1. Lesson # 3 Le Chatelier’s Principle
Chemical Equilibrium
Lesson # 3
Le Chatelier’s Principle

2. Definitions
Le Châtelier’s Principle: when a chemical system at equilibrium is disturbed by a change in a property, the system adjusts in a way that opposes the change.
This principle allows chemists to predict the qualitative effects of changes in concentration, pressure, and temperature on a chemical reaction system in equilibrium.
The goal behind the use of this principle is often to maximize the yield of products in a equilibrium system, which for a long time was solely based on trial and error.

3. 1. Changing Concentration
Changing the concentration of reactants and products which forces a change of equilibrium is called an equilibrium shift.
One way to shift the equilibrium is to add more reactants to the reaction vessel. If we increase the concentration of the reactants, then Le Châtelier’s principle says that equilibrium will move towards the products to oppose this change.
This makes sense as adding more reactants (increasing concentration of reactants means more collisions) should mean more creation of products.

4. Changing Concentration

5. Changing Concentration
We can also apply Le Châtelier’s principle if we removed products as they were formed.
If we decrease the concentration of the products, then the equilibrium will shift to the right towards the products to counteract the lowered concentration (the concentration of reactants will be high compared to products, so moves to right).

6. Concentration & Engineering
This concentration change effect is often used by chemical engineers when designing industrial processes based on reversible reactions.
Often productions involve the continuous addition of reactants or removal of products, which prevents the chemical system from ever reaching equilibrium, and so the formation of products will always be favoured.

7. Formation of Nitric Acid
Nitric acid is commonly made from nitrogen dioxide and water, and is used in the synthesis of fertilizers, explosives, dyes, and perfumes.
3 NO2 (g) + H2O (l) = 2 HNO3 (aq) + NO (g).
Nitric oxide is not particularly useful, so it is often removed from the system to react with oxygen to produced more nitrogen dioxide.
The removal of NO forces the equilibrium shift towards the products and so more nitric acid is formed along with NO.

8. Haemoglobin & Oxygen
In the body, haemoglobin in your blood binds to dissolved oxygen.
Hb (aq) + O2 (aq) = HbO2 (aq).
Oxygen is absorbed by the lungs where the concentration is high, so the equilibrium shifts to the right, toward the Hb-O complex.
The Hb-O complex is pumped into your body cells where the oxygen concentration is low. In other words, the concentration of this reactant is low in our body cells.
To compensate, the equilibrium shifts to the left, in favour of the reactants. As a result, the oxygen is released and is available for use by your body’s cells.

9. 2. Changing Energy
The equilibrium will shift when the temperature is increased or decreased.
You must determine whether the reaction is exothermic or endothermic to discover the shift.

10. Endothermic Reactions
reactants + energy = products
If we cool an endothermic reaction (remove energy), it is like one of our reactants has been removed. The equilibrium will shift left towards the reactants to counteract the change, and energy will be released.
If we add heat an endothermic reaction, then it is like our concentration of reactants has increased, and the equilibrium will shift to the right and absorb energy to create products.

11. Exothermic Reactions reactants = products + energy
If we remove thermal energy from an exothermic reaction, the equilibrium will shift towards the products to counteract the change, and energy will be released.
If we add energy to an exothermic reaction, then the equilibrium will shift to the left to counteract the change, and energy will be used to convert products to reactants.
In summary, increasing temperature will always shift towards the direction that makes the reaction endothermic (heat absorbing), and decreasing temperature will always shift towards the exothermic direction (heat releasing)

12. 3. Changing Gas Volume
Changing the volume of a container also changes the concentration of gases in the container.
We must make the assumption that our gas behaves like an ideal gas to apply the gas laws we learned in grade 11
An ideal gas is composed of particles that have no size, travel in straight lines, and have attraction to each other (no intermolecular forces) and obeys all gas laws.

13. 3. Changing Gas Volume
We know from Boyle’s law that as the pressure of a container increases, the volume decreases, as long as temperature remains constant.
We also know that when a container has a mixture of gases, each gas exerts its own partial pressure (which is the same pressure that it would exert if it was the only gas in the container), and the sum of all these is the total pressure.
When the volume of a mixture decreases, the concentration (number of particles in a given volume) increases proportionally to the increase of their partial pressures.

14. 3. Changing Gas Volume For example:N2 (g) + 3 H2 (g) = 2 NH3 (g)
When we reduce the volume by one half, the total pressure of the system will double, and the number of molecules of each gas in a given volume will also double.
If you look at the balanced equation, there are 4 particles on the reactants side, and 2 on the products side.
This means that the change in concentration of the reactants is greater than the change for the products; Therefore the equilibrium will shift to the right to reduce the total number of particles per unit volume in the container.
In general, the equilibrium will always shift towards the side that has fewer particles when increasing the concentration.

15. Changing an Equilibrium System without Affecting Equilibrium Position
1. Catalysts
We know that catalysts provide an alternative path for a chemical reaction that has a lower activation energy.
In a reversible reaction, catalysts increase the reaction rates of both forward and reverse reactions equally, since both reactants and products can form by the lower-energy path.
Therefore, adding a catalyst does not change the equilibrium position, and the final equilibrium concentrations of reactants and products are not altered. The reaction will simply proceed faster.

16. Changing an Equilibrium System without Affecting Equilibrium Position
2. Inert Gas
Inert gases are not reactive, and so will not enter into a chemical reaction.
If you add an inert gas to an equilibrium system and the volume is kept constant, the total number of particles in the space increases therefore the pressure increases.
Since collisions with inert gases will not result in a chemical reaction, inert molecules only force collisions to happen around them, which means they will happen less frequently.
The combination of increasing total pressure and slowing down collision possibilities at the same time cause no overall shift in the equilibrium.

17. Changing an Equilibrium System without Affecting Equilibrium Position
3. State of Reactants
If there is more than one state of matter in a chemical reaction, equilibrium is affected only by changes in concentration of particles that are in the same state as the ones involved in the chemical reaction system.
For example: H2 (g) + I2 (s) = 2 HI (g)
Adding solid iodine to the reaction would change nothing else in the system except itself, so no equilibrium shift would occur.
Adding more of hydrogen or hydrogen iodide would, however, alter the equilibrium.

18. Changes to Keq
Concentration and changes in pressure or volume do not alter Keq as the ratio of products to reactants is the same, but temperature does alter Keq as it is based on energy.
Endothermic reactions where there is an increase in temperature causes Keq to increase
Exothermic reactions where there is an increase in temperature decreases Keq
VIDEO (4:15 onward)