18.2 Reversible Reactions and Equilibrium
Standards: 8b, 9a, 9b, 9c
Vocabulary
reversible reaction, chemical equilibrium, equilibrium position, Le Chatelier’s principle, equilibrium constant
Connection to your world
For years, scientist tried to produce nitrogen compounds for fertilizers to increase the amount of available foodstuffs. Unfortunately, none of these efforts proved to be commercially successful. Finally, in the early 1900s, two German chemists, Fritz Haber and Karl Bosch, refined the process of making ammonia from elemental nitrogen and hydrogen. Their success came from controlling the temperature and pressure under which two gases are reacted. In this section, you will learn how changing the reaction conditions can influence the yield of a chemical reaction.
Reversible Reaction
From what you have already learned, you may have inferred that chemical reaction go completely from reactants to products, just as the chemical equation indicates this is not usually the case. Some reactions are reversible. A reversible reaction is one in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously. One example of a reversible reaction is the following.
Forward reaction: 2SO2(g) + O2(g) à 2SO3(g)
Reverse reaction: 2SO2(g) + O2(g) ß 2SO3(g)
In the first section, which is read from left to right, sulfur dioxide and oxygen produce sulfur trioxide? In the second, which id read from right to left, sulfur trioxide decomposes into oxygen and sulfur dioxide. The first section is called the forward reaction. The second is called the reverse reaction. The two equations can be combined into one using a double arrow. The double arrow tells you that this reaction is reversible.
2SO2(g) + O2(g) 2SO3(g)
Sulfur Oxygen Sulfur
dioxide trioxide
PCl5(g) + heat PCl3(g) + Cl2(g)
These graphs show how the concentrations of A, B, C, and D over time
What actually happens when sulfur dioxide and oxygen gases are mixed in a sealed chamber? The two reactants begin to react to form sulfur trioxide at a particular rate. Because no sulfur trioxide is present at the beginning of the reaction, the initial rate of the reverse reaction is zero. As sulfur trioxide begins to form, however, the decomposition of sulfur trioxide begins. This reverse reaction proceeds slowly at first, but its rate increased as the concentration of sulfur trioxide increase. Simultaneously, the rate of the forward reaction decreases because sulfur dioxide and oxygen are begin used up. Eventually sulfur trioxide is decomposing to sulfur trioxide and oxygen as fast as sulfur dioxide and oxygen are forming sulfur trioxide. When the rates of the forward and reverse reaction are equal, the reaction has reached a state of balance called chemical equilibrium. Changes in concentrations of the three components during the course of the reaction are shown in the graphs in figure 18.10. The graph on the left shows the process of the reaction that start with specific concentration of SO2 and O2 but with zero concentration of SO3. The graph on the right shows concentration for a reaction that begins with an initial concentration of SO3 and zero concentration for SO2 and O2 . Notice that after a certain time, all concentration remain constant. At chemical equilibrium, no net change occur s in the actual amounts of the components of the system. The amount of SO3 in the equilibrium mixture is the maximum amount that can be produced by this reaction under the conditions of the reaction.
The unchanging amounts of SO2, O2 and SO3 in the reaction mixture at equilibrium might cause you to think both reactions have stopped. This is not the case. Chemical equilibrium is a dynamic state. Both the forward and reverse reactions continue, but because their rates are equal, no net change occur in their concentrations. The escalators in Figure 18.11 are like the double arrows in a dynamic equilibrium equation. The number of the people using the up escalator must equal the number of the people using the down escalator for the number of the people on both floors to remain constant
√Checkpoint: Why is chemical equilibrium called a dynamic state?
Although the rates of the forward and reverse reactions are equal at chemical equilibrium, the concentrations of the components on both sides of the chemical equation are not necessarily the same. In fact, they can be dramatically different. Figure 18.10 shows the equilibrium concentrations of SO2, O2, and SO3. The concentration of SO3 is significantly greater then the concentration of SO2, and O2. The relative concentrations of the reactants and products at equilibrium constitute the equilibrium position of a reaction. The equilibrium position indicates whether the reactants or products are favored in a reversible reaction. If A reacts to give B and the equilibrium mixture contains significantly more of B—say 1% A and 99% B—then the formation of B is said to be favored.
A B
1% 99%
On the other hand, if the mixture contains 99% A and 1% B at equilibrium, then the formation of A is favored.
A B
99% 1%
Notice that the equilibrium arrows are not of equal length; the longer of the two arrows indicates the favored direction of the reaction.
In principle, almost all reactions are reversible to some extent under the right conditions. In practice, one set of components is often so favored at equilibrium that the other set cannot be detected. If one set of components (reactants) is completely converted to new substances (products), you can say that the reaction has gone to completion, or is irreversible. When you mix chemicals expecting to get a reaction but no products can be detected, you can say that there is no reaction. Reversible reactions occupy a middle ground between the theoretical extremes of irreversibility and no reaction.
A catalyst speeds up both the forward and the reverse reactions equally because the reverse reaction is exactly the opposite of the forward reaction. The catalyst lowers the activation energy of the reaction by the same amount in both the forward and reverse directions. Catalysts do not affect the amounts of reactants and products present at equilibrium; they simply decrease the time it takes to establish equilibrium.
√Checkpoint: What is chemical equilibrium?
Factors Affecting Equilibrium: Le Chatelier’s principle
A delicate balance exists in a system at equilibrium. Changes of almost any kind can disrupt this balance. When the equilibrium of a system is disrupted, the system makes adjustments to restore equilibrium. However, the equilibrium position of the restored equilibrium is different from the original equilibrium position; that is, the amount of products ad reactants may have increased or decreased. Such a change is called a shift in the equilibrium position.
The French chemist Henri Le Chaterlier (1850-1936) study how the equilibrium position shifts as a result of changing conditions. He proposed what has come to be called Le Chatelier’s Principle: if a stress is applied to a system in dynamic equilibrium, the system changes in a way tat relives the stress. Stresses that upset the equilibrium of a chemical system include changes in the concentration of reactants or products, changes in temperature, and changes in pressure.
The following examples of applications of Le Chatelier’s principle all involve reversible reactions. For simplicity and clarity, the components to the left of the reaction arrow will be considered the reactants and the components to the right of the reaction arrow will be considered the products. Blue arrows indicate the shifts resulting for additions to or removals from the system. The arrows always point in the direction of the resulting shift in the equilibrium position-that is, toward the favored side.
Concentration Changing the amount, or concentration, of any reactant or product in a system at equilibrium disturbs the equilibrium. The system adjusts to minimize the effects of the change. Consider the decomposition of carbonic acid (h2co3) in aqueous solution to for the products carbon dioxide and water. The system has reached equilibrium. At equilibrium the amount of carbonic acid is less than 1%.
Add CO2
ß Direction of shift
H2CO3(aq) CO2(aq) + H2O (l)
Remove CO2
<1% Direction shift à >99%
Adding more carbon dioxide disrupts the equilibrium. It increases the concentration of CO2 in the mixture and cause the rate of the reverse reaction to increase. As more reactant (H2CO3) is formed, the rate of the forward reaction also begins to increase. In time, the rates of the forward and reverse reaction again become equal, and a new equilibrium is established with a higher concentration of reactant (H2CO3). Adding a product to a reaction at equilibrium pushes a reversible reaction in the direction of reactants.
If, on the other hand, carbon dioxide is remove, the concentration of CO2 decreases. This causes the rate of the reverses reaction to decrease. As less reactant (H2CO3) is being formed, the rated of the forward reaction also begins to decrease. When the rates of the forward and reverse reaction again become equal, equilibrium is restored but as different equilibrium position. Removing a product always pushes a reversible reaction in the direction of products.
Farmers use this technique to increase the yield of eggs laid by hens. Hens lay eggs and then proceed to hatch them. If the eggs are removed after they are laid (removing the product), the hen will lay more eggs (increasing the yield). Similarly as products are removed from a reaction mixture, they system continually changes to restore equilibrium by producing more products. But because the products are being removed, the reactants can never reestablish equilibrium. The reaction continues to produce products until the reactants are used up. Another example of this concept is found in the body. Blood contains dissolved carbonic acid in equilibrium with carbon dioxide and water. The body uses the removal of products to keep the concentration of carbonic acid within a safe range. When the athletes exhale carbon dioxide, the equilibrium shifts toward carbon dioxide and water, thus reducing the amount of carbonic acid. The same principle applies to adding or removing reactants. When a reactant is added to a system at equilibrium, the reaction shirts in the direction of the formation of products. When a reactant is removed, the reaction shifts in the direction of formation of reactants.
√Checkpoint According to Le Chatelier’s principle, how does a system at equilibrium respond to a stress?
Temperature Increasing the temperature causes the equilibrium position of a reaction to shift in the direction that absorbs heat. The heat absorption reduces the applied temperature stress. For example, consider the following exothermic reaction that occurs when SO3 and O2.
Add heat
ß Direction of shift
2SO2 (g) + O2 2SO3 (g) + heat
Remove heat (cool)
Direction of shift à
Heat can be considered to be a product, just like SO3. Heating the reaction mixture at equilibrium pushes the equilibrium position to the left, which favors the reactants. As a result, the product yield decreases. Cooling, or removing heat, pulls the equilibrium to the right, and the product yield increases.
Pressure A change in the pressure on a system affects only gaseous equilibria that have unequal number of moles of reactants and products. An example is the reaction you read about is hydrogen and nitrogen react too form ammonia. Imagine that the three gases are at equilibrium in a cylinder that has a piston attached to a plunger-similar to a bicycle pump by with the hose sealed. A catalyst has been including to speed up the reaction. What happens to the pressure when you push the plunger down? The pressure on the gas momentarily increases because the same number of molecules is contained in a smaller volume. The system immediately relieves some of the pressure increase by reducing the number of gas molecules. For every two molecules of ammonia made, four molecules of reactant are used up (three molecules of hydrogen and one of nitrogen). Therefore, the equilibrium position shifts to make more ammonia. There are then fewer molecules in the system. The pressure decreases, although it will not decrease all the way to the original pressure. As you can see, increasing the pressure on the system results in a shift in the equilibrium position that favors the formation of product.
Increase pressure
Direction of shift à
N2(g) + 3H2 2NH3 (g)
Reduce pressure
ß Direction of shift
The equilibrium position for this reaction can be made to factor the reactants instead of the product. Imagine pulling the plunger of the piston device back up so the volume containing the gases increases. The increase in volume decreases the pressure on the system. To restore the higher starting pressure, the system can produce more gas molecules by decomposition of some ammonia molecules. Decomposition of two molecules of gaseous NH3 produces four molecules of reactants (three H2 and one N2). Pressure at the new equilibrium is higher than when the pressure was first decreased, but not as high as it was at the starting equilibrium. Lowering the pressure on the system thus results in a shift of the equilibrium to favor the reactants.
CONCEPTUAL PROBLEM 18.2
1. How do the amounts of reactants and products change after a reaction has reached chemical equilibrium?
2. What are three stresses that can upset the equilibrium of a chemical system?
3. What does the value of the equilibrium constant tell you about the amounts of reactants and products present at equilibrium?
4. How can a balanced chemical equation be used to write an equilibrium-constant expression?
5. Can a pressure change shift the equilibrium position in every reversible reaction? Explain.