Chapter 3 — Metals And Non Metals

What this topic covers
This overview introduces the basic physical and chemical differences between metals and non‑metals, how metals occur and are extracted, the reactivity series, and important processes such as corrosion and alloying.

Physical properties

• Metals: generally lustrous, good conductors of heat and electricity, malleable and ductile, usually solid at room temperature (except Hg), high melting and boiling points (many).
• Non‑metals: lack metallic luster, poor conductors (except graphite), brittle if solid, wide range of melting/boiling points, may be gases, liquids or solids.

Chemical behaviour

• Metals tend to lose electrons to form cations (e.g., M → M^{n+} + ne^-). Their oxides are usually basic; many react with oxygen, water and acids.
• Non‑metals tend to gain electrons to form anions or share electrons (covalent bonding). Non‑metal oxides are often acidic (e.g., CO2, SO2).

Occurrence and extraction
Metals are found as native metals (e.g., gold) or in ores — minerals from which metals are obtained. General steps in extraction: concentration of ore (removal of gangue), conversion to oxide (if needed), reduction of oxide to metal (by chemical reduction with carbon or by electrolytic reduction), and refining (e.g., electrolytic refining) to get pure metal.

Reactivity series (summary)
Metals are arranged by ease of losing electrons. Very reactive metals (K, Na, Ca, Mg, Al) react with water/air; moderately reactive (Zn, Fe) react with acids; less reactive (Cu, Ag, Au) do not react easily. Reactivity determines extraction method and displacement reactions.

Important applications and processes
Alloys (mixtures of metals and sometimes non‑metals) are made to improve properties (e.g., steel = Fe + C). Corrosion (e.g., rusting of iron) is the chemical deterioration of metals — prevented by painting, galvanizing, sacrificial protection. Amphoteric behaviour: some metal oxides/hydroxides (e.g., ZnO, Al2O3) react with both acids and bases.

How this helps you
Understanding these basics lets you predict reactions (metal + acid → salt + H2), choose extraction/refinement methods, and apply everyday concepts like why aluminium doesn’t rust easily or why iron objects need protection.

Definition: Metals are elements that generally show a characteristic set of physical properties such as lustre, malleability, ductility, thermal and electrical conductivity, and high density. These properties arise from their metallic bonding in which positive metal ions are surrounded by a 'sea' of delocalised electrons.

Why metals show these properties (Electron-sea model): In a metallic lattice each metal atom donates one or more valence electrons to a shared pool of delocalised electrons. The positive metal ions are held together by electrostatic attraction to this sea of electrons. Mobile electrons allow metals to conduct electricity and heat; the non-directional bonding allows layers of atoms to slide past one another (giving malleability and ductility); and the free electrons reflect light (giving lustre).

• Lustre (Shiny appearance): Free electrons absorb and re-emit visible light, so freshly cut metal surfaces appear shiny. Example: gold, silver, aluminium foil.
• Malleability: Metals can be beaten into thin sheets because metal ions can slide over each other without breaking the metallic bond. Example: aluminium is made into foil; gold beaten into gold leaf.
• Ductility: Metals can be drawn into wires as atoms slip in layers under tensile stress. Example: copper and aluminium wires in electrical wiring.
• Electrical conductivity: Due to mobile (delocalised) electrons, metals conduct electric current. Conductivity is high in copper, silver and aluminium. (Mercury is a metal but liquid at room temperature.)
• Thermal conductivity: Free electrons transfer kinetic energy rapidly through the lattice, so metals are good conductors of heat. Example: cooking pans are made of copper or aluminium.
• Sonorous: Metals produce sound when struck because the regular lattice transmits vibrations well. Example: bells and musical cymbals.
• High density, melting and boiling points: Most metals have high densities and relatively high melting/boiling points due to strong metallic bonds (exceptions exist: sodium, potassium have low melting points; mercury is liquid).
• Crystalline and opaque: Most metals are crystalline solids at room temperature and are opaque (they do not transmit light).

Typical exceptions & remarks: Mercury (Hg) is a metal that is liquid at room temperature. Some metals (alkali metals) are soft and have low melting points. Alloys (e.g., brass, bronze, steel) modify physical properties — e.g., steel is harder and less ductile than pure iron.

Summary (cause → effect): Metallic bonding (delocalised electrons + positive ion lattice) → free electrons & flexible/non-directional bonds → electrical/thermal conductivity, lustre, malleability, ductility, sonorous nature, etc.

Overview: Metals are elements that tend to lose electrons and form positive ions (cations). Their chemical properties are governed by this tendency to be oxidized. Metals react with oxygen, water, acids, salts and non-metals to form ionic compounds (oxides, hydroxides, salts) and show characteristic displacement and redox behaviour.

Chemical properties (with concise explanations)

• Reaction with oxygen: Most metals combine with oxygen to form metal oxides. These oxides are usually basic (basic oxides). Example general reaction: 2M + O2 → 2MO. Some metals (like Al, Zn) form oxides that are amphoteric.
• Reaction with water: Very reactive metals (alkali metals, some alkaline earth metals) react with cold water to give hydroxides and hydrogen. Less reactive metals react only with steam (hot water vapor) to form metal oxides and hydrogen.
  • Cold water: M + H2O → M(OH)n + H2 (for reactive metals)
  • Steam: M + H2O(g) → MO + H2 (for metals like Mg)
• Reaction with acids: Metals above hydrogen in the reactivity series displace hydrogen from dilute acids to give a salt and hydrogen gas.
  • General: M + acid → M salt + H2  (e.g., Zn + H2SO4 → ZnSO4 + H2)
• Displacement (single displacement) reactions): A more reactive metal displaces a less reactive metal from its salt solution.
  • General: M1 + M2X → M1X + M2  (e.g., Zn + CuSO4 → ZnSO4 + Cu)
• Reaction with non-metals (halogens): Metals react with halogens to form ionic halides (salts). Example: 2Na + Cl2 → 2NaCl.
• Corrosion and tarnishing: Metals like iron corrode (rust) in presence of moisture and oxygen to form hydrated iron(III) oxide (rust). Silver tarnishes by forming silver sulfide on exposure to H2S in air. Noble metals (gold, platinum) are very unreactive.
• Amphoteric behaviour: Some metal oxides/hydroxides (e.g., Al2O3, ZnO, PbO) react with both acids and bases to give salts and complex ions. This shows they are amphoteric.
• Redox view: Chemical reactions of metals are oxidation reactions (metal atoms lose electrons): M → M^{n+} + ne−. The electrons reduce another species (O2, H+, metal ion).

Importance (CBSE context)

These chemical properties explain reactivity trends (reactivity series), methods of metal extraction, storage/handling (prevent corrosion), and common laboratory tests (hydrogen evolution, displacement reactions) used in Class 10 experiments.

Definition: Non-metals are elements that generally lack the typical physical properties of metals. They are found on the right side of the periodic table and occur in all three states: gases, liquids and solids.

General physical properties:

• States: Non-metals exist as solids (carbon, sulfur, iodine), liquids (bromine) and gases (oxygen, nitrogen, chlorine).
• Appearance and lustre: Most non-metals are dull (non-lustrous). Exceptions: iodine and graphite show metallic shine.
• Hardness: Wide range: diamond (a form of carbon) is extremely hard; many non-metal solids (sulfur, phosphorus) are soft and brittle.
• Malleability and ductility: Non-metals are generally not malleable or ductile; they are brittle and break or powder on hammering.
• Density: Usually lower densities than metals (many non-metals are light gases or low-density solids), though there are exceptions.
• Melting and boiling points: Generally low for molecular non-metals (e.g., O2, N2, Cl2), but covalent-network solids (diamond, graphite) have very high melting points.
• Electrical conductivity: Most non-metals are poor conductors of electricity. Important exception: graphite conducts electricity due to delocalised electrons in its layers.
• Thermal conductivity: Non-metals are generally poor conductors of heat (insulators).
• Magnetism and sonority: Non-metals are not magnetic and are not sonorous (do not produce sound when struck).

Why these properties occur (brief explanation): Non-metals have high ionisation energies and high electronegativities; they tend to gain electrons and form covalent bonds. Many exist as small molecules (weak intermolecular forces → low melting/boiling points) or as covalent networks (strong directional bonds → high hardness and high melting points). Graphite’s electrical conductivity arises from delocalised pi electrons within its layered structure.

Common comparisons with metals: Compared with metals, non-metals are poor conductors of heat/electricity, usually brittle (not malleable/ductile), often low-density and often found in gaseous form at room temperature.

Important exceptions to note (Class 10 level): Diamond (a non-metal) is extremely hard and has very high melting point; graphite (a non-metal) is lustrous, soft, slippery and conducts electricity.

Non-metals show characteristic chemical behaviour that is different from metals. Most non-metals have high ionization energies and high electronegativities, so they tend to gain or share electrons rather than lose them. Their common chemical properties include the formation of covalent hydrides, reaction with oxygen to form acidic oxides, reaction with metals to form ionic or polar compounds, displacement reactions among halogens, and reactions with water and dilute acids. Exceptions (e.g., graphite conducting electricity) arise from special structures.

1. Reaction with hydrogen

• Non-metals combine with hydrogen to form covalent hydrides (e.g., HCl, NH3, CH4). These hydrides can be acidic, basic or neutral depending on the element (HCl is acidic, NH3 is basic, CH4 is neutral).

2. Reaction with oxygen (combustion / oxidation)

• Non-metals burn in oxygen to form oxides which are generally acidic (acidic oxides). Examples: C + O2 → CO2, S + O2 → SO2. These oxides react with water to give acids (CO2 + H2O → H2CO3, SO2 + H2O → H2SO3).

3. Reaction with metals

• Non-metals react with metals to form ionic or polar compounds (metal halides, metal oxides, sulfides). Example: 2Na + Cl2 → 2NaCl. Non-metal atoms gain electrons and form anions in such reactions.

4. Reaction with water and dilute acids/bases

• Some non-metal oxides dissolve in water to form acids (acidic behaviour). Chlorine disproportionates in water: Cl2 + H2O → HCl + HClO. CO2 dissolves to form carbonic acid (H2CO3).

5. Displacement reactions (halogen reactivity)

• More reactive halogens displace less reactive halogens from their salts: Cl2 + 2KBr → 2KCl + Br2. A halogen reactivity order (typical): F2 > Cl2 > Br2 > I2.

6. Acidic character of oxides

• Non-metal oxides are generally acidic (react with bases to form salts). Example: SO2 + 2NaOH → Na2SO3 + H2O; CO2 + 2NaOH → Na2CO3 + H2O.

7. Electrical conductivity and bonding

• Most non-metals are poor conductors of electricity in solid state due to localized electrons and covalent bonding (exceptions: graphite conducts due to delocalized electrons in layers).

Why these behaviours occur (brief theory)

• High electronegativity and electron affinity: non-metals attract electrons, which leads to formation of anions or to sharing (covalent bonds).
• Small atomic size and high ionization enthalpy discourage formation of cations, so non-metals prefer to gain or share electrons.
• Oxides formed have acidic character because non-metal atoms bonded to oxygen stabilize higher oxidation states and yield oxyacids in water.

Points to remember (Class 10 focus)

• Non-metals form acidic oxides; metals form basic oxides.
• Non-metals form covalent compounds with other non-metals and ionic/polar compounds with metals.
• Halogen displacement tests are useful to determine relative reactivity.
• Real-life uses (bleaching, disinfecting, fertilizers, preservatives) arise from their chemical reactivity.

Definition: The reactivity series of metals is an arrangement of metals in order of decreasing chemical reactivity. The most reactive metals are placed at the top and the least reactive (noble) metals at the bottom. Hydrogen is included to compare metals with non-metal reactivity toward acids.

Common CBSE Class 10 series (high to low):

• K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Ag > Au > Pt

Key points and significance:

• Displacement reactions: A metal higher in the series can displace a metal lower from its salt solution. (Example: Zn displaces Cu2+ from CuSO4.)
• Reaction with water and acids: Very reactive metals (K, Na, Ca) react with cold water; moderately reactive (Mg, Al) react with steam or on heating; metals above H react with dilute acids to produce hydrogen gas.
• Extraction of metals: The position in the series helps decide the extraction method. Metals above carbon (C) are usually obtained by electrolysis of their molten compounds; metals below carbon can often be extracted by reduction with carbon (smelting).
• Corrosion and protection: More reactive metals corrode more readily. Some reactive metals (e.g., Al) form a protective oxide layer. Less reactive metals (Cu, Ag, Au) are used where corrosion resistance is needed.
• Practical uses: Metals low in the series (Cu, Ag, Au) are used for electrical wiring, jewellery and coins; very reactive metals are used where sacrificial protection or strong reducing behavior is needed (e.g., Mg anodes).

Rules and quick tests:

• If metal A is above metal B in the series, then A will displace B from a solution of B's salts: A + Bx+ <–> displacement.
• Metals above H react with dilute acids to give H2; metals below H do not.
• Choice of extraction: if metal > C (more reactive than carbon) → electrolysis; if metal < C → reduction by carbon possible.

Exceptions & notes: Aluminium is high in the series but its oxide layer prevents easy reaction unless oxide is removed; noble metals (Au, Pt) show very low chemical reactivity.

Mineral: A naturally occurring inorganic substance with a definite chemical composition and crystalline structure (e.g., quartz, pyrite, bauxite).

Ore: A mineral or mixture of minerals from which a metal can be economically extracted. An ore contains the metal-bearing material plus impurities called gangue (e.g., hematite Fe2O3 and magnetite Fe3O4 are ores of iron).

Key terms

• Gangue: Non-metallic impurities (silica, sand, clay).
• Concentration (enrichment): Removal of gangue; methods depend on physical/chemical properties of the ore.
• Calcination: Heating carbonate ores in limited air to remove volatile impurities and CO2 (MCO3 → MO + CO2).
• Roasting: Heating sulphide ores in excess air to convert them to oxides and produce SO2 (2MS + O2 → 2M O + SO2).
• Flux and slag: Flux (e.g., CaCO3) combines with gangue (e.g., SiO2) to form slag (e.g., CaSiO3), which is removed.

Types of minerals/ores (common classes)

• Native elements: Metals found in elemental form (gold, silver, platinum).
• Oxides: Hematite (Fe2O3), magnetite (Fe3O4), rutile (TiO2).
• Sulphides: Galena (PbS), sphalerite (ZnS), chalcopyrite (CuFeS2).
• Carbonates: Calcite (CaCO3), malachite (CuCO3·Cu(OH)2).
• Hydroxides and ores like bauxite (Al2O3·2H2O) for aluminium.

Concentration methods (brief)

• Gravity (or hydraulic) separation: For dense minerals (gold, tin) using flowing water.
• Magnetic separation: For magnetic ores (magnetite).
• Froth flotation: For fine sulphide ores (copper, lead, zinc); collectors make ore particles hydrophobic so they attach to air bubbles and float.
• Leaching: Using chemical solvents to dissolve the metal (e.g., cyanide leaching for gold, ammonia for copper).

From concentrated ore to metal

• Roasting/calcination to convert ore to oxide or remove volatile impurities.
• Reduction: Less reactive metals are reduced using carbon (C) or carbon monoxide (CO) in furnaces (e.g., iron in blast furnace). Highly reactive metals require electrolysis (e.g., aluminium).
• Refining: Purification of extracted metal (electrorefining for copper).

Which extraction method is used? It depends on the metal's reactivity (reactivity series):

• Low reactivity metals (below carbon): reduction by carbon.
• Moderately reactive: chemical reduction/roasting then reduction.
• Highly reactive metals (above carbon): electrolytic reduction of molten compounds.

Environmental & practical notes

• Roasting of sulphide ores releases SO2 → acid rain; cleaning/processing of gases is required.
• Mining and ore processing create waste (tailings, slag) that must be managed.

Summary (class-10 focus): A mineral is a natural substance; an ore is a mineral containing a metal that can be profitably extracted. Concentration (gravity, magnetic, froth flotation) removes gangue. Calcination and roasting prepare ores for reduction. The method of extraction depends on the metal's reactivity; fluxes are used to form removable slag.

Definition: Concentration (enrichment) of an ore is the process of removing impurities (gangue) from the ore to increase the percentage of the metal compound before further extraction.

Why it is needed: Raw ores contain valuable metal-bearing minerals mixed with unwanted materials (sand, clay, rock). Concentration reduces the amount of material to be processed, lowers cost, saves energy, and increases metal recovery.

Key terms:

• Ore: Rock that contains a useful metal compound in sufficient amounts to make extraction worthwhile.
• Gangue: Non-metallic impurities associated with the ore.
• Grade: Percentage of metal present in the ore.

Methods of concentration (CBSE Class 10 scope):

• Hand picking – Manual separation of large and visible impurities from the ore. Used when ore and gangue are distinguishable by sight or size. Example: coal, copper-tin ores in traditional mines.
• Winnowing – Uses wind (air current) to remove lighter impurities from powdered/heavy ore. Mostly used in agriculture and in few ore-processing cases where particles differ in density/size.
• Gravity separation (hydraulic washing / panning) – Uses flowing water to carry away lighter gangue; heavier ore particles settle. Used for alluvial gold, tin (cassiterite), and coal. Example: panning for gold in riverbeds.
• Magnetic separation – Magnetic ore (e.g., magnetite, Fe3O4) is separated from non-magnetic gangue by magnets. Widely used for iron ores containing magnetite.
• Froth flotation – Used for fine particles of sulphide ores (e.g., chalcopyrite CuFeS2, galena PbS, sphalerite ZnS). The powdered ore is mixed with water and reagents (collectors like xanthates, frothers like pine oil). Sulphide particles attach to air bubbles and rise as froth which is skimmed off. Non-sulphide gangue sinks.
• Leaching – Uses a solvent to dissolve the desired metal from the ore. The metal is later recovered from the solution. Common for low-grade or soluble ores (e.g., cyanide leaching of gold, heap leaching of copper). Example reaction for gold cyanidation: 4Au + 8NaCN + O2 + 2H2O → 4Na[Au(CN)2] + 4NaOH.

General steps in concentration: Crushing → Grinding → Size separation (sieves) → Method-specific treatment (gravity, magnetic, flotation) → Concentrate collection.

Practical considerations and limitations: Choice of method depends on ore type (oxide, sulphide), particle size, physical properties (density, magnetic susceptibility), and economics. Some methods require chemical reagents and generate waste (tailings) that must be managed.

Relationship to metal extraction: Concentration is the first step in metallurgy. After enrichment, the concentrated ore is subjected to chemical processes (roasting, reduction, electrolysis) to obtain the metal.

Overview
The extraction of iron in a blast furnace is a continuous high-temperature process in which iron is obtained from iron ores (mainly haematite, Fe₂O₃) using coke (carbon) as reducing agent and limestone (CaCO₃) as a flux. Raw materials are charged at the top and a hot blast of air is blown in near the bottom; gases move up while solids move down (counter-current operation).

Raw materials

• Iron ore: chiefly haematite (Fe₂O₃) or magnetite (Fe₃O₄).
• Coke: acts as fuel and reducing agent (mainly carbon).
• Limestone (CaCO₃): acts as flux to remove silica and other impurities as slag.
• Hot air blast: pre-heated air injected near the bottom to maintain high temperatures.

Furnace zones and temperatures

• Top (cooler) zone: rough temperature a few hundred °C; ore drying and pre-heating occur.
• Middle (reduction) zone: around 500–900 °C where gaseous reduction mainly occurs (Fe₂O₃ → FeO → Fe).
• Bottom (combustion/temperature) zone: up to ~1500 °C; coke combusts and molten iron & slag collect in the hearth.

Key steps and reactions

• 1. Combustion of coke to generate heat:
  • C + O₂ → CO₂
  • CO₂ + C → 2CO (formation of carbon monoxide)
• 2. Reduction of iron oxides by CO (primary reducing agent):
  • Fe₂O₃ + 3CO → 2Fe + 3CO₂
• 3. Direct reduction by carbon at high temperature (minor):
  • Fe₂O₃ + 3C → 2Fe + 3CO
• 4. Fluxing to remove silica impurities:
  • CaCO₃ → CaO + CO₂ (decomposition of limestone)
  • CaO + SiO₂ → CaSiO₃ (formation of molten slag)

Products
The molten heavier iron (pig iron) collects at the bottom and is tapped periodically. The lighter molten slag (calcium silicate) floats on top and is removed separately. Off-gases—rich in CO—are cleaned and often used as fuel elsewhere in the plant.

Important features & practical notes

• The process is continuous and operates at very high temperatures; efficiency depends on ore quality, coke quality, blast temperature and furnace design.
• Almost all commercial iron is first produced as pig iron in the blast furnace and then converted to steel (e.g., in a basic oxygen furnace or by electric methods).
• Environmental considerations: CO and CO₂ emissions; modern plants recover waste gases and treat effluents.

Introduction: Aluminium is an abundant metal obtained commercially from the ore bauxite. Pure aluminium metal is produced by two main stages: (1) refining bauxite to give pure aluminium oxide (alumina) by the Bayer process, and (2) electrolytic reduction of alumina to aluminium metal by the Hall–Héroult process.

1. Occurrence and ore: Bauxite is the principal ore of aluminium and contains hydrates/oxides of aluminium such as gibbsite (Al(OH)3), boehmite (AlO(OH)) and diaspore, together with impurities (Fe2O3, SiO2, TiO2).

2. Bayer process (to obtain pure Al2O3):

1. Crushing and grinding: Bauxite is crushed and ground to a fine powder.
2. Digestion: Powdered bauxite is treated with concentrated NaOH at high temperature and pressure. Aluminium compounds dissolve as sodium aluminate.
3. Clarification: Insoluble impurities (red mud) are removed by settling/filtration.
4. Precipitation: Aluminium hydroxide is precipitated from sodium aluminate solution by seeding with Al(OH)3 crystals.
5. Calcination: Precipitated Al(OH)3 is heated (900–1000°C) to remove water and give anhydrous alumina (Al2O3).

Why these steps? NaOH selectively dissolves aluminium compounds but not iron oxides; precipitation and calcination yield pure, dry Al2O3 suitable for electrolysis.

3. Hall–Héroult process (electrolytic reduction):

Pure Al2O3 has a very high melting point (~2072°C) and is not practical to melt alone. It is therefore dissolved in molten cryolite (Na3AlF6) which: lowers the working temperature to ~950°C, increases conductivity and reduces energy consumption.

1. Electrolytic cell: The cell contains molten cryolite with dissolved alumina. Carbon (graphite) blocks serve as the anodes; the cathode is the carbon-lined cell. A strong current is passed through the molten mixture.
2. Reactions: At the cathode Al3+ ions gain electrons and deposit as liquid aluminium. Oxygen ions migrate to the carbon anode where they react with carbon to form CO2, consuming the anode.

Key points: Electrolysis is energy-intensive (roughly 13–15 kWh per kg Al). Carbon anodes are consumed and emit CO2. Recycled aluminium requires only ~5% of the energy needed for primary extraction, so recycling is environmentally and economically important.

Environmental and economic aspects: Hall–Héroult emits CO2 from anode consumption; red mud disposal from Bayer process is an environmental challenge. Large-scale aluminium production therefore focuses on energy efficiency, anode alternatives and recycling.

Overview
Extraction of metals means obtaining a metal in pure form from its ore. Metals are extracted by different methods depending on their reactivity and the type of ore (oxide, carbonate or sulfide). Copper is an important example because it occurs both as sulfide and oxide ores and uses almost all classical steps: concentration, roasting, smelting and refining.

Occurrence and concentration

• Most metals are found as oxides, sulphides or carbonates. Copper common ores: chalcopyrite (CuFeS₂), cuprite (Cu₂O), malachite (CuCO₃. Cu(OH)₂).
• Concentration (removal of gangue): methods include gravity separation (for heavy minerals), magnetic separation (for magnetic impurities) and froth flotation (common for sulphide ores). Froth flotation: powdered ore + water + collector + frothing agent -> hydrophobic sulphide particles attach to air bubbles and float.

Extraction of copper from sulfide ores (typical industrial route)

• Concentration (froth flotation) to get enriched sulphide concentrate.
• Roasting: partial oxidation of sulphide to oxide and release of SO₂. Example (for CuS): 2CuS + 3O₂ → 2CuO + 2SO₂. (Industrial roasting of chalcopyrite is more complex but gives oxides, sulphates and SO₂.)
• Smelting: heated with flux (silica) to separate iron as slag. Iron oxides form silicate slag (FeO + SiO₂ → FeSiO₃) and a copper-containing matte (sulphides) is produced.
• Conversion (or reduction): Cu₂S (matte) is oxidised to give impure copper (blister copper). Simplified: Cu₂S + O₂ → 2Cu + SO₂.
• Refining (electrolytic refining): impure copper is made anode, pure copper deposited at cathode. Anode reaction: Cu (anode) → Cu²⁺ + 2e^-. Cathode reaction: Cu²⁺ + 2e^- → Cu (cathode). This gives high-purity copper.

Extraction of copper from oxide/carbonate ores (hydrometallurgy)

• Oxide ores (CuO, Cu₂O) are leached with dilute sulphuric acid: CuO + H₂SO₄ → CuSO₄ + H₂O.
• The copper sulfate solution can be purified and subjected to electrolysis or extracted by solvent extraction and electrowinning (SX–EW).
• Displacement method (laboratory): CuSO₄ + Fe → FeSO₄ + Cu (iron displaces copper from solution).

Electrolytic refining details
In the electrolytic cell: impure copper anode, pure copper cathode, and copper sulphate solution as electrolyte. Impurities less reactive than Cu (like Ag, Au) fall off as anode mud and are recovered; more reactive impurities remain in solution.

Extraction of some other metals (brief)

• Iron: Extracted from haematite (Fe₂O₃) in a blast furnace. Reduction by CO: Fe₂O₃ + 3CO → 2Fe + 3CO₂. Limestone (CaCO₃) acts as flux to remove SiO₂ as CaSiO₃ slag.
• Aluminium: Bauxite (Al₂O₃·xH₂O) is converted to pure Al₂O₃ (Bayer process), then reduced by electrolysis (Hall–Héroult) in molten cryolite. Cathode: Al³⁺ + 3e^- → Al.
• Zinc: Roasting of ZnS to ZnO, then reduction (pyrometallurgy) or hydrometallurgy. Example: ZnO + C → Zn + CO.

Key points to remember

• Choice of extraction method depends on metal reactivity: highly reactive metals (Na, K, Al) are obtained by electrolysis; moderately reactive metals (Zn, Fe, Cu) by reduction; less reactive metals by displacement from solutions.
• Fluxes (like silica, limestone) help remove gangue as slag.
• Electrolytic refining gives very pure metal and recovers noble impurities.
• Pollution issues: roasting and smelting release SO₂ which must be converted to H₂SO₄ or managed to prevent acid rain.

What is refining? Refining (or purification) of metals is the process of removing impurities from crude or impure metal obtained after extraction, to obtain metal of required purity and properties. Impurities change physical and chemical properties (conductivity, malleability, colour, strength), so refining is essential before commercial use.

Main methods of refining (Class 10 level)

1. Electrolytic refining

   Principle: Electrolysis. An impure metal piece is used as the anode and a pure sheet of the same metal as the cathode. The electrolyte is a solution of a salt of that metal. On passing current the metal from anode goes into solution and gets deposited on the cathode as pure metal. Insoluble impurities fall off and collect as anode mud (often containing precious metals).

   Example: Copper—used to produce 99.9% pure copper for electrical wires. Main reactions (copper):

   At anode: Cu(s) → Cu2+(aq) + 2e−
   At cathode: Cu2+(aq) + 2e− → Cu(s)

2. Distillation

   Principle: Separation by differences in boiling points. Metals (or their vapours) with low boiling points are vaporized and condensed separately from less volatile impurities.

   Examples: Mercury (from cinnabar historically) and zinc (in some processes) can be purified by distillation/condensation because they vaporize at relatively low temperatures.

3. Liquation

   Principle: Melting-point difference. Used when one component has a much lower melting point; the low-melting metal is melted off or runs away from higher-melting impurities.

   Typical application: Separation of low-melting metals (like tin or lead) from alloys/impure mixtures in traditional metallurgy.

4. Zone refining (brief advanced note)

   Principle: A narrow molten zone is moved along a solid rod; impurities concentrate in the melt and are carried along, leaving behind an ultrapure solid. Widely used to obtain extremely pure silicon and germanium for semiconductors.

Why electrolytic refining is widely used: It gives very high purity, is controllable, and recovers valuable impurities (in anode mud). It is especially important for metals used in electrical and electronic applications (e.g., copper, silver).

Practical notes & safety: Distillation of mercury is hazardous (toxic vapour) and now restricted; modern industries follow strict environmental and safety norms and often use alternative processes.

Summary: Choose the refining method based on the nature of the metal and impurity (volatility, melting point, desired purity). Electrolytic refining for high purity solids (e.g., copper), distillation for volatile metals (mercury, zinc), liquation for low-melting components, and zone refining for semiconductor-grade purity.

Alloys

Definition: An alloy is a homogeneous mixture (solid solution or a mixture of phases) of two or more elements, where at least one is a metal. Alloys are designed to combine desirable properties of their constituents.

Why alloys are made: To improve strength, hardness, toughness, corrosion resistance, wear resistance, melting point control, or to reduce cost and weight compared with pure metals.

How alloys are prepared

Alloys are generally prepared by melting the base metal, adding the alloying elements in required proportions, mixing thoroughly and allowing the mixture to solidify. Rapid cooling or controlled cooling can change the final structure (microstructure) and properties.

Types of alloys

• Substitutional alloys: Solute atoms replace host metal atoms in the crystal lattice (e.g., brass: Cu and Zn of similar atomic size).
• Interstitial alloys: Smaller atoms occupy interstitial spaces between host atoms (e.g., steel: C atoms in Fe lattice).
• Ferrous alloys: Based on iron (Fe) — steels and cast irons.
• Non-ferrous alloys: Based on other metals — brass, bronze, duralumin, etc.

Key features and effects

• Alloys are usually harder and stronger than their component pure metals because impurity or alloying atoms distort the lattice and hinder dislocation motion.
• Electrical and thermal conductivity usually decrease on alloying (compared to pure metals) because electron scattering increases.
• Melting behaviour: Alloys can have a range of melting temperatures (solidus to liquidus). Some mixtures have a eutectic composition with a sharp melting point lower than either pure component.
• Properties depend on composition and cooling history (microstructure: grains, phases, precipitates).

Design rules (qualitative)

Hume-Rothery rules guide the formation of substitutional solid solutions: similarity in atomic radii, crystal structure, electronegativity and valency favors extensive solubility of one metal in another.

Practical notes (Class 10 level)

• Steel is iron with controlled carbon and other alloying elements. Small % of carbon (0.02–2%) changes hardness and ductility markedly.
• Stainless steel contains chromium (≈11–20%) and often nickel; chromium forms a passive oxide layer that prevents corrosion.
• Brass (Cu + Zn) and bronze (Cu + Sn) are important copper alloys with different mechanical and corrosion properties.

Corrosion is the gradual destructive change in a metal caused by chemical or electrochemical reaction with its environment. It applies to many metals (e.g., copper, aluminium, iron). Rusting is a specific kind of corrosion that occurs only for iron and its alloys (steel) in presence of oxygen and moisture.

Why and how rusting occurs (electrochemical mechanism)

• On an iron surface small regions become anodic and others cathodic. At the anodic sites iron is oxidised: Fe → Fe2+ + 2e−.
• The released electrons travel through the metal to cathodic sites where oxygen is reduced in presence of water: O2 + 2H2O + 4e− → 4OH−.
• Fe2+ ions combine with OH− to form ferrous hydroxide: 2Fe2+ + 4OH− → 2Fe(OH)2. This is further oxidised to hydrated iron(III) oxide (Fe2O3·xH2O) — the flaky brown substance called rust.
• Overall (simplified) reaction: 4Fe + 3O2 + xH2O → 2Fe2O3·xH2O (rust).

Factors accelerating rusting

• Presence of moisture/water and dissolved ions (electrolytes) — e.g., salt (NaCl) increases conductivity and speeds rusting.
• Acidic conditions (low pH) increase corrosion rates.
• Temperature — higher temperature generally increases reaction rates.
• Mechanical stress, impurities and dissimilar metal contact (galvanic corrosion).

Differences: corrosion vs rusting

• Corrosion: general term for deterioration of metals by chemical/electrochemical reactions (affects copper, aluminium, zinc, etc.).
• Rusting: specific to iron/steel and produces hydrated iron(III) oxide (rust).

Methods to prevent corrosion and rusting

• Coatings: paint, enamel, varnish prevent contact with air and water.
• Oiling/greasing: forms a water-proof film (used for tools, machinery parts).
• Galvanization: coating iron with zinc. Zinc corrodes preferentially (sacrificial anode) protecting iron.
• Sacrificial protection: attaching a more reactive metal (e.g., magnesium or zinc) that corrodes instead.
• Cathodic protection: applying an external impressed current to make the metal the cathode.
• Alloying and passivation: stainless steel (Cr) forms an inert oxide layer; aluminium forms a protective Al2O3 layer.
• Electroplating: depositing a protective metal layer (e.g., chromium plating).

Practical importance

Rust weakens structures (bridges, buildings, vehicles), increases maintenance costs, causes leaks in pipelines, and can cause safety hazards. Controlling corrosion is essential in engineering, transport, and everyday life.

What the topic covers
Representative chemical reactions and equations describe typical ways in which metals and non‑metals react and how to write balanced chemical equations (including ionic and net ionic forms) that represent those reactions. The emphasis is on types of reactions (combination, decomposition, displacement, double displacement, redox), predicting products using the reactivity series, and showing electron transfer where relevant.

Key ideas

• Types of reactions: combination (synthesis), decomposition, single displacement, double displacement (metathesis), and redox reactions.
• Reactivity series: ranks metals by chemical reactivity and predicts whether a metal will displace another from a compound or react with water/acids.
• Balancing equations: atoms and charge must be conserved; show physical states (s, l, g, aq) and catalysts if present.
• Ionic and net ionic equations: useful for aqueous reactions — separate strong electrolytes into ions and cancel spectator ions to show the actual chemical change.
• Oxidation and reduction: oxidation = loss of electrons (increase in oxidation number), reduction = gain of electrons (decrease in oxidation number). Metal atoms commonly oxidize to form cations; non‑metals often reduce to form anions.

Representative reaction classes for metals & non‑metals

• Metal + oxygen → metal oxide (combination; often exothermic). Example: 2Mg(s) + O2(g) → 2MgO(s).
• Metal + water → metal hydroxide (or oxide) + H2 (reactivity depends on metal). Example: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g).
• Metal + acid → salt + H2 (single displacement/redox). Example: Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g).
• Metal displacement (single displacement): A + BC → AC + B when A is higher in reactivity than B. Example: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s).
• Double displacement (precipitation): AB(aq) + CD(aq) → AD + CB, often forms an insoluble precipitate. Example: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq).
• Non‑metal reactions (e.g., halogen displacement): Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(l).
• Decomposition of carbonates and hydroxides on heating (thermal decomposition): MCO3(s) → MO(s) + CO2(g); 2Cu(OH)2(s) → CuO(s) + H2O(l) + 1/2 O2(g) (balanced appropriately).

Writing ionic and net ionic equations

• Write the full balanced molecular equation with states.
• For aqueous strong electrolytes, split into ions (complete ionic equation).
• Cancel spectator ions to obtain the net ionic equation, which shows the actual chemical change (e.g., Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)).

Practical points & safety

• Predicting products relies on the reactivity series: very reactive metals (like Na, K) react with water; moderate metals (Mg, Zn) react with acids; less reactive metals (Cu, Ag, Au) do not displace H+ from acids and often do not react with water.
• Many reactions release heat or may be violent (combustion of alkali metals in water), so perform demonstrations only with proper safety.

Introduction: Metals and non-metals are two broad classes of elements that we use every day. Metals (like iron, copper, aluminium) are typically lustrous, good conductors of heat and electricity, malleable and ductile. Non-metals (like oxygen, nitrogen, sulphur, chlorine) are usually poor conductors and show a wide variety of physical states and chemical behaviour.

Common properties (brief):

• Metals: high electrical and thermal conductivity, malleability, ductility, metallic lustre, usually solid at room temperature (except mercury).
• Non-metals: poor conductors (insulators), brittle if solid, varied states (gases like O2, solids like S, liquids like Br2), often gain electrons in reactions.

Uses in daily life: Metals and non-metals are chosen for uses based on their physical and chemical properties.

• Metals:
  • Iron/Steel: construction (beams, reinforcement), automobiles, tools due to strength and toughness.
  • Copper: electrical wiring and motors because of excellent electrical conductivity.
  • Aluminium: cooking utensils, aircraft parts, foils — light weight and corrosion-resistant oxide layer.
  • Gold/Silver/Platinum: jewellery, electronics (gold contacts) — corrosion resistance and conductivity.
  • Alloys: brass (Cu + Zn) for musical instruments, bronze (Cu + Sn) for bearings, stainless steel (Fe + Cr + Ni) for cutlery.
• Non-metals:
  • Oxygen: respiration, medical oxygen, oxidation reactions in industry.
  • Nitrogen: inert atmosphere for food packaging, ammonia production for fertilisers.
  • Carbon (graphite): electrodes, lubricants, pencil leads; carbon (diamond) in cutting tools.
  • Chlorine: water disinfection (swimming pools), production of many compounds (PVC).
  • Phosphorus and sulphur: fertilizers, matches, dyes, sulphuric acid — industrial chemicals.

Important chemical behaviours (why used):

• Good electrical conductors (metals like Cu, Al) are used in wiring and electronics.
• Reactive metals (Na, K) are used in chemical syntheses; less reactive metals (Au, Pt) are used where corrosion resistance is essential.
• Non-metals such as chlorine and oxygen participate in important disinfecting and oxidation processes.

Environmental aspects:

• Extraction impacts: Mining for metal ores disturbs land, causes habitat loss, soil erosion and generates large waste (tailings). Many ores require energy-intensive refining (e.g., aluminium from bauxite consumes large electricity).
• Air and water pollution: Smelting and ore processing release sulphur dioxide (SO2), particulate matter and heavy metals into air and water. Acid mine drainage (sulphide oxidation) lowers pH and releases toxic metals into waterways.
• Heavy metal toxicity: Lead, mercury, cadmium, and arsenic accumulate in organisms (bioaccumulation and biomagnification), causing neurological, renal and developmental damage.
• Corrosion and waste: Metal corrosion (e.g., rusting of iron) shortens lifespan of infrastructure; improper disposal of metal-containing products (batteries, electronics) leads to soil and water contamination.
• Non-metal pollution: Combustion of fossil fuels releases CO2, SO2, and NOx, causing climate change, acid rain and respiratory problems. Chlorinated organics and other non-metal compounds can be persistent pollutants.
• Energy footprint: Production of some metals (especially aluminium) is energy-intensive, contributing to greenhouse gas emissions unless low-carbon electricity is used.

Mitigation and sustainable practices:

• Recycling metals (steel, aluminium, copper) saves energy and reduces mining. Example: recycling aluminium uses ~95% less energy than producing from bauxite.
• Proper disposal and treatment of electronic waste to recover precious metals and prevent toxic leaks.
• Pollution control at smelters (filters, scrubbers) to capture SO2 and particulates; treatment of acid mine drainage (neutralisation, constructed wetlands).
• Substitution and green chemistry: using less toxic materials, designing products for reuse and easier recycling.
• Regulation and monitoring: emission standards, safe limits for heavy metals in water and soil, public awareness.

Summary: Metals and non-metals are indispensable in daily life due to complementary properties — metals for strength and conduction, non-metals for biological roles and chemical reactivity. However, their extraction, use and disposal present environmental challenges that can be reduced by recycling, cleaner production methods and responsible policy.