In this chapter, you will learn
- —Compare physical properties of metals and non-metals (lustre, hardness, malleability, conductivity, melting/boiling points)
- —Describe chemical properties of metals — reactions with oxygen, water, acids, and salt solutions
- —Understand the reactivity series of metals and its applications
- —Explain how metals are extracted from ores (reduction methods based on reactivity)
- —Understand ionic bond formation through electron transfer
- —Learn about corrosion and its prevention methods
Physical Properties of Metals and Non-metals
Metals and non-metals differ significantly in their physical properties:
| Property | Metals | Non-metals |
|---|---|---|
| Lustre | Shiny (metallic lustre) | Dull (except iodine, diamond) |
| Hardness | Generally hard (except Na, K — soft) | Generally soft (except diamond — hardest) |
| Malleability | Can be beaten into thin sheets (gold is most malleable) | Not malleable — break when hammered (brittle) |
| Ductility | Can be drawn into wires (gold is most ductile) | Not ductile |
| Conductivity | Good conductors of heat and electricity (silver is best, then copper) | Poor conductors (except graphite — conducts electricity) |
| Melting/Boiling point | Generally high (except Ga, Cs — low; tungsten — highest 3422°C) | Generally low |
| Sonority | Produce a ringing sound when struck | Not sonorous |
| State at room temp | Solid (except mercury — liquid) | Solid, liquid, or gas (bromine — liquid) |
Key Exceptions to Remember:
- Mercury — only metal that is liquid at room temperature
- Gallium and Caesium — metals with very low melting points (melt in your hand)
- Sodium and Potassium — soft metals, can be cut with a knife
- Diamond — non-metal but the hardest natural substance
- Graphite — non-metal but conducts electricity
- Iodine — non-metal but has lustre
Exam Tip
Exceptions are very commonly asked: mercury (liquid metal), graphite (non-metal conductor), diamond (hardest non-metal), Na/K (soft metals). Memorise all of these.
Common Mistake
Students say 'all metals are hard' or 'all non-metals are poor conductors'. Always remember the exceptions — Na, K are soft; graphite conducts electricity.
Chemical Properties of Metals
Metals show characteristic chemical reactions:
1. Metals + Oxygen → Metal Oxides
2Mg(s) + O₂(g) → 2MgO(s) (burns with dazzling white flame)
4Al(s) + 3O₂(g) → 2Al₂O₃(s)
2Cu(s) + O₂(g) → 2CuO(s) (copper turns black on heating)
- Most metal oxides are basic (react with acids): MgO + 2HCl → MgCl₂ + H₂O
- Some are amphoteric (react with both acid and base): Al₂O₃, ZnO
- Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (acts as base)
- Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (acts as acid)
2. Metals + Water → Metal Hydroxide/Oxide + Hydrogen
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)↑ (vigorous, catches fire)
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)↑ (violent reaction)
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)↑ (sinks, steady bubbles)
Mg(s) + 2H₂O(hot) → Mg(OH)₂(aq) + H₂(g)↑ (reacts with hot water)
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)↑ (reacts with steam only)
- Na and K react vigorously with cold water — stored in kerosene
- Ca reacts with cold water but less vigorously
- Mg reacts with hot water
- Fe, Zn react with steam only
- Cu, Ag, Au do not react with water at all
3. Metals + Dilute Acids → Salt + Hydrogen
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)↑
Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)↑
Cu, Ag, Au do not react with dilute acids (below hydrogen in reactivity series)
4. Metals + Salt Solutions → Displacement Reaction
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
A more reactive metal displaces a less reactive metal from its salt solution
Exam Tip
The order of reactivity with water is critical: Na/K (cold water, vigorous) > Ca (cold water) > Mg (hot water) > Fe/Zn (steam only) > Cu/Ag/Au (no reaction). This is tested every year.
Common Mistake
Students write that all metals react with dilute acids. Metals below hydrogen in the reactivity series (Cu, Ag, Au, Pt) do NOT react with dilute HCl or H₂SO₄.
Reactivity Series of Metals
The reactivity series is a list of metals arranged in decreasing order of their reactivity.
K > Na > Ca > Mg > Al > Zn > Fe > Ni > Sn > Pb > [H] > Cu > Hg > Ag > Au > Pt
Most reactive → Least reactive
Applications of the Reactivity Series:
- A more reactive metal can displace a less reactive metal from its salt solution
- Metals above hydrogen react with dilute acids to produce H₂; metals below hydrogen do not
- Metals at the top (K, Na, Ca) react vigorously with cold water
- Metals in the middle (Mg, Al, Zn, Fe) react with steam or hot water
- Metals at the bottom (Cu, Ag, Au) do not react with water at all
- The method of extraction of a metal depends on its position in the reactivity series
Memorisation Aid:
King Nathan Came More Active Zealous In New Shiny Land [H] Could He Secure Gold Platinum
(K, Na, Ca, Mg, Al, Zn, Fe, Ni, Sn, Pb, [H], Cu, Hg, Ag, Au, Pt)
Exam Tip
You MUST memorise the reactivity series — it's the backbone of this chapter. Many questions on displacement, extraction, and reactions with water/acid depend on it.
Common Mistake
Students forget that hydrogen is placed in the reactivity series as a reference. Metals above H react with dilute acids to give H₂; metals below H do not.
Extraction of Metals from Ores
Metals are found in nature as minerals. Minerals from which metals can be profitably extracted are called ores.
Steps in Extraction:
Step 1: Enrichment / Concentration of Ore
- Removal of impurities (gangue) from the ore
- Methods: hand-picking, washing, magnetic separation, froth floatation (for sulphide ores)
Step 2: Extraction — depends on position in reactivity series:
| Reactivity | Metals | Extraction Method |
|---|---|---|
| High (top) | K, Na, Ca, Mg, Al | Electrolytic reduction (electrolysis of molten ore) |
| Medium (middle) | Zn, Fe, Ni, Sn, Pb | Reduction with carbon (coke) after roasting/calcination |
| Low (bottom) | Cu, Hg, Ag, Au | Heating alone (self-reduction) or in air |
Important Reactions:
Roasting (heating sulphide ore in air): 2ZnS + 3O₂ → 2ZnO + 2SO₂
Calcination (heating carbonate ore in absence of air): ZnCO₃ → ZnO + CO₂
Reduction with carbon: ZnO + C → Zn + CO
Heating cinnabar (HgS): 2HgS + 3O₂ → 2HgO + 2SO₂, then 2HgO → 2Hg + O₂
Thermite Reaction:
Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat
Highly exothermic. Used for welding railway tracks and cracked machine parts (thermite welding).
Step 3: Refining — Electrolytic refining
- Anode: Impure metal
- Cathode: Pure metal (thin strip)
- Electrolyte: Salt solution of the same metal
- Impure metal dissolves from anode, pure metal deposits at cathode. Impurities settle as anode mud.
Exam Tip
Extraction method depends on reactivity: top → electrolysis, middle → carbon reduction, bottom → heating. The thermite reaction (Fe₂O₃ + 2Al) is a board exam favourite.
Common Mistake
Students confuse roasting and calcination. Roasting = heating SULPHIDE ore in AIR (gives oxide + SO₂). Calcination = heating CARBONATE ore WITHOUT air (gives oxide + CO₂).
Ionic Bonding — Formation of Ionic Compounds
Metals and non-metals combine by transfer of electrons to form ionic bonds (electrovalent bonds). The resulting compounds are called ionic compounds.
Example — Formation of NaCl:
Na (2, 8, 1) → Na⁺ (2, 8) + e⁻ (metal loses electron)
Cl (2, 8, 7) + e⁻ → Cl⁻ (2, 8, 8) (non-metal gains electron)
Na⁺ + Cl⁻ → NaCl (electrostatic attraction holds them together)
Example — Formation of MgCl₂:
Mg (2, 8, 2) → Mg²⁺ (2, 8) + 2e⁻
2Cl + 2e⁻ → 2Cl⁻
Mg²⁺ + 2Cl⁻ → MgCl₂
Properties of Ionic Compounds:
- Physical state: Solid at room temperature (strong electrostatic force of attraction)
- Melting and boiling points: Very high (requires large energy to break ionic bonds)
- Solubility: Generally soluble in water, insoluble in organic solvents (kerosene, petrol)
- Electrical conductivity: Do NOT conduct electricity in solid state (ions are fixed). Conduct electricity in molten state or when dissolved in water (ions are free to move)
Exam Tip
The formation of NaCl by electron transfer and properties of ionic compounds (high MP, conduct in molten/solution but not solid) are asked almost every year.
Common Mistake
Students say ionic compounds conduct electricity in solid state. They do NOT — ions are locked in fixed positions. They conduct only when molten or dissolved in water (ions become free).
Corrosion and Its Prevention
Corrosion is the slow deterioration of a metal by chemical action of air, moisture, or chemicals on its surface.
Rusting of Iron:
Iron + Oxygen + Water → Hydrated iron(III) oxide (rust)
4Fe + 3O₂ + xH₂O → 2Fe₂O₃·xH₂O (reddish-brown, flaky)
Requires BOTH moisture and oxygen
Other examples of corrosion:
- Silver: Black coating of Ag₂S (tarnishing — reacts with H₂S in air)
- Copper: Green coating of CuCO₃ (patina — reacts with CO₂ and moisture)
- Aluminium: Forms Al₂O₃ layer — but this layer is protective (prevents further corrosion)
Prevention of Corrosion:
- Painting / Oiling / Greasing: Prevents contact with air and moisture
- Galvanisation: Coating iron with zinc. Even if scratched, zinc corrodes first (sacrificial protection)
- Electroplating: Coating with chromium, nickel, or tin
- Alloying: Making stainless steel (Fe + Cr + Ni) — resistant to corrosion
- Anodising: Making a thick oxide layer on aluminium by electrolysis
Alloys — Important Examples:
| Alloy | Composition | Property/Use |
|---|---|---|
| Stainless steel | Fe + Cr + Ni + C | Rust-resistant, utensils |
| Brass | Cu + Zn | Decorative items, taps |
| Bronze | Cu + Sn | Statues, medals, coins |
| Solder | Pb + Sn | Welding electrical wires (low MP) |
| Amalgam | Hg + other metals | Dental fillings |
Key: Pure gold (24 carat) is very soft. Jewellery is made from 22 carat gold (alloyed with copper or silver) to make it harder.
Exam Tip
Alloy compositions are frequently asked: brass (Cu+Zn), bronze (Cu+Sn), stainless steel (Fe+Cr+Ni+C), solder (Pb+Sn). Also know: 24 carat = pure gold, 22 carat = jewellery gold.
Common Mistake
Students think aluminium corrodes easily because it's reactive. Aluminium forms a tough protective oxide layer (Al₂O₃) that PREVENTS further corrosion — unlike iron rust which flakes off.