A solution is a homogeneous mixture of two or more substances. A solution can either be in the gas phase, the liquid phase, the solid phase, or in a combination of these phases. The enthalpy change of solution refers to the overall amount of heat which is released or absorbed during the dissolving process (at constant pressure). The enthalpy of solution can either be positive (endothermic reaction) or negative (exothermic reaction). The enthalpy of solution is commonly referred to as ΔH_solution.

When understanding the enthalpy of solution, it is easiest to think of three processes happening between two substances. One substance is the solute, let’s call that . The other substance is the solvent, let’s call that ?. The first process that happens deals only with the solvent, ?; ? has to break apart from the intramolecular forces holding it together. This means the solute molecules separate from each other. The enthalpy of this process is called ΔH₁. This is an endothermic process because energy is required for this reaction, so ΔH₁ > 0.

The second process is very similar to the first step. Much like the solvent, ?, needed to break apart from itself, the solute, ?, needs to overcome the intermolecular forces holding it together. This means the solvent molecules separate from each other. The enthalpy of this process is called ΔH₂. Like the first step, this reaction is an endothermic one and ΔH₂ > 0 because energy is required to break the forces between the ? molecules.

At this point, let us visualize what has happened so far. The solute, ?, has broken from the intermolecular forces holding it together and the solvent, ?, has broken from the intermolecular forces holding it together as well. It is at this time that the third process happens. We also have two values ΔH₁ and ΔH₂. Both of these values are greater than zero (again, because both processes are endothermic).

The third process is when substance ? and substance ? mix. The separated solute molecules and the separated solvent molecules join together to form a solution. The mixing of these two will create a solution, one in which one mole of this solution contains molecules from the solute, ?, and molecules from the solvent, ?. The enthalpy of this process is called ΔH₃. This is an exothermic reaction, because energy is given off as the two substances bond together; therefore, ΔH₃ < 0.

The final value for the enthalpy of solution can either be endothermic or exothermic (so ΔH_solution can either be greater than zero or less than zero), depending on how much energy is required or given off in each step. The enthalpy of solution can be written as a formula; ΔH_solution = ΔH₁ + ΔH₂ + ΔH₃. If ΔH_solution = 0, then these solutions are called ideal solutions. If ΔH_solution > 0 or ΔH_solution < 0, then these solutions are called nonideal solutions. The diagrams below can be used as visuals to help facilitate the understanding of this concept. Figure 1 is for an endothermic reaction, where ΔH_solution > 0. Figure 2 is for an exothermic reaction, where ΔH_solution < 0. Figure 3 is for an ideal solution, where ΔH_solution = 0.

The enthalpy of solution depends on the intermolecular forces of the solute and solvent. If the solution is ideal, and ΔH_solution = 0, then that means ΔH₁ added to ΔH₂ is equal to ΔH₃. This means the forces of attraction between like (the solute-solvent) and unlike (the solute-solute and solvent-solvent) molecules are the same (Figure 3). If the solution is nonideal, then either ΔH₁ added to ΔH₂ is greater than ΔH₃ or ΔH₃ is somewhat greater than the sum of ΔH₁ and ΔH₂. The first case means the forces of attraction of unlike molecules is greater than the forces of attraction between like molecules. The second case means the forces of attraction between like molecules is greater than the forces of attraction between unlike molecules (Figure 2). If ΔH₃ is much greater than the sum of ΔH₁ and ΔH₂, then a heterogeneous mixture occurs. Dissolution occurs to a very small extent (Figure 1).

Enthalpy change is part of the driving force in the formation of solutions. The effect of enthalpy change should be combined with the entropy change to get the Gibbs free energy change, ΔG. Gibbs free energy change is the ultimate driving force in the formation of solutions (at constant temperature and pressure).

An example of this concept is the dissolution of salts. NaCl (table salt) dissolves readily in water. In solid NaCl, the positive sodium ions are attracted to the negative chloride ions. The same is true of the solvent, water; the partially positive hydrogen atoms are attracted to the partially negative oxygen atoms. While NaCl dissolves in water, the positive sodium cations and chloride anions are being stabilized by the water molecule electric dipoles. Thus, the intermolecular bonds between NaCl are broken and the salt is dissolved. The overall chemical equation for this reaction is as follows: NaCl (s)--> Na⁺ (aq) + Cl^- (aq). An external link provided below can help with the understanding of the dissolution of salts.