Le Chatelier's Principle
| Le Chatelier's Principle | Changes in Concentration |
| Changes in Pressure | Changes in Temperature |
In 1884 the French chemist and engineer Henry-Louis LeChatelier proposed one of the central concepts of chemicalequilibria. Le Chatelier's principle can bestated as follows: A change in one of the variables thatdescribe a system at equilibrium produces a shift in the positionof the equilibrium that counteracts the effect of this change.
Le Chatelier's principle describes what happens to a systemwhen something momentarily takes it away from equilibrium. Thissection focuses on three ways in which we can change theconditions of a chemical reaction at equilibrium:
(1) changing the concentration of one of the components of thereaction
(2) changing the pressure on the system
(3) changing the temperature at which the reaction is run.
To illustrate what happens when we change the concentration ofone of the reactants or products of a reaction at equilibrium,let's consider the following system at 500oC.
| N2(g) | + | 3 H2(g) |  | 2 NH3(g) | Kc = 0.040 | |||
| Initial: | 0.100 M | 0.100 M | 0 | |||||
| Equilibrium: | 0.100 - C | 0.100 - 3 C | 2 C |
We obtain the following results when we solve this problem.
[NH3] = 2 C  0.0020 M |
| [N2] = 0.100 - C 0.099 M |
| [H2] = 0.100 - 3 C 0.097 M |
The fact that C is small compared with the initialconcentrations of N2 and H2 makes thiscalculation relatively easy to do. But it implies that verylittle ammonia is actually produced in the reaction. According tothis calculation, only 1% of the nitrogen present initially isconverted into ammonia.
What would happen if we add enough N2 to increasethe initial concentration by a factor of 10? The reaction can'tbe at equilibrium any more because there is far too much N2in the system. Adding an excess of one of the reactants thereforeplaces a stress on the system. The system responds by minimizingthe effect of this stress 
by shifting the equilibrium toward theproducts. The reaction comes back to equilibrium when theconcentrations of the three components reach the followingvalues.
| [NH3] = 2 C 0.0055 M |
| [N2] = 1.00 - C 1.00 M |
| [H2] = 0.10 - 3 C 0.092 M |
By comparing the new equilibrium concentrations with thoseobtained before excess N2 was added to the system, wecan see the magnitude of the effect of adding the excess N2.
| Before | After | |
| [NH3] 0.0094 M | [NH3] 0.026 M | |
| [N2] 0.095 M | [N2] 0.99 M | |
| [H2] 0.29 M | [H2] 0.26 M |
Increasing the amount of N2 in the system by afactor of 10 leads to an increase in the amount of NH3at equilibrium by a factor of about 3. Adding an excess of one ofthe products would have the opposite effect; it would shift theequilibrium toward the reactants.
The effect of changing the pressure on a gas-phase reactiondepends on the stoichiometry of the reaction. We can demonstratethis by looking at the result of compressing the followingreaction at equilibrium.
| N2(g) | + | 3 H2(g) | | 2 NH3(g) |
Let's start with a system that initially contains 2.5 atm of N2and 7.5 atm of H2 at 500oC, where Kpis 1.4 x 10-5, allow the reaction to come toequilibrium, and then compress the system by a factor of 10. Whenthis is done, we get the following results.
| Before Compression | After Compression | |
| PNH3 = 0.12 atm | PNH3 = 8.4 atm | |
| PN2 = 2.4 atm | PN2 = 21 atm | |
| PH2 = 7.3 atm | PH2 = 62 atm |
Before the system was compressed, the partial pressure of NH3was only about 1% of the total pressure. After the system iscompressed, the partial pressure of NH3 is almost 10%of the total.
These data provide another example of Le Chatelier'sprinciple. A reaction at equilibrium was subjected to a stress an increasein the total pressure on the system. The reaction then shifted inthe direction that minimized the effect of this stress. Thereaction shifted toward the products because this reduces thenumber of particles in the gas, thereby decreasing the totalpressure on the system, as shown in the figure below.
| N2(g) | + | 3 H2(g) | | 2 NH3(g) |
Changes in the concentrations of the reactants or products ofa reaction shift the position of the equilibrium, but do notchange the equilibrium constant for the reaction.
Similarly, a change in the pressure on a gas-phase reactionshifts the position of the equilibrium without changing themagnitude of the equilibrium constant. Changes in the temperatureof the system, however, affect the position of the equilibrium bychanging the magnitude of the equilibrium constant for thereaction.
Chemical reactions either give off heat to their surroundingsor absorb heat from their surroundings. If we consider heat to beone of the reactants or products of a reaction, we can understandthe effect of changes in temperature on the equilibrium.Increasing the temperature of a reaction that gives off heat isthe same as adding more of one of the products of the reaction.It places a stress on the reaction, which must be alleviated byconverting some of the products back to reactants.
The reaction in which NO2dimerizes to form N2O4 provides an exampleof the effect of changes in temperature on the equilibriumconstant for a reaction. This reaction is exothermic.
| 2 NO2(g) | | N2O4(g) | Ho = -57.20 kJ |
Thus, raising the temperature of this system is equivalent toadding excess product to the system. The equilibrium constanttherefore decreases with increasing temperature.
Predict the effect of the following changes on the reaction in which SO3 decomposes to form SO2 and O2.
2 SO3(g) 
2 SO2 (g) + O2 (g) 
Ho = 197.78 kJ
(a) Increasing the temperature of the reaction.
(b) Increasing the pressure on the reaction.
(c) Adding more O2 when the reaction is at equilibrium.
(d) Removing O2 from the system when the reaction is at equilibrium.
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