A solution is formed when one substance disperses uniformly throughout another. With the exception of gas mixtures, all solutions involve substances in a condensed phase. We learned in Chapter 11 that the molecules or ions of substances in the liquid and solid states experience intermolecular attractive forces that hold them together. Intermolecular forces also operate between solute particles and the solvent molecules that surround them.
Any of the various kinds of intermolecular forces that we discussed in Chapter 11 can operate between solute and solvent particles in a solution. Solutions form when the attractive forces between solute and solvent are comparable in magnitude with those that exist between the solute particles themselves or between the solvent particles themselves. For example, the ionic substance NaCl dissolves readily in water because the attractive interaction between the ions and the polar H2O molecules overcomes the lattice energy of NaCl(s). Let's examine this solution process more closely, paying attention to these attractive forces.
When NaCl is added to water (Figure 13.1), the water molecules orient themselves on the surface of the NaCl crystals. The positive end of the water dipole is oriented toward the Cl– ions, and the negative end of the water dipole is oriented toward the Na+ ions. The ion-dipole attractions between Na+ and Cl– ions and water molecules are sufficiently strong to pull these ions from their positions in the crystal.
Figure 13.1 A schematic illustration of the solution process of an ionic solid (a) in water. The solid substance is hydrated by water molecules (b), with the oxygen atoms of the water molecules oriented toward the cations and the hydrogens oriented toward the anions. (c) As the solution process proceeds, the individual ions are removed from the solid surface and become completely separate, hydrated species in solution.
Once separated from the crystal, the Na+ and Cl– ions are surrounded by water molecules, as shown in Figures 13.1(b and c) and 13.2. Such interactions between solute and solvent molecules are known as solvation. When the solvent is water, the interactions are known as hydration.
Figure 13.2 Hydrated Na+ and Cl– ions. The negative ends of the water dipole point toward the positive ion, and the positive ends point toward the negative ion.
Sodium chloride dissolves in water because the water molecules have a sufficient attraction for the Na+ and Cl– ions to overcome the attraction of these two ions for one another in the crystal. To form an aqueous solution of NaCl, water molecules must also separate from one another to form spaces in the solvent that will be occupied by the Na+ and Cl– ions. Thus, we can think of the overall energetics of solution formation as having three components, illustrated schematically in Figure 13.3. The overall enthalpy change in forming a solution, Hsoln, is the sum of these three terms:
Figure 13.3 Depiction of the three enthalpic contributions to the overall heat of solution of a solute. As noted in the text, H1 and H2 represent endothermic processes, requiring an input of energy, whereas H3 represents an exothermic process.
Figure 13.4 depicts the enthalpy change associated with each of these components. Separation of the solute particles from one another requires an input of energy to overcome their attractive interactions (for example, separating Na+ and Cl– ions). The process is therefore endothermic (H1 > 0). Separation of solvent molecules to accommodate the solute also requires energy ( H2 > 0). The third component arises from the attractive interactions between solute and solvent, leading to an exothermic process ( H3 < 0).
Figure 13.4 Analysis of the enthalpy changes accompanying the solution process. The three processes are illustrated in Figure 13.3. The diagram on the left illustrates a net exothermic process (Hsoln < 0); that on the right shows a net endothermic process (Hsoln > 0).
As shown in Figure 13.4, the three enthalpy terms in Equation 13.1 can add together to give either a negative or a positive sum. Thus, the formation of a solution can be either exothermic or endothermic. For example, when magnesium sulfate, MgSO4, is added to water, the resultant solution gets quite warm: Hsoln = -91.2 kJ/mol. In contrast, the dissolution of ammonium nitrate, NH4NO3, is endothermic: Hsoln = 26.4 kJ/mol. These particular substances have been used to make instant heat packs and ice packs that are used to treat athletic injuries (Figure 13.5). These packs consist of a pouch of water and a dry chemical, MgSO4 for hot packs and NH4NO3 for cold packs. When the pack is squeezed, the seal separating the solid from the water is broken and a solution forms, either increasing or decreasing the temperature.
In Chapter 5 we learned that the enthalpy change in a process can provide information about the extent to which a process will occur. Processes that are exothermic tend to proceed spontaneously. A solution will not form if Hsoln is too endothermic. The solvent-solute interaction must be strong enough to make H3 comparable in quantity to H1 + H2. This is why ionic solutes like NaCl do not dissolve in nonpolar liquids such as gasoline. The nonpolar hydrocarbon molecules of the gasoline would experience only weak attractive interactions with the ions, and these interactions would not go very far toward compensating for the energies required to separate the ions from one another.
By similar reasoning, we can understand why a polar liquid such as water does not form solutions with a nonpolar liquid such as octane, C8H18. The water molecules experience strong hydrogen-bonding interactions with one another. These attractive forces must be overcome to disperse the water molecules throughout the nonpolar liquid. The energy required to separate the H2O molecules is not recovered in the form of attractive interactions between H2O and C8H18 molecules.
When two nonpolar substances such as CCl4 and hexane, C6H14, are mixed, they readily dissolve in one another in all proportions. The attractive forces between molecules in both of these substances are London dispersion forces. The two substances have similar boiling points: CCl4 boils at 77°C, and C6H14 boils at 69°C. It is therefore reasonable to suppose that the magnitudes of the attractive forces between molecules are comparable in the two substances. When the two are mixed, there is little or no energy change. Yet the dissolving process occurs spontaneously; that is, it occurs to an appreciable extent without any extra input of energy from outside the system. Two distinct factors are involved in processes that occur spontaneously. The most obvious is energy; the other is disorder.
If you let go of a book, it falls to the floor because of gravity. At its initial height, it has a higher potential energy than it has when it is on the floor. Unless it is restrained, the book falls and so loses energy. This fact leads us to the first basic principle identifying spontaneous processes and the direction they take: Processes in which the energy content of the system decreases tend to occur spontaneously. Spontaneous processes tend to be exothermic. Change tends to occur in the direction that leads to a lower energy for the system.
However, we can think of processes that do not result in lower energy for a system or that may even be endothermic and still occur spontaneously. For example, NH4NO3 readily dissolves in water, even though the solution process is endothermic. The mixing of CCl4 and C6H14 provides another simple example. All such processes are characterized by an increase in the disorder, or randomness, of the system. Suppose that we could suddenly remove a barrier that separates 500 mL of CCl4 from 500 mL of C6H14, as in Figure 13.6(a). Before the barrier is removed, each liquid occupies a volume of 500 mL. We know that we can find all the CCl4 molecules in the 500 mL to the left of the barrier and all the C6H14 molecules in the 500 mL to the right. When equilibrium has been established after removal of the barrier, the two liquids together occupy a volume of about 1000 mL. Formation of a homogeneous solution has resulted in increased disorder, or randomness, in that the molecules of each substance are now mixed and distributed in a volume twice as large as that which they occupied before mixing. This example illustrates our second basic principle: Processes in which the disorder of the system increases tend to occur spontaneously.
Figure 13.6 Formation of a homogeneous solution between CCl4 and C6H14 upon removal of a barrier separating the two liquids. The solution in (b) is more disordered, or random, in character than the separate liquids before solution formation (a).
When molecules of different types are brought together, mixing and hence an increase in disorder occur spontaneously unless the molecules are restrained by sufficiently strong intermolecular forces or by physical barriers. Thus, gases spontaneously mix and expand unless restrained by their containers; in this case intermolecular forces are too weak to restrain the molecules. However, because strong bonds hold sodium and chloride ions together, sodium chloride does not spontaneously dissolve in gasoline.
We will discuss spontaneous processes again in Chapter 19. At that time we will consider the balance between the tendencies toward lower energy and toward increased disorder in greater detail. For the moment we need to be aware that in most cases formation of solutions is favored by the increase in disorder that accompanies mixing. Consequently, a solution will form unless solute-solute or solvent-solvent interactions are too strong relative to the solute-solvent interactions.
In all our discussions of solutions, we must be careful to distinguish the physical process of solution formation from chemical processes that lead to a solution. For example, nickel metal is dissolved on contact with hydrochloric acid solution. Due to the following chemical reaction:
In this instance the chemical form of the substance being dissolved is changed. If the solution is evaporated to dryness, Ni(s) is not recovered as such; instead, NiCl2 6H2O(s) is recovered (Figure 13.7). In contrast, when NaCl(s) is dissolved in water, it can be recovered by evaporation of its solution to dryness. Our focus throughout this chapter is on solutions from which the solute can be recovered unchanged from the solution.