Question

When ammonium chloride dissolves in water, the solution becomes colder. (a) Is the solution process exothermic or endothermic? (b) Why does the solution form?

Short answer

(a) The solution process is endothermic because the ammonium chloride solution becomes colder when it dissolves in water, indicating that heat is absorbed from the surroundings. (b) The solution forms because the force of attraction between the water molecules and ammonium chloride ions overcomes the solute-solute and solvent-solvent interactions, allowing the ammonium chloride to dissolve in water.

Step-by-step solution

(a) Identify the type of process from temperature change)

(If the ammonium chloride solution becomes colder when it dissolves in water, it implies heat is absorbed from the surroundings. This characteristic is indicative of an endothermic process, where heat energy is absorbed by the system from the surroundings.)

(b) Understand why the solution forms)

(Based on the principles of intermolecular forces, solutions form when the solvent-solute interaction is stronger or approximately equal to the solute-solute and solvent-solvent interactions. In the case of ammonium chloride dissolving in water, the force of attraction between the water molecules and ammonium chloride ions overcomes the solute-solute and solvent-solvent interactions, which allows the ammonium chloride to dissolve in water, forming a solution.)

Key concepts

Solution Formation

Understanding why certain substances dissolve in solvents, like salt in water or carbon dioxide in soda, begins with the fundamental process of solution formation. In the exercise provided, when ammonium chloride dissolves in water to form a solution, it is an intriguing phenomenon that involves the distribution of the solute (ammonium chloride) throughout the solvent (water).

To comprehend this process, picture a group of solute particles being embraced and steadily pulled apart by solvent molecules. As these particles disperse, they interact with and are surrounded by solvent molecules, resulting in a uniform mixture – a solution. The intriguing part of our textbook exercise reveals that this process is governed by the delicate balance of interactions between solute and solvent, which we will further explore in the next section.

Solvent-Solute Interaction

A key player in solution formation is the interaction between solvent and solute molecules. Our example of ammonium chloride dissolving in water illustrates this beautifully. At a microscopic level, when the ionic compound (ammonium chloride) encounters water, an intricate dance begins. Water molecules, being polar, are attracted to the positively charged ammonium ions and the negatively charged chloride ions.

This attraction leads to the solvent molecules tugging at the solute ions, effectively pulling them away from the solid structure they were in. The stronger or comparable interaction between solvent and solute compared to that within the solute or the solvent itself, the more likely a solution will form. If the solute particles are fond enough of the solvent molecules, they prefer to associate with them, leading to a solution, just like in our homework problem.

Intermolecular Forces

Let's dig deeper into the intimate forces that dictate the dance of molecules during solution formation – the intermolecular forces. These are the forces that hold molecules together and also pull them apart in different scenarios. There are several types of these forces, including hydrogen bonds, dipole-dipole interactions, and London dispersion forces.

These forces not only keep the solvent and solute molecules cozy in their own company but also play matchmaker to create new interactions. In our example of ammonium chloride in water, it's the electrostatic attraction – a type of intermolecular force – that enables water molecules to engage with and dissolve the ammonium chloride. The mastery of how these forces operate is essential to predicting the solubility of different substances in various solvents and is the cornerstone of understanding solution chemistry, as highlighted in the exercise.