Question

Which of the following orders are correct? (1) Thermal stability \(\mathrm{BeCO}_{3}<\mathrm{MgCO}_{3}<\mathrm{CaCO}_{3}<\mathrm{SrCO}_{3}<\mathrm{BaCO}_{3}\) (2) Basic nature \(\mathrm{LiOH}>\mathrm{NaOH}>\mathrm{KOH}>\mathrm{RbOH}>\mathrm{CsOH}\) (3) Solubility in water \(\mathrm{BeSO}_{4}<\mathrm{MgSO}_{4}<\mathrm{CaSO}_{4}<\mathrm{SrSO}_{4}<\mathrm{BaSO}_{4}\) (4) Melting point \(\mathrm{NaCl}>\mathrm{KCl}>\mathrm{RbCI}>\mathrm{CsCl}>\mathrm{LiCl}\) (a) (1), (4) (b) \((1),(2),(4)\) (c) \((2),(3)\) (d) (1), (2) (3), (4)

Short answer

Option (a) (1), (4) is correct as only the thermal stability order is given correctly.

Step-by-step solution

Thermal stability order

The thermal stability of carbonates generally increases down the group in the periodic table because larger cations stabilize the carbonate ion more effectively. Thus, the correct order is \(\mathrm{BeCO}_{3}<\mathrm{MgCO}_{3}<\mathrm{CaCO}_{3}<\mathrm{SrCO}_{3}<\mathrm{BaCO}_{3}\), so this statement is correct.

Basic nature order

The basic nature of hydroxides also increases down the group due to the reduction in lattice energy variations and increase in ionic radius. Therefore, the correct order for basicity should be \(\mathrm{LiOH}<\mathrm{NaOH}<\mathrm{KOH}<\mathrm{RbOH}<\mathrm{CsOH}\), opposite to the given order. Hence, this statement is incorrect.

Solubility in water order

The solubility of sulfates decreases down the group due to the increasing size of the cations, which reduces the lattice energies less effectively compared to hydration energies. Thus, the correct order should be \(\mathrm{BeSO}_{4}>\mathrm{MgSO}_{4}>\mathrm{CaSO}_{4}>\mathrm{SrSO}_{4}>\mathrm{BaSO}_{4}\), opposite to the given order, so this statement is incorrect.

Melting point order

For halide salts, the melting point generally decreases as you move down the group with larger cations leading to weaker lattice energies. However, \(\mathrm{LiCl}\) has a particularly high melting point due to its covalent character caused by small ion size and high polarization. Therefore, \(\mathrm{NaCl} > \mathrm{KCl} > \mathrm{RbCl} > \mathrm{CsCl} > \mathrm{LiCl}\) is incorrect, due to the incorrect placement of \(\mathrm{LiCl}\).

Key concepts

Thermal stability of carbonates

In Group 2 elements, the thermal stability of carbonates increases as you move down the group. This pattern is due to the increasing size of the metal cations from beryllium to barium. Larger cations have a stronger ability to stabilize the carbonate ion.When heating carbonates, they decompose into metal oxides and carbon dioxide: \[ \text{MCO}_3 (s) \rightarrow \text{MO} (s) + \text{CO}_2 (g) \] Here, the stability of the carbonate is linked to the strength of the bonds between the carbonate ion and the metal cation in the compound. As you move down the group, these bonds become more stable.- Larger cations like Ba²⁺ have a lower charge density, which means the electrostatic attraction they exert on carbonate ions is more spread out, leading to increased thermal stability.- Smaller cations like Be²⁺ result in less stable carbonates which decompose at lower temperatures.

Basicity of hydroxides

The basic nature of Group 1 and 2 metal hydroxides generally increases down the group. This increase is attributed to two main factors: decreasing lattice energies and increasing ionic radii. - As the ions become larger down the group, the lattice energies decrease. This means it takes less energy to dissociate the hydroxide into its component ions—resulting in more free OH⁻ ions, which contribute to basicity. - Larger ionic radii typically mean that the bond strength between the metal cation and the hydroxide decreases. This weaker bond aids the release of hydroxide ions into the solution, making the solution more basic. This concept is reversed for Group 1 metal hydroxides such as LiOH, NaOH, etc., where the ionic nature is a critical determinant of basicity. Remember, more free hydroxide ions in solution translates to stronger basic properties.

Solubility of sulfates

The solubility of sulfates in water decreases as you go down Group 2. This trend is closely related to the balance between lattice and hydration energies. - **Lattice Energy:** This is the energy required to separate the ions in a compound. Down the group, as cation size increases, lattice energy decreases. - **Hydration Energy:** This is the energy released when ions are solvated by water molecules. Larger ions have lower hydration energies as they do not interact as strongly with water molecules. For Group 2 sulfates: - BeSO₄ is highly soluble due to its high hydration energy overcoming its lattice energy. - BaSO₄, on the other hand, becomes less soluble as its lattice energy is much higher compared to its hydration energy. This results in a decrease in solubility as we move from BeSO₄ to BaSO₄.

Melting points of halides

The melting points of halide salts associated with Group 1 metals usually decrease as you move down the group. - **Larger Ionic Size:** As the metal cations become larger, the distance between the cations and anions in the crystal lattice increases, which weakens the force of attraction. - **Lattice Energy:** Lower lattice energies result in lower melting points because the ions require less energy to be separated in a molten state. An interesting exception within these salts is lithium chloride (LiCl), which exhibits a relatively high melting point. This is because LiCl has a significant degree of covalent character due to its small ion size and high polarization ability. Such covalent interactions lead to higher melting points. However, for other halides like NaCl, KCl, and CsCl, the expected trend of decreasing melting point with increasing ionic size holds true.