IP Chemistry – Chemical Equilibrium and Le Chatelier’s Principle

First of all, this topic on Chemical Equilibrium is not being tested for GCE O-Level Pure Chemistry (syllabus code 6092) in Singapore. It is only being tested in IP Chemistry, IGCSE Chemistry and IB Chemistry.

At the end of the blog post, you will be able to answer three important questions:

• Q1: What is meant by ‘dynamic equilibrium’?
• Q2: How is equilibrium being maintained?
• Q3: How can experimental conditions be manipulated to obtain the maximum yield?

Just like all the Chemistry tuition classes and holiday revision workshops which I have been personally conducting over the last 16+ years, I would like you to have an overview in what you be learning today.

The key points in this post are:

• A) Reversible Reactions
• B) Dynamic Equilibrium
• C) Le Chatelier’s Principle & Position of Equilibrium
• D) Haber Process (Case Study)

A) Reversible Reactions

Many chemical reactions can proceed in one direction only. i.e. they cannot be reversed and they go towards completion.

e.g. Neutralisation reaction between potassium hydroxide & hydrochloric Acid

KOH(aq) + HCl(aq) → KCl(aq) + H₂O(l)

Some chemical reactions can be reversible .i.e. reactions can go either directions and they reached an equilibrium, instead of going towards completion.

At equilibrium, the forward and backward reactions do not stop; they continue, but at the same speed. Hence, there is no overall change in the amounts of reactants and products.

At the end of reaction, a mixture of reactants and products is present and they are known to be in equilibrium.

e.g. Haber Process:           N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

By altering conditions of temperature, pressure and  use of catalysts, these reactions can be adjusted to favour more reactants or more products.

Whether a reaction is reversible or irreversible depends on activation energy.

If the activation energy of the reverse reaction (i.e. E_b) is exceptionally high, then this reaction will be unfavourable and the reaction is described as irreversible.

B) Dynamic Equilibrium

Let’s consider a reversible reaction whereby:

A + B ⇌ C + D

When the mixture of A and B reacts to become C and D, the concentrations of A and B (reactants) decrease with time, while the concentrations of C and D (products) increases with time.

As the concentrations of reactants decrease, the rate of the forward reaction also decreases with time

At the start, the rate of backward reaction is zero because there is no C and D.

As the reaction proceeds, concentrations of C and D increase. Hence, rate of backward reaction also increases.

After a period of time, Dynamic Equilibrium is reached.

Dynamic Equilibrium refers to a reversible reaction in which the rates of forward and reverse reactions have become equal and there is no change in the concentrations of the reactants and the products.

C) Le Chatelier’s Principle & Position of Equilibrium

Summaries the effect of external factors (changes in temperature, concentration or pressure) on a system at equilibrium.

Le Chatelier’s Principle states that if a change is made to a system in equilibrium, the system reacts in such a way as to tend to oppose the change, and a new equilibrium is formed.

Basically, whatever is done to a system in equilibrium, the system does the opposite to it!

• If something is added to a system at equilibrium, the system will behave as to remove it

• If something is removed from the system, the system will behave so as to put it back

Let’s consider 4 external factors: concentration, pressure, temperature and addition of a catalyst.

C.1 Changes in Concentration

If reactants are added (or products removed) in an equilibrium, a new equilibrium is produced which contains a higher proportion of the products i.e. position of equilibrium shifts to the right hand side.

At the new equilibrium, concentrations of the reactants and products are not the same as those in the previous equilibrium.

C.2 Changes in Pressure

Only applicable for systems containing gas(es) only.

If the pressure of an equilibrium mixture is increased (or volume decreased), the mixture will try to reduce the pressure by reducing the number of moles (or molecules) of gases.

• ↑ Pressure favours reaction which produces  fewer moles of gas

• ↓ Pressure favours reaction which produces  more  moles of gas

C.3 Changes in Temperature

This is linked to the concepts you have learned in the topic of Energy Changes (also known as Chemical Energetics or Thermodynamics).

• If temperature is increased, system tries to decrease it  by ‘absorbing’ extra heat energy i.e. favours Endothermic reactions

• If temperature is decreased, system tries to increase it  by ‘producing’ extra heat energy i.e. favours Exothermic reactions

C.4 Presence of Catalyst

This is linked to the concepts on catalysts which you have learned in the topic of Energy Changes as well as Rate of Reaction.

• When a catalyst is added to an equilibrium, it increases both the forward and reverse reaction rates by the same extent

• The activation energies for both the forward and reverse reactions are lowered by the same extent

• Catalysts shortens the time needed to attain the same final equilibrium concentrations

• It does not affect the position of equilibrium nor the equilibrium composition

• Only increase the rate of reaction so that equilibrium is reached faster

D) Haber Process (Case Study)

N₂ (g) + 3H₂( g) ⇌ 2NH₃(g)                      ∆H = – 92 kJmol^-1

Note that we have covered Manufacturing of Ammonia using Haber Process previously in this blog, which touch on the key concepts and associated keywords necessary for students taking GCE O-Level Pure Chemistry (syllabus code: 6092). We have described how nitrogen gas (from fractional distillation of liquefied air), and hydrogen gas (from cracking of crude oil), as raw materials for producing ammonia in the Haber Process. We have also looked into the essential /optimum conditions for the process. We shall not talk about those here today, so, it’s really good that you refer to that previous blogpost.

Instead, we shall look at how chemical engineers and chemists can use Le Chatelier’s Principle to maximise the yield of ammonia in the Haber Process, by manipulating the conditions in it.

Conditions (Optimum)

1. Temperature: 450 ^oC

• Le Chatelier’s Principle predicts that lower temperature gives a higher yield of NH₃ (Adv)

• If temperature is too low, reaction rate will be too slow (Disadv)

2. Pressure: 200 atm

• Le Chatelier’s Principle predicts that higher pressure gives a higher yield of NH₃ (Adv)

• Too high a pressure is dangerous and involve higher cost of maintaining equipment (Disadv)

3. Catalyst: Finely divided Fe catalyst

• Increased the rate of production of NH₃ but not the yield of NH₃

4. Continual removal of ammonia

• Removal of NH₃ shifts position of equilibrium to the RHS i.e. increasing yield of NH₃

• Done by cooling the reaction mixture to -50^oC to liquefy NH₃  formed

• BP of NH₃ is around -50^oC, while N₂ & H₂ have lower BPs; will remain as gases. Only NH₃ will be liquefied.

I hope you find the content easy for your understanding and if you have any questions, leave me a comment below. Feel free to share this blog post with your friends.

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Related Articles:

• O Level Chemistry – Rate of Reaction Mini Series Part 3
• O Level Chemistry – Rate of Reaction Mini Series Part 2
• O Level Chemistry: Energy Changes – Exothermic & Endothermic Reactions
• O Level Chemistry: Energy Changes (Exo/Endo) & Bond Energy
• O Level Chemistry – Rate of Reaction Mini Series Part 1