Class 10 Science Chapter 3 Notes: Metals and Non-metals (2026-2027)

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H1 Title: Class 10 Science Chapter 3 Notes: Metals and Non-metals (Ultimate 2026-2027 Board Revision Guide)

Navigating through the CBSE Class 10 Science curriculum requires an in-depth understanding of elemental classification, electronic configurations, coordination lattices, and metallurgical extraction principles. Chapter 3 of Class 10 Science, “Metals and Non-metals”, forms the foundation of modern structural engineering, industrial metallurgy, semiconductor electronics, and electrochemistry. It investigates the physical and chemical contrasts between metallic and non-metallic elements; explores the reactivity series governing displacement energetics and oxide amphoterism; and details the extraction of pure metals from mineral ores alongside the synthesis, structure, and ionic bonding of electrovalent salts. To help students master every aspect of this high-weightage chapter, this comprehensive guide offers textbook-accurate, highly structured, and pedagogically sound responses strictly aligned with the latest CBSE evaluation standards.

Every concept presented in the official NCERT textbook—ranging from line-by-line theoretical elucidations and an exhaustive dissection of all 14 practical NCERT activities to zero-mark-loss balanced chemical equations, authentic Wikimedia visual assets, and a master revision cheat sheet—has been detailed with precision. Key scoring terms, physical state symbols, reaction conditions, and examiner cautions have been highlighted to ensure students secure maximum marks in their CBSE Board Examinations.

Chapter 3: Metals and Non-metals

Master Concept Overview & Thematic Architecture

Metals are electropositive elements that readily lose valence electrons to form positive cations, displaying characteristic lustre, malleability, ductility, and high electrical conductivity. Non-metals are electronegative elements that gain or share electrons to form anions or covalent bonds, exhibiting high ionisation energies and forming acidic or neutral oxides.

Analytical Dimension	Metallic Elements	Non-Metallic Elements	Notable Chemical / Physical Exceptions
Electronic Tendency	Electropositive ($\text{M} \longrightarrow \text{M}^{n+} + ne^-$)	Electronegative ($\text{X} + ne^- \longrightarrow \text{X}^{n-}$)	Hydrogen acts as electropositive under specific conditions
Physical State at $25^\circ\text{C}$	Solid crystalline lattices	Solids, gases, and one liquid	Mercury ($\text{Hg}$) is liquid; Bromine ($\text{Br}_2$) is liquid non-metal
Thermal & Electrical Conductivity	Exceptionally high (delocalised sea of electrons)	Extremely poor insulators	Graphite (allotrope of Carbon) conducts electricity
Mechanical Properties	Malleable, ductile, high tensile strength	Brittle solids, non-ductile	Sodium ($\text{Na}$), Potassium ($\text{K}$) are soft enough to cut with a knife
Melting & Boiling Points	High melting and boiling points	Generally low melting and boiling points	Gallium ($\text{Ga}$) and Caesium ($\text{Cs}$) melt on palm; Diamond has high m.p.
Oxide Character	Basic oxides; some amphoteric ($\text{Al}_2\text{O}_3, \text{ZnO}$)	Acidic oxides ($\text{SO}_2, \text{CO}_2$); some neutral ($\text{CO}, \text{H}_2\text{O}$)	Metallic oxides are predominantly insoluble in water

🧠 Examiner’s Secret: When defining amphoteric oxides, always write both balanced equations (reaction with an acid AND reaction with a base). Stating only the definition without showing the formation of sodium aluminate [$\text{NaAlO}_2$] or sodium zincate [$\text{Na}_2\text{ZnO}_2$] results in a direct 1-mark deduction in CBSE Section C/D evaluations.

Concept Map: Structural Overview of Metallurgy & Bonding

[mermaid]
graph TD
A[Ores of Metals] –> B[Enrichment / Concentration of Ore]
B –> C[Metals of High Reactivity: K, Na, Ca, Mg, Al]
B –> D[Metals of Medium Reactivity: Zn, Fe, Pb, Cu]
B –> E[Metals of Low Reactivity: Hg, Ag, Au]

C --> C1[Electrolysis of Molten Chloride/Oxide]
C1 --> C2[Pure Metal at Cathode]

D --> D1[Carbonate Ore]
D --> D2[Sulphide Ore]
D1 --> D3[Calcination: Absence of Air]
D2 --> D4[Roasting: Excess Air]
D3 & D4 --> D5[Metal Oxide]
D5 --> D6[Reduction with Carbon / Al]
D6 --> D7[Purification / Electrolytic Refining]

E --> E1[Sulphide Ores: e.g., HgS, Cu2S]
E1 --> E2[Roasting in Air]
E2 --> E3[Thermal Reduction to Metal]
E3 --> E4[Refining]

[/mermaid]

Physical Properties of Metals and Non-metals

Physical Properties of Metals

Physical properties of metals are the intrinsic observable characteristics arising from metallic bonding, where positive metal ions reside within a sea of delocalised valence electrons.

1. Metallic Lustre:

Metals, in their pure state, possess a bright, shiny reflecting surface. When exposed to atmospheric gases ($\text{O}_2, \text{H}_2\text{O}, \text{CO}_2$), they tarnish due to the formation of thin surface layers of oxides, carbonates, or sulphides. Rubbing the surface with sandpaper removes this passivating coating and restores the bright metallic lustre.

2. Hardness:

Most metals are hard and possess high tensile strength. The hardness varies significantly across elements:

Hard Metals: Iron ($\text{Fe}$), Copper ($\text{Cu}$), and Aluminium ($\text{Al}$).
Exceptions (Alkali Metals): Lithium ($\text{Li}$), Sodium ($\text{Na}$), and Potassium ($\text{K}$) have low densities, weak metallic bonding, and can be cut effortlessly with an ordinary lab knife.

3. Malleability:

Malleability is the physical property of a metal that allows it to be beaten, hammered, or rolled into extremely thin foils without fracturing.

Gold ($\text{Au}$) and Silver ($\text{Ag}$) are the two most malleable metals known. Gold can be beaten into transparent foils less than $0.0001\text{ mm}$ thick.

4. Ductility:

Ductility is the mechanical property of a metal that allows it to be drawn or stretched into thin, continuous wires without snapping.

Gold is the most ductile metal: a single gram of pure gold can be drawn into a wire approximately $2\text{ kilometres}$ long.

5. Thermal Conductivity:

Metals are efficient thermal conductors because heat energy is rapidly transferred through lattice vibrations and free-electron mobility.

Best Conductors: Silver ($\text{Ag}$) followed closely by Copper ($\text{Cu}$).
Poorest Conductors: Lead ($\text{Pb}$) and Mercury ($\text{Hg}$).

6. Electrical Conductivity:

Under an applied electric potential, free valence electrons drift through the metallic lattice, conducting electric current. Wires carrying domestic electrical currents are coated with Polyvinylchloride (PVC) or vulcanised rubber, which act as insulating dielectrics to prevent electric shocks.

7. Melting and Boiling Points:

Metals generally display exceptionally high melting and boiling points (e.g., Tungsten [$\text{W}$] melts at $3422^\circ\text{C}$, making it ideal for incandescent lamp filaments).

Notable Exceptions: Gallium ($\text{Ga}$) ($m.p. = 30^\circ\text{C}$) and Caesium ($\text{Cs}$) ($m.p. = 28.5^\circ\text{C}$) have melting points so low that they melt into liquid when rested on the palm of a human hand.

8. Sonorous Property:

Metals emit a distinct, resonant ringing sound when struck forcefully against an unyielding rigid surface. This property is exploited in manufacturing acoustic instruments, temple bells, and school bells.

Physical Properties of Non-metals

Non-metals are elements that lack metallic characteristics, exhibiting high electronegativity, structural brittleness, and an absence of delocalised electron seas.

1. States of Matter:

At room temperature ($25^\circ\text{C}$), non-metals exist either as brittle solids (e.g., Carbon, Sulphur, Phosphorus, Iodine) or as discrete gases (e.g., Oxygen, Nitrogen, Hydrogen, Chlorine).

Liquid Exception: Bromine ($\text{Br}_2$) is the only non-metal that exists as a reddish-brown liquid at room temperature.

2. Absence of Lustre:

Non-metals typically display dull surfaces that absorb or scatter incident light rather than reflecting it.

Exception: Iodine ($\text{I}_2$) is a non-metal that forms lustrous, purplish-black crystals with a metallic sheen.

3. Allotropic Diversity of Carbon:

Carbon displays allotropy, existing in distinct structural forms that possess widely divergent physical properties:

Diamond: A three-dimensional tetrahedral lattice of $sp^3$-hybridised carbon atoms. It is the hardest naturally occurring substance known and exhibits an exceptionally high melting point ($> 3500^\circ\text{C}$) and high thermal conductivity, though it remains an electrical insulator.
Graphite: A planar hexagonal layered lattice of $sp^2$-hybridised carbon atoms held together by weak van der Waals forces. Because each carbon atom retains one unbonded, mobile delocalised electron, graphite is a good conductor of electricity and serves as a dry industrial lubricant.

[👉 Also Read: Class 10 Science Chapter 2 Acids, Bases and Salts Notes]

Chemical Properties of Metals

Reaction of Metals with Oxygen (Combustion)

Almost all metals combine chemically with atmospheric oxygen upon heating to synthesize corresponding solid basic metallic oxides.

$$\text{Metal} + \text{Oxygen} \longrightarrow \text{Metal Oxide}$$

Differential Reactivity Patterns with Oxygen:

Potassium ($\text{K}$) and Sodium ($\text{Na}$):
React vigorously with atmospheric oxygen and moisture at ambient temperature, catching fire spontaneously:
$$4\text{Na}(s) + \text{O}_2(g) \longrightarrow 2\text{Na}_2\text{O}(s)$$
$$4\text{K}(s) + \text{O}_2(g) \longrightarrow 2\text{K}_2\text{O}(s)$$
Safety Storage Rule: To prevent accidental ignition from atmospheric exposure, sodium and potassium are stored fully immersed in dry kerosene oil.
Magnesium ($\text{Mg}$), Aluminium ($\text{Al}$), Zinc ($\text{Zn}$), and Lead ($\text{Pb}$):
Do not catch fire at room temperature. Instead, their surfaces oxidise to form an impermeable, non-porous passivation layer that shields the underlying metal from ongoing oxidation:
$$4\text{Al}(s) + 3\text{O}_2(g) \longrightarrow 2\text{Al}_2\text{O}_3(s)$$
Iron ($\text{Fe}$):
Massive iron bars do not burn upon heating. However, finely divided iron filings sprinkled into a Bunsen burner flame combust with bright scintillating sparks:
$$3\text{Fe}(s) + 2\text{O}_2(g) \xrightarrow{\Delta} \text{Fe}_3\text{O}_4(s) [\text{Iron(II,III) Oxide}]$$
Copper ($\text{Cu}$):
Does not burn in air, but prolonged heating coats its exterior with a black layer of Copper(II) Oxide:
$$2\text{Cu}(s) + \text{O}_2(g) \xrightarrow{\Delta} 2\text{CuO}(s) [\text{Black}]$$
Silver ($\text{Ag}$) and Gold ($\text{Au}$):
Completely unreactive noble metals that do not combine with oxygen even at high temperatures.

Nature of Metal Oxides:

Most metal oxides are basic anhydrides that are insoluble in water. Some dissolve to form alkalis:
$$\text{Na}_2\text{O}(s) + \text{H}_2\text{O}(l) \longrightarrow 2\text{NaOH}(aq)$$
$$\text{K}_2\text{O}(s) + \text{H}_2\text{O}(l) \longrightarrow 2\text{KOH}(aq)$$

Amphoteric Oxides:

Amphoteric oxides are metal oxides that react chemically with both strong acids and strong bases to synthesize salt and water.

$$\text{Oxide} + \text{Acid} \longrightarrow \text{Salt} + \text{Water}$$
$$\text{Oxide} + \text{Base} \longrightarrow \text{Complex Salt} + \text{Water}$$

Reactions of Aluminium Oxide ($\text{Al}_2\text{O}_3$):
- Acting as a Base (Reacting with Acid):
  $$\text{Al}_2\text{O}_3(s) + 6\text{HCl}(aq) \longrightarrow 2\text{AlCl}_3(aq) + 3\text{H}_2\text{O}(l)$$
- Acting as an Acid (Reacting with Base):
  $$\text{Al}_2\text{O}_3(s) + 2\text{NaOH}(aq) \longrightarrow 2\text{NaAlO}_2(aq) [\text{Sodium Aluminate}] + \text{H}_2\text{O}(l)$$
Reactions of Zinc Oxide ($\text{ZnO}$):
- With Acid:
  $$\text{ZnO}(s) + 2\text{HCl}(aq) \longrightarrow \text{ZnCl}_2(aq) + \text{H}_2\text{O}(l)$$
- With Base:
  $$\text{ZnO}(s) + 2\text{NaOH}(aq) \longrightarrow \text{Na}_2\text{ZnO}_2(aq) [\text{Sodium Zincate}] + \text{H}_2\text{O}(l)$$

💡 Did You Know? (Anodising): Anodising is an industrial electrolytic process used to produce a thick, corrosion-resistant oxide layer on aluminium articles. The cleaned aluminium component serves as the anode in an electrolytic bath of dilute sulphuric acid. The oxygen gas liberated at the anode reacts with the aluminium surface to form an integrated, scratch-resistant coating of aluminium oxide ($\text{Al}_2\text{O}_3$), which can be dyed in vibrant colours.

Reaction of Metals with Water

Metals react with water to yield a metal oxide (or metal hydroxide) while liberating flammable hydrogen gas.

$$\text{Metal} + \text{Water} \longrightarrow \text{Metal Oxide} + \text{Hydrogen Gas}\uparrow$$
$$\text{Metal Oxide} + \text{Water} \longrightarrow \text{Metal Hydroxide} \quad (\text{if oxide is soluble})$$

Differential Reactivity with Water:

Violent Reaction with Cold Water ($\text{K}, \text{Na}$):
The reaction is violent and highly exothermic. The evolved hydrogen gas catches fire immediately due to the heat of the reaction:
$$2\text{K}(s) + 2\text{H}_2\text{O}(l) \longrightarrow 2\text{KOH}(aq) + \text{H}_2(g)\uparrow + \text{Heat Energy}$$
$$2\text{Na}(s) + 2\text{H}_2\text{O}(l) \longrightarrow 2\text{NaOH}(aq) + \text{H}_2(g)\uparrow + \text{Heat Energy}$$
Controlled Reaction with Cold Water ($\text{Ca}$):
Calcium reacts less violently. The heat evolved is insufficient to ignite the hydrogen gas. The metal begins floating because adhering bubbles of hydrogen gas buoy it toward the surface:
$$\text{Ca}(s) + 2\text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{H}_2(g)\uparrow$$
Reaction with Hot Water ($\text{Mg}$):
Magnesium does not react with cold water; it requires boiling hot water. It also floats as hydrogen gas bubbles stick to its surface:
$$\text{Mg}(s) + 2\text{H}_2\text{O}(l) [\text{Hot}] \longrightarrow \text{Mg(OH)}_2(aq) + \text{H}_2(g)\uparrow$$
Reaction Strictly with Steam ($\text{Al}, \text{Fe}, \text{Zn}$):
These metals react neither with cold nor with hot liquid water, but react with gaseous steam to form metal oxides:
$$2\text{Al}(s) + 3\text{H}_2\text{O}(g) [\text{Steam}] \longrightarrow \text{Al}_2\text{O}_3(s) + 3\text{H}_2(g)\uparrow$$
$$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) [\text{Steam}] \longrightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)\uparrow$$
$$\text{Zn}(s) + \text{H}_2\text{O}(g) [\text{Steam}] \longrightarrow \text{ZnO}(s) + \text{H}_2(g)\uparrow$$
Metals That Do Not React With Water Under Any Conditions:
Lead ($\text{Pb}$), Copper ($\text{Cu}$), Silver ($\text{Ag}$), and Gold ($\text{Au}$).

Reaction of Metals with Acids

Metals situated above hydrogen in the activity series displace hydrogen from dilute mineral acids to produce a metal salt and hydrogen gas.

$$\text{Metal} + \text{Dilute Acid} \longrightarrow \text{Metal Salt} + \text{Hydrogen Gas}\uparrow$$

Reactivity Gradient with Dilute $\text{HCl}$:
$$\text{Mg} > \text{Al} > \text{Zn} > \text{Fe}$$
- $\text{Mg}(s) + 2\text{HCl}(aq) \longrightarrow \text{MgCl}_2(aq) + \text{H}_2(g)\uparrow \quad (\text{Fastest, highest temp rise})$
- $2\text{Al}(s) + 6\text{HCl}(aq) \longrightarrow 2\text{AlCl}_3(aq) + 3\text{H}_2(g)\uparrow$
- $\text{Zn}(s) + 2\text{HCl}(aq) \longrightarrow \text{ZnCl}_2(aq) + \text{H}_2(g)\uparrow$
- $\text{Fe}(s) + 2\text{HCl}(aq) \longrightarrow \text{FeCl}_2(aq) + \text{H}_2(g)\uparrow$
Non-Reactivity of Copper:
$$\text{Cu}(s) + \text{HCl}(aq) \longrightarrow \text{No Reaction}$$

The Nitric Acid ($\text{HNO}_3$) Anomaly:

Hydrogen gas is not evolved when most metals react with nitric acid. Nitric acid is a powerful oxidising agent that oxidises the evolved $\text{H}_2$ into liquid water ($\text{H}_2\text{O}$), while reducing itself to nitrogen oxides ($\text{NO}_2, \text{NO}, \text{N}_2\text{O}$).

Exceptions: Only Magnesium ($\text{Mg}$) and Manganese ($\text{Mn}$) react with extremely dilute ($1%$) $\text{HNO}_3$ to liberate hydrogen gas:
$$\text{Mg}(s) + 2\text{HNO}_3(\text{very dil.}) \longrightarrow \text{Mg(NO}_3)_2(aq) + \text{H}_2(g)\uparrow$$
$$\text{Mn}(s) + 2\text{HNO}_3(\text{very dil.}) \longrightarrow \text{Mn(NO}_3)_2(aq) + \text{H}_2(g)\uparrow$$

Aqua Regia (Royal Water):

Aqua Regia is a fuming, corrosive liquid prepared by mixing concentrated Hydrochloric acid and concentrated Nitric acid in a precise volumetric ratio of $3 : 1$:
$$\text{Aqua Regia} = 3\text{ parts Conc. HCl} : 1\text{ part Conc. HNO}_3$$
It dissolves noble metals like Gold ($\text{Au}$) and Platinum ($\text{Pt}$) by generating reactive nascent chlorine:
$$\text{HNO}_3 + 3\text{HCl} \longrightarrow \text{NOCl} + 2\text{H}_2\text{O} + 2[\text{Cl}]$$

The Reactivity Series (Activity Series) of Metals

The reactivity series is an empirical, vertically ordered arrangement of metals in descending order of their chemical reactivity and electrode potentials.

$$\begin{aligned}
\text{K} \quad &\text{(Potassium)} && \text{[Most reactive metal; burns in cold water]} \
\text{Na} \quad &\text{(Sodium)} && \text{[Violent reaction with water; stored in kerosene]} \
\text{Ca} \quad &\text{(Calcium)} && \text{[Reacts with cold water; floats]} \
\text{Mg} \quad &\text{(Magnesium)} && \text{[Reacts with hot water; floats]} \
\text{Al} \quad &\text{(Aluminium)} && \text{[Passivates; reacts with steam]} \
\text{Zn} \quad &\text{(Zinc)} && \text{[Reacts with steam; amphoteric oxide]} \
\text{Fe} \quad &\text{(Iron)} && \text{[Reacts with steam; non-combustible mass]} \
\text{Pb} \quad &\text{(Lead)} && \text{[Unreactive with water]} \
\mathbf{[H]} \quad &\mathbf{(Hydrogen)} && \mathbf{[\text{Non-metal reference standard}]} \
\text{Cu} \quad &\text{(Copper)} && \text{[Unreactive with dil. HCl; forms black oxide]} \
\text{Hg} \quad &\text{(Mercury)} && \text{[Extracted by simple roasting of HgS]} \
\text{Ag} \quad &\text{(Silver)} && \text{[Noble metal; tarnishes in H}_2\text{S]} \
\text{Au} \quad &\text{(Gold)} && \text{[Most inert noble metal; native state]}
\end{aligned}$$

Exhaustive NCERT Laboratory Activities Vault

Activity 3.1: Demonstration of Metallic Lustre

Aim: To observe the appearance and lustre of clean metal surfaces.
Setup: Collect samples of iron, copper, aluminium, and magnesium. Inspect their surface appearance before and after rubbing with fine-grit sandpaper.
Observations: The initial samples appear dark and dull due to surface oxidation and corrosion. Once rubbed with sandpaper, the passivating layers are abraded away, revealing shiny, light-reflecting surfaces.
Board Conclusion: Pure metals display metallic lustre, which can be obscured by exposure to atmospheric moisture and gases.

Activity 3.2: Assessment of Hardness in Metals

Aim: To demonstrate variations in hardness among metals.
Setup: Attempt to cut iron, copper, and aluminium with a utility knife. Next, retrieve a small piece of sodium metal from kerosene using forceps, dry it between folds of filter paper on a watch glass, and attempt to slice it with the knife.
Observations: Iron, copper, and aluminium resist cutting and cannot be sliced with a knife. In contrast, sodium slices easily like cold butter.
Board Conclusion: Most transition and structural metals possess high hardness, whereas alkali metals (Sodium, Potassium, Lithium) are soft solids due to their single valence electron and weak metallic bonding.

Activity 3.3: Malleability Test of Metals

Aim: To test the malleability of various metallic specimens.
Setup: Place clean samples of iron, copper, aluminium, and lead on an anvil or heavy iron block and strike each repeatedly with a hammer.
Observations: The metal samples do not crack or shatter; instead, they flatten and expand into thin sheets.
Board Conclusion: Metals are malleable. They deform plastically under compressive stress because their atomic layers can slide past one another without disrupting the metallic bond.

Activity 3.4: Ductility Test

Aim: To identify metals that can be drawn into wires.
Setup: Survey commercial applications of metals such as copper, aluminium, iron, lead, and gold.
Observations: Copper and aluminium are drawn into electrical cables; iron is drawn into fencing wires; gold is drawn into fine jewellery threads. Lead, however, fractures when drawn under tension.
Board Conclusion: Ductility is a characteristic property of metals, with gold being the most ductile element.

Activity 3.5: Thermal Conductivity of Metals

Aim: To demonstrate high thermal conductivity and evaluate the melting point of a metal wire.
Setup: Clamp an aluminium or copper wire horizontally on an iron stand. Fix a small iron pin to the free end of the wire using molten wax. Heat the wire with a Bunsen burner near the clamped end.
Observations:
1. Thermal energy conducts rapidly from the heated region along the length of the metal wire toward the free end.
2. The wax melts, and the iron pin drops to the benchtop.
3. The metal wire itself does not melt in the flame.
Board Conclusion: Metals are efficient conductors of heat and possess high melting points. Silver and copper have the highest thermal conductivities, whereas lead and mercury are poor thermal conductors.

Activity 3.6: Electrical Conductivity Circuit

Aim: To verify the electrical conductivity of metallic test specimens.
Setup: Construct an open circuit comprising a $6\text{ V}$ battery, an indicator bulb, a switch, and two test clips ($A$ and $B$). Insert samples of various metals and non-metals between the clips.
Observations: When metals (copper, aluminium, iron) bridge the clips, the circuit closes and the bulb glows brightly. When non-metals (sulphur, plastic) are inserted, the bulb remains unlit. Inserting graphite allows the bulb to glow.
Board Conclusion: Metals conduct electricity due to the movement of free valence electrons. Non-metals act as insulators, with graphite serving as a notable conductive exception.

Activity 3.7: Physical Properties of Non-metals

Aim: To evaluate the physical properties of non-metallic elements.
Setup: Test samples of carbon (coal, graphite), sulphur, and iodine for lustre, hardness, malleability, ductility, and electrical conduction.
Observations: Sulphur and coal crush into powder when struck (brittle); they cannot be drawn into wires; they do not conduct electricity. Iodine shows no malleability but displays a distinct crystalline lustre.
Board Conclusion: Non-metals are non-malleable, non-ductile, brittle, and poor conductors of heat and electricity.

Activity 3.8: Combustion and Oxide Nature of Magnesium vs. Sulphur

Aim: To compare the chemical nature of metallic oxides and non-metallic oxides.
Setup:
1. Burn magnesium ribbon in air, collect the white ash ($\text{MgO}$), dissolve it in water, and test with red and blue litmus paper.
2. Burn sulphur powder in a deflagrating spoon, collect the evolved gas ($\text{SO}_2$) in an inverted gas jar, add water, shake to dissolve, and test with litmus paper.
Observations:
- Magnesium oxide solution turns red litmus paper blue (no change on blue litmus).
- Sulphur dioxide solution turns blue litmus paper red (no change on red litmus).
Balanced Chemical Equations:
- Metallic Oxide Sequence:
  $$2\text{Mg}(s) + \text{O}_2(g) \longrightarrow 2\text{MgO}(s)$$
  $$\text{MgO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Mg(OH)}_2(aq) [\text{Basic Solution}]$$
- Non-Metallic Oxide Sequence:
  $$\text{S}(s) + \text{O}_2(g) \longrightarrow \text{SO}_2(g)$$
  $$\text{SO}_2(g) + \text{H}_2\text{O}(l) \longrightarrow \text{H}_2\text{SO}_3(aq) [\text{Sulphurous Acid, Acidic}]$$
Board Conclusion: Metals burn in oxygen to yield basic oxides, whereas non-metals burn in oxygen to yield acidic oxides.

Activity 3.9: Heating Metals in Air

Aim: To observe the relative rates of oxidation of different metals.
Setup: Hold samples of magnesium, aluminium, copper, and iron filings in a burner flame using tongs.
Observations: Magnesium catches fire with a dazzling white flame; aluminium forms a dull grey surface oxide film; iron filings ignite in bright sparks; copper does not burn, forming a black surface layer of copper(II) oxide.
Reactivity Order Toward Oxygen:
$$\text{Mg} > \text{Al} > \text{Fe} > \text{Cu}$$

Activity 3.10: Action of Metals with Water

Aim: To classify metals based on their reactions with cold water, hot water, and steam.
Setup: Place small pieces of various metals into beakers containing cold water. For unreactive metals, heat the water. For metals that remain unreactive in hot water, pass steam over the heated metal sample in a glass-tube setup with delivery to an inverted water-displacement trough.
Observations:
- $\text{Na}$ and $\text{K}$ react violently with cold water, skittering across the surface and igniting in yellow and lilac flames.
- $\text{Ca}$ reacts steadily with cold water and floats as hydrogen bubbles stick to it.
- $\text{Mg}$ reacts only with hot water, floating as gas bubbles evolve.
- $\text{Al}, \text{Zn},$ and $\text{Fe}$ react only with steam, releasing hydrogen gas that collects over water.
Balanced Chemical Equations:
$$\text{Ca}(s) + 2\text{H}_2\text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{H}_2(g)\uparrow$$
$$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \longrightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)\uparrow$$

Activity 3.11: Action of Metals with Dilute Acids

Aim: To monitor temperature shifts and hydrogen gas evolution during metal-acid reactions.
Setup: Add equal volumes of dilute hydrochloric acid to test tubes containing equal masses of magnesium, aluminium, zinc, iron, and copper. Insert thermometers into each tube to record temperature changes.
Observations:
- Magnesium shows the fastest effervescence and highest temperature rise.
- Copper shows no effervescence, no bubbles, and no temperature change.
Reactivity Gradient:
$$\text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Cu}$$

Activity 3.12: Mutual Displacement Reactions

Aim: To confirm the relative positions of iron and copper in the reactivity series.
Setup: Hang a clean iron nail in blue copper sulphate solution (Tube A), and place a clean copper wire in green ferrous sulphate solution (Tube B) for $20\text{ minutes}$.
Observations:
- Tube A: The blue colour fades to pale green, and a reddish-brown copper deposit forms on the iron nail.
- Tube B: No colour change or reaction occurs.
Balanced Chemical Equation:
$$\text{Fe}(s) + \text{CuSO}_4(aq) [\text{Blue}] \longrightarrow \text{FeSO}_4(aq) [\text{Pale Green}] + \text{Cu}(s) [\text{Red-Brown}]$$
$$\text{Cu}(s) + \text{FeSO}_4(aq) \longrightarrow \text{No Reaction}$$
Board Conclusion: Iron is more reactive than copper and displaces it from its salt solution, whereas copper cannot displace iron.

Activity 3.13: Properties of Ionic Compounds

Aim: To evaluate the physical state, melting point, solubility, and conductivity of ionic salts.
Setup: Test samples of sodium chloride ($\text{NaCl}$) and magnesium chloride ($\text{MgCl}_2$) on a burner spatula, test their solubility in water versus kerosene, and test their electrical conductivity in solid, aqueous, and molten forms.
Observations:
1. The salts do not melt over a standard Bunsen flame (confirming high melting points).
2. They dissolve readily in water, but remain insoluble in kerosene and petrol.
3. Solid salt crystals do not conduct electricity; aqueous and molten forms conduct current and light up the test bulb.
Board Conclusion: Ionic compounds are crystalline solids with high melting points that conduct electricity only when dissolved in water or melted, due to the release of mobile ions.

Activity 3.14: Necessary Conditions for Rusting of Iron

Aim: To prove that both oxygen and water are required for the rusting of iron.
Setup: Place clean iron nails in three separate test tubes:
- Tube A: Half-filled with ordinary tap water (contains both air and moisture).
- Tube B: Filled with boiled distilled water (air removed) covered with an impermeable $1\text{ cm}$ layer of oil.
- Tube C: Contains dry air with anhydrous calcium chloride ($\text{CaCl}_2$) to absorb all moisture, sealed with a rubber stopper.
Observations (after several days):
- Tube A: Iron nails show heavy, flaky reddish-brown rust.
- Tube B: Iron nails remain completely rust-free.
- Tube C: Iron nails remain completely rust-free.
Balanced Chemical Equation:
$$4\text{Fe}(s) + 3\text{O}_2(g) + 2x\text{H}_2\text{O}(l) \longrightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}(s) [\text{Hydrated Iron(III) Oxide}]$$
Board Conclusion: Rusting requires the simultaneous presence of both oxygen and water.

How Metals and Non-metals React: Ionic Bonding

Ionic Bonding

An ionic bond is an electrostatic force of attraction that binds oppositely charged cations and anions formed by the complete transfer of one or more electrons from a metal to a non-metal.

Atoms react to achieve a stable electronic configuration matching that of the nearest noble gas (an octet of 8 valence electrons, or a duplet of 2 for Helium). Metals achieve this by losing valence electrons to form positive cations, while non-metals achieve it by gaining those electrons to form negative anions.

1. Formation of Sodium Chloride ($\text{NaCl}$):

Sodium Atom ($\text{Na}$): Electronic Configuration $= 2, 8, 1$. It loses its single valence electron:
$$\text{Na} \longrightarrow \text{Na}^+ + e^- \quad [2, 8 \text{ (Neon core)}]$$
Chlorine Atom ($\text{Cl}$): Electronic Configuration $= 2, 8, 7$. It gains that electron:
$$\text{Cl} + e^- \longrightarrow \text{Cl}^- \quad [2, 8, 8 \text{ (Argon core)}]$$
Electrostatic Combination:
$$\text{Na}^+ + \text{Cl}^- \longrightarrow \text{NaCl}$$

2. Formation of Magnesium Chloride ($\text{MgCl}_2$):

Magnesium Atom ($\text{Mg}$): Electronic Configuration $= 2, 8, 2$. It loses two valence electrons:
$$\text{Mg} \longrightarrow \text{Mg}^{2+} + 2e^- \quad [2, 8 \text{ (Neon core)}]$$
Two Chlorine Atoms: Each gains one electron:
$$2\text{Cl} + 2e^- \longrightarrow 2\text{Cl}^- \quad [2(2, 8, 8)]$$
Electrostatic Combination:
$$\text{Mg}^{2+} + 2\text{Cl}^- \longrightarrow \text{MgCl}_2$$

Characteristic Properties of Ionic Compounds

1. Physical Nature:

Ionic compounds are hard, rigid solids due to the strong electrostatic forces holding their alternating positive and negative ions in a crystal lattice. They are brittle and shatter into fragments when struck because external mechanical shear forces bring like-charged ions into alignment, causing repulsive cleavage.

2. High Melting and Boiling Points:

A large amount of thermal energy is needed to overcome the strong electrostatic attractions holding the ionic crystal lattice together:

Ionic Compound	Melting Point ($\text{K}$)	Boiling Point ($\text{K}$)
$\text{NaCl}$	$1074$	$1686$
$\text{LiCl}$	$887$	$1600$
$\text{CaCl}_2$	$1045$	$1900$
$\text{CaO}$	$2850$	$3120$
$\text{MgCl}_2$	$987$	$1685$

3. Solubility:

Ionic compounds dissolve readily in polar solvents like water ($\text{H}_2\text{O}$), where polar water molecules hydrate and separate individual ions. They are insoluble in non-polar organic solvents such as kerosene, benzene, and petrol.

4. Electrical Conductivity:

In Solid State: Ionic compounds do not conduct electricity because their ions are held rigidly in fixed lattice positions and cannot move.
In Molten and Aqueous States: The electrostatic attractions are broken by heat or hydration, allowing free-floating ions to move toward oppositely charged electrodes and carry electric current.

Metallurgy: Extraction of Metals from Natural Ores

Basic Terminology of Metallurgy

Minerals: Naturally occurring inorganic chemical compounds or elements found within the Earth’s crust.
Ores: Minerals containing a high percentage of a particular metal that can be extracted profitably on an industrial scale (e.g., Bauxite [$\text{Al}_2\text{O}_3 \cdot 2\text{H}_2\text{O}$] is an ore of aluminium). All ores are minerals, but not all minerals are ores.
Gangue (Matrix): The unwanted rocky, sandy, silty, and earthy impurities associated with mined ore bodies that must be removed prior to extraction.

Enrichment / Concentration of Ore

Mined ores are crushed and processed through techniques such as gravity separation, froth flotation, and magnetic separation, selected based on differences in the physical or chemical properties of the metal ore versus its gangue impurities.

Extraction of Metals at the Bottom of the Activity Series (Low Reactivity)

Metals low in the activity series are relatively unreactive. Their oxides can be reduced to pure metal simply by heating in air.

1. Extraction of Mercury from Cinnabar ($\text{HgS}$):

Step 1 (Roasting sulphide to oxide):
$$2\text{HgS}(s) [\text{Cinnabar}] + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{HgO}(s) + 2\text{SO}_2(g)\uparrow$$
Step 2 (Thermal reduction of oxide):
$$2\text{HgO}(s) \xrightarrow{\Delta > 300^\circ\text{C}} 2\text{Hg}(l) + \text{O}_2(g)\uparrow$$

2. Extraction of Copper from Copper Glance ($\text{Cu}_2\text{S}$):

Partial roasting in air followed by auto-reduction:
$$2\text{Cu}_2\text{S}(s) + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{Cu}_2\text{O}(s) + 2\text{SO}_2(g)\uparrow$$
$$2\text{Cu}_2\text{O}(s) + \text{Cu}_2\text{S}(s) \xrightarrow{\Delta} 6\text{Cu}(s) + \text{SO}_2(g)\uparrow$$

Extraction of Metals in the Middle of the Activity Series (Medium Reactivity)

Metals in the middle of the reactivity series ($\text{Fe}, \text{Zn}, \text{Pb}$) typically occur as sulphides or carbonates. Because metals are more easily reduced from oxides, these ores are first converted into metal oxides using Roasting or Calcination.

Roasting vs. Calcination:

Operational Parameter	Roasting Process	Calcination Process
Applicable Ore Type	Applied primarily to Sulphide Ores	Applied primarily to Carbonate & Hydrated Ores
Thermal Atmosphere	Heated strongly in the presence of excess air	Heated strongly in limited air or absence of air
Gaseous Byproduct	Releases toxic Sulphur Dioxide Gas ($\text{SO}_2\uparrow$)	Releases Carbon Dioxide Gas ($\text{CO}_2\uparrow$)
Prototypical Reaction	$2\text{ZnS}(s) + 3\text{O}_2(g) \xrightarrow{\Delta} 2\text{ZnO}(s) + 2\text{SO}_2(g)\uparrow$	$\text{ZnCO}_3(s) \xrightarrow{\Delta} \text{ZnO}(s) + \text{CO}_2(g)\uparrow$

Reduction of Metal Oxides:

Once converted to an oxide, the metal is reduced to its free state using a chemical reducing agent.

Reduction Using Carbon (Coke):
Zinc oxide is reduced by heating with powdered coke:
$$\text{ZnO}(s) + \text{C}(s) \xrightarrow{\Delta} \text{Zn}(s) + \text{CO}(g)\uparrow$$
Reduction Using Reactive Metals (Displacement Reduction):
More reactive metals like Aluminium ($\text{Al}$), Sodium ($\text{Na}$), or Calcium ($\text{Ca}$) can displace less reactive metals from their oxides. These reactions are highly exothermic, releasing enough heat to produce the extracted metal in a molten state:
$$3\text{MnO}_2(s) + 4\text{Al}(s) \xrightarrow{\Delta} 3\text{Mn}(l) + 2\text{Al}_2\text{O}_3(s) + \text{Heat Energy}$$

The Thermite Reaction:

The thermite reaction is the exothermic reduction of iron(III) oxide by aluminium powder:
$$\text{Fe}_2\text{O}_3(s) + 2\text{Al}(s) \xrightarrow{\text{Ignition}} 2\text{Fe}(l) [\text{Molten Iron}] + \text{Al}_2\text{O}_3(s) + \text{Heat Energy}$$
The heat released melts the produced iron, which flows directly into gaps to weld together broken railway tracks, cracked machine frames, and heavy gears on site.

Extraction of Metals at the Top of the Activity Series (High Reactivity)

Metals high in the reactivity series ($\text{K}, \text{Na}, \text{Ca}, \text{Mg}, \text{Al}$) have a greater affinity for oxygen than carbon does, meaning their oxides cannot be reduced using carbon or carbon monoxide.

These metals are extracted using Electrolytic Reduction (Electrolysis) of their molten chlorides or oxides:

Extraction of Molten Sodium from Fused $\text{NaCl}$:

At the Cathode (Negative Electrode – Reduction):
Sodium cations gain electrons:
$$\text{Na}^+ + e^- \longrightarrow \text{Na}(l)$$
At the Anode (Positive Electrode – Oxidation):
Chloride anions lose electrons to evolve chlorine gas:
$$2\text{Cl}^- \longrightarrow \text{Cl}_2(g) + 2e^-$$

🧠 Examiner’s Secret: Aqueous sodium chloride solution ($\text{brine}$) cannot be used to extract metallic sodium. Because water is reduced more easily than sodium ions, electrolysing aqueous $\text{NaCl}$ yields hydrogen gas at the cathode and sodium hydroxide in solution (the Chlor-Alkali process). Metallic sodium can only be extracted from anhydrous molten (fused) $\text{NaCl}$.

Refining of Metals: Electrolytic Refining of Copper

Electrolytic refining is an electrolytic process used to purify crude, blistered metals by dissolving them at an anode and redepositing pure metal onto a cathode.

Electrochemical Cell Setup:

Anode (Positive Electrode): A thick block of impure crude copper.
Cathode (Negative Electrode): A thin strip of high-purity refined copper.
Electrolyte: An aqueous solution of copper sulphate ($\text{CuSO}_4$) acidified with dilute sulphuric acid.

Electrode Reactions:

At the Anode: Metallic copper oxidises and dissolves into the electrolyte as copper ions:
$$\text{Cu}(\text{impure}, s) \longrightarrow \text{Cu}^{2+}(aq) + 2e^-$$
At the Cathode: An equivalent amount of copper ions from the electrolyte deposit onto the cathode as pure copper metal:
$$\text{Cu}^{2+}(aq) + 2e^- \longrightarrow \text{Cu}(\text{pure}, s)$$

Disposition of Impurities:

Soluble Impurities: Metals more electropositive than copper dissolve into the acidic electrolyte solution as ions.
Insoluble Impurities (Anode Mud): Less electropositive noble metals (such as Silver [$\text{Ag}$], Gold [$\text{Au}$], and Platinum [$\text{Pt}$]) do not oxidise. They settle beneath the anode as an insoluble sludge known as anode mud, which is recovered industrially to offset refining costs.

Corrosion and Its Prevention

Mechanisms of Corrosion

Corrosion is the natural electrochemical deterioration of metals caused by chemical reactions with atmospheric gases, moisture, and industrial pollutants.

Rusting of Iron:
Requires the simultaneous presence of both oxygen and water:
$$4\text{Fe}(s) + 3\text{O}_2(g) + 2x\text{H}_2\text{O}(l) \longrightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}(s) [\text{Hydrated Iron(III) Oxide}]$$
Rust is a non-adherent, porous material that flakes away and exposes fresh metal to ongoing corrosion.
Tarnishing of Silver:
Silver reacts with trace hydrogen sulphide ($\text{H}_2\text{S}$) in the air to form a black surface layer of Silver Sulphide:
$$2\text{Ag}(s) + \text{H}_2\text{S}(g) \longrightarrow \text{Ag}_2\text{S}(s) [\text{Black}] + \text{H}_2(g)$$
Corrosion of Copper:
Copper reacts slowly with moist atmospheric carbon dioxide to form a protective green patina composed of Basic Copper Carbonate:
$$2\text{Cu}(s) + \text{H}_2\text{O}(l) + \text{CO}_2(g) + \text{O}_2(g) \longrightarrow \text{CuCO}_3 \cdot \text{Cu(OH)}_2(s) [\text{Green Patina}]$$

Methods for Preventing Corrosion

Barrier Protection: Coating the surface with paint, oil, grease, or enamel to block contact with air and moisture.
Galvanisation: Coating iron or steel with a thin protective layer of molten Zinc ($\text{Zn}$). Even if the surface is scratched, the zinc oxidises preferentially, continuing to protect the exposed iron through sacrificial protection.
Alloying: Combining metals with other elements to alter their chemical properties:
- Stainless Steel: An alloy of Iron ($\text{Fe}$), Nickel ($\text{Ni}$), and Chromium ($\text{Cr}$) that resists rusting.
Anodising: Electroless or electrolytic thickening of surface aluminium oxide layers.

Alloys: Engineering Material Properties

An alloy is a homogeneous metallic mixture of two or more metals, or a metal and a non-metal, synthesized by melting the primary component and dissolving the other elements in specific proportions.

Objectives of Alloying:

Increase mechanical hardness and tensile strength.
Lower electrical conductivity and melting points.
Improve resistance to chemical corrosion.

High-Yield CBSE Alloys Reference:

Steel: Pure iron is soft and stretches easily when hot. Adding $0.05%$ Carbon ($\text{C}$) produces steel, which is hard and strong.
Stainless Steel: Composed of Iron ($\text{Fe}$), Chromium ($\text{Cr}$), and Nickel ($\text{Ni}$). It is hard and resistant to rusting.
Brass: An alloy of Copper ($\text{Cu}$, $70%$) and Zinc ($\text{Zn}$, $30%$). It is malleable, resistant to corrosion, and has a lower electrical conductivity than pure copper.
Bronze: An alloy of Copper ($\text{Cu}$, $90%$) and Tin ($\text{Sn}$, $10%$). It is hard, sonorous, and a poor electrical conductor; used for medals, statues, and ship propellers.
Solder: An alloy of Lead ($\text{Pb}$, $50%$) and Tin ($\text{Sn}$, $50%$). It has a low melting point, making it suitable for soldering electrical connections.
Amalgam: Any alloy that contains Mercury ($\text{Hg}$) as one of its constituents (e.g., dental silver amalgam used for tooth fillings).
Purity of Gold (22-Carat Gold): Pure gold is $24\text{ carats}$, but it is too soft to hold shape in jewellery. It is alloyed as $22\text{ carats}$ ($22\text{ parts pure gold}$ mixed with $2\text{ parts}$ of either Copper or Silver) to increase its hardness.

💡 Did You Know? (The Delhi Iron Pillar): The rust-resistant Iron Pillar near the Qutub Minar in New Delhi was forged over $1600\text{ years ago}$. Weighing over $6\text{ tonnes}$ and standing $8\text{ metres}$ tall, it has resisted monsoon rains and weather without rusting due to a thin, protective surface film of crystalline iron hydrogen phosphate formed from high phosphorus levels during ancient forging.

[👉 Also Read: Class 10 Science Chapter 4 Carbon and its Compounds Notes]

Master Revision Cheat Sheet

Material / System	Physical State / Appearance	Key Governing Chemical Reaction	Real-World Application
Magnesium Ribbon	Silvery-white flexible metal	$2\text{Mg}(s) + \text{O}_2(g) \longrightarrow 2\text{MgO}(s)$	Burns with a dazzling white flame
Aluminium Oxide	White amphoteric solid	$\text{Al}_2\text{O}_3 + 2\text{NaOH} \longrightarrow 2\text{NaAlO}_2 + \text{H}_2\text{O}$	Amphoteric oxide character proof
Thermite Blend	$\text{Fe}_2\text{O}_3$ and $\text{Al}$ powder	$\text{Fe}_2\text{O}_3 + 2\text{Al} \xrightarrow{\Delta} 2\text{Fe}(l) + \text{Al}_2\text{O}_3$	In-situ welding of railway tracks
Cinnabar Ore	Dark reddish-brown sulphide	$2\text{HgS} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{HgO} + 2\text{SO}_2$	Extraction of metallic Mercury
Zinc Blende	Insoluble mineral sulphide	$2\text{ZnS} + 3\text{O}_2 \xrightarrow{\Delta} 2\text{ZnO} + 2\text{SO}_2$	Roasting stage of zinc extraction
Calamine Ore	Natural zinc carbonate ore	$\text{ZnCO}_3 \xrightarrow{\Delta} \text{ZnO} + \text{CO}_2$	Calcination stage of zinc extraction
Solder Alloy	Silvery-grey malleable wire	$\text{Lead } (\text{Pb}) + \text{Tin } (\text{Sn})$	Low-melting joining of electronics
Anode Mud	Insoluble sludge under anode	Precious insoluble residue ($\text{Ag}, \text{Au}, \text{Pt}$)	Secondary recovery in copper refining
Rust Residue	Flaky reddish-brown solid	$4\text{Fe} + 3\text{O}_2 + 2x\text{H}_2\text{O} \longrightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$	Corrosive destruction of iron
Aqua Regia	Fuming, orange-yellow liquid	$3\text{HCl}(\text{conc}) : 1\text{HNO}_3(\text{conc})$	Dissolving noble metals ($\text{Au}, \text{Pt}$)