Le Chatelier's Principle (2024)

Le Chatelier's Principle

Le Chatelier's Principle Changes in Concentration
Changes in Pressure Changes in Temperature

LeChatelier's Principle

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.

Changes inConcentration

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) Le Chatelier's Principle (2) 2 NH3(g) Kc = 0.040
Initial: 0.100 M 0.100 M 0
Equilibrium: 0.100 -Le Chatelier's Principle (3)C 0.100 - 3 Le Chatelier's Principle (4)C 2 Le Chatelier's Principle (5)C

We obtain the following results when we solve this problem.

[NH3] = 2 Le Chatelier's Principle (6)C Le Chatelier's Principle (7) 0.0020 M
[N2] = 0.100 - Le Chatelier's Principle (8)C Le Chatelier's Principle (9) 0.099 M
[H2] = 0.100 - 3 Le Chatelier's Principle (10)C Le Chatelier's Principle (11) 0.097 M

The fact that Le Chatelier's Principle (12)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 Le Chatelier's Principle (13)by shifting the equilibrium toward theproducts. The reaction comes back to equilibrium when theconcentrations of the three components reach the followingvalues.

[NH3] = 2 Le Chatelier's Principle (14)C Le Chatelier's Principle (15)0.0055 M
[N2] = 1.00 - Le Chatelier's Principle (16)C Le Chatelier's Principle (17)1.00 M
[H2] = 0.10 - 3 Le Chatelier's Principle (18)C Le Chatelier's Principle (19) 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] Le Chatelier's Principle (20) 0.0094 M [NH3] Le Chatelier's Principle (21) 0.026 M
[N2] Le Chatelier's Principle (22) 0.095 M [N2] Le Chatelier's Principle (23) 0.99 M
[H2] Le Chatelier's Principle (24) 0.29 M [H2] Le Chatelier's Principle (25) 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.

Changes inPressure

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) Le Chatelier's Principle (27) 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 Le Chatelier's Principle (28) 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.

Le Chatelier's Principle (29)

N2(g) + 3 H2(g) Le Chatelier's Principle (30) 2 NH3(g)

Changes inTemperature

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) Le Chatelier's Principle (32) N2O4(g) Le Chatelier's Principle (33)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.

Practice Problem 7:

Predict the effect of the following changes on the reaction in which SO3 decomposes to form SO2 and O2.

2 SO3(g) Le Chatelier's Principle (34) 2 SO2 (g) + O2 (g) Le Chatelier's Principle (35)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.

Click here to check your answer to Practice Problem 7

Le Chatelier's Principle (2024)

FAQs

What is Le Chatelier's principle in simple words? ›

Le Châtelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change to reestablish an equilibrium.

What is the Le Chatelier's principle of shifting equilibrium? ›

If the concentration of a substance is increased, the reaction that consumes that substance is favored, and the equilibrium shifts away from that substance. If the concentration of a substance is decreased, the reaction that produces that substance is favored, and the equilibrium shifts toward that substance.

What are the 3 factors that affect the Le Chatelier's principle? ›

Temperature, pressure, and concentration all affect the state of equilibrium in a system.

How do you answer Le Chatelier's principle? ›

Correct answer:

Le Chatelier's principle states that changes in pressure are attributable to changes in volume. If we increase the volume, the reaction will shift toward the side that has more moles of gas. If we decrease the volume, the reaction will shift toward the side that has less moles of gas.

How to remember Le Chatelier's principle? ›

So the simple way to remember this is AA TT. For temperature and pressure, the thing to remember is that the reaction will always try to shift back to equilibrium. For example, if you have an exothermic reaction and you increase the temperature, the reaction needs to cool down to get back to equilibrium.

What is an example of Le Chatelier's principle in real life? ›

Everyday tasks like drying clothes are also examples of Le Chatelier's principle and chemical equilibrium in real life. On a windy day, the water vapors are carried away faster, now to establish an equilibrium, the water from the clothes starts drying, hence drying the clothes faster.

Why is the Le Chatelier's principle important? ›

Le Chatelier's principle is important because it allows us to weigh up input and output in order to find the most profitable combination of reactants and conditions. Without it, many of our industrial processes would be much more inefficient.

What is an example of Le Chatelier's principle concentration? ›

Le Chatelier's Principle Examples

Concentration: In a system where the reaction A + B ⇌ C + D is currently in equilibrium, increasing the concentration of one of the reactants, e.g., 2 A + B ⇌ C + D , will cause a shift in equilibrium to the right.

What is the Chatelier's principle of heat? ›

Using Le Châtelier's principle, heat can be seen as either a reactant or a product depending on if the reaction is endothermic or exothermic respectively. So adding or removing heat will shift the reaction left or right, or speed up either the forward or reverse reactions.

What causes equilibrium to shift to the right? ›

Raising the temperature of the system is akin to increasing the amount of a reactant, and so the equilibrium will shift to the right. Lowering the system temperature will likewise cause the equilibrium to shift left.

What is the Le Chatelier's principle of pH? ›

Based on Le Châtelier's principle, increasing the pH means more hydroxide ions are present, which will shift the equilibrium of the reaction to the left. When the pH is reduced, more hydrogen ions are present, which will shift the equilibrium of the reaction to the right.

References

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