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Chemical Equilibrium
Introduction

In the laboratory portion of your chemistry course you have had the opportunity to observe a number of chemical reactions. In some cases you have been asked to calculate the amounts of products formed assuming that the reactions go to completion, meaning that the limiting reactants are used up. In fact, many reactions do not go to completion but rather approach an equilibrium state in which both reactants and products are present. Thus, after a certain amount of time, these reactions appear to “stop”—colors stop changing, gases stop evolving, and so forth—before the reaction is complete, leading to a mixture of reactants and products.

As an example, consider N2O4 and NO2 (Figure 15.1), which readily interconvert. When pure frozen N2O4 is warmed above its boiling point (21.2°C), the gas in the sealed tube turns progressively darker as colorless N2O4 gas dissociates into brown NO2 gas (Figure 15.2). Eventually the color change stops even though there is still N2O4 in the tube. We are left with a mixture of N2O4 and NO2 in which the concentrations of the gases no longer change.

The condition in which the concentrations of all reactants and products in a closed system cease to change with time is called chemical equilibrium. Chemical equilibrium occurs when opposing reactions are proceeding at equal rates: The rate at which the products are formed from the reactants equals the rate at which the reactants are formed from the products. For equilibrium to occur, neither reactants nor products can escape from the system.

We have already encountered several equilibrium processes. For example, the vapor above a liquid is in equilibrium with the liquid phase. (Section 11.5) The rate at which molecules escape from the liquid into the gas phase equals the rate at which molecules in the gas phase strike the surface and become part of the liquid. In a saturated solution of sodium chloride, moreover, the solid sodium chloride is in equilibrium with the ions dispersed in water. (Section 13.2) The rate at which ions leave the solid surface equals the rate at which other ions are removed from the liquid to become part of the solid. Both of these examples involve a pair of opposing processes. At equilibrium these opposing processes are occurring at the same rate.

Chemical equilibria explain a great many natural phenomena, and they play important roles in many industrial processes. In this and the next two chapters we will explore chemical equilibria in some detail. Here we will learn how to express the equilibrium position of a reaction in quantitative terms, and we will study the factors that determine the relative concentrations of reactants and products at equilibrium. We begin by exploring the relationship between the rates of opposing reactions and how this relationship leads to chemical equilibrium.



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