Le Chatelier's Principle
Dr. Davinder Bhachu
Teacher
Contents
What is Dynamic Equilibrium?
Some chemical reactions go to completion – meaning that nearly all the reactants are used up to form products. These reactions are shown with a single arrow (→) in the chemical equation, indicating that the reaction is one-way.
Examples of Irreversible Reactions
Combustion:
Neutralisation:
Precipitation:
In these cases, the reactions "go to completion" – there's no meaningful reverse reaction.
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Dynamic Equilibrium
Many chemical reactions don’t go all the way to completion. Instead, they reach a state of equilibrium, where both reactants and products are present at the same time. At this point, no overall change in the amounts of substances occurs, even though the reaction hasn’t stopped.
A familiar example is the reaction between ethanoic acid and ethanol, which produces an ester (ethyl ethanoate) and water. This reaction is reversible:
Once the system reaches equilibrium, all four substances are present in the mixture – the original acid and alcohol, and the ester and water formed.
On the molecular scale, particles are still reacting – ethanoic acid and ethanol continue combining to make ester and water, while at the same time, ester and water are breaking back down into acid and alcohol. Because these two processes happen at equal rates, the concentrations of each substance remain unchanged overall.
From a macroscopic point of view, nothing appears to change – but in reality both reactions are still occurring. This is why we call it dynamic equilibrium. It’s shown in chemical equations using the symbol , which represents the continuous but balanced movement in both directions.
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Position of Equilibrium
In any reversible reaction, the position of equilibrium describes the relative amounts of reactants and products when equilibrium is reached. Depending on how the reaction is set up – including the initial concentrations, temperature, and pressure – the final mixture can vary.
If the system is disturbed by changing one of these conditions, it will shift to restore equilibrium. This results in a new position of equilibrium, where the concentrations of substances have adjusted accordingly.
If more products are made, we say the equilibrium has shifted to the right (towards the products, which are usually written on the right-hand side of the equation).
If more reactants are formed instead, the equilibrium has shifted to the left (towards the left-hand side of the equation).
The position of equilibrium can be influenced by changes in:
the concentration of substances involved (if they’re in solution),
the pressure (only relevant if gases are involved),
the temperature of the system.
In the late 1800s, a scientist called Henri Le Chatelier studied how different reactions responded to changes in conditions. From this, he developed a set of rules to predict how equilibrium systems respond to changes. These ideas are known today as Le Chatelier’s Principle, which can be summarised as:
If an external condition is changed, the equilibrium will shift in the direction that helps oppose that change.
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Applying Le Chatelier’s Principle
Changing Concentration
Orange-brown bromine gas () reacts with colourless hydrogen gas () to form colourless hydrogen bromide (HBr(g)) in a reversible reaction:
This is a great example to show how changing the concentration of reactants affects equilibrium. If you increase the amount of either or the system responds by forming more . According to Le Chatelier’s Principle, the equilibrium shifts to the right, reducing the concentration of the added species.
If you remove from the system – for example, by letting it react in a side reaction – the system compensates by shifting the equilibrium to the left, producing more and reducing the amount of .
In industry, reactions are often designed to shift equilibrium towards the products. One common strategy is to remove a product as it forms, so the system keeps producing more.
A good example of this is ester formation, such as:
(ethanoic acid + ethanol ⇌ ethyl ethanoate + water)
Here, if you remove the ester (for example, by distilling it off as it forms), the equilibrium keeps shifting to the right, encouraging more ester to be produced. This technique is often used in laboratories to improve yield in reversible reactions.
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Changing the Pressure
Pressure can affect the position of equilibrium for reactions involving gases. According to Le Chatelier’s Principle, if you increase the pressure of a system, the equilibrium will shift in the direction that produces fewer gas molecules – this helps reduce the pressure and oppose the change.
Increasing pressure: Shifts equilibrium towards the side with fewer gas molecules (to reduce pressure).
Decreasing pressure: Shifts equilibrium to the side with more gas molecules (to increase pressure).
No effect if the number of gas molecules is the same on both sides.
Example:
Example
What is the effect of increasing pressure on the yield of ethanol in the following equilibrium?
Answer:
On the left there are 2 moles of gas, on the right only 1 mole.
Increasing pressure causes equilibrium to shift right.
This results in a higher yield of ethanol.
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Tip
High pressures give faster rates and better yields (if fewer gas moles on the product side), but they are costly and dangerous to maintain.
Hence, a compromise pressure is often used in industry.
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Changing the Temperature
For every reversible reaction, the forward and backward reactions have opposite energy changes – if one is exothermic, the other is endothermic but the same magnitude. When the temperature of a system is increased, the equilibrium shifts in the direction of the endothermic reaction to absorb the added heat and minimise the disturbance.
If temperature increases:
The system shifts to oppose the change, favouring the endothermic direction (absorbing heat).
If temperature decreases:
The system shifts in the exothermic direction (releasing heat) to raise the temperature.
Example
What happens to the yield of sulfur trioxide if temperature is increased in this reaction?
Answer:
The forward reaction is exothermic.
Increasing temperature shifts equilibrium in the endothermic (backwards) direction to oppose the increase.
Therefore, yield of SO₃ decreases as equilibrium shifts left.
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Tip:
Low temperatures favour high yield but slow the rate – so industrial processes use a compromise temperature.
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Effect of Catalysts on Equilibrium
Catalysts increase the rate at which equilibrium is reached by speeding up both the forward and backward reactions equally.
Key Point:
Catalysts do not affect the position of equilibrium – they simply allow equilibrium to be reached faster.
Example 1
Contact Process – Making Sulfuric Acid
Stage 1:
Stage 2:
Conditions: ~450°C, ~2 atm, catalyst =
Low temp = better yield but slower reaction
High pressure gives marginally better yield but isn't economical.
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Example 2
Making Methanol
Conditions: ~400°C, ~50 atm, catalyst = Cr/Zn oxides
Low temp = good yield, slow rate → compromise temp
High pressure = better yield but expensive equipment needed
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Example 3
Hydration of Propene to Make Propanol
Conditions: ~300°C, ~70 atm, catalyst = conc
Low temp favours yield, but reaction is slow → compromise
High pressure = faster rate but costly and can cause polymerisation
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Summary Points
Catalysts increase rate, allow for lower temperatures and reduced energy costs, but do not shift the equilibrium.
High pressure boosts rate and yield in some cases, but adds cost due to expensive pumps and strong equipment needed to withstand high pressure.
Recycling unreacted reactants helps improve overall yield in all processes.
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