Sulphates of Be and Mg are readily soluble in water but sulphates of Ca, Sr and Ba are insoluble. This is due to the fact

Options

A. the greater hydration enthapies of `Be^(2+) and Mg^(2+)` overcome the lattice enthalpy
B. high lattice enthalpy of `Be^(2+) and Mg^(2+)` makes them soluble in water
C. solubility decreases from `BeSO_(4)` to `BaSO_(4)` due to increase in ionic size
D. `BeSO_(4) and MgSO_(4)` are ionic in nature while other sulphates are covalent

Explanation

To solve the question regarding the solubility of sulphates of beryllium (Be) and magnesium (Mg) compared to those of calcium (Ca), strontium (Sr), and barium (Ba), we can follow these steps:

### Step 1: Understand the Concept of Solubility
The solubility of ionic compounds in water is influenced by two main factors: lattice enthalpy and hydration enthalpy.

- **Lattice Enthalpy**: This is the energy required to separate one mole of a solid ionic compound into its gaseous ions. Higher lattice enthalpy means the ions are held together more strongly in the solid state.

- **Hydration Enthalpy**: This is the energy released when gaseous ions are surrounded by water molecules and solvated. Higher hydration enthalpy means the ions interact more favorably with water.

### Step 2: Analyze the Sulphates of Be and Mg
Beryllium and magnesium sulphates are soluble in water. This can be attributed to the following:

- The **hydration enthalpy** of Be²⁺ and Mg²⁺ ions is significantly higher than their respective lattice enthalpy. This means that when these ions are dissolved in water, the energy released from hydration is sufficient to overcome the energy required to break the ionic bonds in the solid sulphate.

### Step 3: Analyze the Sulphates of Ca, Sr, and Ba
On the other hand, the sulphates of calcium, strontium, and barium are insoluble in water. The reasons include:

- As we move down the group from Ca to Ba, the size of the ions increases, which leads to a decrease in hydration enthalpy. The larger the ion, the less energy is released during hydration.

- The lattice enthalpy of these sulphates does not decrease significantly enough to compensate for the lower hydration enthalpy, resulting in their insolubility in water.

### Step 4: Conclusion
The difference in solubility between the sulphates of Be and Mg versus those of Ca, Sr, and Ba can be summarized as follows:

- **Beryllium and Magnesium Sulphates**: High hydration enthalpy > Lattice enthalpy → Soluble
- **Calcium, Strontium, and Barium Sulphates**: Hydration enthalpy < Lattice enthalpy → Insoluble

Thus, the main reason for the difference in solubility is that the hydration enthalpy of Be²⁺ and Mg²⁺ ions is greater than their lattice enthalpy, while for Ca²⁺, Sr²⁺, and Ba²⁺, the hydration enthalpy is not sufficient to overcome the lattice enthalpy.

### Final Answer
The sulphates of Be and Mg are soluble in water due to the higher hydration enthalpy of Be²⁺ and Mg²⁺ ions compared to their lattice enthalpy, while the sulphates of Ca, Sr, and Ba are insoluble because their hydration enthalpy cannot overcome their lattice enthalpy.

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