In this explainer, we will learn how to explain the effect of changing the temperature, concentration, or pressure on an equilibrium according to Le Chatelier’s principle.
There are many proverbs among gardeners about pruning trees and shrubs. Some gardeners say that to make a rose bush grow twice as fast, we must prune the branches first. The idea is that the plant is resisting the change that has been imposed upon it. To replace what was lost, the rose grows back rapidly. In chemistry, there is a similar idea that concerns chemical reactions in equilibrium.
Definition: Dynamic Equilibrium
A dynamic equilibrium is reached in a reversible chemical reaction when the rate of the forward reaction is equal to the rate of the reverse reaction. The concentrations of reactants and products will remain constant. The forward reaction rate and the reverse reaction rate are not zero. The reaction must be contained within a closed system.
When chemical reactions have reached dynamic equilibrium, there are microscopic, or molecular level, changes still taking place. This is different from static equilibrium. These microscopic changes are in balance with each other, so there is no overall change to the system to the outside observer.
For example, a chemical reaction in dynamic equilibrium, in a closed system, may remain a constant color. Of course, this would depend on the color of the reactants and products involved in the overall reaction. An example of a system in equilibrium is the conversion between nitrogen dioxide and dinitrogen tetroxide:
Nitrogen dioxide is a dark-brown, or reddish-brown, gas and dinitrogen tetroxide is a colorless gas. At equilibrium, both gases are present in the mixture and a constant pale brown color may be observed. If a change is suddenly made to this gas mixture, when it is in equilibrium, we will observe a color change.
The changes that can be made to this system may include temperature, pressure, and concentration. Any of these changes may cause the position of the equilibrium to shift.
A catalyst speeds up the forward and reverse reactions equally, so it does not affect the position of an equilibrium. Catalysts simply reduce the time taken for the system to reach equilibrium. The equilibrium will shift either in the direction of the reactants or in the direction of the products. A shift in the direction of the reactants is often referred to as a shift to the left, or left-hand side. Conversely, a shift in the direction of the products is often referred to as a shift to the right, or right-hand side. When a system in equilibrium is suddenly disturbed in this way, it will respond until the equilibrium is restored. The French chemist Henri Louis Le Chatelier was one of the first people to investigate the effects of different factors on equilibria. This led to Le Chatelier’s principle.
Definition: Le Chatelier’s Principle
For a dynamic equilibrium, if the conditions change (concentration, temperature, or pressure), the position of equilibrium will move to counteract the change.
We can explore how Le Chatelier’s principle works by applying the idea to various systems in equilibrium. We will explore how concentration changes affect the position of an equilibrium first.
If iron(III) ions are added to a solution of thiocyanate ions (), a bloodred color develops rapidly as the thiocyanatoiron(III) ion is formed. The color is caused by a complex ion with the formula . At relatively low initial concentrations of reactants, the bloodred color is not too intense and an equilibrium is rapidly established:
If we suddenly change the concentration of one of the species in the equilibrium mixture, we can observe how the equilibrium responds. The response will be according to Le Chatelier’s principle. For example, if we add a soluble iron(III) salt to the mixture, such as iron(III) nitrate, we will raise the concentration of the iron(III) ions in the solution. To respond to this change, the equilibrium will shift to the side of the products, or to the right-hand side. This shift in the position of the equilibrium opposes the original change. Some of the extra iron(III) ions added to the mixture will react with available ions. The concentration of iron(III) ions will then drop, as will the concentration of ions. The concentration of complex ions will rise as equilibrium is restored. We will observe an increase in the intensity of the bloodred color in the mixture.
Conversely, we could remove some of the dissolved ions from the equilibrium mixture. This could be achieved by precipitating them out by adding a few drops of sodium hydroxide solution. The equilibrium would oppose the change and shift to the left-hand side in order to replace the ions that were removed from the solution. We would observe a fading of the bloodred color and the mixture would become a paler red color.
Example 1: Determining How a Change in Concentration Affects the Equilibrium Position
Iodine trichloride is a bright yellow interhalogen compound formed in the following reaction:
What effect will removing chlorine have on the position of the equilibrium?
- The equilibrium will move to the right.
- The equilibrium will move to the left.
- The equilibrium will not move.
Chlorine lies on the left-hand side of the equilibrium equation. If some of the chlorine is removed, the equilibrium responds to oppose the change. This is according to Le Chatelier’s principle. The equilibrium will shift to the left-hand side to replace the chlorine that has been removed. The concentration of iodine trichloride in the mixture will fall accordingly. The bright yellow color of iodine trichloride will be replaced by the darker red/brown color of iodine monochloride.
The answer is therefore B. The equilibrium will move to the left.
When a reaction system in equilibrium involves gaseous particles, the particles are constantly colliding with each other. The particles are also colliding with the walls of the container that ensures there is a sealed system. These collisions with the container exert pressure on the container. The total pressure in the container is related directly to the number of gas particles that it contains.
The equilibrium position in a gas-phase system can be influenced by the total pressure of the system. If the total pressure is suddenly increased, the gas particles are forced closer together. To relieve the pressure, the equilibrium can shift to oppose the change. This is in accordance with Le Chatelier’s principle. Take, for example, the reaction of nitrogen gas with hydrogen gas to produce ammonia gas:
In this reaction, one molecule of nitrogen reacts with three molecules of hydrogen to make two molecules of ammonia. On the left-hand side of the equation, there are four molecules, and on the right-hand side, there are just two. If the equilibrium were to move in the forward direction, forming more ammonia, then there would be a decrease in the total number of molecules in the system. This would result in a decrease in the total pressure of the system. If the total pressure of the system were to be increased, the pressure change would be opposed by a shift to the right-hand side. More ammonia would be formed as equilibrium is restored.
In general terms, an increase in pressure will favor the reaction that produces fewer molecules of gas, or fewer moles of gas. A decrease in pressure will favor the reaction that produces more moles of gas. In the reaction of nitrogen with hydrogen, a higher yield of ammonia is obtained if the reaction is run at higher pressure.
Some gas-phase reactions have the same number of moles of gas particles on each side of the balanced equation. The reaction of hydrogen with gaseous iodine, to produce hydrogen iodide gas, is an example:
In this case, a change in the total pressure of the equilibrium mixture will have no effect on the position of the equilibrium. Pressure changes cannot be opposed here by shifting the equilibrium position to the left or right.
Example 2: Determining How a Change in Pressure Affects the Equilibrium Position
The following reaction is part of the Ostwald process that produces nitric acid:
Which of the following statements explains why the percentage of produced may decrease as the pressure increases?
- Increasing the pressure moves the equilibrium to the right.
- The forward reaction is favored by an increase in pressure.
- The total volume of gas increases for the backward reaction.
- Fewer moles of gas are on the reactant’s side.
In this reaction, there are a total of nine gas moles on the reactant side. This is composed of four moles of ammonia plus five moles of oxygen. On the product side of the equation, there are a total of ten gas moles. This is composed of four moles of and six moles of water vapor. If we increase the total pressure of the system, the equilibrium shifts to the side with less gas moles. This is to oppose the change that has been made. A shift to the left-hand side of the equilibrium would be expected. Some of the will react with some of the water vapor and more ammonia and more oxygen will be produced. The percentage of in the new equilibrium mixture will therefore be reduced. The reason for this reduction as the pressure increases is because there are fewer moles of gas on the reactant’s side, and so the correct answer is D.
We will now explore the effect of temperature on the position of an equilibrium. We will use the reaction of nitrogen dioxide in equilibrium with its dimer, dinitrogen tetroxide, as an example:
The forward reaction is exothermic in this case. The reverse reaction process must therefore be endothermic.
Definition: Exothermic Reaction
In an exothermic reaction, heat energy is released to the surroundings. The heat content of the products is less than the heat content of the reactants.
Definition: Endothermic Reaction
In an endothermic reaction, heat energy is absorbed from the surroundings. The heat content of the products is greater than the heat content of the reactants.
If the equilibrium shifts to the right-hand side, more dinitrogen tetroxide is produced. This gas is a pale color, much paler than the dark-brown nitrogen dioxide. Since the forward reaction is exothermic, a release of heat energy will accompany the formation of dinitrogen tetroxide.
If the equilibrium shifts to the left-hand side, more of the dark-brown nitrogen dioxide is produced. This is an endothermic process and heat would be absorbed from the surroundings.
If an equilibrium mixture of these gases was placed into hot water, heat energy is added to the system. Le Chatelier’s principle suggests that the system will oppose the change. To absorb the added heat, the equilibrium will shift in the endothermic direction. The mixture will become a darker color as more nitrogen dioxide is formed.
If the equilibrium mixture was placed in ice-cold water, the equilibrium will shift in the exothermic direction to oppose this change. This favors the production of dinitrogen tetroxide, which is a pale gas. The color of the gas mixture will become paler as less of the dark-brown nitrogen dioxide is present in the new equilibrium mixture.
Example 3: Determining How a Change in Temperature Affects the Equilibrium Position
is a brown gas that can exist in equilibrium with colorless dinitrogen tetroxide, as shown in the following reaction:
Is the enthalpy change of the forward reaction exothermic or endothermic? Which ampoule in the picture is at the highest temperature?
- The forward reaction is endothermic, and the ampoule on the left is the hottest.
- The forward reaction is exothermic, and the ampoule on the left is the hottest.
- The forward reaction is endothermic, and the ampoule on the right is the hottest.
- The forward reaction is exothermic, and the ampoule on the right is the hottest.
The forward reaction here involves the combination of two identical molecules, which results in bond formation. We can see that the enthalpy change, quoted for the forward reaction, is negative. This means that the forward reaction is exothermic. The high temperature will therefore favor the endothermic reverse reaction. The equilibrium would shift to the left-hand side. When the temperature is high, we would expect there to be more in the equilibrium mixture. is a brown gas, so the mixture will darken. The ampoule on the right of the picture is the darkest as it contains the most nitrogen dioxide. This indicates that the ampoule on the right is the hottest, and so the correct answer is D.
Example 4: Determining How a Change in Temperature Affects the Equilibrium Position
The industrial production of ethanol uses ethene gas and steam, as shown:
Considering that the forward reaction is exothermic, what effect will increasing the temperature have on the position of the equilibrium?
- The equilibrium will not move.
- The equilibrium will move to the left.
- The equilibrium will move to the right.
Since the forward reaction process is exothermic, heat will be released if the equilibrium shifts in this direction. Adding more heat to the system, by raising the temperature, favors the endothermic process. If the temperature were raised, the equilibrium would oppose the change and shift in the endothermic direction. This involves a shift to the left. The yield of ethanol in the mixture would decrease.
The correct answer is therefore B. The equilibrium will move to the left.
- A dynamic equilibrium is established in a sealed system when the forward rate of a reversible reaction is equal to the rate of the reverse reaction.
- If a change is made to the system in equilibrium, the equilibrium position will shift to oppose the change that was made. This is the basis of Le Chatelier’s principle.
- The changes that may affect the position of the equilibrium can include increases or decreases in concentration, total pressure, or temperature.
- Adding a catalyst does not influence the position of an equilibrium. It reduces the time taken for the system to reach equilibrium.
- Increasing the temperature of a system in equilibrium favors the endothermic process. This could be the forward reaction or the reverse reaction, depending on the enthalpy changes of these respective processes.