12.3: Enthalpy of Solution

There are two criteria that favor, but do not guarantee, the spontaneous formation of a solution:

1. A decrease in the internal energy of the system (an exothermic change, as discussed in the previous chapter on thermochemistry)
2. An increased dispersal of matter in the system (which indicates an increase in the entropy of the system, as you will learn about in the later chapter on thermodynamics)

In the process of dissolution, an internal energy change often, but not always, occurs as heat is absorbed or evolved. An increase in matter dispersal always results when a solution forms from the uniform distribution of solute molecules throughout a solvent.

Spontaneous solution formation is favored, but not guaranteed, by exothermic dissolution processes. While many soluble compounds do, indeed, dissolve with the release of heat, some dissolve endothermically. Ammonium nitrate (NH₄NO₃) is one such example and is used to make instant cold packs for treating injuries. A thin-walled plastic bag of water is sealed inside a larger bag with solid NH₄NO₃. When the smaller bag is broken, a solution of NH₄NO₃ forms, absorbing heat from the surroundings (the injured area to which the pack is applied) and providing a cold compress that decreases swelling. Endothermic dissolutions such as this one require greater energy input to separate the solute species than is recovered when the solutes are solvated, but they are spontaneous nonetheless due to the increase in disorder that accompanies the formation of the solution.

This text is adapted from Openstax, Chemistry 2e, Section 11.1: The Dissolution Process.

Dissolving a solute in a solution is either an exothermic or an endothermic process.

When sodium hydroxide dissolves in water, heat is transferred from the solution to the surrounding water causing the temperature of the water to increase. This is an exothermic process.

In endothermic processes, such as dissolving ammonium chloride in water, heat is absorbed by the solution causing the temperature of the water to decrease.

At constant pressure, the heat released or absorbed is called the enthalpy change.

Solution formation has three steps, each associated with a corresponding enthalpy change.

Step one is the separation of the solute particles. This requires an input of energy to overcome the attractive forces between the solute particles.

Step two is the separation of the solvent particles. This is also an endothermic step since energy is required to disrupt the attractive forces between the solvent particles.

Step three occurs when the solute and solvent particles mix. This step is exothermic because the attractive interactions between solute particles and solvent particles release energy.

For a stepwise process, Hess’s Law states that the net enthalpy change is the sum of the enthalpy changes in each step. The sign of the net enthalpy depends on the magnitudes of the enthalpies of the components.

If the sum of the component enthalpies is less than the enthalpy of mixing, the net enthalpy change is negative, and the dissolution process is exothermic.

If the sum of the component enthalpies is greater than the enthalpy of mixing, the enthalpy change is positive, and the dissolution process is endothermic.

If the two are equal, heat is neither released nor absorbed.

Solution formation is different from a chemical reaction. When a solute is dissolved in a solvent, the change is physical. Upon evaporating the solution, the solute can be recovered.

On the contrary, a chemical reaction alters the properties of the reactants. When copper hydroxide is dissolved in hydrochloric acid, evaporating the solution will not return copper hydroxide.

Instead, we will obtain the product, copper chloride.