Properties of Ionic Compounds – Long Answer Questions (CBSE Class 10 Science)
Medium Level (Application & Explanation)
Q1. Explain why ionic compounds have high melting and boiling points. Compare NaCl and MgO to support your answer.
Answer:
Ionic compounds have strong electrostatic forces of attraction between oppositely charged ions arranged in a regular lattice. These ionic bonds require a lot of energy to break.
As a result, ionic compounds have high melting and boiling points. For example, NaCl melts around ~800°C and boils at ~1413°C.
The strength of the ionic bond depends on ionic charge and ionic size. Higher charge and smaller ions mean stronger attractions.
MgO has Mg²⁺ and O²⁻ ions, both with double charges, leading to stronger ionic bonds than NaCl (Na⁺ and Cl⁻). Hence, MgO has a much higher melting point (~2852°C) than NaCl.
Therefore, due to their strong ionic bonds, ionic compounds need large amounts of heat energy to melt or boil, unlike many covalent compounds (like wax or sugar) which melt at much lower temperatures.
Q2. Why do most ionic compounds dissolve in water but not in kerosene or oil? Give examples and exceptions.
Answer:
Water is a polar solvent with partial positive (H) and partial negative (O) ends. These ends surround and pull apart ions from the crystal lattice, a process called hydration.
As ions separate and disperse, the ionic compound dissolves. Examples: NaCl (table salt) dissolves readily; KNO₃ is highly soluble; CuSO₄ gives a blue solution.
Non-polar solvents like kerosene, petrol, oils lack charges, so they cannot attract ions or separate them from the lattice. Thus, ionic compounds are insoluble in such solvents.
The rule of thumb is: “Like dissolves like.” Polar solvents dissolve ionic/polar substances; non-polar solvents dissolve non-polar substances.
Exceptions exist: BaSO₄ and PbCl₂ are sparingly soluble or almost insoluble in water due to very strong ionic lattices that water cannot effectively separate.
Q3. Describe how ionic compounds conduct electricity in molten or aqueous states but not in solid state. Include the role of ions.
Answer:
In the solid state, ions in an ionic compound are fixed in place within a rigid ionic lattice. There are no free ions to carry electric charge, so solids do not conduct electricity.
When the compound is molten (melted) or dissolved in water, the lattice breaks down. Ions become free to move.
Mobile ions act as charge carriers: cations move towards the cathode and anions move towards the anode, allowing electric current to pass.
Example: Solid NaCl is a non-conductor; molten NaCl and NaCl(aq) are good conductors. A CuSO₄ solution conducts due to Cu²⁺ and SO₄²⁻ ions.
Therefore, state of the substance controls conductivity: no movement in solids, but free-moving ions in molten and aqueous states enable electrical conduction.
Q4. Design an experiment to compare the conductivity of solid NaCl, NaCl solution, and sugar solution. Write observations and conclusions.
Set up a simple conductivity circuit: battery → wire → electrode 1 → solution → electrode 2 → wire → bulb → battery.
Test 1 (solid): Place dry NaCl between electrodes (no water). Observe the bulb.
Test 2 (ionic solution): Dissolve NaCl in water. Dip electrodes without touching; note bulb.
Test 3 (covalent solution): Dissolve sugar in water and repeat.
Observations:
Solid NaCl: bulb does not glow.
NaCl(aq): bulb glows (free ions conduct).
Sugar(aq): bulb does not glow (no ions).
Conclusion:
Ionic compounds conduct only when molten or in aqueous solution due to mobile ions.
Covalent compounds like sugar form neutral molecules in water and do not conduct.
Q5. “Strength of ionic bonds depends on charge and size of ions.” Explain this statement with examples from NaCl, KCl, and MgO.
Answer:
The strength of ionic bonds increases with:
Higher ionic charge: more attraction between ions.
Smaller ionic size: ions can come closer, increasing attraction.
NaCl vs KCl: Both have ions with ±1 charges, but Na⁺ is smaller than K⁺. So NaCl has slightly stronger bonds and a higher melting point than KCl.
MgO vs NaCl: MgO has Mg²⁺ and O²⁻ (higher charges), producing much stronger attraction than NaCl (Na⁺, Cl⁻). Thus, MgO has a much higher melting point (~2852°C) than NaCl (~800°C).
Therefore, charge magnitude and ionic size determine how strongly ions bind, which affects melting/boiling points, hardness, and thermal stability of ionic compounds.
High Complexity (Analytical & Scenario-Based)
Q6. A white crystalline solid melts at ~770°C, dissolves in water, does not conduct as a solid but conducts when molten and in solution. Identify the type of compound and justify with tests.
Answer:
The data indicates a typical ionic compound:
High melting point (~770°C) suggests strong ionic bonds (similar to KCl).
Soluble in water implies interaction with a polar solvent and hydration of ions.
Non-conducting in solid state but conducting when molten/aqueous confirms free ions are required for conduction.
Identification Tests:
Perform a conductivity test on the aqueous solution—the bulb should glow.
Try dissolving in kerosene/petrol—it should be insoluble (ionic → not soluble in non-polar).
If available, perform flame test or specific ion tests (e.g., for K⁺, Na⁺, Cl⁻) to narrow down the salt.
Conclusion: The substance is ionic (likely a simple salt). The combination of thermal and electrical properties decisively supports this.
Q7. During extraction of aluminium, why is molten Al₂O₃ used for electrolysis instead of solid Al₂O₃? Explain in terms of ionic movement and conductivity.
Answer:
Electrolysis requires mobile charge carriers. In ionic compounds, these are free ions.
Solid Al₂O₃ has immobile ions locked in a lattice; hence it is a non-conductor and unsuitable for electrolysis.
When Al₂O₃ is molten, the lattice breaks and Al³⁺ and O²⁻ ions become free to move, allowing current to flow.
Practically, pure Al₂O₃ has a very high melting point, so it is often dissolved in a medium to lower the effective melting point and make the process energy-efficient (industrial detail).
At the electrodes, Al³⁺ ions move to the cathode to gain electrons (forming aluminium metal), while O²⁻ ions move to the anode and release oxygen.
Thus, molten/dissolved state is essential for ionic conduction and electrolytic extraction of aluminium.
Q8. Salt is spread on icy roads in winter. Analyse how this works using the properties of ionic compounds and predict limitations.
Answer:
Common salt (NaCl) is an ionic compound that dissolves in water (from surface melt), forming an aqueous ionic solution.
This solution has a lower freezing point than pure water, a phenomenon called freezing point depression. As a result, ice melts even when the air temperature is below 0°C.
The presence of ions in the liquid phase disrupts the formation of the ice lattice, slowing down refreezing.
Limitations:
At very low temperatures (significantly below 0°C), NaCl becomes less effective; other salts (like CaCl₂) may be needed.
Salt can cause corrosion of vehicles/bridges and harm plants/soil.
Excess salt can increase salinity in nearby water bodies.
Conclusion: The solubility and ionic nature of salt enable freezing point reduction, making roads safer, but with environmental and temperature limitations.
Q9. A powder melts around 180°C, is insoluble in water but dissolves in petrol, and does not conduct in any state. Could it be ionic? Justify carefully.
Answer:
The given properties suggest it is not ionic:
Low melting point (~180°C) is typical of covalent molecular substances (e.g., wax, naphthalene), not of ionic compounds, which have high melting points.
Insolubility in water but solubility in petrol/kerosene indicates a non-polar substance. Ionic compounds dissolve in polar solvents like water, not in non-polar solvents.
No electrical conductivity in any state indicates absence of free ions. Ionic compounds conduct when molten or dissolved.
Therefore, the powder is likely a covalent, non-polar compound with weak intermolecular forces (van der Waals).
Conclusion: Based on melting point, solubility behavior, and conductivity, the substance is not ionic.
Q10. BaSO₄ and PbCl₂ are ionic but almost insoluble in water. Analyse this “exception” using lattice and hydration ideas in simple terms.
Answer:
Although BaSO₄ and PbCl₂ are ionic, they are sparingly soluble in water due to a balance of forces:
Inside the crystal, ions are held by strong ionic (lattice) forces.
For dissolution, water must separate and stabilize the ions using hydration (polar water molecules surround ions).
If the lattice attraction is very strong (e.g., large, highly charged ion pairs with good fit), and the hydration effect is not strong enough to compensate, the compound remains insoluble.
In BaSO₄ and PbCl₂, the attraction within the lattice is so significant that water **c...
