A chemical reaction reaches a state of chemical equilibrium when the rate of the forward reaction (reactants turning into products) equals the rate of the reverse reaction (products turning back into reactants). This balance is dynamic, meaning the reactions never truly stop; they continue at the molecular level, but the overall concentrations of all substances in the system remain constant over time. This constant composition is the definitive characteristic of a system at equilibrium, setting the stage for how external forces, such as pressure, can influence the outcome.
The Principle Guiding Equilibrium Shifts
The response of a system at equilibrium to an external change is governed by a foundational scientific concept. If a system experiences a disturbance, such as a change in pressure, temperature, or concentration, the system will adjust itself to partially oppose or counteract the applied stress. The reaction will temporarily favor either the forward or the reverse direction until a new equilibrium state is established under the altered conditions.
The Mechanism: Pressure, Volume, and Moles
The effect of pressure on equilibrium is directly tied to changes in the volume of a reaction container, which primarily impacts gaseous substances. Pressure is the result of gas molecules colliding with the walls of their container; therefore, the total pressure is proportional to the total number of gas molecules present. When the volume of a closed container is decreased, the gas molecules are forced closer together, causing an immediate increase in the system’s pressure.
The system counteracts this imposed pressure increase by shifting its equilibrium position to the side of the reaction that contains a smaller total number of gas molecules, or moles. By converting a larger number of gas molecules into a smaller number, the system effectively reduces the total number of collisions with the container walls. This molecular adjustment serves to partially relieve the high-pressure stress.
Conversely, if the pressure is decreased, often by increasing the container’s volume, the system will favor the side of the reaction with the greater number of moles of gas. This shift increases the total number of gas particles, working to restore the pressure that was initially lost.
Conditions for Pressure-Induced Shifts
Pressure changes only influence a chemical equilibrium under two specific and necessary conditions. First, at least one of the reacting substances or products must exist in the gaseous state, as the compression or expansion of solids and liquids results in only negligible changes to their concentrations. Second, the total number of moles of gaseous reactants must be different from the total number of moles of gaseous products. If the count of gas moles is the same on both sides of the reaction arrow, a pressure change will have no effect on the equilibrium position, even if gases are involved.
In a reaction where the total moles of gas on the reactant side is greater than the total moles of gas on the product side, an increase in pressure will push the reaction toward the products. For instance, if four moles of gas react to form two moles of gas, compressing the system shifts the balance to create the two-mole side. A decrease in pressure for the same reaction would cause a shift back toward the four-mole reactant side.
Inert Gas Addition
A common point of confusion involves the addition of an inert gas, such as argon, to a system at equilibrium. If this inert gas is added without changing the container’s volume, the total pressure increases, but the partial pressures—the individual pressures exerted by the reacting gases—remain unchanged. Since the chemical equilibrium constant is defined by the partial pressures of the reacting species, no shift occurs. Only a change that alters the concentration or partial pressure of the reacting gases will induce a shift in the equilibrium position.
Industrial Application of Pressure Control
The industrial synthesis of ammonia, known as the Haber Process, provides a clear example of how pressure control is used to maximize product yield. This reaction combines one mole of nitrogen gas with three moles of hydrogen gas (four total moles) to produce two moles of ammonia gas.
To optimize the production of ammonia, industrial chemists deliberately apply high pressures, typically around 200 atmospheres. Increasing the pressure forces the equilibrium to shift to the side with fewer gas molecules, which is the product side containing ammonia. This high-pressure condition pushes the reaction forward to maximize the conversion of reactants into the desired ammonia product. The resulting increase in yield demonstrates the commercial utility of manipulating the effect of pressure on chemical equilibrium.