Le Chatelier's Principle

 

The Basics

According to Le Chatelier himself, the principle means: 

 

any system in stable chemical equilibrium, subjected to the influence of an external cause which tends to change either its temperature or its condensation (pressure, concentration, number of molecules in unit volume), either as a whole or in some of its parts, can only undergo such internal modifications as would, if produced alone, bring about a change of temperature or of condensation of opposite sign to that resulting from the external cause.

 

In simpler terms, systems in dynamic equilibrium counteract disturbances to their conditions by shifting the equilibrium to counteract and "balance" the system and restore equilibrium. The physical application for this principle is that if a chemical reaction in equilibrium experiences a change in pressure, temperature, or concentration of products or reactants, then the equilibrium shifts in the opposite direction  in order to offset the change. Le Chatelier's principle is extremely useful in determining what happens when conditions in a reaction at dynamic equilibrium are changed; however, the principle does not address why these things happen molecularly. Le Chatelier's principle answers "what" not "how." 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Via http://img.docstoccdn.com/

 

 

It is useful to think of Le Chatelier's principle as a see-saw or balance while keeping in mind that the system wants to keep balanced so as not to tip over. In this way, we can see that an addition of reactants or products would tip the scale, so logically, the system must favor the opposite in order to keep itself from "tipping over."

 

Changing Concentrations

According to Le Chatelier's Principle, if the concentration of one reacting species is increased, the system must favor the reaction in which that species is consumed. Essentially, Le Chatelier's Principle states that a system at dynamic equilibrium does not want the concentrations of any products or reactants to change, and the system will do whatever it needs to do to return the concentrations to an equilibrium state. This is due the relationship between the reaction quotient and the equilibrium constant. If more reactants are added to the system, the reaction will proceed to the right because the reaction quotient is less than the equilibrium constant: Q<K. The inverse is true when more products are added to the system. The reaction quotient will be greater than the equilibirum constant and the reaction will proceed to the left: Q>K. 

 

 

 

 

 

 

 

 

 

     

   Via http://chemwiki.ucdavis.edu/

     Consider the reaction: A+2B⇌C+D

 

Sticking to the "see-saw" or "balance" train of thought, we can consider the collective concentrations on each side of the reaction as objects being "balanced." Therefore, if we add A or B, the reaction must produce more C and D by using up A and B in order to stay balanced. The same is true for adding C or D. Similarly, if we reduce the concentration of C or D, the reaction must also proceed in favor of them in order to stay balanced. Think of the concentrations as weights being balanced on a scale or on each end of a see-saw, and it is much easier to determine which side needs more (which way the reaction will proceed) to maintain equilibrium. 

 

Practice Problem #1: The reaction (HC2H3O2 <----> H+ + C2H3O2-) is in dynamic equilibrium. If the concentration of H+ is increased, the system will shift to re-establish equilibrium. What will be a result of this shift?

 

A.) More C2H3O2- will be produced

 

B.) Less HC2H3O2 will be produced

 

C.) More HC2H3O2 will be produced

 

D.) Less C2H3O2- will be produced

 

Changing Pressure

Le Chatelier's Principle applies to pressure changes when gaseous species are involved in a reaction; however, not all species need to be a gas for the principle to apply. In a system at equilibrium, pressure changes will make the reaction favor either products or reactants depending upon which will keep the pressure the same. For instance, if pressure is increased, the system will favor less moles of gas and the reaction will proceed in whatever way that facilitates that. The inverse is true for a decrease in pressure. If pressure is decreased, the system will favor more moles of gas. In other terms, a change in pressure can be described as a change in volume, with an increased volume equating to a decreased pressure and a decreased volume equating to an increased pressure. This is because the system is in a confined space and the volume of the container will directly correlate to the pressure of the system when gaseous species are involved. 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

  Via s-owl.cengage.com

 

As you can see, as the volume was decreased and pressure increased, the system favored less moles of gas, and as a result, more NO2 was converted to N2O4. This is due to the fact that 2 moles of NO2 are present for every mole of N2O4. This shift allows the system to maintain the same pressure it had at equilibrium. Keep in mind, the system simply wants to maintain the pressure it had at equilibrium. As a result, it will shift in order to obtain the number of moles of gas that it needs to do that based on the conditions imposed upon the system. 

 

Practice Problem #2: The reaction (N2 (g) + 3H2 (g) <---> 2NH3 (g)) is in dynamic equilibrium in a sealed piston. If the piston is compressed, in which direction will the reaction shift and why?

 

A.) The reaction will shift to the left because there are more moles of gas and the system needs to raise its pressure

 

B.) The reaction will shift to the right because there are less moles of gas and the system needs to lower its pressure

 

C.) The reaction will not shift at all

 

D.) The reaction will shift to the right because there are less moles of gas and the system needs to raise its pressure

 

Changing Temperature

It is a bit more complicated to examine the effect of temperature change on a system in equilibrium using Le Chatelier's Principle than concentration or pressure. This is due to the fact that the sign of reaction enthalpy must be known (endothermic or exothermic). While incorrect technically speaking, it is convenient to think of heat as either a product (exothermic, ΔH<0  ) or a reactant (endothermic, ΔH>0 ). By doing this, we can once again use the see-saw or balance method. If the temperature of a reaction is increased, the reaction will respond by shifting in whatever way that decreased the temperature again. The inverse is true for if the temperature is decreased.

 

 

 

 

 

 

 

 

 

 

   Via http://faculty.sdmiramar.edu/

Even though the chart provides more information than is needed, we can see the effects of temperature change on a system in dynamic equilibrium depending on the enthalpy of reaction. 

 

Practice Problem 3: The reaction (H2 (g) + I2 (g) <---> 2HI (g)    ΔH = -51.0 kJ) is in dynamic equilibrium and occurs at 100°C. If the reaction is heated an additional 100°C, how will the system respond?

 

A.) The reaction will shift to the left in order to account for the increased temperature 

 

B.) The reaction will shift to the right because there is more energy in the products now

 

C.) The reaction will shift to the left in order to account for the energy added to the products

 

D. The reaction will shift to the right in order to account the increased temperature

 

 

Traps to Avoid

Overall, Le Chatelier's Principle is an easy concept to grasp; however, there are some aspects of it that can easily trick you. First of all, inert gases should not be considered when determining the affect a pressure change has on a system. While they will increase the overall pressure of the system, they do not contribute to the reaction; they are simply present. Secondly, if the moles of gas on each side of a reaction are equal, the reaction will not shift to favor one side when the pressure or volume of the reaction is varied. Finally, catalysts have no effect on the position of equilibrium of a reaction, and therefore, Le Chatelier's Principle does not apply. 

 

Practice Problem #4: The reaction (PCl3 (g) + Cl2 (g) <---> PCl5 (g)) is in dynamic equilibrium. It is then released into a larger, sealed container with Ne (g) and PCL5 (g) so that the reaction becomes

(Ne (g) + PCl3 (g) + Cl2 (g) <--->  PCl5 (g) +Ne(g)) and the concentration of PCl5 (g) is greater than it was previously. How will the system respond in order to re-establish equilibrium? 

 

A.) It is impossible to tell because there are too many variables

 

B.) The reaction will proceed to the left in order to increase pressure and account for the addition of PCl5 (g)

 

C. The reaction will proceed to the right in order to increase pressure and account for the addition of PCl5 (g)

 

D. The reaction will proceed to the left in order to decrease pressure and account for the addition of Ne (g)

 

The Man Behind the Principle

Born October 8, 1850, Henry-Louis Le Chatelier would some day change the world of chemistry. Le Chatelier was priveleged to come from a burgeous Roman Catholic family in Paris, France. As a result, he recieved a good education. He was able to attend the Collège Rollin in Paris. He would go on to earn two undergraduate degrees here in 1867 and 1868 before enrolling in École Polytechnique in 1869. The very next year, Le Chatelier entered the mining engineer program at École des Mines, earning a degree there in 1873. He would go on to marry Geneviève Nicolas with whom  he had seven children with. 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

         Via http://amandalee6.blogspot.com

              Henry-Louis Le Chatelier as a young man

 

Le Chatelier spent two years as a mining engineer before returning to École des Mines in 1877 to teach Chemistry. Initially, his academic endeavours were geared towards way to improve mining conditions, safety, and effectiveness; however, his early work led him down the path of experimental thermodynamics. In 1884, Le Chatelier released a general principle that would define him for the rest of history onto the chemistry world. He would go on to contribute directly to industrial development, become president of the Société d’Encouragement pour l’Industrie Nationale, teach at Sorbonne and École des Mines, and sit as a scientific expert on a variety of governmental committees. Even with his impressive list of accomplishments, "Le Chatelier's Principle" will forever define his life, career, and legacy.