Clackamas Community College

CH 105: INTRODUCTORY CHEMISTRY

 

 

Contact instructor:

Eden Francis

Physical Science
19600 Molalla Avenue
Oregon City, OR 97045
(503) 594-3352
TDD (503) 650-6649

Lesson 7: Equilibria & Le Chatelier's Principle

In many of the acid-base reactions we have written we’ve used double arrows to indicate that the reaction can proceed in both the forward and reverse directions.

Such reactions are often called reversible reactions. Many reactions are not reversible, however. When was the last time you saw carbon dioxide, water, and fireplace ash combine to form a piece of Douglas fir? Actually, it happens all the time, albeit quite slowly, but the process is not the reverse of the one in which the fir burns to form carbon dioxide, water, and ash. In a reversible reaction, the reverse process proceeds by exactly the same path as the forward.

In a typical reversible reaction the rate of the forward reaction depends on the concentration of the reactants, while the rate of the reverse reaction depends on the concentration of products.

The forward reaction occurs as a result of collisions involving the reactant molecules. The higher their concentration, the more frequently they collide and the higher the rate of the forward reaction. Similarly, the reverse reaction occurs because the product ions collide with each other and stick together. The higher their concentration, the more frequently they collide and the more rapid the reverse reaction.

 

The rate of the forward reaction is highest at the beginning of the reaction, when the concentration of reactants is highest. As the reactant molecules are used up, the forward reaction slows down.

 

Picture the burning piece of Douglas fir again: the fire dies out slowly (the reaction rate decreases) as the fuel is used up. Not all reactions follow this simple pattern, however. Even the fire seems to burn at a constant rate for a long time before it starts to die. This is partly because there is so much oxygen in the air that its concentration does not change appreciably while the fire is burning while, as the log burns, new wood is continually exposed to keep the concentration of available fuel roughly constant until the log gets pretty small.

Meanwhile, the rate of the reverse reaction starts at zero because there are no products around to react. As products are slowly formed, however, the reate of the reverse reaction slowly increases.

This assumes, of course, that the reverse reaction is capable of occurring at all. Many are not, in which case the forward reaction is said to be irreversible. This is the case with the burning wood example.

 

 

As the forward reaction slows and the slow reverse reaction speeds up, there finally comes a point when the two reactions are proceeding at equal rates.

The point at which this equalization of rate occurs depends on the intrinsic rates of the forward and reverse reactions, the theory of which is called chemical kinetics.

As you might already suspect, however, the point at which the forward and reverse rates are the same does not necessarily – in fact almost never – occurs when the concentrations of reactants and products are the same.

 

When the forward and reverse rates are equal, there is no further change in the concentrations of reactants and products and the reaction is said to have reached equilibrium.

If the forward reaction is intrinsically fact and the reverse reaction slow, equilibrium will occur when the concentration of reactants is very low and that of products is very high. Equilibrium is said to lie to the right or to favor products.

If the forward reaction is intrinsically slow and the reverse reaction fact equilibrium will occur when the concentration of reactants is still very high and that of products is very low. Equilibrium is said to lie to the left of to favor reactants.

Although, at equilibrium, there is no further change in concentrations, both the forward and reverse reactions are still taking place, the system is not static and so the equilibrium is called a dynamic equilibrium.

Chemists can demonstrate that a system at equilibrium is dynamic by adding to it reactants “labeled” with radioactive isotopes and showing that over time the labeled atoms begin to appear in product molecules. Similar experiments with labeled product molecules demonstrate that the reverse reaction is occurring as well.

 

Let’s now explore what happens to a system at equilibrium when we disturb it by adding or removing either reactants or products.

Strictly speaking we must change the concentration of reactant (or product) molecules in order to disturb the equilibrium. If we were to change both the number of molecules and the volume of the solution so that the concentrations did not change, the equilibrium would not be affected.

 

If we add only reactant molecules, the rate of the forward reaction would increase, but the rate of the reverse reaction would stay the same because we would not have changed the concentration of product molecules.

By adding reactant, we do change slightly the total volume occupied by the reactants and products and therefore, by diluting them, slightly lower the concentration of product molecules. This has the effect of slowing the reverse reaction a little. For this discussion, we can ignore this small effect.

 

As a result, the concentration of reactants would begin once more to decrease, while the concentration of products would begin to increase and we say that the equilibrium shifts to the right.

You may have noticed that the response of the system countered the change we imposed: we added reactant molecules to the mixture and the equilibrium naturally shifted to the right, using them up.

And you can probably figure out that if we had increased the product concentration, the reaction would have shifted to the left, decreasing the product concentration and increasing the reactant concentration. What would have happened if we have decreased the product concentration? If we had decreased the reactant concentration?

This general principle, that any change imposed on a system at equilibrium will cause the equilibrium to shift in the direction that counteracts the change, is called “Le Chatelier’s Principle” and it applies to all systems in dynamic equilibrium.

Thus, equilibrium will shift to the right whenever the concentration of a reactant is increased or the concentration of a product is decreased.

It will shift to the left whenever the concentration of a reactant is decreased or that of a product is increased.

If a change is imposed which the equilibrium cannot counteract by shifting in one direction or the other, the equilibrium will not be affected by the change.

 

For acid-base equilibria, the pH measures the H+ ion concentration and, indirectly, the OH- ion concentration.

Recall that the pH = -log [H+], so that as the concentration of H+ increases the as the pH decreases. Since the concentrations of H+ and OH- are inversely proportional, as one goes up, the other goes down.

So for an acid-base equilibrium, such as the one shown, decreasing the pH amounts to increasing the H+ concentration, which would cause the equilibrium to shift to the left.

The opposite occurs if the pH is raised – and the H+ concentration lowered – by adding a base. The H+ ions aren’t really “removed” from solution, but the OH- ions in the base react with the H+ ions to form water, which effectively takes them out of the acid-base equilibrium reaction and does, in fact, lower the concentration of the H+ ions.

 

In each square in the table to the right, think about whether imposing the change at the left will increase the concentration of the species at the top or if imposing the change will decrease the concentration.

For example, in the first box you should think about what will happen to the concentration of HA when adding HCl. Remember Le Chatelier’s principle. The equilibrium will shift in the direction that counteracts the change imposed. For each change imposed, figure out which direction the equilibrium must shift in order to counteract it, and then decide what effect shifting in that direction will have on the concentrations of the other species in the reaction.

Let's walk through one together. Adding HA will increase the concentration of a reactant. The equilibrium will therefore shift to the right to reduce it. This will cause the concentrations of H3O+ and A- to increase. Since the H3O+ ion concentration increases, the pH decreases. Although the concentration of HA will also go down when the equilibrium shifts, if will not decrease all the way back to what it was before the additional HA was introduced. So, adding HA does, in fact, increase the concentration of HA, but, because of the resulting equilibrium shift, it does not increase by as much as was added.

change

[HA]

[H3O+]

[A-]

pH

Add HCl

increase (equilibrium is shifting left to use up additional H3O+)

increase (HCl is a strong acid; source of H3O+)

decrease

decrease (gets more acidic)

Add NaOH

decrease (equilibrium is shifting right to make up for less H3O+)

decrease (reacts with the OH-)

increase

increase (gets more basic)

Add HA

increase

increase

increase

decrease

Add A-

increase (equilibrium is shifting left to use up additional A-)

decrease

increase

increase

Remove A-

decrease (equilbrium is shifting right to make up for less A-)

increase

decrease

decrease

 

Next try Exercise 16 before moving on to the last section in this lesson.

Answers to Exercise 16:

16.a. the position of equilibrium will shift to the left
        the pH will increase
        [HX] will increase
        [H3O+] will decrease

    b. [NH3] will increase
        [NH4+] will decrease

    c. pH will decrease
        [HCO3-] will increase

    d. Whenever more acid dissociates, increasing the acidity of a solution, the concentrations of both the hydronium ion and the conjugate base increase.   The conjugate base doesn't make the solution more basic because it is a weaker base than water.

 

In our final section, we'll apply Le Chatelier's principle to a type of a solution called a buffer.

 

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