CBSE Class 10 Science Metals and non metals Sure Shot Questions Set 04

This CBSE Class 10 Science Sure Shot page for Chapter 3, Metals and Non-metals, collects 19 of the most frequently asked board-style questions with step-by-step answers. It covers ionic bond formation, roasting and calcination, extraction and electrolytic refining of copper, rusting and its prevention, alloys, and the reactions of metals with oxygen, water and acids.

Short Answer Questions

Question 1. Write the balanced chemical equations for the reactions taking place when: (i) zinc carbonate is calcined, (ii) zinc sulphide is roasted, (iii) zinc oxide is reduced to zinc, (iv) cinnabar is heated in air, (v) manganese dioxide is heated with aluminium powder.
Answer:
(i) ZnCO₃ → ZnO + CO₂ (heat, calcination)
(ii) 2ZnS + 3O₂ → 2ZnO + 2SO₂ (roasting)
(iii) ZnO + C → Zn + CO (reduction by carbon)
(iv) 2HgS + 3O₂ → 2HgO + 2SO₂, then 2HgO → 2Hg + O₂
(v) 3MnO₂ + 4Al → 3Mn + 2Al₂O₃ + heat
In simple words: Carbonate ores are calcined, sulphide ores are roasted, and the oxide formed is then reduced to the metal.

Teacher's Note:
1. Balance each equation by counting atoms of every element on both sides.
2. Remember the order: ore to oxide first, then oxide to metal.

Question 2. Give reasons why metals conduct electricity. Write two advantages of making alloys.
Answer:
Conduction: metal atoms easily lose their valence electrons, so a metal has a sea of free (mobile) electrons. When a potential difference is applied these electrons move through the metal, which carries the current.
Advantages of alloys:
1. An alloy with the required properties (hardness, strength, low melting point) can be made.
2. Alloys resist corrosion better and are stronger than the pure metals, e.g. stainless steel.
In simple words: Free electrons carry current in a metal, and mixing metals gives a stronger, rust-resistant material.

Teacher's Note:
1. Do not say that the metal ions carry current; only the free electrons move.
2. Always give one example of an alloy with its use.

Teacher's Note:
1. Bronze contains tin, not zinc; many students mix up bronze and brass.
2. An alloy of mercury with another metal is called an amalgam.

Question 4. (a) All ores are minerals but all minerals are not ores. Justify. (b) What is galvanisation? (c) Explain roasting with a reaction. (d) What is an amalgam?
Answer:
(a) A mineral is any natural substance containing a metal. An ore is the mineral from which the metal can be extracted profitably. So every ore is a mineral, but a mineral from which extraction is not economical is not an ore. For example, bauxite is an ore of aluminium but clay is not.
(b) Galvanisation is coating iron or steel with a thin layer of zinc to protect it from rusting.
(c) Roasting is strong heating of a sulphide ore in excess of air: 2ZnS + 3O₂ → 2ZnO + 2SO₂.
(d) An amalgam is an alloy in which one of the metals is mercury.
In simple words: Ores are the profitable minerals; galvanisation is zinc coating; roasting converts a sulphide ore into its oxide.

Teacher's Note:
1. Use the bauxite and clay pair as your example for ore and mineral.
2. Zinc protects iron even when the coating is scratched, because zinc is more reactive.

Question 5. (a) What is an ionic bond? (b) How is it formed? (c) Show the formation of magnesium chloride.
Answer:
(a) The bond formed by the transfer of electrons from one atom to another is called an ionic (electrovalent) bond.
(b) A metal atom with 1, 2 or 3 valence electrons loses them to attain a noble gas configuration and a non-metal atom with 5, 6 or 7 valence electrons gains them. The oppositely charged ions are held by strong electrostatic attraction.
(c) Magnesium (2, 8, 2): Mg → Mg²⁺ + 2e⁻ (2, 8). Chlorine (2, 8, 7): Cl + e⁻ → Cl⁻ (2, 8, 8), so two chlorine atoms accept the two electrons.
Mg²⁺ + 2Cl⁻ → MgCl₂
In simple words: Magnesium gives two electrons, two chlorine atoms take one each, and the ions stick together as MgCl₂.

Teacher's Note:
1. Show the electron transfer with arrows and write the charges on the ions.
2. The formula MgCl₂ comes from balancing +2 with two −1 charges.

Teacher's Note:
1. Galvanisation protects only the surface, while an alloy changes the bulk material.
2. Learn one example of each: galvanised iron pipes and stainless steel.

Question 7. How is the method of extraction of metals high up in the reactivity series different from that for metals in the middle? Why can the same process not be applied for them? Name the process used for extraction of these metals.
Answer:
Metals high up (Na, Mg, Ca, Al): they form very stable oxides and have a greater affinity for oxygen than carbon has. So carbon cannot reduce their oxides. They are obtained by electrolysis of their molten chlorides or oxides; this process is called electrolytic reduction.
Metals in the middle (Zn, Fe, Pb): they usually occur as sulphides or carbonates. These are first converted into oxides by roasting or calcination, and the oxides are then reduced to the metal by heating with carbon (or aluminium).
The same process cannot be applied because the reactivities, and so the stability of the oxides, are different.
In simple words: Very reactive metals need electricity to be reduced, but middle metals can be reduced with carbon after converting the ore to oxide.

Teacher's Note:
1. Link the method of extraction to the position of the metal in the reactivity series.
2. Write one example for each: sodium by electrolysis, zinc by carbon reduction.

Question 8. Give reasons: (i) Most metals conduct electricity well. (ii) The reaction of iron(III) oxide with heated aluminium is used to join cracked machine parts.
Answer:
(i) Metals have free electrons which move easily through the metal and carry current.
(ii) The thermite reaction Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat is highly exothermic. The heat produced is so large that the iron is formed in molten state, and this molten iron is poured into the crack to weld it.
In simple words: The thermite reaction gives molten iron, which fills and welds the crack.

Teacher's Note:
1. Name the reaction (thermite) and say aluminium acts as the reducing agent.
2. This is a displacement reaction, since aluminium is more reactive than iron.

Long Answer Questions

Teacher's Note:
1. Split the answer into the four parts and write each one clearly.
2. Roasting uses plenty of air, calcination uses little or none.

Question 10. How does rusting take place? Write the conditions necessary for rusting. How can rusting be prevented?
Answer:
Rusting is the slow oxidation of iron in the presence of air (oxygen) and moisture to form a brown, flaky substance called rust, which is hydrated iron(III) oxide, Fe₂O₃.xH₂O.
Conditions: 1. Presence of air (oxygen). 2. Presence of moisture (water). Iron does not rust in dry air or in boiled water.
Prevention: 1. Painting, oiling or greasing the surface. 2. Galvanisation, i.e. coating iron with zinc. 3. Alloying, e.g. making stainless steel with chromium and nickel. 4. Electroplating with tin or chromium.
In simple words: Iron rusts when it meets oxygen and water together, so keep it from air and water or coat it with a protective layer.

Teacher's Note:
1. Write the conditions as two separate points; examiners look for both.
2. Galvanised iron stays protected even if the zinc layer is scratched.

Question 11. What is meant by refining of metals? Describe the electrolytic refining of copper with a neat labelled diagram.

Answer:
Refining is the process of removing impurities from the metal obtained by reduction to get pure metal. The most widely used method is electrolytic refining.
1. The impure copper is made the anode and a thin strip of pure copper is made the cathode.
2. Acidified copper sulphate solution is the electrolyte.
3. On passing current, copper from the anode dissolves in the solution and an equal amount of pure copper is deposited on the cathode.
At anode: Cu → Cu²⁺ + 2e⁻; at cathode: Cu²⁺ + 2e⁻ → Cu.
4. Soluble impurities remain in the solution and insoluble impurities settle below the anode as anode mud.
In simple words: Impure copper dissolves from the anode and pure copper plates on the cathode, leaving the impurities behind.

Teacher's Note:
1. Label the anode, cathode, electrolyte and anode mud in the diagram.
2. Remember that the anode is the impure block and the cathode is the pure strip.

Question 12. Discuss the physical properties of non-metals.
Answer:
1. Non-metals are brittle; they break into pieces when hammered and are not malleable.
2. They are not ductile and cannot be drawn into wires.
3. They are bad conductors of heat and electricity because they have no free electrons; graphite is an exception.
4. They are generally non-lustrous; iodine and graphite are lustrous.
5. They are generally soft; diamond is the hardest natural substance.
6. They generally have low melting and boiling points because the forces between their molecules are weak; graphite and diamond are exceptions.
7. They have low densities. Non-metals are solids or gases at room temperature, except bromine which is a liquid.
In simple words: Non-metals are brittle, dull, poor conductors, soft and light, with a few exceptions such as diamond, graphite and iodine.

Teacher's Note:
1. Learn the exceptions together with the general rules.
2. A table of metals versus non-metals gives full marks quickly.

Question 13. In what forms are metals found in nature? Explain, with examples and equations, how metals react with oxygen, water and dilute acids.
Answer:
Metals are found in the free (native) state, e.g. gold and silver, and in the combined state as oxides, sulphides and carbonates.
With oxygen: metals form metal oxides, e.g. 2Cu + O₂ → 2CuO.
With water: sodium reacts with cold water, 2Na + 2H₂O → 2NaOH + H₂; magnesium reacts with hot water, Mg + 2H₂O → Mg(OH)₂ + H₂; iron reacts with steam, 3Fe + 4H₂O → Fe₃O₄ + 4H₂.
With dilute acids: metals above hydrogen form a salt and hydrogen gas, e.g. Mg + 2HCl → MgCl₂ + H₂.
In simple words: Unreactive metals occur free; reactive metals occur as compounds and react with oxygen, water and acids.

Teacher's Note:
1. Write one equation for each type of reaction.
2. The oxide formed by iron and steam is Fe₃O₄.

Question 14. (i) Design an activity to show the conditions needed for iron nails to rust. (ii) How can rusting of iron be prevented?
Answer:
(i) Take three test tubes A, B and C with clean iron nails.
1. Tube A: add some water and cork it (air and water both present).
2. Tube B: add boiled distilled water, then pour about 1 mL of oil on it and cork it (no air; the oil stops air from dissolving).
3. Tube C: add anhydrous calcium chloride and cork it (dry air, no moisture).
Leave for a few days. The nail rusts only in tube A, so both air and water are necessary for rusting.
(ii) Rusting is prevented by painting, oiling, greasing, galvanising or alloying so that air and moisture do not touch the iron.
In simple words: The nail rusts only when both air and water are present.

Teacher's Note:
1. Tell clearly what each tube is meant to prove: A has both, B only water, C only dry air.
2. Boiling removes dissolved air from the water.

Question 15. A metal M is found in nature as its sulphide ore (M₂S) and is used in making electric wires. (i) Identify M. (ii) Write the balanced chemical equations involved in its extraction. (iii) Draw a labelled diagram of the electrolytic refining of the metal.

Answer:
(i) Metal M is copper (ore: copper glance, Cu₂S).
(ii) 1. Roasting: 2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂
2. Self-reduction: 2Cu₂O + Cu₂S → 6Cu + SO₂
(iii) The impure copper is the anode, a strip of pure copper is the cathode and acidified copper sulphate solution is the electrolyte (see diagram). Pure copper collects on the cathode and the impurities settle as anode mud.
In simple words: Heating copper sulphide ore in air gives copper, which is then purified by electrolysis.

Teacher's Note:
1. Name the two steps: roasting and then self-reduction.
2. Use the labelled diagram of the electrolytic cell for part (iii).

Question 16. Write the names and symbols of the two most reactive metals. Show, by drawing the electronic structure, how any one of them reacts with a halogen. State two physical properties of the compound formed.
Answer:
The two most reactive metals are potassium (K) and sodium (Na).
Sodium and chlorine: Na (2, 8, 1) loses one electron to become Na⁺ (2, 8); Cl (2, 8, 7) gains one electron to become Cl⁻ (2, 8, 8). The ions combine to form NaCl.
Na → Na⁺ + e⁻; Cl + e⁻ → Cl⁻; Na⁺ + Cl⁻ → NaCl
Properties of NaCl: 1. It is a hard crystalline solid with high melting and boiling points. 2. It is soluble in water and conducts electricity in the molten or aqueous state, not in the solid state.
In simple words: Sodium passes one electron to chlorine and the two ions stay together as common salt.

Teacher's Note:
1. Draw the electron arrangement of both atoms before and after the transfer.
2. State the type of bond: ionic.

Question 17. (i) Write the chemical name of the coating that forms on silver and copper articles left in moist air. (ii) What is galvanisation and why is it done? (iii) Define an alloy. How are alloys prepared? How do the properties of iron change when (a) a small quantity of carbon, (b) nickel and chromium are mixed with it?
Answer:
(i) Silver gets a black coating of silver sulphide, Ag₂S. Copper gets a green coating of basic copper carbonate, CuCO₃.Cu(OH)₂.
(ii) Galvanisation is coating iron with a layer of zinc. It is done to prevent rusting.
(iii) An alloy is a homogeneous mixture of two or more metals, or of a metal and a non-metal. It is made by melting the primary metal, adding the other elements in a definite proportion and cooling.
(a) Iron with a small amount of carbon becomes hard and strong (steel). (b) Iron with nickel and chromium becomes stainless steel, which is hard, strong and does not rust.
In simple words: Tarnish layers form on silver and copper, zinc coating stops rust, and adding other elements improves iron.

Teacher's Note:
1. Silver tarnishes with sulphur compounds in air, while copper turns green with CO₂ and moisture.
2. Pure iron is soft; carbon makes it harder.

Question 18. Two ores A and B were heated. Ore A gives CO₂ gas and ore B gives SO₂ gas. Write the steps to convert these ores into metals, giving chemical equations.
Answer:
Ore A gives CO₂, so it is a carbonate ore (ACO₃).
1. Calcination: heating strongly in limited air gives the oxide: ACO₃ → AO + CO₂.
2. Reduction: AO + C → A + CO.
Ore B gives SO₂, so it is a sulphide ore (BS).
1. Roasting: heating strongly in excess air gives the oxide: 2BS + 3O₂ → 2BO + 2SO₂.
2. Reduction: BO + C → B + CO.
In simple words: A carbonate ore is calcined and a sulphide ore is roasted, and the oxide is then reduced with carbon.

Teacher's Note:
1. CO₂ on heating points to a carbonate; SO₂ points to a sulphide.
2. Write both steps with balanced equations to score full marks.

Question 19. Discuss the exceptions in the properties of metals and non-metals.
Answer:
1. All metals are solids at room temperature except mercury, which is a liquid. Gallium and caesium have very low melting points and melt on the palm.
2. Iodine is a non-metal but it is lustrous.
3. Carbon, a non-metal, has allotropes. Diamond is the hardest natural substance with very high melting and boiling points, and graphite is a good conductor of electricity.
4. Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife, and they have low densities and low melting points.
In simple words: A few metals and non-metals do not follow the general rules, for example mercury, iodine, diamond and graphite.

Teacher's Note:
1. Learn the exceptions as a short list; they are asked often.
2. Bromine is the only liquid non-metal and mercury the only liquid metal.