CHEMISTRY THE CENTRAL SCIENCE
13 PROPERTIES OF SOLUTIONS
A TROPICAL BEACH. Ocean water is a complex aqueous solution of many dissolved substances with sodium chloride having the highest concentration. Dana Edmunds/PacificStock.com
13.1 THE SOLUTION PROCESS
We begin by considering what happens at the molecular level when a substance dissolves, paying particular attention to the role of intermolecular forces. Two important aspects of the solution process are the natural tendency of particles to mix and changes in energy.
13.2 SATURATED SOLUTIONS AND SOLUBILITY
We learn that when a saturated solution is in contact with undissolved solute, the dissolved and undissolved solutes are in equilibrium. The amount of solute in a saturated solution defines the solubility of the solute, the extent to which a particular solute dissolves in a particular solvent.
13.3 FACTORS AFFECTING SOLUBILITY
We next consider the major factors affecting solubility. The nature of the solute and solvent determines the kinds of intermolecular forces among solute and solvent particles and strongly influences solubility. Temperature also affects solubility: Most solids are more soluble in water at higher temperatures, whereas gases are less soluble in water at higher temperatures. The solubility of gases increases with increasing pressure.
13.4 EXPRESSING SOLUTION CONCENTRATION
We examine several common ways of expressing concentration, including mole fraction, molarity, and molality.
13.5 COLLIGATIVE PROPERTIES
We observe that some physical properties of solutions depend only on concentration and not on the identity of the solute. These colligative properties include the extent to which the solute lowers the vapor pressure, increases the boiling point, and decreases the freezing point of the solvent. The osmotic pressure of a solution is also a colligative property.
We close the chapter by investigating colloids, mixtures that are not true solutions but consist of a solute-like phase (the dispersed phase) and a solvent-like phase (the dispersion medium). The dispersed phase consists of particles larger than typical molecular sizes.
IN CHAPTERS 10, 11, AND 12, we explored the properties of gases, liquids, and solids. Most of the discussion focused on pure substances. However, the matter that we encounter in our daily lives, such as air, seawater, and sand, is usually composed of mixtures. In this chapter we examine homogeneous mixtures. As we noted in earlier chapters, homogeneous mixtures are called solutions. (Sections 1.2 and 4.1)
When we think of solutions, we usually think of liquids, such as a solution of a salt in water, like the seawater shown in this chapter's opening photograph. Solutions, however, can also be solids or gases. For example, sterling silver is a homogeneous mixture of about 7% copper in silver and so is a solid solution. The air we breathe is a homogeneous mixture of several gases, making air a gaseous solution. Because liquid solutions are the most common, however, we focus our attention on them in this chapter.
Each substance in a solution is called a component of the solution. As we saw in Chapter 4, the solvent is normally the component present in the greatest amount, and all the other components are called solutes. In this chapter we compare the physical properties of solutions with the properties of the components in their pure form. We will be particularly concerned with aqueous solutions, which contain water as the solvent and either a gas, liquid, or solid as a solute.
13.1 THE SOLUTION PROCESS
A solution is formed when one substance disperses uniformly throughout another. The ability of substances to form solutions depends on two factors: (1) the natural tendency of substances to mix and spread into larger volumes when not restrained in some way and (2) the types of intermolecular interactions involved in the solution process.
The Natural Tendency toward Mixing
Suppose we have O2(g) and Ar(g) separated by a barrier, as in FIGURE 13.1. If the barrier is removed, the gases mix to form a solution. The molecules experience very little in the way of intermolecular interactions and behave like ideal gas particles. As a result, their molecular motion causes them to spread through the larger volume, and a gaseous solution is formed.
The mixing of gases is a spontaneous process, meaning it occurs of its own accord without any input of energy from outside the system. The extent of spreading of the molecules and their associated kinetic energies is related to a thermodynamic quantity called entropy. We will examine spontaneous processes and entropy in Chapter 19. For our discussion of solutions, it is sufficient to merely recognize that the mixing that occurs when solutions form is associated with an entropy increase. Furthermore, it is the balance of the tendencies of systems to increase their entropy and decrease their energy (or enthalpy) that determines whether a process is spontaneous. Thus, the formation of solutions is favored by the increase in entropy that accompanies mixing.
When molecules of different types are brought together, mixing occurs spontaneously unless the molecules are restrained either by sufficiently strong intermolecular forces or by physical barriers. Thus, gases spontaneously mix unless restrained by their containers because with gases intermolecular forces are too weak to restrain the molecules. However, when the solvent or solute is a solid or liquid, intermolecular forces become important in determining whether or not a solution forms. For example, although ionic bonds hold sodium and chloride ions together in solid sodium chloride (Section 8.2), the solid dissolves in water because of the compensating strength of the attractive forces between the ions and water molecules. Sodium chloride does not dissolve in gasoline, however, because the intermolecular forces between the ions and the gasoline molecules are too weak.
GIVE IT SOME THOUGHT
Which two thermodynamic quantities determine whether or not a process is spontaneous?
The Effect of Intermolecular Forces on Solution Formation
Any of the various intermolecular forces discussed in Chapter 11 can operate between solute and solvent particles in a solution. These forces are summarized in FIGURE 13.2. Dispersion forces, for example, dominate when one nonpolar substance, such as C7H16, dissolves in another, such as C5H12, and ion-dipole forces dominate in solutions of ionic substances in water.
FIGURE 13.1 Spontaneous mixing of two gases to form a homogeneous mixture (solution).
Why does the oxygen atom in H2O point toward Na+ in the ion–dipole interaction?
FIGURE 13.2 Intermolecular interactions involved in solutions.
Three kinds of intermolecular interactions are involved in solution formation:
1. Solute-solute interactions between solute particles must be overcome in order to disperse the solute particles through the solvent.
2. Solvent-solvent interactions between solvent particles must be overcome to make room for the solute particles in the solvent.
3. Solvent-solute interactions between solvent and solute particles occur as the particles mix.
The extent to which one substance is able to dissolve in another depends on the relative magnitudes of these three types of interactions. Solutions form when the magnitudes of the solvent-solute interactions are either comparable to or greater than the solute-solute and solvent-solvent interactions. For example, heptane (C7H16) and pentane (C5H12) dissolve in each other in all proportions. For this discussion, let's arbitrarily call heptane the solvent and pentane the solute. Both substances are nonpolar, and the magnitudes of the solvent-solute interactions (attractive dispersion forces) are comparable to the solute-solute and the solvent-solvent interactions. Thus, no forces impede mixing, and the tendency to mix (increase entropy) causes the solution to form spontaneously.
Solid NaCl dissolves readily in water because the attractive solvent-solute interactions between the polar H2O molecules and the ions are strong enough to overcome the attractive solute-solute interactions between ions in the NaCl(s) and the attractive solvent-solvent interactions between H2O molecules. When NaCl is added to water (FIGURE 13.3), the water molecules orient themselves on the surface of the NaCl crystals with the positive end of the water dipole oriented toward Cl– ions and the negative end oriented toward Na+ ions. These ion-dipole attractions are strong enough to pull the surface ions away from the solid, thus overcoming the solute-solute interactions. In order for the solid to dissolve, some solvent-solvent interactions must also be overcome to create room for the ions to “fit” among all the water molecules.
Once separated from the solid, the Na+ and Cl– ions are surrounded by water molecules. Interactions such as this between solute and solvent molecules are known as solvation. When the solvent is water, the interactions are referred to as hydration.
GIVE IT SOME THOUGHT
Why doesn't NaCl dissolve in nonpolar solvents such as hexane, C6H14?
Energetics of Solution Formation
Solution processes are typically accompanied by changes in enthalpy. For example, when NaCl dissolves in water, the process is slightly endothermic, ΔHsoln = 3.9 kJ/mol. We can use Hess's law to analyze how the solute–solute, solvent–solvent, and solute–solvent interactions influence the enthalpy of solution. (Section 5.6)
How does the orientation of H2O molecules around Na+ differ from that around Cl–?
FIGURE 13.3 Dissolution of an ionic solid in water.
We can consider the solution process as having three components, each with an associated enthalpy change: Solute particles separate from one another (ΔHsolute), solvent particles separate from one another (ΔHsolvent), and solute and solvent particles mix (ΔHmix). The overall enthalpy change, ΔHsoln, is
Separation of the solute particles from one another always requires an input of energy to overcome their attractive interactions. The process is therefore endothermic (ΔHsolute > 0). Likewise, separation of solvent molecules to accommodate the solute also always requires energy (ΔHsolvent > 0). The third component, which arises from the attractive interactions between solute particles and solvent particles, is always exothermic (ΔHmix < 0).
As shown in FIGURE 13.4, the three enthalpy terms in Equation 13.1 can be added 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 solution process is exothermic: ΔHsoln = –91.2 kJ/mol. In contrast, the dissolution of ammonium nitrate (NH4NO3) is endothermic: ΔHsoln = 26.4 kJ/mol. These particular salts are the main components in the instant heat packs and ice packs used to treat athletic injuries (FIGURE 13.5). The packs consist of a pouch of water and the solid salt sealed off from the water—MgSO4(s) for hot packs and NH4NO3(s) 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.
The enthalpy change for a process can provide insight into the extent to which the process occurs. (Section 5.4) Exothermic processes tend to proceed spontaneously. On the other hand, if ΔHsoln is too endothermic, the solute might not dissolve to any significant extent in the chosen solvent. Thus, for solutions to form, the solvent–solute interaction must be strong enough to make ΔHmix comparable in magnitude to ΔHsolute + ΔHsolvent. This fact further explains why ionic solutes do not dissolve in nonpolar solvents. The nonpolar solvent molecules experience only weak attractive interactions with the ions, and these interactions do not compensate for the energies required to separate the ions from one another.
How does the magnitude of ΔHmix compare with the magnitude of ΔHsolute + ΔHsolute for exothermic solution processes?
FIGURE 13.4 Enthalpy changes accompanying the solution process.
By similar reasoning, a polar liquid solute, such as water, does not dissolve in a non-polar liquid solvent, such as octane (C8H18). The water molecules experience strong hydrogen-bonding interactions with one another. (Section 11.2) These attractive forces must be overcome in order to disperse the water molecules throughout the octane solvent. The energy required to separate the H2O molecules from one another is not recovered in the form of attractive interactions between the H2O and C8H18 molecules.
FIGURE 13.5 Ammonium nitrate instant ice pack.
GIVE IT SOME THOUGHT
Label the following processes as exothermic or endothermic:
a. breaking solvent–solvent interactions to form separated particles;
b. forming solvent–solute interactions from separated particles.
Solution Formation and Chemical Reactions
In discussing solutions, we must be careful to distinguish the physical process of solution formation from chemical reactions that lead to a solution. For example, nickel metal dissolves on contact with hydrochloric acid solution because the following reaction occurs:
In this instance the resultant solution is not that of the Ni metal but rather its salt NiCl2. If the solution is evaporated to dryness, NiCl2 • 6 H2O(s) is recovered (FIGURE 13.6). When NaCl(s) is dissolved in water, on the other hand, no chemical reaction occurs. If the solution is evaporated to dryness, NaCl is recovered. Our focus throughout this chapter is on solutions from which the solute can be recovered unchanged from the solution.
FIGURE 13.6 The reaction between nickel metal and hydrochloric acid is not a simple dissolution.
A CLOSER LOOK
Frequently, hydrated ions remain in crystalline salts that are obtained by evaporation of water from aqueous solutions. Common examples include FeCl3 • 6 H2O [iron(III) chloride hexahydrate] and CuSO4 • 5 H2O [copper(II) sulfate pentahydrate]. The FeCl3 • 6 H2O consists of Fe(H2O)63+and Cl– ions; the CuSO4 • 5 H2O consists of Cu(H2O)42+ and SO4(H2O)2– ions. Water molecules can also occur in positions in the crystal lattice that are not specifically associated with either a cation or an anion. BaCl2 • 2 H2O (barium chloride dihydrate) is an example. Compounds such as FeCl3 • 6 H2O, CuSO4 • 5 H2O, and BaCl2 • 2 H2O, which contain a salt and water combined in definite proportions, are known as hydrates. The water associated with them is called water of hydration. FIGURE 13.7 shows a hydrate and the corresponding anhydrous (water-free) substance.
RELATED EXERCISE: 13.4
FIGURE 13.7 A hydrate and its anhydrous salt. The anhydrous salt is the white substance, which turns blue upon addition of water.