Effects of Change in Environmental Factors on Reactions

According to Le Chatelier's principle, if an external stress is applied at equilibrium, the system adjusts to offset the stress and reaches a new equilibrium position. This stress can be a change in the temperature, pressure, volume, or concentration. For example, when more reactants are added to a chemical system or if the products formed are removed, the equilibrium will shift to the right, toward the side of the products. If more products are added or if reactants are removed, the equilibrium will shift to the left, toward the side of reactants.

Thus, reversible reactions are self-directive. That is, even if they are thrown out of balance, they tend to naturally shift and regain their balance. This tendency to regain equilibrium offers an important result that can be implemented in manufacturing of industrial chemicals. Removing products can trigger production of more products. For example, in the contact process, a process that is used in production of sulfuric acid on an industrial scale, sulfur trioxide is the product. Sulfur trioxide is continually removed so that the reaction continues.

Another example is in the Haber-Bosch process, used in the synthesis of ammonia. Nitrogen and hydrogen are made to react to form ammonia gas. This is a reversible reaction with a certain value of equilibrium. The ammonia gas often dissociates again into nitrogen and hydrogen. Since the goal of the process is to produce ammonia, allowing the reaction to reach equilibrium is not beneficial. The goal is to maximize the rate of the forward reaction. One way to keep the reaction moving forward is to continuously remove the ammonia that forms. This change causes a stress in the system, causing more gases to react and form ammonia. Without this change, the yield of ammonia would be very poor.

A temperature change can affect the equilibrium constant. This naturally changes the equilibrium. Depending on whether the reaction is endothermic or exothermic, the left or right side of the reversible reaction may become more favorable. An exothermic reaction releases heat, and an endothermic reaction absorbs heat from the surroundings. A useful way to understand how heat affects a reversible reaction is to treat heat as a reactant or a product. An exothermic reaction has heat as a product. An endothermic reaction has heat as a reactant:

$\rm{A}\rightleftharpoons\rm{B}+\rm{heat}\hspace{25pt}$ Exothermic reaction. Enthalpy $\left( {\Delta H} \right)$, the change in heat in a system, is negative.

$\rm{A}+\rm{heat}\rightleftharpoons\rm{B}\hspace{25pt}$ Endothermic reaction. Enthalpy $\left( {\Delta H} \right)$, the change in heat in a system, is positive.

In an exothermic reaction, increasing temperature will reduce K_c and K_p, favoring the reactants' side. Reducing temperature will increase K_c and K_p, favoring the products' side.

In an endothermic reaction, increasing temperature will increase K_c and K_p, favoring the products' side. Reducing temperature will decrease K_c and K_p, favoring the reactants' side.

The decomposition of dinitrogen tetroxide into nitrogen dioxide is a reversible, endothermic reaction. At a temperature of 100°C, both K_c and K_p values are low. When heat is supplied, K_c and K_p values change dramatically. Heat causes the system to shift toward the product side, and there is a greater yield of nitrogen dioxide.

See the following example reaction:

Pressure and volume changes are often related. An increase in volume without changing the amount of gas in a system will result in a pressure decrease. A decrease in volume will result in a pressure increase. Adding or removing gases from a system may change pressure without changing volume.

If pressure is increased, the equilibrium will shift to the side of the reaction that has fewer moles of gas. In the same way, if the pressure is decreased, the shift will be to the side that favors a greater number of moles of gas. If both sides have the same number of moles of gas, a change in pressure will not affect equilibrium.

Again, consider the synthesis of ammonia by the Haber-Bosch process in a closed system. In this reaction, one mole of nitrogen gas reacts with three moles of hydrogen gas to form two moles of ammonia. The reactants have four moles of gas, while the product has two.

In a closed system, if pressure is increased, the equilibrium shifts toward the side with fewer moles of gas—the products' side, in this example. If pressure is decreased, the equilibrium shifts toward the side with more moles of gas—the reactants' side, in this example.

A catalyst has no effect on either K_c or K_p. When a catalyst is added to a reversible reaction, it will not have any effect on the equilibrium constants because it speeds up both the forward and reverse reactions equally, thereby allowing the system to reach equilibrium faster.