Up to this point, chemical reactions have likely been presented to students as being either "complete" or having Òno reaction.Ó As a consequence, it is probably not surprising that students are puzzled by chemical reactions where constant amounts of reactants and products are present simultaneously. This phenomenon is because chemical reactions can be reversible. Consider the reaction between nitrogen oxide and ozone represented by the equation:
Using a simple collision theory model of reaction, NO and O3 molecules can collide, rearrange their atoms, and form NO2 and O2 molecules as products. But if this is so, why can't NO2 and O2 molecules collide and reform NO and O3 molecules? The answer is that they can; the reaction is probably better represented by the equation
The rate at which NO2 and O2 form varies because of the changing concentrations of NO and O3 as they are used up while being converted to NO 2 and O2. The same is true for the reverse reaction. Eventually, the rates of the forward and reverse reactions become the same, and the consequence is a state of "balance" or "equilibrium" among the amounts of NO, O3, NO2, and O2 present. Although the reaction appears to have ceased on the macroscopic scale, the forward and reverse reactions continue at the molecular level. This is a "dynamic" equilibrium in contrast to a "static" equilibrium as represented, for example, by a balanced see-saw.
Several examples of dynamic equilibria (as physical changes) may already be familiar to your students. When water evaporates in a closed container, a constant "vapor pressure" is reached. This pressure is exerted by the water vapor molecules when the rate of evaporation is equal to the rate of condensation. Also, in the dissolving of a solid in a liquid solvent, a limit is reached where no more solid will dissolve. This saturation point is reached when the rate of dissolution is equal to the rate of crystallization.
In quantitative terms, chemists are interested in the position of equilibrium,
i.e., to what extent the reaction has approached completion when the equilibrium
state is reached (i.e., is the reaction 90% complete, 60% complete?). The
extent of completion for chemical reactions varies with different reactions.
It is usually expressed as the ratio of concentrations of products to reactants
(which is theoretically a constant value, K, at a fixed temperature). Specifically,
for the hypothetical reaction: a A + b B
c
C + d D, where capital letters represent reactants and products and lowercase
letters represent coefficients: 
The position of an equilibrium can be shifted toward reactants or products according to LeChatelier's principle (a stress placed on a system at equilibrium will cause a shift that minimizes the stress). Factors (stresses) that influence chemical equilibrium positions are those that affect the rates of forward reactions differently from the rates of reverse reactions. The principal factors are concentration effects and temperature effects. Pressure effects are often included in this list, but the ideal gas law shows that pressure can only be changed by changing concentration (concentration is equal to n/V; consequently any changes in amount of gas or volume of gas will change the concentration), or by changing the temperature. Equilibrium is usually considered a ÒlaterÓ topic in first-year chemistry. When it is taught, there are three different levels to approaching the content: (1) limiting instruction to descriptive or qualitative issues (LeChatelierÕs principle), (2) interpreting values of K (e.g., weak vs. strong acids), and (3) a complete quantitative approach (calculating values of K). The focus of this module is on the qualitative level of equilibrium, leaving quantitative aspects to Extensions and Projects .
After completing their study of equilibrium, students should be able to:
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