Chapter 3: Metals and Non-Metals

1. Introduction

The Earth’s crust contains different elements. Some of these are metals, some are non-metals, and a few show properties of both (called metalloids).

 Metals and non-metals are important in our daily life — from the utensils we use to the wires that carry electricity and even the gases we breathe.
 Let’s understand their physical and chemical properties, reactions, and uses in detail.

2. Physical Properties of Metals

Property

Description

Examples / Exceptions

1. Lustre (Shiny surface)

Metals have a shining surface.

Gold, silver, copper. (Exception: Tarnished metals lose shine due to oxidation.)

2. Hardness

Most metals are hard and solid.

Sodium and potassium are soft and can be cut with a knife.

3. Malleability

Can be beaten into thin sheets.

Gold and silver are most malleable.

4. Ductility

Can be drawn into wires.

Copper and aluminium. Gold is most ductile.

5. Conductivity

Good conductors of heat and electricity.

Silver and copper are best conductors.

6. Sonority

Produce ringing sound when hit.

Bell metal.

7. Density and Strength

Usually have high density and strong.

Iron is strong and used in construction.

8. Melting & Boiling Points

Generally high.

Mercury is liquid at room temperature (exception).

3. Physical Properties of Non-Metals

Property

Description

Examples / Exceptions

1. State

Usually gases or brittle solids.

Oxygen, nitrogen (gases), sulphur, phosphorus (solids).

2. Lustre

Dull appearance (not shiny).

Iodine is an exception (it is lustrous).

3. Hardness

Soft and brittle.

Diamond (carbon) is the hardest natural substance (exception).

4. Conductivity

Poor conductors of heat and electricity.

Graphite (carbon) conducts electricity.

5. Melting & Boiling Points

Generally low.

Carbon, boron have very high melting points.

4. Chemical Properties of Metals

(a) Reaction with Oxygen

  • Metals react with oxygen to form metal oxides.

Example:
 4Na + O₂ → 2Na₂O
 2Mg + O₂ → 2MgO

Nature of oxides:

  • Basic: Most metal oxides (like Na₂O, CaO).
  • Amphoteric: Some oxides react with both acid and base (like Al₂O₃, ZnO).

Example:
 Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
 Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

(b) Reaction with Water

  • Highly reactive metals (like Na, K) react vigorously with water to form metal hydroxide and hydrogen gas.
    Example: 2Na + 2H₂O → 2NaOH + H₂ ↑

  • Less reactive metals (like Mg) react slowly with hot water.
    Example: Mg + H₂O → Mg(OH)₂ + H₂ ↑

  • Unreactive metals (like Au, Ag, Cu) do not react with water.

Observation: Hydrogen gas evolves; reaction may be vigorous for active metals.

(c) Reaction with Acids

  • Metals react with acids to form salt and hydrogen gas.

Example:
 Zn + 2HCl → ZnCl₂ + H₂ ↑

Observation:
 Effervescence due to hydrogen gas.

Note: Metals like Cu, Ag, Au do not react with dilute acids because they are less reactive.

(d) Reaction with Bases

  • Some metals react with strong bases (like NaOH, KOH) to form complex salts and hydrogen gas.

Example:
 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂ ↑

5. Chemical Properties of Non-Metals

(a) Reaction with Oxygen

  • Non-metals react with oxygen to form non-metal oxides, which are generally acidic in nature.

Example:
 C + O₂ → CO₂
 S + O₂ → SO₂

Nature:

  • Dissolve in water to form acids.
    Example: SO₂ + H₂O → H₂SO₃

(b) Reaction with Water

Non-metals generally do not react with water.
 However, they may form acidic oxides which react with water to give acids.

(c) Reaction with Acids

Non-metals do not react with dilute acids because they do not displace hydrogen from acids.

(d) Reaction with Bases

Non-metals react with strong bases to form complex compounds.
 Example: Cl₂ + NaOH → NaCl + NaClO + H₂O

6. Reactivity Series of Metals

The reactivity series is a list of metals arranged in decreasing order of their reactivity.

| Most Reactive | K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au | Least Reactive |

Key Points:

  • Metals above hydrogen can displace hydrogen from acids.
  • More reactive metals can displace less reactive metals from their compounds.

Example:
 Zn + CuSO₄ → ZnSO₄ + Cu (Displacement reaction)

7. Occurrence of Metals

Metals occur in nature in the form of minerals. Those minerals which contain a large amount of metal are called ores.

Example:

  • Bauxite (Al₂O₃·2H₂O) → Aluminium
  • Haematite (Fe₂O₃) → Iron
  • Zinc blende (ZnS) → Zinc

Metals are extracted from their ores by various processes depending on their reactivity.

8. Extraction of Metals

(a) For Less Reactive Metals (Like Au, Ag, Cu):

  • Found in free or native state.
  • Can be obtained by physical separation methods.

(b) For Moderately Reactive Metals (Like Zn, Fe, Pb):

  • Present as oxides, carbonates, or sulphides.
  • Steps of extraction:
    1. Concentration of ore – Remove impurities.
    2. Roasting (for sulphide ores):
      2ZnS + 3O₂ → 2ZnO + 2SO₂
    3. Calcination (for carbonate ores):
      ZnCO₃ → ZnO + CO₂
    4. Reduction of oxide to metal:
      ZnO + C → Zn + CO

(c) For Highly Reactive Metals (Like Na, K, Ca, Al):

  • Extracted by electrolysis of molten compounds.
    Example:
     Electrolysis of molten NaCl → Na and Cl₂ gas.

9. Refining of Metals

  • After extraction, metals often contain impurities.
     They are purified by electrolytic refining.

    Example: Copper Refining

    • Anode: Impure copper
    • Cathode: Pure copper plate
    • Electrolyte: Acidified copper sulphate solution

     

Process:
 When electric current is passed, pure copper deposits on the cathode and impurities settle below as anode mud.

10. Corrosion

Definition: The slow deterioration of metals when exposed to air, moisture, or chemicals is called corrosion.

Example:
 Iron reacts with oxygen and moisture to form rust (Fe₂O₃·xH₂O).
 2Fe + 3O₂ + xH₂O → Fe₂O₃·xH₂O

Prevention of Corrosion
  1. Painting or greasing – Prevents contact with air and moisture.
  2. Galvanization – Coating of zinc on iron.
  3. Alloying – Mixing metals to improve strength and resistance (e.g. stainless steel = Fe + Cr + Ni).
  4. Electroplating – Coating a less reactive metal like chromium on iron.

11. Uses of Metals

  • Iron: Construction, machinery, tools.
  • Copper: Electrical wires.
  • Aluminium: Utensils, aircraft, packaging.
  • Gold & Silver: Jewellery, coins.
  • Zinc: Galvanization, batteries.

12. Uses of Non-Metals

  • Oxygen: Respiration and combustion.
  • Nitrogen: Fertilizers.
  • Chlorine: Disinfectant.
  • Sulphur: Making sulphuric acid.
  • Phosphorus: Matchsticks, fireworks.
  • Carbon: Fuels and steel industry.

13. Alloys

An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal.

Alloy

Constituents

Uses

Brass

Copper + Zinc

Musical instruments

Bronze

Copper + Tin

Medals, statues

Steel

Iron + Carbon

Construction

Stainless steel

Iron + Cr + Ni

Utensils, tools

Solder

Lead + Tin

Joining electrical wires

Amalgam

Mercury + Metal

Dental fillings

Advantages of Alloys:
    • Greater strength.
    • Resistance to corrosion.
    • Lower melting point.

14. Differences Between Metals and Non-Metals

 

Property

Metals

 

          Non-Metals

1. Lustre

Shiny

 

Dull

2. Conductivity

Good conductors

 

Poor conductors

3. Malleability

Can be beaten into sheets

 

Brittle

4. Ductility

Can be drawn into wires

 

Not ductile

5. Nature of oxides

Basic or amphoteric

 

Acidic

6. Reaction with acids

Produce H₂ gas

 

No reaction

7. Position in periodic table

Left side

 

Right side

8. Examples

Na, Cu, Fe

 

S, C, N, O

15. Ionic Compounds

Metals and non-metals combine by transfer of electrons to form ionic compounds.

Example:
 Na + Cl → Na⁺Cl⁻ → NaCl

Properties:

  • Formed by strong electrostatic attraction between ions.
  • Have high melting and boiling points.
  • Conduct electricity in molten or aqueous state (not in solid).

Generally soluble in water.

16. Summary of Key Reactions

  1. Metal + Oxygen → Metal oxide
     Example: 2Mg + O₂ → 2MgO

  2. Metal + Water → Metal hydroxide + Hydrogen
     Example: 2Na + 2H₂O → 2NaOH + H₂

  3. Metal + Acid → Salt + Hydrogen
     Example: Zn + H₂SO₄ → ZnSO₄ + H₂

  4. Metal oxide + Acid → Salt + Water
     Example: ZnO + 2HCl → ZnCl₂ + H₂O

  5. Non-metal + Oxygen → Non-metal oxide
     Example: S + O₂ → SO₂

Displacement Reaction:
 Zn + CuSO₄ → ZnSO₄ + Cu

17. Summary of the Chapter

  • Metals are electropositive; they lose electrons to form positive ions.
  • Non-metals are electronegative; they gain electrons to form negative ions.
  • Reactivity series helps predict displacement reactions.
  • Corrosion can be prevented by galvanization, painting, or alloying.
  • Ionic compounds have strong forces and conduct electricity in solution.
  • Alloys improve properties of metals for better use.
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