium and lithium are metals with low density.

Chemical properties of Metals

  1. Reaction with oxygen:
  • Almost all metals combine with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

For e.g. 2Cu + O2 → 2CuO

  • Metal oxides are basic in nature. Some metal oxides, such as aluminium oxide, zinc oxide, etc., which react with both acids as well as bases to produce salts and water are known as amphoteric oxides.

For e.g. Al2O3 + 6HCl → 2AlCl3 + 3H2O

Al2O3 + 2NaOH → 2NaAlO2 + H2O

  • Most metal oxides are insoluble in water but some dissolve to form alkalis like sodium and potassium oxides.
  • Metals such as potassium and sodium react so vigorously with oxygen that they catch fire if kept in the open. So to protect them they are kept immersed in kerosene oil.

For e.g. Na2O(s) + H2O(l) → 2NaOH(aq)

  • Anodising: Process of forming a thick oxide layer of aluminium that makes it resistant to further corrosion.
  1. Reaction with water:
  • All metals do not react with water. Those which react form metal oxide and hydrogen gas. Metal oxides that are soluble in water further form metal hydroxide.

Metal + Water → Metal oxide + Hydrogen

Metal oxide + Water → Metal hydroxide

  • Metals like sodium & potassium react with cold water vigorously; metals like magnesium react with hot water. Iron, zinc reacts with steam while lead, silver and gold do not react with water at all.

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

Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)

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  1. Reaction with Dilute Acid:
  • Most metals react with acids to give a salt and hydrogen gas.

Metal + Dilute acid → Salt + Hydrogen

  • Hydrogen gas is not evolved when a metal reacts with nitric acid. It’s a strong oxidizing agent and oxidizes hydrogen produced to water and itself gets reduced to any of the nitrogen oxides (N2O, NO, NO2).

But magnesium (Mg) and manganese (Mn) react with very dilute

HNO3 to evolve H2 gas.

  • Aquaregia:
  • Freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1.
  • Is a highly corrosive, fuming liquid and one of the few reagents able to dissolve gold and platinum.
  • Reactivity of metals / Reactivity Series:

Figure 1: Activity Series : Relative reactivities of metals

  1. Reaction of metals with metal salts:
  • Reactive metals can displace less reactive metals from their compounds in solution or molten form.

How do metals and non-metals react:

  • The compounds formed by the transfer of electrons from a metal to a non-metal are known as ionic or electrovalent compounds.
  • Properties of ionic compound:
  1. Physical Nature: Solid and hard due to strong inter-ionic force of attraction; generally brittle.
  2. Melting and boiling points: High melting and boiling points since a considerable force is required to break the strong inter-ionic attraction.
  3. Solubility: Generally soluble in water but insoluble in solvents such as kerosene, petrol, etc.
  4. Conduction of electricity:
  • Conducts electricity through solution due to involvement of charged particles (ions).
  • As movement of ions is not possible in solid state, due to rigid structure, do not conduct electricity.
  • In molten state this movement is overcome due to heat and thus conducts electricity.

Occurrence of Metals:

  • Mineral: The elements or compounds, which occur naturally in the earth’s crust.
  • Ore: Mineral that contains high percentage of metal that can be extracted profitably from it.
  • Every ore is a mineral but every mineral is not an ore.

Obtaining metal from ore:

  • Different techniques are to be used for obtaining the metals on the basis of their reactivity.
  1. Enrichment of ore:
  • Ores mined from earth contain large amount of impurities such as sand, soil, etc. called gangue.
  • Prior to the extraction of metal, based on the differences between the physical or chemical properties of gangue and the ore, different processes are used to remove gangue.
  1. Extraction of Metal:
  2. Extraction of metals low in the Activity Series:
  • These metals are generally very unreactive.
  • Oxides of these can be reduced to metals by heating alone.

For e.g. 2HgS(s) + 3O2 (g) + Heat → 2HgO(s) + 2SO2 (g)

  1. Extraction of Metals in the middle of the Activity Series:
  • It’s easy to obtain a metal from its oxide compared to its sulphide and carbonate.
  • Roasting is a process of converting sulphide ores into oxides by heating strongly in the presence of excess air.
  • Calcination is a process of converting carbonate ores into oxides by heating strongly in limited air.
  • Roasting

2ZnS(s) + 3O2 (g) → 2ZnO(s) + 2SO2 (g)

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ZnCO3 (s) →ZnO(s) + CO2 (g)

  • Metal oxides are then reduced to corresponding metals using suitable reducing reagents such as coke, aluminum, etc. on basis of their reactivities using displacement reactions.
  • These displacement reactions are highly exothermic; hence metals are produced in molten state.
  • Thermit reaction: Reaction of iron oxide with aluminium used to join railway tracks or cracked machine parts.

Fe2O3(s) + 2Al(s) → 2Fe(l) + Al2O3(s) + Heat

Figure 2: Steps involved in extraction of metals from ores.

  1. Extraction of metals high in the Activity Series:
  • Since these are very reactive metals and thus cannot be obtained by displacement reactions. These metals are obtained by electrolytic refining.
  • They are generally obtained by electrolysis of their molten chlorides. Metals are deposited at cathode (negatively charged), while chlorine is liberated at anode.

At cathode Na+ + e → Na

At anode 2Cl → Cl2 + 2e

  • Aluminium is obtained by electrolytic reduction of aluminium oxide.

Figure 3: Activity Series and related metallurgy.

  1. Electrolytic Refining:
  • Metals obtained by various reduction processes contain impurities. The most widely used method for refining impure metals is electrolytic refining.
  • Apparatus setup:

At Anode – Impure Metal

At Cathode – Pure Metal

Electrolyte – Solution of the metal salt

  • At Anode: Pure metal from anode dissolve into electrolyte.

At Cathode: An equivalent amount of pure metal from electrolyte is deposited at cathode.

  • Soluble impurities go into solution; insoluble impurities settle at the bottom of anode called as anode mud.


  • The eating up of metal by the action of gases, moisture or acids present in air.
  • Corrosion of Iron:

Figure 4 : Investigating the conditions under which iron rusts.

In tube A, both air and water are present. In tube B, there is no air dissolved in the water. In tube C, the air is dry.

  • Prevention of corrosion of Iron:
  1. Painting
  2. Applying grease
  3. Galvanisation : Process of protecting steel and iron from rusting by coating them with thin layer of zinc.
  4. Chromium plating/ tin plating
  5. Alloying: Improve the properties of a metal.
  • Alloy is a homogeneous mixture of a metal and two or more other metals or non-metals.

It has better properties than metals like in case of strength, corrosion and lower electrical conductivity& melting points.

It is prepared by melting primary metal first and then dissolving other in definite proportion and then cooling to room temperature.

  • Some common alloys:
  1. Stainless Steel – Alloy of Fe, Ni and Cr
  2. Brass – Alloy of Cu and Zn
  3. Bronze – Alloy of Cu and Sn
  4. Solder – Alloy of Pb and Sn
  5. Amalgam – Any alloy containing mercury
  6. Alloy of Gold – Contains gold and silver/copper.

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