Metals and Non-metals — CBSE Class 10 Science Important Questions

Answer: (b) Sodium

Sodium is so reactive that it catches fire when exposed to air and reacts violently with moisture. To keep it away from air and water, it is stored under kerosene oil.

Answer: (b) Mercury

Mercury is the only metal that is liquid at room temperature. (Bromine is a liquid but it is a non-metal.)

Answer: (c) $Cu + FeSO_4 \rightarrow CuSO_4 + Fe$Cu + FeSO₄ → CuSO₄ + Fe

Copper is less reactive than iron, so it cannot displace iron from iron sulphate solution — this reaction does not occur. In the other three, the added metal is more reactive than the one in solution, so displacement takes place.

Answer: (b) Aluminium is oxidised and acts as the reducing agent

Aluminium gains oxygen to form $Al_2O_3$Al₂O₃, so it is oxidised and reduces $Fe_2O_3$Fe₂O₃ to iron — hence it is the reducing agent. The reaction is a highly exothermic displacement (redox) reaction; the heat released melts the iron, which is used to weld the tracks.

Answer: (a) Both A and R are true and R is the correct explanation of A

$HNO_3$HNO₃ is a strong oxidising agent. It oxidises the $H_2$H₂ produced to water and is itself reduced to an oxide of nitrogen ($N_2O$N₂O, $NO$NO or $NO_2$NO₂), so hydrogen gas is not given off. (Magnesium and manganese are exceptions: they react with very dilute $HNO_3$HNO₃ to give $H_2$H₂.)

(a) An amphoteric oxide is a metal oxide that reacts with both acids and bases to form salt and water. Examples: aluminium oxide ($Al_2O_3$Al₂O₃) and zinc oxide ($ZnO$ZnO).

(b) Aluminium reacts with oxygen of air to form a thin, hard, protective layer of aluminium oxide on its surface. This layer prevents further contact of the metal with air and moisture, so the aluminium underneath does not corrode easily.

High melting and boiling points: Ionic compounds have strong electrostatic forces of attraction between their oppositely charged ions. A large amount of heat energy is needed to break these forces, so their melting and boiling points are high.

Electrical conductivity: In the solid state the ions are held in fixed positions and cannot move, so no current flows. In the molten (or aqueous) state the ions become free to move and carry charge, so the compound conducts electricity.

Roasting: heating a sulphide ore strongly in the presence of excess air to convert it into the metal oxide.

$2ZnS + 3O_2 \xrightarrow{\Delta} 2ZnO + 2SO_2$2ZnS + 3O₂ → (Δ) 2ZnO + 2SO₂

Calcination: heating a carbonate ore strongly in limited air (absence of air) to convert it into the metal oxide.

$ZnCO_3 \xrightarrow{\Delta} ZnO + CO_2$ZnCO₃ → (Δ) ZnO + CO₂

Difference: roasting is used for sulphide ores and takes place in plenty of air (releasing $SO_2$SO₂), while calcination is used for carbonate ores in limited air (releasing $CO_2$CO₂). Both convert the ore into an oxide, which is easier to reduce to the metal.

(a) Sodium, magnesium and aluminium are very high in the reactivity series and have a greater affinity for oxygen than carbon, so carbon cannot reduce their oxides. These metals are obtained by electrolytic reduction of their molten chlorides — for example, when molten $NaCl$NaCl is electrolysed, sodium is deposited at the cathode and chlorine is liberated at the anode.

(b) Gold and silver are least reactive (noble) metals — they do not corrode or tarnish easily and have a lasting shine (lustre) and are highly malleable, which makes them ideal for jewellery.

(c) In moist air, copper reacts slowly with carbon dioxide, water and oxygen to form a green layer of basic copper carbonate, $Cu(OH)_2 \cdot CuCO_3$Cu(OH)₂ · CuCO₃, which makes the vessel dull and green. This is corrosion of copper.

Magnesium (atomic number 12, electronic configuration 2,8,2) loses its 2 outer electrons; each chlorine (2,8,7) needs 1 electron, so two chlorine atoms are required.

• $Mg \rightarrow Mg^{2+} + 2e^-$Mg → Mg^2+ + 2e^-
• $2Cl + 2e^- \rightarrow 2Cl^-$2Cl + 2e^- → 2Cl^-

The $Mg^{2+}$Mg^2+ and $Cl^-$Cl^- ions attain stable noble-gas configurations and are held together by a strong electrostatic force, forming the ionic compound $MgCl_2$MgCl₂ (i.e. $Mg^{2+}(Cl^-)_2$Mg^2+(Cl^-)₂).

Two properties:

1. It is a hard crystalline solid with a high melting point.
2. It conducts electricity in the molten or dissolved state (free ions), but not as a solid; it is also soluble in water.

(a) Galvanisation is the method of coating iron or steel with a thin layer of zinc. It protects the iron from rusting: the zinc layer keeps out air and moisture, and even if the coating is broken, the more reactive zinc corrodes first, so the iron remains protected.

(b) Pure gold is very soft, so jewellery made from it would bend and wear easily. Alloying it with silver or copper (for example, 22-carat gold) makes it harder and more durable.

(c) Solder — an alloy of lead and tin. It has a low melting point, so it melts easily and is used to join (weld) electrical wires.

Zinc is a metal of moderate reactivity, so it is extracted by converting the ore to its oxide and then reducing the oxide with carbon.

Step 1 - Concentration: The powdered ore is concentrated to remove gangue (impurities).

Step 2 - Roasting: The concentrated sulphide ore is heated strongly in excess air to form zinc oxide:

$2ZnS + 3O_2 \xrightarrow{\Delta} 2ZnO + 2SO_2$2ZnS + 3O₂ → (Δ) 2ZnO + 2SO₂

(If the ore were the carbonate calamine $ZnCO_3$ZnCO₃, calcination would be used: $ZnCO_3 \xrightarrow{\Delta} ZnO + CO_2$ZnCO₃ → (Δ) ZnO + CO₂.)

Step 3 - Reduction: Zinc oxide is reduced to zinc metal by heating with carbon (coke):

$ZnO + C \rightarrow Zn + CO$ZnO + C → Zn + CO

Step 4 - Refining: The impure zinc obtained is purified to get very pure metal by electrolytic refining, in which impure zinc is made the anode, a thin strip of pure zinc the cathode, and a zinc salt solution the electrolyte; pure zinc deposits on the cathode.

(a) With oxygen — metals form basic oxides:
$2Mg + O_2 \rightarrow 2MgO$2Mg + O₂ → 2MgO
(Magnesium burns with a dazzling white flame.)

(b) With water — reactivity varies: sodium reacts with cold water, iron with steam:
$3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$3Fe + 4H₂O → Fe₃O₄ + 4H₂
(Sodium: $2Na + 2H_2O \rightarrow 2NaOH + H_2$2Na + 2H₂O → 2NaOH + H₂, very vigorous.)

(c) With dilute acids — metals above hydrogen release hydrogen gas:
$Zn + 2HCl \rightarrow ZnCl_2 + H_2$Zn + 2HCl → ZnCl₂ + H₂

(d) Reactivity series: an arrangement of metals in the decreasing order of their reactivity (K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au). A metal higher in the series can displace a metal lower down from its salt solution.

(e) Metals below hydrogen in the series — copper, silver and gold — do not react with dilute acids, because they are less reactive than hydrogen and cannot displace it from the acid.

(a) Iron is more reactive than copper, so it displaces copper from copper sulphate solution. The blue $Cu^{2+}$Cu^2+ ions are used up (forming copper metal) and pale green $Fe^{2+}$Fe^2+ ions (iron sulphate) are produced, so the blue colour fades to green.

$Fe + CuSO_4 \rightarrow FeSO_4 + Cu$Fe + CuSO₄ → FeSO₄ + Cu

This is a displacement (redox) reaction.

(b) The reddish-brown deposit on the nails is copper metal ($Cu$Cu) set free from the solution.

(c) In the second test tube, copper is less reactive than iron, so it cannot displace iron from iron sulphate solution. Therefore no reaction takes place and there is no change.