⚙️ Physical Properties
Comparing metals and non-metals
| Property | ||
|---|---|---|
| Lustre | Shiny (metallic lustre) | Dull (Exception: Iodine) |
| Hardness | Generally hard (Exc: Na, K — soft) | Generally soft (Exc: Diamond — hardest) |
| Malleability | Can be beaten into sheets (Au, Al) | Brittle — break when hammered |
| Ductility | Can be drawn into wires (Au, Cu) | Not ductile |
| Conductivity | Good (Ag > Cu > Au > Al) | Poor (Exc: Graphite conducts) |
| Melting Point | Generally high (Exc: Hg, Ga, Cs) | Generally low |
| Sonority | Sonorous (produce ringing sound) | Not sonorous |
| State | Solid at room temp (Exc: Hg — liquid) | Solid, liquid, or gas (Br₂ — liquid) |
💡 CBSE Favourite Exceptions: Mercury (liquid metal), Sodium (soft metal), Iodine (lustrous non-metal), Diamond (hardest non-metal), Graphite (conducting non-metal).
Common Elements — Metals vs Non-Metals
Metalloid
🧪 Chemical Properties of Metals
Reactions with oxygen, water, acids, and salt solutions
1. Reaction with Oxygen
Metal + O₂ → Metal Oxide (basic)
4Al + 3O₂ → 2Al₂O₃
2Cu + O₂ → 2CuO (black)
2Mg + O₂ → 2MgO (white, dazzling)
2Cu + O₂ → 2CuO (black)
2Mg + O₂ → 2MgO (white, dazzling)
Non-metals form acidic/neutral oxides:
C + O₂ → CO₂ (acidic) | S + O₂ → SO₂ (acidic)
2. Reaction with Water
Metal + H₂O → Metal hydroxide/oxide + H₂
2Na + 2H₂O → 2NaOH + H₂↑ (vigorous!)
Ca + 2H₂O → Ca(OH)₂ + H₂↑
3Fe + 4H₂O(steam) → Fe₃O₄ + 4H₂↑
Ca + 2H₂O → Ca(OH)₂ + H₂↑
3Fe + 4H₂O(steam) → Fe₃O₄ + 4H₂↑
No reaction: Cu, Ag, Au (noble metals)
3. Reaction with Acids
Metal + Dilute Acid → Salt + H₂↑
Zn + H₂SO₄ → ZnSO₄ + H₂↑
Mg + 2HCl → MgCl₂ + H₂↑
Fe + 2HCl → FeCl₂ + H₂↑
Mg + 2HCl → MgCl₂ + H₂↑
Fe + 2HCl → FeCl₂ + H₂↑
Note: Cu, Ag, Au do NOT react with dilute acids (below H in series). HNO₃ is different — it's an oxidising acid.
4. Displacement Reactions
More reactive metal displaces less reactive metal
Fe + CuSO₄(blue) → FeSO₄(green) + Cu↓
Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag↓
Zn + CuSO₄ → ZnSO₄ + Cu ✓
Cu + ZnSO₄ → No reaction ✗
Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag↓
Zn + CuSO₄ → ZnSO₄ + Cu ✓
Cu + ZnSO₄ → No reaction ✗
⚠️ Aqua regia (3:1 conc. HCl : HNO₃) can dissolve even gold and platinum! This is a favourite CBSE fact.
📊 Reactivity Series
Metals arranged in decreasing order of reactivity
💡 Mnemonic: "Please Stop Calling Me A Cute Zebra, I Like HydroGen Silver, Copper and Gold are Pretty Tiresome"
K → Na → Ca → Mg → Al → Zn → Fe → Ni → Sn → Pb → [H] → Cu → Hg → Ag → Au → Pt
K → Na → Ca → Mg → Al → Zn → Fe → Ni → Sn → Pb → [H] → Cu → Hg → Ag → Au → Pt
Highly Reactive (K, Na, Ca)
- React with cold water vigorously
- Stored under kerosene
- Extracted by electrolysis
- Cannot be reduced by carbon
Medium Reactive (Mg → Fe)
- React with steam (not cold water)
- React with dilute acids
- Extracted by reduction with carbon
- Exception: Al extracted by electrolysis
Less Reactive (Cu, Ag, Au)
- Do NOT react with water or dilute acids
- Found in free/native state
- Extracted by roasting or self-reduction
- Noble metals — used in jewellery
⛏️ Extraction of Metals
From ore to pure metal — step by step
Step 1: Concentration of Ore
Removing gangue (impurities) from the ore:
- Hydraulic washing — lighter gangue washed away
- Magnetic separation — for magnetic ores
- Froth flotation — for sulphide ores (ore floats with froth)
Step 2: Conversion to Oxide
Roasting: heating sulphide ore in air
2ZnS + 3O₂ → 2ZnO + 2SO₂
Calcination: heating carbonate ore without air
ZnCO₃ → ZnO + CO₂
Step 3: Reduction to Metal
Medium reactive: reduce oxide with carbon
ZnO + C → Zn + CO
Fe₂O₃ + 3CO → 2Fe + 3CO₂ (blast furnace)
Highly reactive: electrolysis of molten ore
Fe₂O₃ + 3CO → 2Fe + 3CO₂ (blast furnace)
2Al₂O₃ → 4Al + 3O₂ (Hall-Héroult process)
Low reactive: simple heating
2HgS + 3O₂ → 2Hg + 2SO₂ (roasting itself gives metal)
Step 4: Refining (Purification)
Electrolytic refining (most common):
- Anode: impure metal block
- Cathode: thin pure metal strip
- Electrolyte: salt solution of same metal
- Pure metal deposits at cathode
- Impurities settle as "anode mud"
💡 Thermite reaction: Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + Heat. Used to join railway tracks! Al reduces iron oxide because Al is more reactive than Fe.
🔥 Corrosion & Prevention
Why metals deteriorate and how to stop it
Rusting of Iron
4Fe + 3O₂ + xH₂O → 2Fe₂O₃·xH₂O (rust)
Conditions needed: BOTH oxygen AND moisture
- Dry air only → No rust ✗
- Boiled water (no O₂) only → No rust ✗
- Moist air (O₂ + H₂O) → Rust ✓
Rust is porous and flaky — does NOT protect the surface (unlike Al₂O₃ which is protective).
Prevention Methods
- Painting / Oiling / Greasing — physical barrier
- Galvanising — coating with zinc (most effective)
- Electroplating — coating with Cr, Ni, Sn
- Alloying — stainless steel (Fe + Cr + Ni)
- Sacrificial protection — attach more reactive metal
Galvanising works even if scratched — Zn oxidises preferentially (sacrificial protection).
Important Alloys
| Alloy | Composition | Properties / Use |
|---|---|---|
| Steel | Fe + C (0.05–1.5%) | Harder than iron; construction |
| Stainless Steel | Fe + Cr + Ni + C | Rust-proof; utensils, surgical tools |
| Brass | Cu + Zn | Golden colour; taps, decorations |
| Bronze | Cu + Sn | Hard; statues, medals, coins |
| Solder | Pb + Sn | Low MP; joining wires/circuits |
| Amalgam | Metal + Hg | Dental fillings |
| 22-carat gold | Au + Cu or Ag | Harder than pure gold; jewellery |
💡 Pure gold is 24 carat (too soft for jewellery). 22K gold = 22 parts Au + 2 parts Cu/Ag. Alloying increases hardness but decreases electrical conductivity.
⚡ Ionic Compounds
Electron transfer, ionic bonds, and properties
Formation of Ionic Bonds
Metals LOSE electrons → Cations (+)
Non-metals GAIN electrons → Anions (−)
Na (2,8,1) → Na⁺ (2,8) + e⁻
Cl (2,8,7) + e⁻ → Cl⁻ (2,8,8)
∴ NaCl = Na⁺Cl⁻ (ionic compound)
Cl (2,8,7) + e⁻ → Cl⁻ (2,8,8)
∴ NaCl = Na⁺Cl⁻ (ionic compound)
Driving force: atoms achieve noble gas configuration (stable octet).
More Examples
Mg (2,8,2) → Mg²⁺ + 2e⁻
O (2,6) + 2e⁻ → O²⁻ (2,8)
∴ MgO = Mg²⁺O²⁻
Ca (2,8,8,2) → Ca²⁺ + 2e⁻
2Cl + 2e⁻ → 2Cl⁻
∴ CaCl₂ = Ca²⁺(Cl⁻)₂
O (2,6) + 2e⁻ → O²⁻ (2,8)
∴ MgO = Mg²⁺O²⁻
Ca (2,8,8,2) → Ca²⁺ + 2e⁻
2Cl + 2e⁻ → 2Cl⁻
∴ CaCl₂ = Ca²⁺(Cl⁻)₂
Properties of Ionic Compounds
- Hard & brittle (rigid crystal lattice)
- High melting & boiling points (strong electrostatic forces)
- Soluble in water (polar solvent)
- Insoluble in organic solvents (kerosene, petrol)
- Conduct electricity when molten/dissolved (free ions)
- Do NOT conduct in solid state (ions fixed in lattice)
⚠️ Why ionic compounds are brittle: When force is applied, like-charged ions come adjacent and repel → crystal shatters. Unlike metals which have a "sea of electrons" allowing layers to slide.