0.17 Phase equilibrium in solutions

Introduction

In our study of phase equilibrium, we have examined only pure materials. However, we will eventually want to study chemical reactions, which will mean understanding solutions with many components that might react with one another. Before considering reactions then, we need to consider what happens to phase equilibrium when there is more than one component present. How does mixing things together change equilibrium? This is actually a quite general question that we will address in many contexts including chemical reactions. For now, we’ll consider the phase equilibrium first, since that is where we have discovered equilibrium and that is where we have developed an understanding based on the concept of dynamic equilibrium.

There are many types of solutions to consider. We can mix together gases with gases, liquids with liquids, gases with liquids, solids with liquids, and so on. Each of these present different challenges, but we will find that there are similarities amongst them as well. In particular, we’ll find that dynamic equilibrium can be applied in each of these cases to understand the phase equilibrium that exists amongst the different components of the solutions.

Foundation

We will assume some understanding of solutions. Recall that the solvent is the major component in a solution, and is typically but not always a liquid. Far more often than not in Chemistry we use water as our solvent. A solute is a minor component of a solution, and in a single solution there may be more than one solute. The solute, in its pure form, can be a solid, another liquid, or even a gas. Once in solution, the solute usually has very different properties and is generally no longer recognizable from its pure form. Think about dissolving sugar or salt in water. The solution formed shows no evidence of the original crystalline solid solute.

Solutions are defined in large part by the concentration of the solute in the solvent. There are many ways to define and measure the concentration. The most common is the “molarity” of the solution, meaning the number of moles of solute per liter of solution. The units of molarity are “molar” with a capital M: a solution with 1.0 mole of solute in 1.0 L of solution is a 1.0 M solution. In this study we will also discuss the mole fraction of the solute. This is simply the number of moles of solute divided by the total number of moles of particles of all types in the solution.

We will lean very heavily on the concept of dynamic equilibrium. The idea will show up in all of our explanations that two competing processes at equilibrium must have the same rate. We will examine several types of processes and the factors that determine their rates.

Observation 1: lowering of the vapor pressure in solution

To begin studying solutions, our first task is to observe what impact, if any, the presence of a solute has on the properties of the solvent. We will begin with a simple two-component solution, with a solvent and a single solute. The type of solute will matter to us, as we will observe different behaviors for different solutes, particularly whether the solute is, in its pure form, a solid, a liquid, or a gas. To start, we will consider solutions formed by dissolving a solid solute into a liquid solvent. This choice is easiest to start with because the solid solute will be assumed to be non-volatile. That is, it does not readily evaporate and therefore has zero vapor pressure. Solids do have a vapor pressure, but for most solids, the vapor pressures are sufficiently small that we can ignore them. As a first guess, then, we might assume that the solution formed from a volatile solvent and a non-volatile solute would have the same vapor pressure as the solvent alone, since the solute seems to contribute nothing to the equilibrium vapor pressure.