CBSE Class 10 Science Metals and Non-metals Notes

Introduction to Metals and Non-metals

Metals and Non-metals is Chapter 3 of Class 10 Science and one of the highest-weightage chapters in board examinations. It covers the physical and chemical properties that distinguish metals from non-metals, how metals react with oxygen, water, acids, and other metal salts, the reactivity series, ionic bonding, and how metals are extracted from their ores (metallurgy).

Metals are everywhere in our daily lives — the iron in your building structure, the copper in electrical wires, the aluminium in foil and utensils, and the gold and silver in jewellery. Non-metals like oxygen, nitrogen, carbon, sulphur, and phosphorus are equally vital for life and industry.

This chapter also explains important phenomena like corrosion, alloy formation, and the role of ionic compounds in everyday chemistry.

Key Topics Covered

•         Physical Properties of Metals and Non-metals

•         Chemical Properties: Reactions with Oxygen, Water, Acids, Salt Solutions

•         Reactivity Series of Metals

•         Formation of Ionic Compounds (Electrovalent Bonding)

•         Metallurgy: Extraction of Metals from Ores

•         Refining of Metals

•         Corrosion — Causes and Prevention

•         Alloys — Definition, Examples, and Uses

1.  Physical Properties of Metals and Non-metals

Metals and non-metals have distinctly different physical properties. These differences arise from the nature of metallic bonding (electron sea model) in metals versus covalent/ionic bonding in non-metals.

Exceptions: Mercury (Hg) is the only liquid metal. Iodine is a non-metal with lustre. Diamond (non-metal) is the hardest natural substance. Graphite (non-metal) conducts electricity.

2.  Chemical Properties of Metals

Metals undergo characteristic chemical reactions with oxygen, water, acids, and salt solutions. The extent and nature of these reactions depends on the reactivity of the metal.

Reaction with Oxygen

Metals react with oxygen to form metal oxides, which are generally basic in nature.

Metal + O2  ->  Metal Oxide (basic)

4Na + O2  ->  2Na2O      (very vigorous; burns in air)

2Mg + O2  ->  2MgO       (bright white flame; white powder)

4Al + 3O2  ->  2Al2O3   (Al2O3 layer prevents further corrosion)

2Fe + O2  ->  2FeO      (slow; iron rusts; forms Fe2O3.xH2O)

Au, Pt, Ag  -->  do NOT react with oxygen  (least reactive)

Amphoteric oxides (Al2O3, ZnO) react with BOTH acids and bases. They are neither purely acidic nor purely basic.

Reaction with Water

Different metals react with water at different rates, depending on their position in the reactivity series.

2Na + 2H2O  ->  2NaOH + H2  (very vigorous; catches fire; floats)

2K  + 2H2O  ->  2KOH  + H2  (explosive; K burns violet)

Mg  + 2H2O  ->  Mg(OH)2 + H2  (only with hot water/steam)

3Fe + 4H2O  ->  Fe3O4 + 4H2   (only with steam; not cold water)

Cu, Ag, Au  -->  do NOT react with water at all

Reaction with Dilute Acids

Metals above hydrogen in the reactivity series displace hydrogen from dilute acids, forming a salt and hydrogen gas.

Metal + Dilute Acid  ->  Salt + Hydrogen gas

Zn + H2SO4  ->  ZnSO4 + H2  (moderate rate; bubbles)

Fe + 2HCl  ->  FeCl2 + H2   (moderate rate)

Mg + 2HCl  ->  MgCl2 + H2  (very fast; vigorous)

Cu, Ag, Au  -->  do NOT react with dilute acids

Copper does not displace hydrogen from acids because Cu is below H in the reactivity series. No H2 gas is produced.

Reaction with Salt Solutions (Displacement)

A more reactive metal displaces a less reactive metal from its salt solution. This is the basis of single displacement reactions.

More reactive Metal + Salt solution  ->  New salt + Less reactive Metal

Fe + CuSO4  ->  FeSO4 + Cu   (iron nails turn copper-coloured)

Zn + CuSO4  ->  ZnSO4 + Cu   (blue colour of CuSO4 fades)

Cu + ZnSO4  -->  NO REACTION   (Cu less reactive than Zn)

3.  Reactivity Series of Metals

The reactivity series (also called the activity series) arranges metals in decreasing order of their reactivity. It is one of the most important tools in predicting chemical behaviour of metals.

Reactivity series mnemonic: Please Stop Calling Me A Cute Zebra Instead Like Him — (K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au, Pt)

4.  Ionic Compounds — Formation and Properties

Metals lose electrons to form positive ions (cations); non-metals gain electrons to form negative ions (anions). The electrostatic attraction between these oppositely charged ions forms an ionic (electrovalent) bond.

Formation of Ionic Bond — Example: NaCl

Sodium (Na) has 1 electron in its outermost shell. Chlorine (Cl) has 7 electrons and needs 1 more to complete its octet. Transfer of 1 electron from Na to Cl forms Na+ and Cl- ions held together by ionic bonding.

Na  ->  Na+  +  e-             (loses 1 electron; becomes cation)

Cl + e-  ->  Cl-              (gains 1 electron; becomes anion)

Na+  +  Cl-  ->  NaCl          (ionic compound; table salt)

Properties of Ionic Compounds

•         Physical state: Solid crystalline structures at room temperature (strong electrostatic forces).

•         Melting and boiling points: Very high (large amount of energy needed to break ionic bonds).

•         Solubility: Generally soluble in water (polar solvent) but insoluble in organic solvents like petrol.

•         Electrical conductivity: Conduct electricity only in molten state or in aqueous solution (ions free to move), NOT in solid state.

•         Brittleness: Brittle — layers shift when struck, causing like charges to face each other and repel, breaking the crystal.

Ionic compounds conduct electricity in molten or dissolved state because ions become free to move. In solid state, ions are fixed in lattice — no conduction.

5.  Occurrence of Metals and Metallurgy

Most metals occur in nature as compounds (ores) because they react with oxygen, sulphur, and other elements. Very few unreactive metals like gold and platinum occur in the free (native) state.

Important Terminology

•         Mineral: A naturally occurring element or compound present in the Earth's crust.

•         Ore: A mineral from which a metal can be extracted profitably. All ores are minerals but not all minerals are ores.

•         Gangue: The impurities (sand, rocks) present along with the ore in the earth.

•         Flux: A substance added during smelting to remove gangue (e.g., limestone CaCO3 added to remove SiO2).

•         Slag: The product formed when flux reacts with gangue (e.g., CaSiO3).

Steps in Extraction of Metals (Metallurgy)

1.       Mining and concentration of ore: Remove gangue by physical methods (froth flotation, gravity separation, magnetic separation, leaching).

2.       Conversion of concentrated ore to metal oxide: Calcination (heating in absence of air) or Roasting (heating in presence of air).

3.       Reduction of metal oxide to metal: Using carbon (coke), CO, Al, or electrolysis depending on reactivity.

4.       Refining/purification of the metal: Electrolytic refining, liquation, distillation, or vapour phase refining.

Reduction Based on Reactivity

Thermite reaction (Fe2O3 + 2Al -> Al2O3 + 2Fe): Al is more reactive than Fe, so Al reduces Fe2O3. Produces enormous heat — used for welding railway tracks.

Electrolytic Refining

Electrolytic refining is the most common method to obtain pure metals. It uses electrolysis to transfer metal from an impure anode to a pure cathode.

•         Anode: Impure metal (dissolves into solution during electrolysis).

•         Cathode: Thin strip of pure metal (pure metal deposits here).

•         Electrolyte: Solution of a salt of the same metal (e.g., CuSO4 solution for copper refining).

•         Anode mud: Insoluble impurities (gold, silver) collect below the anode.

6.  Corrosion and Alloys

Corrosion

Corrosion is the gradual deterioration of metals due to reaction with moisture, oxygen, CO2, acids, or other chemicals in the environment. It is an oxidation process. The most well-known example is rusting of iron.

4Fe + 3O2 + xH2O  ->  2Fe2O3 . xH2O   (rust)

•         Silver tarnishes: forms black Ag2S when exposed to H2S in air.

•         Copper corrodes: forms green patina (basic copper carbonate) in moist air.

•         Conditions for rusting: presence of BOTH oxygen and water is essential; either alone is not sufficient.

Prevention of Corrosion

•         Painting or greasing: Creates a barrier between metal and environment.

•         Galvanising: Coating iron with zinc (Zn). Zinc protects iron even if coating is scratched (sacrificial protection).

•         Electroplating: Depositing a protective metal layer (e.g., nickel, chromium) using electrolysis.

•         Alloying: Making stainless steel (Fe + Cr + Ni) is highly resistant to corrosion.

•         Sacrificial protection: Attaching a more reactive metal (Mg or Zn) to the structure; Mg/Zn corrodes instead of Fe.

Alloys — Definition and Examples

An alloy is a homogeneous mixture of two or more metals (or a metal and a non-metal). Alloys have improved properties compared to pure metals — harder, stronger, more resistant to corrosion, or lower melting point.

Pure gold (24 carat) is too soft for jewellery. It is alloyed with Ag or Cu to make 22K or 18K gold, which is harder and more durable.

7.  Key Reactions Summary

All important reactions from this chapter tested in CBSE Class 10 board examinations:

8.  Board Exam Practice Questions

These questions cover all question types from CBSE Class 10 Science (Chapter 3) board examinations.

1 Mark Questions

1.       Name the metal that is the best conductor of electricity.

2.       What is an alloy? Give one example.

3.       Name the method used to extract highly reactive metals like sodium.

4.       What is gangue?

5.       Which metal is stored under kerosene and why?

3 Mark Questions

1.       Explain why metals are good conductors of electricity. How does this differ from ionic compounds in solid state?

2.       State what happens when iron nails are dipped in copper sulphate solution. Write the equation and name the type of reaction.

3.       What is thermite reaction? Write the equation. Give its practical application.

4.       Differentiate between calcination and roasting with one example each.

5 Mark Questions

1.       Describe the process of electrolytic refining of copper with a labelled diagram description. What is anode mud? Name two metals found in it.

2.       Explain the extraction of metals of moderate reactivity from their sulphide ores with relevant chemical equations (take ZnS as example).

3.       Compare the physical and chemical properties of metals and non-metals under five headings with one example each.