Select your language

Suggested languages for you:
Log In Start studying!
StudySmarter - The all-in-one study app.
4.8 • +11k Ratings
More than 3 Million Downloads
Free
|
|

All-in-one learning app

  • Flashcards
  • NotesNotes
  • ExplanationsExplanations
  • Study Planner
  • Textbook solutions
Start studying

Application of Le Chatelier's Principle

Save Save
Print Print
Edit Edit
Sign up to use all features for free. Sign up now
Application of Le Chatelier's Principle

If you are a business working in industry, then yield is of the utmost importance. Ideally, you want to maximise your yield whilst minimising your costs in order to achieve the greatest profits. Following Le Chatelier's Principle, we know that by changing the conditions of an equilibrium, we can favour one reaction or the other.

  • This article is about industrial applications of Le Chatelier's principle.
  • We'll begin by exploring why applying Le Chatelier's principle is important. This will include defining compromise conditions.
  • We'll then look at four specific examples of how Le Chatelier's principle is applied to real-life reactions.

The importance of applications of Le Chatelier's principle

Industrial processes, such as fertiliser production or making sulphuric acid (H₂SO₄), are all about profit. This comes down to yield versus the cost of input. Le Chatelier's principle is useful because it allows us to increase the yields of certain products formed in reversible reactions and thus increase profit.

However, yield isn’t the only consideration when it comes to running a chemical reaction. For example, a lower temperature might increase the yield of the desired product but slow the rate of reaction down too much to be economically useful. Or the opposite might be true - a higher pressure or temperature could increase the yield. However, it would be costly to build, run and maintain production plants that can cope with these extremes. When it comes to choosing reaction conditions, price must be factored in as well. This is why industrial equilibrium reactions often use compromise conditions.

Compromise conditions are conditions that don’t necessarily give the greatest yield of the product, but are the most economical when it comes to balancing factors like cost and rate of reaction.

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. In this article, we're going to explore some real-life examples of applications of Le Chatelier's principle in industry.

Applications of Le Chatelier's Principle in the chemical industry

We’ll now look at four different examples of compounds formed in industrial reactions that involve Le Chatelier's principle:

  • Methanol
  • Ethanol
  • Sulphuric acid. This reaction is known as the Contact process.
  • Ammonia. This reaction is known as the Haber process.

Here's how you make them.

Le Chatelier's principle and methanol production

Methanol is made by reacting synthesis gas, which is a mixture of carbon monoxide and hydrogen, with a copper catalyst. It has the following equation:

CO(g) + 2H2(g) ⇌ CH3OH(g) ΔH = -91 kJ mol-1

You should now be able to predict the effect of certain conditions on the yield of methanol:

  • The forward reaction is exothermic. This means a lower temperature shifts the equilibrium to the right and increases the yield of methanol. However, a temperature that is too low slows down the rate of reaction and so a compromise temperature of 500 K is used.
  • The forward reaction produces fewer moles of gas. This means that increasing the pressure shifts the equilibrium to the right and increases the yield of methanol. However, maintaining a high pressure is expensive, and so a compromise pressure of 10,000 kPa is used.

33 million tonnes of methanol are produced every year. Most of it is used to make methanal, an aldehyde further transformed into many types of plastics. However, methanol is also seeing a surge in popularity as a fuel. It can be used in typical diesel and petrol cars with little modification to their existing engines and is even being tested in boats.

Next, we’ll take a look at making another alcohol, ethanol.

Le Chatelier's principle and ethanol production

Ethanol is made in two different ways:

  • Fermentation
  • Hydration of ethene

Out of the two methods, the hydration of ethene is a reversible reaction, and so we'll focus our attention on it here.

You’ll compare fermentation and the hydration of ethene in more detail in the article “Production of Ethanol”.

Hydrating ethene uses a phosphoric acid catalyst. It has the following equation:

C2H4(g) + H2O(g) ⇌ C2H5OH(g) ΔH = -46 kJ mol-1

From the equation, we can infer the following:

  • The forward reaction is exothermic. This means a lower temperature shifts the equilibrium to the right and increases the yield of ethanol. However, a temperature that is too low slows down the rate of reaction and so a compromise temperature of 570 K is used.
  • The forward reaction produces fewer moles of gas. This means that increasing the pressure shifts the equilibrium to the right and increases the yield of ethanol. However, maintaining a high pressure is expensive, and so a compromise pressure of 6,500 kPa is used.
  • Adding excess steam shifts the equilibrium to the right and increases the yield of ethanol. However, too much steam dilutes the catalyst and so slows down the rate of reaction. Instead, the ethanol is removed as it is formed, decreasing its concentration and therefore favouring the forward reaction, and the ethene and steam are repeatedly recycled over the catalyst.

As well as being one of the main components of alcoholic drinks, ethanol also plays an important role as an antimicrobial agent. It destroys microorganisms by disrupting their lipid bilayer membrane and denaturing their proteins.

Le Chatelier's principle and the Contact process

Another example of Le Chatelier's principle is the industrial production of sulphuric acid. This process is called the Contact process and takes place in a number of steps. First, sulphur dioxide is transformed into sulphur trioxide. This uses a vanadium (V) oxide catalyst and is a reversible reaction:

2SO2(g) + O2(g) ⇌ 2SO3(g) ΔH = -196 kJ mol-1

The sulphur trioxide is then turned into sulphuric acid. We first dissolve it in a small amount of sulphuric acid, and then react the resulting solution with water:

H2SO4(l) + SO3(g) → H2S2O7(l)

H2S2O7(l) + H2O(l) → 2H2SO4(l)

As with the reversible reactions we've looked at so far, we can change the conditions of the first reaction in order to increase the yield of sulphur trioxide. This in turn increases the yield of sulphuric acid.

  • The forward reaction is exothermic. This means a lower temperature shifts the equilibrium to the right and increases the yield of sulphur trioxide. However, a temperature that is too low slows down the rate of reaction and so a compromise temperature of 670 K is used.
  • The forward reaction produces fewer moles of gas. This means that increasing the pressure shifts the equilibrium to the right and increases the yield of sulphur trioxide. However, maintaining a high pressure is expensive, and so a compromise pressure of 200 kPa is used (which is 2 atm).
  • Adding excess oxygen shifts the equilibrium to the right and increases the yield of sulphur trioxide. However, adding too much oxygen in turn lowers the concentration of sulphur dioxide, which decreases the rate of reaction - the oxygen has got nothing to react with! Instead, a 1:1 ratio of sulphur dioxide to oxygen has been shown to produce the highest yield.

Once sulphur trioxide has been produced, it can then be converted into sulphuric acid.

Most of the sulphuric acid produced in industry is used in fertilisers. However, it is also used for the production of detergents, resins, pigments and pharmaceuticals.

Finally, let’s have a look at the production of ammonia.

Le Chatelier's principle and the Haber process

Ammonia is made in a reversible reaction called the Haber process, using an iron catalyst:

N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ mol-1

We can say the following:

  • The forward reaction is exothermic. This means a lower temperature shifts the equilibrium to the right and increases the yield of ammonia. However, a temperature that is too low slows down the rate of reaction and so a compromise temperature of 670 K is used.
  • The forward reaction produces fewer moles of gas. This means that increasing the pressure shifts the equilibrium to the right and increases the yield of ammonia. However, maintaining a high pressure is expensive, and so a compromise pressure of 20,000 kPa is used.

As in the production of ethanol, the product is removed and the unreacted nitrogen and hydrogen gases are recycled back over the catalyst. This increases the yield.

Over 80% of ammonia manufactured industrially each year is used to make fertilisers. The rest forms products like plastics, dyes and even explosives!

Comparing applications of Le Chatelier's principle

Here’s a handy table to help you compare the conditions needed for methanol, ethanol, sulphuric acid and ammonia production. When it comes to sulphuric acid production in the Contact process, we've only included the reversible part of the reaction in order to keep things simple.

Product

Equation

Temperature (K)

Pressure (kPa)

Catalyst

Methanol

CO(g) + 2H2(g) ⇌ CH3OH(g)

500

10,000

Copper

Ethanol

C2H4(g) + H2O(g) ⇌ C2H5OH(g)

570

6,500

Phosphoric acid

Sulphuric acid

2SO2(g) + O2(g) ⇌ 2SO3(g)

670

200

Vandadium(V) oxide

Ammonia

N2(g) + 3H2(g) ⇌ 2NH3(g)

670

20,000

Iron

Applications of Le Chatelier's Principle - Key takeaways

  • Le Chatelier's principle tells us that the position of an equilibrium in a reversible reaction is affected by factors such as concentration, pressure and temperature. We can use this to increase the yield of a reaction.
  • Many industrial processes use Le Chatelier’s principle to help increase yield but often use compromise conditions. These balance high yield with cost and rate of reaction.
  • Methanol, ethanol, sulphuric acid and ammonia production are all examples of reversible reactions in industry that use Le Chatelier's principle.

Frequently Asked Questions about Application of Le Chatelier's Principle

We can use Le Chatelier’s principle to increase the profits and yields of many industrial reversible reactions by looking at the effect of changing conditions on the position of equilibrium. For example, the reaction’s equation might tell you that increasing the pressure increases the equilibrium yield. We can therefore apply this to the reaction in industry in order to maximise profit.

Le Chatelier’s principle allows us to change the conditions of an equilibrium in order to shift its position. It is important in industry because it helps to increase yield and maximise profit.


By using Le Chatelier’s principle, we can predict how changing the conditions of an equilibrium reaction affects the position of the equilibrium and influences yield. For example, the Haber process is used to make ammonia and involves a reversible reaction. Le Chatelier’s principle tells us that the useful forward reaction is favoured by a higher pressure, and so this is taken into consideration when considering the reaction conditions in industry.

Examples of Le Chatelier’s principle include synthesising methanol and ammonia. They both involve reversible reactions, and Le Chatelier’s principle helps us find the best conditions to balance cost and yield.

Le Chatelier’s principle is used in real life to increase the profitability and yield of reversible reactions. One example of this is the Haber process, used to make ammonia. Le Chatelier’s principle tells us that increasing the pressure increases the yield of ammonia, and so this is taken into consideration when choosing reaction conditions.

Final Application of Le Chatelier's Principle Quiz

Question

What are compromise conditions?

Show answer

Answer

Compromise conditions are conditions that don’t necessarily give the greatest yield of the product, but are the most economical when it comes to balancing factors like cost and rate of reaction.

Show question

Question

Name the reactants used to make methanol.

Show answer

Answer

Synthesis gas (a mixture of carbon monoxide and hydrogen)

Show question

Question

Give the reactants used to make ethanol via hydration.

Show answer

Answer

Ethene (C2H4) and steam (H2O).

Show question

Question

Give the reactants used to make sulphuric acid in the Contact process.

Show answer

Answer

Sulphur dioxide (SO2), oxygen (O2) and a small amount of sulphuric acid (H2SO4).

Show question

Question

Give the reactants used to make ammonia in the Haber process.

Show answer

Answer

Nitrogen (N2) and hydrogen (H2).

Show question

Question

Many industrial reversible reactions are exothermic. Explain how compromise temperatures are used to increase the yield. 

Show answer

Answer

The forward reaction is exothermic, so decreasing the temperature favours the forward reaction. However, decreasing the temperature too much decreases the overall rate of reaction, and so a compromise temperature that is somewhere in the middle is used instead.

Show question

Question

Many industrial reversible reactions are favoured by a high pressure. Explain how compromise temperatures are used in these cases.

Show answer

Answer

Increasing the pressure favours the forward reaction and so increases the yield. However, maintaining a high pressure is expensive and so a compromise pressure that is somewhere in the middle is used instead.

Show question

Question

Why are catalysts used in industrial reversible reactions?

Show answer

Answer

They increase the rate of reaction and so increase yield.

Show question

More about Application of Le Chatelier's Principle
60%

of the users don't pass the Application of Le Chatelier's Principle quiz! Will you pass the quiz?

Start Quiz

Discover the right content for your subjects

No need to cheat if you have everything you need to succeed! Packed into one app!

Study Plan

Be perfectly prepared on time with an individual plan.

Quizzes

Test your knowledge with gamified quizzes.

Flashcards

Create and find flashcards in record time.

Notes

Create beautiful notes faster than ever before.

Study Sets

Have all your study materials in one place.

Documents

Upload unlimited documents and save them online.

Study Analytics

Identify your study strength and weaknesses.

Weekly Goals

Set individual study goals and earn points reaching them.

Smart Reminders

Stop procrastinating with our study reminders.

Rewards

Earn points, unlock badges and level up while studying.

Magic Marker

Create flashcards in notes completely automatically.

Smart Formatting

Create the most beautiful study materials using our templates.

Just Signed up?

Yes
No, I'll do it now

Sign up to highlight and take notes. It’s 100% free.