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:
$$ CO(g)+2H_2(g)\rightleftharpoons CH_3OH(g)\qquad \Delta H^\circ = -91 kJ mol^{-1} $$
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:
$$ C_2H_4(g)+H_2O(g) \rightleftharpoons C_2H_5OH(g)\quad \Delta H= -46kJ\space mol^{-1} $$
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:
$$ 2SO_2(g)+O_2(g)\rightleftharpoons 2SO_3(g)\quad \Delta H= \space -196\space kJ\space mol^{-1} $$
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:
$$ H_2SO_4(l)+SO_3(g)\rightarrow H_2S_2O_7(l) $$
$$ H_2S_2O_7(l)+H_2O(l)\rightarrow 2H_2SO_4(l) $$
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:
$$ N_2(g)+3H_2(g)\rightleftharpoons 2NH_3(g)\quad \Delta H=\space -92\space kJ\space mol^{-1} $$
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 | \(CO(g)+2H_2(g)\rightleftharpoons CH_3OH(g)\) | 500 | 10,000 | Copper |
Ethanol | \(C_2H_4(g)+H_2O(g)\rightleftharpoons C_2H_5OH(g)\) | 570 | 6,500 | Phosphoric acid |
Sulphuric acid | \(2SO_2(g)+O_2(g)\rightleftharpoons 2SO_3(g)\) | 670 | 200 | Vandadium(V) oxide |
Ammonia | \(N_2(g)+3H_2(g)\rightleftharpoons 2NH_3(g)\) | 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.
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