Previous Year Questions: Metals and Non-metals - Class 10 PDF Download

Ans: (b) Calcium and magnesium
Explanation:

• Highly reactive metals like sodium, calcium, magnesium, and aluminium are extracted by electrolytic reduction of their molten chlorides.
• Gold and silver are less reactive, often found native or extracted by other methods (e.g., cyanide process).
• Iron is extracted by reduction with carbon in a blast furnace.
• Calcium and magnesium fit the description as they require electrolysis (e.g., NaCl or CaCl₂).

Ans: (a) Mg · O → Mg²⁺[O²⁻]
Explanation:

• Magnesium (Mg, atomic number 12: 2,8,2) loses 2 electrons to form Mg²⁺.
• Oxygen (O, atomic number 8: 2,6) gains 2 electrons to form O²⁻.
• The ionic compound is MgO (Mg²⁺[O²⁻]), with a 1:1 ratio.
• Option (a) correctly shows the electron transfer and ionic bond formation.

Ans: (a) Reduction with carbon

Explanation:

• Metals in the middle of the reactivity series (e.g., iron, zinc) are commonly extracted by reducing their oxides with carbon 
  (e.g., Fe₂O₃ + 3C → 2Fe + 3CO).
• Hydrogen reduction is less common, used for specific metals like tungsten.
• Aluminium reduction (thermite) is specific to certain metals (e.g., chromium).
• Electrolytic reduction is for highly reactive metals (e.g., sodium, aluminium).

Ans: (c) It has weak electrostatic forces of attraction between its oppositely charged ions.
Compound C is ionic (A loses electrons, B gains, forming A⁺ and B⁻).
Ionic compounds have:

• High melting points due to strong electrostatic forces (a).
• High solubility in water for many ionic compounds (b).
• Conductivity in molten or aqueous state due to mobile ions (d).
• Weak electrostatic forces (c) are not characteristic, as ionic bonds are strong.

Ans: (b) When it is heated with iron (III) oxide, molten iron is obtained.

• Thermite reaction: 2Al + Fe₂O₃ → 2Fe + Al₂O₃ + Energy.
• Aluminium reduces iron oxide to molten iron, which is used to join railway tracks.
• The reaction is exothermic (A is partially correct but less specific).
• Molten aluminium oxide (c) is a byproduct, not used for joining.
• No alloy is formed (D is incorrect).

Ans: (b) Al₂O₃ and MgO

• Aluminium burns in air: 4Al + 3O₂ → 2Al₂O₃ (aluminium oxide).
• Magnesium burns in air: 2Mg + O₂ → 2MgO (magnesium oxide).
• Al₃O₄ and MgO₂ are not chemically correct formulas.

Ans: (d) Zinc

• Reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu.
• Metal M displaces Fe from FeSO₄, so M is above Fe (Mg, Al, Zn).
• M does not displace Al from Al₂(SO₄)₃, so M is below Al.
• Zinc (Zn) fits: Zn + FeSO₄ → ZnSO₄ + Fe; no reaction with Al₂(SO₄)₃.

Ans: (b) Pale green

• Reaction: Fe + CuSO₄ → FeSO₄ + Cu.
• CuSO₄ solution is blue (Cu²⁺ ions).
• FeSO₄ solution is pale green (Fe²⁺ ions).
• After 1 hour, Cu²⁺ is replaced by Fe²⁺, changing the color to pale green.

Ans: (b) Ductility

• Ductility allows metals to be drawn into wires.
• Malleability is for hammering into sheets.
• Rigidity refers to stiffness, and resistivity is electrical resistance.

Ans: (a) Aluminium

• Aluminium forms a protective Al₂O₃ layer, preventing further corrosion.
• Copper forms a green patina (Cu₂(OH)₂CO₃), not just oxide.
• Silver and gold are resistant but do not rely on oxide layers.

Ans: (c) Assertion (a) is true, but Reason (R) is false.

• A: Most metals do not produce H₂ with HNO₃ because it’s a strong oxidizing agent, oxidizing H₂ to H₂O.
• R: Nitric acid is an oxidizing agent, not a reducing agent.
• Thus, A is true, R is false.

Ans: (b) Both Assertion (a) and Reason (R) are true, but Reason (R) is not the correct explanation of Assertion (a).

• A: Highly reactive metals (e.g., Na, K, Al) have strong bonds in oxides, requiring electrolysis, not carbon reduction.
• R: Displacement reactions (e.g., thermite) are used for some metals, but this does not explain A.

Ans: (d) Assertion (a) is false, but Reason (R) is true.

• A: Cooking utensils require malleability (shaping into sheets) and thermal conductivity, not ductility (wire drawing).
• R: Copper is both ductile and malleable, so R is true.

Ans: (c) Assertion (a) is true, but Reason (R) is false.

• A: Brass is made by melting copper and adding zinc (not tin).
• R: Copper is the primary metal in brass, but A mentions tin, making it incorrect. Corrected A would make R true, but as stated, R does not explain A.

Ans: (a) Both Assertion (a) and Reason (R) are true and Reason (R) is the correct explanation of the Assertion (a).

• A: AgCl turns grey due to formation of silver metal.
• R: 2AgCl → 2Ag + Cl₂ (photolytic decomposition), explaining the grey color.

Ans: (c) (i) and (ii)

• (i) Mg burns with a dazzling white flame, correct.
• (ii) Forms white MgO powder, correct.
• (iii) Mg does not significantly vaporize; it reacts with O₂.
• (iv) MgO forms Mg(OH)₂ in water, which is basic and turns red litmus blue, not blue to red.

Ans: (c) It is an endothermic reaction.

• 2Mg + O₂ → 2MgO is exothermic, releasing heat and light.
• (a), (b), and (d) are true: white flame, white MgO powder, combination reaction.

Ans: (a) Only (i)

• Reactivity series: Mg > Al > Zn > Fe > Pb > Cu.
• (i) Mg displaces Cu: Mg + CuSO₄ → MgSO₄ + Cu.
• (ii) Pb is below Fe, no reaction.
• (iii) Al is below Ca, no reaction.
• (iv) Ca is below Zn, no reaction.

Ans:
Safety Measures:

1. Wear safety goggles to protect eyes from the dazzling flame.
2. Use tongs to hold the ribbon to avoid burns.

Observations:

1. Burns with a dazzling white flame.
2. Forms a white powder (MgO).

Explanation:

• The bright flame can harm eyes, and the hot ribbon can cause burns.
• Reaction: 2Mg + O₂ → 2MgO, producing intense light and white ash.

Ans:
(a) Type: Displacement reaction.
Equation: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag
(b) Method: Electrolytic refining.

Explanation:

• Cu, more reactive than Ag, displaces Ag from AgNO₃.
• Electrolytic refining purifies silver using a silver anode and cathode in an electrolyte.

Ans:
Metal X: Mercury (Hg).
Equations:

1. 2HgS + 3O₂ → 2HgO + 2SO₂
2. 2HgO → 2Hg + O₂

Explanation:

• Cinnabar (HgS) is roasted to form mercury(II) oxide (HgO).
• HgO decomposes on heating to yield mercury metal.

Ans:
(i) Free State: Gold (Au), low in reactivity series (below Cu).
(ii) Compound: Aluminium (Al, as bauxite Al₂O₃), high in reactivity series (above Zn).

Explanation:

• Gold, being unreactive, exists native.
• Aluminium, reactive, forms stable compounds like oxides.

Ans: (a)

Formation:

• Mg (2,8,2) → Mg²⁺ (2,8) + 2e⁻
• Cl (2,8,7) + e⁻ → Cl⁻ (2,8,8) [×2 for 2Cl]
• Mg + 2Cl → MgCl₂ (Mg²⁺[Cl⁻]₂)

Cation: Mg²⁺; Anion: Cl⁻

Explanation:

• Mg donates 2 electrons to two Cl atoms, forming ionic MgCl₂.

Ans: (b)
Statement: Aluminium oxide (Al₂O₃) is amphoteric, reacting with both acids and bases.
Equations:

• With acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• With base: Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

Explanation:

• Al₂O₃ reacts with HCl to form a salt (like a base) and with NaOH to form sodium aluminate (like an acid), proving amphoteric nature.

Ans: (a)

Activity: Place three iron nails in test tubes:

• Tube A: Iron nail in water with air.
• Tube B: Iron nail in boiled water (no oxygen) with oil layer.
• Tube C: Iron nail in dry air (no moisture).

Observations:

• Tube A: Rust (Fe₂O₃·nH₂O) forms; requires water and oxygen.
• Tube B: No rust; oxygen absent.
• Tube C: No rust; moisture absent.

Conclusion: Iron rusts only with both water and oxygen.

Explanation:

• Rusting is a chemical reaction: 4Fe + 3O₂ + 2nH₂O → 2Fe₂O₃·nH₂O.
• Moisture and oxygen are essential for corrosion.

Ans:
(a) Justification:
Displacement reactions reduce metal oxides using more reactive metals.
Examples:

1. Fe₂O₃ + 2Al → 2Fe + Al₂O₃ (thermite, extracts iron).
2. Cr₂O₃ + 2Al → 2Cr + Al₂O₃ (extracts chromium).

(b) Reason: Highly reactive metals (e.g., Na, K, Al) form very stable oxides due to strong metal-oxygen bonds, which carbon cannot break. Electrolysis is required.

Explanation:

• Middle-series metals (Fe, Zn) have less stable oxides, reducible by Al.
• High-series metals need stronger reducing conditions (electrolysis).

Ans:
Electron-dot structures:

• Sodium (Na): 2,8,1 → Na·
• Oxygen (O): 2,6 → :Ö:

Formation of Na₂O:

• 2Na· → 2Na⁺ + 2e⁻
• :Ö: + 2e⁻ → [Ö]²⁻
• 2Na⁺ + [Ö]²⁻ → Na₂O

Cation: Na⁺; Anion: O²⁻

• Each Na donates 1 electron to O, forming ionic Na₂O.

Ans:
Activity for Conductivity:

• Take a metal rod (e.g., copper) and fix thumbtacks with wax at equal intervals.
• Heat one end with a burner.

Observation: Wax melts sequentially from the heated end, showing heat conduction.

Activity for Melting Point:

• Heat small pieces of copper and plastic in separate crucibles.

Observation: Plastic melts quickly; copper requires much higher heat, indicating a high melting point.

Explanation:

• Metals conduct heat due to free electrons.
• High melting points result from strong metallic bonds.

Ans:
(i) Least Reactive: D
(ii) Observation: C displaces Cu from CuSO₄ (CuSO₄ + C → CSO₄ + Cu).
(iii) Order: B > C > A > D

Explanation:

• Reactivity series: More reactive metal displaces less reactive one.
• D: No reactions, least reactive.
• A: Displaces Cu, below Zn and Al.
• C: Displaces Fe, below Al, likely displaces Cu.
• B: Displaces Fe, Cu, Zn, most reactive.

Ans:
Method: Electrolytic refining.
Description:

• Impure metal is the anode, pure metal is the cathode.
• Electrolyte is a solution of the metal’s salt (e.g., CuSO₄ for copper).
• On passing current, metal ions from the anode dissolve, deposit as pure metal on the cathode.
• Impurities (sludge) settle at the bottom.

Explanation:

• Used for metals like Cu, Al, ensuring high purity.

Ans: (b)
(a) Formation of AlN:

• Al (2,8,3) → Al³⁺ (2,8) + 3e⁻
• N (2,5) + 3e⁻ → N³⁻ (2,8)
• Al³⁺ + N³⁻ → AlN

(b) Reason: Ionic compounds have strong electrostatic forces forming a rigid lattice, but pressure displaces ions, causing like charges to repel, leading to brittleness.
Explanation:

• Al donates 3 electrons to N, forming ionic AlN.
• Ionic lattices break when lattice alignment is disrupted.

Ans:
(a) Reactivity Series: A list of metals arranged in order of decreasing reactivity.

• Development: Based on displacement reactions and tendency to lose electrons.

• Order: Ca > Al > Pb > Cu

(b) Equation: Fe₂O₃ + 2Al → 2Fe + Al₂O₃

Explanation:

• Reactivity is tested by observing which metals displace others from solutions.
• Thermite reaction reduces Fe₂O₃ to iron.

Ans: (a) (i)

• I: Ag (silver, unreactive).
• II: Al (forms protective Al₂O₃).
• III: K (catches fire spontaneously).
• IV: Cu (forms black CuO on heating).

(ii) Amphoteric Oxides: React with both acids and bases.

• Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O
  

(iii) Alkalis: Water-soluble bases. Example: NaOH.

Explanation:

• Al₂O₃ shows dual reactivity, confirming amphoteric nature.
• Alkalis dissolve in water, producing OH⁻ ions.

Ans:
(I) Changes:

• Electrical conductivity decreases due to disruption of metallic lattice by alloying element.
• Melting point usually decreases as alloys have less ordered structures.
  (II) Alloy: Solder; Constituents: Lead, Tin.
  (III)(a) Alloys: Mixtures of two or more metals (or metal and non-metal) to enhance properties.
• Brass Preparation: Melt copper, add zinc in specific proportions, and cool.

Explanation:

• Alloys improve strength and corrosion resistance.
• Solder’s low melting point aids welding.
• Brass (Cu + Zn) is harder than pure copper.