CBSE Class 10 Chemistry Solubility and Electrolysis MCQs Set 16

Answer: (a) kilogram of solvent

Explanation:
1. Molality m = moles of solute / mass of solvent in kg.

Teacher's Note:
1. Molality uses the mass of solvent, so it does not change with temperature.
2. Molarity uses the volume of solution.

Answer: (b) 0.50 m

Explanation:
1. 0.5 mole of the hydrate contains 5 mole of water = 90 g.
2. Total solvent = 910 + 90 = 1000 g = 1 kg.
3. Molality = 0.5 mol / 1 kg = 0.50 m.

Teacher's Note:
1. Add the water of crystallisation to the solvent.
2. A solution made with a known mass of solvent is described by molality, not molarity.

Answer: (c) 1.00

Explanation:
1. Molality = 0.10 mol / 0.100 kg = 1.00 m.

Teacher's Note:
1. Convert grams of solvent to kilograms.
2. 100 g = 0.1 kg.

Answer: (d) 0.500

Explanation:
1. Molality = 0.100 mol / 0.200 kg = 0.500 m.

Teacher's Note:
1. Divide moles by kilograms of solvent.
2. 200 g = 0.2 kg.

Answer: (a) 1.00 m

Explanation:
1. Molality = 0.100 mol / 0.100 kg = 1.00 m.
2. The values of density and K_f given are not needed for molality.

Teacher's Note:
1. Ignore data that the question does not ask for.
2. The freezing point constant would be needed only for the depression in freezing point.

Answer: (b) 0.500 m

Explanation:
1. 0.5 mole of the hydrate contains 9 mole of water = 162 g.
2. Total solvent = 838 + 162 = 1000 g = 1 kg.
3. Molality = 0.500 mol / 1 kg = 0.500 m.

Teacher's Note:
1. Treat the water in the crystal as solvent once it dissolves.
2. Volume is not given, so the answer cannot be a molarity.

Answer: (c) the pressure of the gas above the liquid is increased

Explanation:
1. By Henry's law the solubility of a gas is proportional to its pressure above the liquid.
2. Raising the temperature lowers gas solubility, and the amount of liquid or stirring does not change the solubility itself.

Teacher's Note:
1. Gas solubility rises with pressure and falls with temperature.
2. Soda water is bottled under pressure for this reason.

Answer: (d) its solubility increases with increasing temperature

Explanation:
1. An endothermic dissolution absorbs heat.
2. By Le Chatelier's principle raising the temperature favours the process, so solubility increases.

Teacher's Note:
1. Heat acts like a reactant in an endothermic process.
2. Hydration energy is smaller than lattice energy here.

Answer: (a) moles per 1000 g of solvent

Explanation:
1. Molality = moles of solute in 1000 g (1 kg) of solvent.

Teacher's Note:
1. Do not confuse it with molarity (moles per litre of solution).
2. The unit of molality is mol/kg.

Answer: (b) Ca(OH)₂

Explanation:
1. The solubility of KCl and KNO₃ rises steeply with temperature and that of NaCl rises slightly.
2. Calcium hydroxide is more soluble in cold water, so its solubility falls as temperature rises.

Teacher's Note:
1. Dissolving Ca(OH)₂ releases heat, so heating reduces solubility.
2. Slaked lime solution is therefore prepared with cold water.

Answer: (c) the colour of the solute

Explanation:
1. Nature of solute and solvent, temperature and (for gases) pressure affect solubility.
2. Colour has no influence.

Teacher's Note:
1. Fine powder dissolves faster, although the final solubility is the same.
2. Like dissolves like.

Answer: (d) decreases; increases

Explanation:
1. Gas solubility falls as temperature rises.
2. It rises with pressure (Henry's law).

Teacher's Note:
1. Gases escape faster from warm liquids.
2. A warm fizzy drink loses its gas quickly.

Answer: (a) molarity

Explanation:
1. Molarity M = moles of solute per litre of solution.

Teacher's Note:
1. Molarity depends on volume and so on temperature.
2. Molality depends only on masses.

Answer: (b) 0.01 M C₁₂H₂₂O₁₁ (sucrose)

Explanation:
1. Sucrose is a non-electrolyte and gives no ions.
2. Acetic acid ionises slightly and NaCl ionises fully.

Teacher's Note:
1. Conductivity depends on the number of free ions.
2. Non-electrolytes do not conduct.

Answer: (c) water is a better proton acceptor than methylbenzene

Explanation:
1. In water HCl transfers a proton to H₂O giving H₃O⁺ and Cl⁻ ions.
2. Methylbenzene cannot accept protons, so HCl stays as molecules.

Teacher's Note:
1. Ions form only if the solvent can accept the proton.
2. Pure water itself is a very poor conductor.

Answer: (d) Al³⁺

Explanation:
1. Heat of hydration increases with ionic charge and decreases with ionic size.
2. Al³⁺ is small and carries three charges, so it is hydrated most strongly.

Teacher's Note:
1. Charge density = charge / size decides hydration energy.
2. Na⁺ and Ba²⁺ have lower charge density.

Answer: (a) Acetic acid

Explanation:
1. Acetic acid ionises in water to give H⁺ and CH₃COO⁻ (a weak electrolyte).
2. Benzene, urea and CCl₄ do not form ions.

Teacher's Note:
1. Electrolytes form ions in solution or in the molten state.
2. Most organic compounds are non-electrolytes, but organic acids are exceptions.

Answer: (b) CCl₄

Explanation:
1. Chlorides in solution give chloride ions that precipitate white AgCl.
2. CCl₄ is covalent, gives no Cl⁻ ions and shows no precipitate.

Teacher's Note:
1. A precipitate needs free Cl⁻ ions.
2. Test for chloride: curdy white precipitate soluble in ammonia.

Answer: (c) solid NaCl

Explanation:
1. Copper conducts through free electrons.
2. Fused and aqueous NaCl conduct through free ions.
3. Solid NaCl has ions locked in the crystal and cannot conduct.

Teacher's Note:
1. Ionic compounds conduct only when molten or dissolved.
2. Metals conduct in the solid state.

Answer: (d) NH₄OH

Explanation:
1. NaOH, KOH and HNO₃ ionise almost completely in water.
2. NH₄OH is a weak base and ionises only slightly.

Teacher's Note:
1. Strong acids and strong bases are strong electrolytes.
2. Ammonium hydroxide is a weak base.

Answer: (a) Carbon tetrachloride

Explanation:
1. NaCl, BaCl₂ and FeCl₃ are ionic and give ions.
2. CCl₄ is a covalent liquid that does not ionise.

Teacher's Note:
1. Non-electrolytes are covalent and do not ionise.
2. Examples: sugar, urea, alcohol, benzene.