What is Le Chatelier’s Principle and Chemical equilibrium?
Le Chatelier’s principle and chemical equilibrium is a concept that is usually discussed in Chemistry.
Some chemical reactions are reversible and can attain equilibrium. Chemical equilibrium is a state of a system (i.e chemical reaction) when the rate of forward reaction equals the rate of backward reaction.
Please recall that a reversible reaction is a reaction that occurs in both directions i.e has both the forward and backward reaction.
The French chemist Henry-Louis Le Chatelier studied the reactions in chemical equilibrium and he discovered that equilibrium can be altered by external constraints.
There is always an upset in chemical equilibrium by altering any of these three conditions; temperature, pressure, and concentration.
Le Chatelier’s principle states that if an external constraint such as a change in temperature, pressure, or concentration is imposed on a system in equilibrium, the equilibrium will shift so as to annul or neutralize the constraint.
Le Chatelier
Actually, what Le Chatelier is saying is that there is a way a system or chemical reaction in equilibrium will behave.
The system will either shift to the right or left when an external constraint is imposed on the system.
Where the system will shift is what we will look at right now.
Recall the Le Chateliers principle mentioned three external constraints that can alter equilibrium position.
Effect of Temperature
Temperature generally increases the rate of chemical reactions according to collision theory. However rate of reaction is not the same as a chemical equilibrium. In chemical equilibrium, we look at which rate is favoured though both rates will be affected but one will be affected more..
The general rule is that increase in temperature will favour endothermic reactions but a decrease in temperature favors exothermic reactions.
The simple reason is that endothermic reactions absorb heat therefore increasing temperature will excite the particles while exothermic reactions give out heat, therefore, decreasing temperature will suit the reaction.
Illustration of effect on temperature on the equilibrium position
- Consider the reaction below;
CaCO3 –> CaO + CO2
If the forward reaction is endothermic, what is the effect of increasing temperature?
Solution
Hint: if the forward reaction is endothermic, then the backward reaction will definitely be exothermic.
Increasing temperature will favour forward reaction shifting equilibrium position to the right.
Decreasing temperature will favour backward reaction shifting equilibrium position to the left.
- Consider the reaction below;
2SO2 + O2 à 2SO3
If the forward reaction is exothermic, what is the effect of temperature on equilibrium position?
Solution
Hint: If the forward reaction is exothermic, then the backward reaction is endothermic.
Increasing temperature will favour backward reaction shifting equilibrium position to the left
Decreasing temperature will favour forward reaction shifting equilibrium position to right.
Effect of Concentration
Basically concentration can be discussed in two ways;
Take a look at this equation
2HCl + CaCO3 CaCl2 + H2O + CO2
For solids, concentration simply means adding more of the solid solute e.g increasing the mass of CaCO3 from 20g to 50g.
For liquids or aqueous solutions, concentration means increasing concentration from 2.0 mol/dm3 to 3.5 mol/dm3.
In general, if you increase the concentration of one side, the equilibrium position will shift to the other side you did not increase.
Also if you reduce or decrease the concentration of one side, the equilibrium will still shift to the side you decreased.
Example 1
CaCO3 à CaO + CO2
What’s the effect of adding more CaCO3 on equilibrium position?
Solution
Adding more CaCO3 will shift the equilibrium position to the right favouring the forward reaction.
Example 2
Consider the ionic equation below;
Fe3+ + SCN– –> [Fe(SCN)]2+
What is the effect of increasing Fe3+ and removing [Fe(SCN)]2+ on equilibrium position?
Solution
Adding more of Fe3+ will shift the equilibrium position to the right favouring forward reaction.
Removing [Fe(SCN)]2+ as it is formed will shift the equilibrium position to the right favouring forward reaction.
Example 3
Zn + 2HCl à ZnCl2 + HCl
What is the effect of increasing concentration from 3.0 mol/dm3 to 5.5 mol/dm3?
Solution
The equilibrium position will shift to the right favouring the forward reaction.
Effect of Pressure
If the pressure should affect a system, then there should be gas on either side of the equation.
If there is a gaseous product and a gaseous reactant, then the number of moles of the gaseous reactant and product should not be the same.
The principle is that pressure can only affect the reaction when the volume of the gaseous reactant and product vary.
An increase in pressure will favour the side with a lower volume (lower number of moles)
A decrease in pressure will favour the side with a higher volume (higher number of moles)
Example 1
Consider the equation below;
2SO2 + O2 à 2SO3
What is the effect of increasing pressure?
Solution
First, analyze the number of moles of reactants and products.
2 = 1 =3 vs 2
An increase in pressure will usually favour the side with lower number of moles.
So increase in pressure will favour the forward reaction shifting the equilibrium position to the right.
Consequently, a decrease in pressure will favour the backward reaction shifting the equilibrium position to the left.
Example 2
N2 + 3H2 à 2NH3
What is the effect of decreasing the pressure of the system?
Solution
1 + 3 = 4 vs 2
A decrease in pressure will favour the backward reaction shifting the equilibrium position to the left.
What is Equilibrium Constant?
.Equilibrium constant is equal to the ratio of the concentration of products to reactants
For example the equilibrium constant of the reaction below
PCl5 à PCl3 + Cl2
K = [PCl3][Cl2]/[PCl5]
I have a detailed post on the Equilibrium constant and its calculations.
In conclusion, chemical equilibrium in reactions can be predicted and determined by understanding the concept of the Le Chatelier principle.