EncyclopediaLe Châtelier's principle
Le Châtelier's principle, chemical principle that states that if a system in equilibrium is disturbed by changes in determining factors, such as temperature, pressure, and concentration of components, the system will tend to shift its equilibrium position so as to counteract the effect of the disturbance (see chemical equilibrium). For example, at a given temperature a covered beaker partly filled with water constitutes a system in which the liquid water is in equilibrium with the water vapor that forms above the surface of the liquid. While some molecules of liquid are absorbing heat and evaporating to become vapor, an equal number of vapor molecules are giving up heat and condensing to become liquid. If stress is put on the system by raising the temperature, then according to Le Châtelier's principle the rate of evaporation will exceed the rate of condensation until a new equilibrium is established. At the new equilibrium point a greater proportion of molecules will exist in the vapor phase. Le Châtelier's principle is evident in chemical systems, as in the common-ion effect and in buffer solutions (see also separate article on pH). Le Châtelier's principle can be used to encourage formation of a desired product in chemical reactions. In the Haber process for the industrial synthesis of ammonia, nitrogen gas and hydrogen gas react to form ammonia gas in the reaction N2+3H2→2NH3 ; the process is exothermic, i.e., one that gives off heat. Since four molecules—three of hydrogen and one of nitrogen—react to form two molecules of ammonia, the reactants have a higher gas pressure than the products. When the reaction is run under high external pressure, up to 1000 atmospheres, and relatively low temperature, about 500°C (932°F), the system favors formation of the substance that will result in a lower total number of molecules, i.e., the ammonia. Running the reaction at relatively low temperature causes it to go far to completion, although if the temperature is too much below 500°C the rate of reaction is too slow.
The Columbia Electronic Encyclopedia, 6th ed. Copyright © 2007, Columbia University Press. All rights reserved.
