Answer — key differences (physical):

• Lustre: Metals are generally lustrous (shiny); non-metals are dull (except graphite).
• Conductivity: Metals conduct heat and electricity well due to free electrons; most non-metals are poor conductors (graphite is an exception for electricity).
• Mechanical properties: Metals are malleable and ductile (can be hammered or drawn into wires); non-metals (solid) are brittle and break when hammered.

These distinctions are important for identifying substances and for practical uses (e.g., metals in wiring, non-metals as insulators).

Answer — chemical differences:

• Oxides: Metals form basic oxides (e.g., MgO — basic); non-metals form acidic oxides (e.g., SO₂ — acidic).
• Reaction with acids: Metals (above H) react with dilute acids to produce H₂ (e.g., Zn + 2 HCl → ZnCl₂ + H₂); non-metals do not displace hydrogen.
• Electron behaviour: Metals tend to lose electrons to form cations (e.g., Na → Na⁺ + e⁻); non-metals gain electrons to form anions (e.g., Cl + e⁻ → Cl⁻).

These chemical behaviours guide reactions, bonding and compound formation in inorganic chemistry.

Answer — reactivity series & significance:

• Definition: The reactivity series ranks metals from most to least reactive (e.g., K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au).
• Practical uses:
  • Predict displacement reactions: a more reactive metal displaces a less reactive metal from its salt (e.g., Zn + CuSO₄ → ZnSO₄ + Cu).
  • Choose extraction/reduction methods: highly reactive metals (e.g., Na, K) are obtained by electrolysis; less reactive metals (e.g., Fe) by reduction with carbon.

Answer — equations & oxide nature:

• Sodium: 4 Na + O₂ → 2 Na₂O — forms sodium oxide, a strongly basic oxide.
• Magnesium: 2 Mg + O₂ → 2 MgO — magnesium oxide is basic.
• Copper: 2 Cu + O₂ → 2 CuO — copper(II) oxide is amphoteric/less basic; many non-metal oxides are acidic, but CuO shows limited amphoteric behaviour in some reactions.

Basic oxides react with acids to form salts (e.g., MgO + 2 HCl → MgCl₂ + H₂O).

Answer — metals with water:

• Sodium (very reactive): 2 Na + 2 H₂O → 2 NaOH + H₂ (vigorous; produces caustic sodium hydroxide and hydrogen gas; sodium often moves on water surface).
• Magnesium (less reactive with cold water): Mg + 2 H₂O → Mg(OH)₂ + H₂ (slow with cold water; reacts more readily with steam to form MgO + H₂).
• Difference: Reactivity and ease of oxidising cause sodium to react violently, while magnesium is slower with cold water but reacts with steam due to higher activation energy.

Answer — metals with dilute acids:

• General reaction: Metal + Acid → Salt + Hydrogen gas (if metal above H in reactivity series).
• Example: Zn + 2 HCl → ZnCl₂ + H₂.
• Gas test: Hydrogen is identified by the 'pop' test — bring a burning splint near the gas; a characteristic pop confirms H₂.

Answer — amphoteric oxides:

• Definition: Amphoteric oxides react with both acids and bases.
• Examples: Zinc oxide (ZnO), aluminium oxide (Al₂O₃).
• Reactions:
  • ZnO + 2 HCl → ZnCl₂ + H₂O (reaction with acid).
  • ZnO + 2 NaOH + H₂O → Na₂[Zn(OH)₄] (reaction with base forming complex).

Answer — ore preparation steps:

• Concentration: Removal of gangue (impurities) by physical methods — e.g., crushing, washing, magnetic separation; increases the percentage of metal compound in ore.
• Roasting: Heating sulfide ores in excess air to convert sulfides to oxides and release SO₂ (e.g., 2 ZnS + 3 O₂ → 2 ZnO + 2 SO₂).
• Calcination: Heating carbonate ores in limited air to remove CO₂ and give oxides (e.g., CaCO₃ → CaO + CO₂), often used for ores unsuitable for roasting.
• Purpose: Both roasting and calcination produce metal oxides, which are easier to reduce to the metal in the subsequent reduction step.

Answer — reduction by carbon:

• Principle: Carbon (or carbon monoxide) removes oxygen from metal oxides, reducing them to metals as it is oxidised to CO or CO₂.
• Blast furnace iron reduction: Fe₂O₃ + 3 CO → 2 Fe + 3 CO₂ (CO acts as reducing agent generated from coke).
• Notes: Slag forming agents (limestone) remove acidic impurities; the process operates at high temperatures with multiple redox steps.

Answer — Hall–Héroult process:

• Feedstock: Aluminium oxide (Al₂O₃) is obtained from bauxite by refining (Bayer process).
• Electrolysis: Al₂O₃ is dissolved in molten cryolite and electrolysed; at cathode Al³⁺ + 3 e⁻ → Al (liquid), at anode O²⁻ → O₂ + 4 e⁻.
• Reason for electrolysis: Aluminium is highly reactive and cannot be reduced by carbon; electrolysis provides the electrical energy to force a non-spontaneous reduction.

Answer — corrosion & rusting:

• Definition: Corrosion is the gradual destruction of metals by chemical or electrochemical reaction with the environment (e.g., rusting of iron).
• Electrochemical rusting: In presence of water and oxygen, iron forms anodic and cathodic regions; Fe oxidises to Fe²⁺ at anode and oxygen reduces at cathode, producing hydrated iron(III) oxide (rust): 4 Fe + 3 O₂ + x H₂O → 2 Fe₂O₃·x H₂O.
• Prevention:
  • Coatings/painting to exclude air and moisture.
  • Galvanisation — zinc coating provides sacrificial protection.

Answer — why copper for wiring:

• High electrical conductivity: Copper has low resistivity allowing efficient current flow.
• Ductility: Can be drawn into thin wires without breaking.
• Corrosion resistance & mechanical strength: Reasonably resistant to corrosion and mechanically robust for long-term installations.

Answer — storage & safety:

• Storage: Stored under dry hydrocarbon oil (paraffin/dry mineral oil) to prevent contact with air and moisture.
• Reason: Both metals react violently with water and oxidise in air, producing hydrogen and sometimes flames/explosions.
• Safety concerns: Risk of fire/explosion on exposure to water; handle with gloves and tongs, store in labeled containers, and keep away from moisture and heat sources.

Answer — alloys & advantages:

• Concept: Alloys are mixtures of metals (and sometimes non-metals) where the combined properties (strength, hardness, corrosion resistance) are superior to pure metals.
• Examples:
  • Steel (Fe + C): stronger and harder than pure iron; used in construction, tools, vehicles.
  • Brass (Cu + Zn): corrosion-resistant and malleable; used for musical instruments, fittings.
• Note: Alloying tailors mechanical and chemical properties for specific applications.

Answer — gas tests:

• Hydrogen (H₂): Test — bring a burning splint to the gas; Observation: a short 'pop' sound. Basis: H₂ is combustible and reacts with O₂ producing a small flame/sound.
• Oxygen (O₂): Test — insert a glowing splint; Observation: the splint relights. Basis: O₂ supports combustion, restoring glow to a smoldering splint.
• Carbon dioxide (CO₂): Test — bubble gas through limewater (Ca(OH)₂); Observation: limewater turns milky (forms CaCO₃). Basis: CO₂ + Ca(OH)₂ → CaCO₃↓ + H₂O.

Answer — environmental impacts & mitigation:

• Impacts: Air pollution (SO₂ from roasting), water contamination (heavy metals, acidic effluents), habitat destruction and soil erosion.
• Mitigation:
  • Install effluent treatment and scrubbers to remove SO₂ and neutralize acidic waters before discharge.
  • Reclaim mined land (rehabilitation), enforce waste management and reduce environmental footprint through cleaner technologies.

Answer — POP preparation & use:

• Preparation: Gypsum is calcium sulfate dihydrate (CaSO₄·2 H₂O). On heating to about 150 °C it loses water to form hemihydrate (plaster of Paris): CaSO₄·2 H₂O → CaSO₄·½ H₂O + 1½ H₂O (lost).
• Use: On mixing POP with water it sets to form hard solid — widely used for making casts in orthopaedics, moulding ornaments and building plaster.

Answer — graphite vs diamond conductivity:

• Graphite: Each carbon is bonded to three others in planar hexagonal sheets. One electron per carbon is delocalised across layers, allowing electron mobility → electrical conduction.
• Diamond: Each carbon forms four strong covalent bonds in a tetrahedral network; no delocalised electrons are available → electrical insulator.
• Conclusion: Difference in bonding and electron delocalisation explains contrasting electrical properties.

Answer — removing temporary hardness:

• Temporary hardness: Caused by bicarbonates of Ca²⁺ and Mg²⁺ in water.
• Action of Na₂CO₃: Na₂CO₃ reacts with Ca²⁺/Mg²⁺ to form insoluble carbonates which precipitate:
  • Ca²⁺ + CO₃²⁻ → CaCO₃↓
• Result: Removal of Ca²⁺/Mg²⁺ ions softens the water; this method treats temporary (carbonate) hardness effectively.

Answer — lab preparation & crystallisation:

• Neutralisation: Mix stoichiometric amounts of dilute HCl and NaOH — HCl + NaOH → NaCl + H₂O (use indicator to check neutral point).
• Purification: Filter if insoluble impurities; evaporate gently to concentrate solution without splattering.
• Crystallisation: Allow saturated solution to cool slowly; crystals form. Filter, wash crystals with small amount of cold distilled water and dry in desiccator or oven at low temperature.
• Note: Slow cooling yields larger crystals suitable for analysis.

Answer — native metals vs ores:

• Native metals: Metals that are chemically unreactive (noble) resist oxidation and weathering — e.g., gold (Au), platinum — so they remain in elemental form in nature.
• Reactive metals: Readily form compounds (oxides, sulfides, carbonates) and thus occur as ores requiring extraction (e.g., iron as Fe₂O₃).
• Implication: Chemical reactivity determines natural occurrence; noble metals are often costly but easily recoverable as nuggets, while reactive metals need processing from ores.

Answer — why not change subscripts:

• Meaning of subscripts: Subscripts define the chemical identity and proportion of atoms in a molecule (e.g., H₂O has 2 H atoms and 1 O atom).
• Effect of changing subscripts: Changing a subscript changes the substance itself (e.g., H₂O → H₂O₂ becomes hydrogen peroxide, a different compound).
• Correct method: Use coefficients to balance atoms: 2 H₂ + O₂ → 2 H₂O (coefficients change quantity, not identity).

Answer — catalysts in industry:

• Function: Catalysts lower activation energy and increase reaction rate without being consumed.
• Industrial example: In the Haber process (ammonia synthesis), an iron catalyst speeds up combination of N₂ and H₂ to produce NH₃ at feasible rates and conditions.
• Metallurgical relevance: Catalysts are used in processes that require controlled reaction rates and selectivity (e.g., catalytic converters, hydrogenation steps in chemical industries).

Answer — sacrificial protection:

• Concept: A more reactive metal (anode) is attached to the metal to be protected (cathode). The sacrificial metal corrodes preferentially, protecting the main structure.
• Application: For ship hulls or pipelines, blocks of magnesium or zinc are attached; these corrode instead of iron/steel, preventing rusting of the protected metal.
• Maintenance: Sacrificial anodes need periodic replacement after they are consumed.

Answer — five exam-crucial points:

1. Reactivity series: Memorise order and use for predicting displacement and extraction methods.
2. Key reactions: Metal with O₂, water, acids — write balanced equations and know observations.
3. Extraction basics: Know concentration, roasting/calcination, reduction (by carbon or electrolysis) and blast furnace concept.
4. Corrosion & prevention: Methods like painting, galvanisation and sacrificial anodes — reasons and examples.
5. Laboratory tests & salts: Gas tests (pop, glowing splint, limewater), precipitate tests (AgCl, BaSO₄), and common salts (NaCl, Na₂CO₃, CaOCl₂).

Focusing on these points with balanced equations and practical examples will greatly help board-exam answers.