6.13. SOLUBILITY EQUILIBRIA OF SPARINGLY SOLUBLE SALTS

We have already known that the solubility of ionic solids in water varies a great deal. Some of these (like calcium chloride) are so soluble that they are hygroscopic in nature and even absorb water vapour from atmosphere. Others (such as lithium fluoride) have so little solubility that they are commonly termed as insoluble.

Solubility refers to the ability of a substance to dissolve in a solvent, particularly in water for ionic solids. The solubility of ionic compounds can differ significantly. For instance, calcium chloride dissolves readily in water, which means it has high solubility and can attract moisture from the air, making it hygroscopic. In contrast, lithium fluoride dissolves poorly in water, classifying it as insoluble due to its low solubility.

Think of a sponge (highly soluble salt) that easily soaks up water, compared to a rock (insoluble salt) under water that doesn’t dissolve. While the sponge absorbs and holds water, the rock remains unchanged in the water.

The solubility depends on a number of factors important amongst which are the lattice enthalpy of the salt and the solvation enthalpy of the ions in a solution.

Several factors influence the solubility of salts in water. Lattice enthalpy is the energy required to separate ions in a solid crystal lattice. A higher lattice enthalpy means stronger attractions between ions, making the salt less soluble. Solvation enthalpy is the energy released when ions interact with the solvent. For dissolution to occur, the energy released via solvation must overcome the energy required to separate the ions, signifying that effective solvation can increase solubility.

Imagine trying to break apart a tightly held group of friends (lattice enthalpy). If someone comes along (water) and offers them something exciting to attract their attention (solvation), they are more likely to separate and engage with that offer. Without sufficient attraction, they won't separate easily.

Each salt has its characteristic solubility which depends on temperature. We classify salts on the basis of their solubility in the following three categories: Category I: Soluble Solubility > 0.1M Category II: Slightly Soluble 0.01M<Solubility< 0.1M Category III: Sparingly Soluble Solubility < 0.01M

Salts can be categorized based on their solubility in water. Category I includes salts that dissolve greatly in water (greater than 0.1M), Category II includes slightly soluble salts (between 0.01M and 0.1M), and Category III includes sparingly soluble salts (less than 0.01M). This classification helps in understanding how much of a specific salt can be dissolved in a solution under given conditions.

Think of it like different kinds of drinks. A soda (soluble) dissolves easily in water, whereas a syrup (slightly soluble) mixes but remains thick. Some compounds, like a pinch of salt in a glass of water (sparingly soluble), dissolve very little. This helps us understand how well or poorly components can mix.

We shall now consider the equilibrium between the sparingly soluble ionic salt and its saturated aqueous solution.

Sparingly soluble salts reach a state of equilibrium in a saturated solution where the rate of dissolution equals the rate of precipitation. In a saturated solution of a salt like barium sulfate, it exists in a dynamic balance where some solid salt keeps dissolving while an equal amount precipitates back. This creates a stable concentration of ions in the solution, allowing the salt to remain in equilibrium with its solid form.

Imagine a busy coffee shop where customers continually enter and exit. People come in (dissolve) to buy coffee while others leave (precipitate) after enjoying it. Eventually, the number of customers remains stable while the processes of entering and leaving continue simultaneously.

Let us now have a solid like barium sulphate in contact with its saturated aqueous solution. The equilibrium between the undisolved solid and the ions in a saturated solution can be represented by the equation: BaSO4(s) Ba2+(aq) + SO42–(aq). The equilibrium constant is given by the equation: K = {[Ba2+][SO42–]} / [BaSO4].

The solubility product constant, Ksp, quantifies the extent of solubility of sparingly soluble salts. For barium sulfate, Ksp is expressed in terms of the concentrations of its dissociation products, barium ions and sulfate ions, while the concentration of the solid is omitted, as it remains constant. Therefore, Ksp equals the product of the concentrations of the ions at equilibrium.

Think of Ksp as a recipe for a perfect balance. To make a cake, you need specific amounts of flour (Ba2+) and sugar (SO42–), but no matter how much cake you have, it doesn't change the amount of flour itself in the recipe.

It is expected from Le Chatelier’s principle that if we increase the concentration of any one of the ions, it should combine with the ion of its opposite charge and some of the salt will be precipitated.

The common ion effect describes how the solubility of a salt decreases when a common ion is added to the solution. By increasing the concentration of one of the ions involved in the solubility equilibrium, the system shifts equilibrium according to Le Chatelier’s principle—favoring the formation of the solid phase over the dissolved ions, thus leading to precipitation.

Imagine a concert where everyone is clapping and cheering (ions in solution), but if someone starts a loud chant (adding a common ion), it shifts focus and people start standing at attention instead of clapping—that’s similar to how precipitation occurs in solution as equilibrium shifts.

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