Kaplan MCAT General Chemistry Review

MCAT Expertise

Le Châtelier’s principle applies to a wide variety of systems and, as such, appears as a fundamental concept in both MCAT science sections.

CHANGES IN CONCENTRATION OF A REACTANT SPECIES

When you add or remove reactants or products from a reaction in equilibrium, you are causing the reaction to be no longer at its energy minimum state. Metaphorically, you have pushed the reaction, like a ball, up a little ways along the slope of the energy hills on either side of the energy valley. The actual effect is that with the change in concentration of one or another of the chemical species, you have caused the system to have a ratio of products to reactants that is not equal to the equilibrium ratio. In other words, changing the concentration of either a reactant or a product results in Q_c ≠ K_eq. By adding reactant or removing product, you have created a situation in which Q_c < K_eq, and the reaction will spontaneously move in the forward direction, increasing the value of Q_c until Q_c = K_eq. By removing reactant or adding product, you have created a situation in which Q_c > K_eq, and the reaction will spontaneously react in the reverse direction, thereby decreasing the value of Q_c until once again Q_c = K_eq. A simple way to remember this is: The system will always react in the direction away from the added species or toward the removed species.

Bridge

Remember this equation:

CO₂ + H₂O HCO₃^- + H⁺

In the tissues, there is a lot of CO₂, and the reaction shifts to the right. In the lungs, CO₂ is lost, and the reaction shifts to the left. Note that blowing off CO₂ (hyperventilation) is used as a mechanism of dealing with acidosis (excess H⁺).

We often take advantage of this particular tendency of chemical reactions in order to improve the yield of chemical reactions. For example, where possible in the industrial production of chemicals, products of reversible reactions are removed as they are formed so as to prevent the reactions from ever reaching their equilibrium states. The reaction will continue to go in the forward direction, producing more and more product (assuming continual replenishment of reactants as they are consumed in the reaction). You could also drive a reaction forward by starting with higher concentrations of reactants. This will lead to an increase in the absolute quantities of products formed, but the reaction would still eventually reach its equilibrium state, unless product was removed as it formed.

CHANGES IN PRESSURE (BY CHANGING VOLUME)

Because liquids and solids are essentially incompressible, only chemical reactions that involve at least one gas species will be affected by changes to the system’s volume and pressure. When you compress a system, its volume decreases, and the total pressure increases. The increase in the total pressure is associated with an increase in the partial pressures of all the gases in the system, and this results in the system no longer being in the equilibrium state, such that Q_p ≠ K_eq. The system will move either forward or in reverse but always toward whichever side has the lower total number of moles of gas. This result is a consequence of the ideal gas law, which tells us that there is a direct relationship between the number of moles of gas and the pressure of the gas. If you increase the pressure on a system, it will respond by decreasing the pressure by means of decreasing the number of gas moles present. (In this case, the volume of the system was decreased and then held constant while the system returned to its equilibrium state.) When you expand a system, its volume increases, and the total pressure and partial pressures decrease. The system is no longer in its equilibrium state and will react in the direction of the side with the greater number of moles of gas.

Consider the following reaction:

N₂ (g) + 3H₂ 2NH₃ (g)

The left side of the reaction has a total of four moles of gas molecules, whereas the right side has only two moles. When the pressure of this system is increased, the system will react in the direction that produces fewer moles of gas. In this case, that direction is to the right: More ammonia will form. However, if the pressure is decreased, the system will react in the direction that produces more moles of gas; the favored reaction will be the reverse, and more reactants will re-form.

CHANGE IN TEMPERATURE