Trends in Solubility of Hydroxides, Sulfates, and Carbonates of Group 2 Elements

This section explores how the solubility of hydroxides, sulfates, and carbonates changes as we move down Group 2 (beryllium, magnesium, calcium, strontium, and barium).

Solubility of Hydroxides

The solubility of hydroxides (M(OH)₂) in water increases significantly as you go down the group. This trend holds true regardless of the specific data set used.

Here’s the trend:

• Beryllium Hydroxide (Be(OH)₂): Be(OH)₂ is nearly insoluble in water. While it may appear insoluble at first glance, a closer look reveals some solubility. Shaking a suspension of Be(OH)₂ in water, filtering it, and testing the pH will show a slightly basic solution. This indicates the presence of a small amount of dissolved Be(OH)₂ and hydroxide ions (OH⁻).

• Magnesium Hydroxide (Mg(OH)₂): Mg(OH)₂ is considered slightly soluble in water. Similar to Be(OH)₂, shaking Mg(OH)₂ with water and testing the filtrate reveals a slightly basic solution, indicating some degree of solubility.

• Calcium Hydroxide (Ca(OH)₂): Ca(OH)₂ is a common example of a moderately soluble hydroxide. One liter of water can dissolve around 1 gram of Ca(OH₂) at room temperature, forming a solution known as “limewater.”

• Strontium Hydroxide (Sr(OH)₂) and Barium Hydroxide (Ba(OH)₂): These hydroxides are highly soluble in water. Barium hydroxide solutions can reach concentrations of around 0.1 mol dm⁻³ at room temperature.

Explanation for the Trend:

This trend is related to the increasing atomic radius down the group. As the size of the metal cation (M²⁺) increases, its attraction to the hydroxide ion (OH⁻) weakens. This weaker attraction allows the crystal lattice of the hydroxide to break down more easily in water, resulting in higher solubility.

Solubility of Sulfates

In contrast to hydroxides, the solubility of sulfates (MSO₄) in Group 2 elements generally decreases as you move down the group.

• Beryllium Sulfate (BeSO₄): BeSO₄ is highly soluble in water.

• Magnesium Sulfate (MgSO₄): MgSO₄ is also highly soluble in water. Epsom salts, a common bath soak, is the hydrated form of magnesium sulfate (MgSO₄⋅7H₂O).

• Calcium Sulfate (CaSO₄): CaSO₄ is sparingly soluble in pure water. However, its solubility increases in the presence of other salts or at higher temperatures.

• Strontium Sulfate (SrSO₄) and Barium Sulfate (BaSO₄): These sulfates are very insoluble in pure water. Barium sulfate’s low solubility makes it a valuable tool for medical imaging procedures.

Explanation for the Trend:

This trend is less straightforward than the hydroxide trend. It’s a balancing act between the size of the cation and the sulfate ion (SO₄²⁻). While the increasing cation size down the group can potentially weaken the ionic bond, the large size of the sulfate ion also plays a role. As the size difference between the cation and sulfate ion decreases down the group, the lattice energy (energy required to break the crystal) can increase, leading to lower solubility.

Solubility of Carbonates

The solubility of carbonates (MCO₃) in water generally decreases as you move down the group. However, there’s an important caveat.

• All Group 2 carbonates are insoluble in pure water. This means they exist as solid minerals in nature.

• Carbonates can dissolve in water containing carbon dioxide (CO₂) due to the formation of bicarbonates (MHCO₃). The reaction with CO₂ helps to overcome the initial low solubility.

• Example: CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq) (Calcium carbonate reacts with carbon dioxide and water to form soluble calcium bicarbonate)

Explanation for the Trend:

Similar to sulfates, the trend in carbonate solubility involves a complex interplay between cation size and the size and charge of the carbonate ion (CO₃²⁻). Additionally, the formation of bicarbonate bridges between carbonate ions can further decrease their solubility in pure water.