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Equilibria and Reversible Reactions

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What are Reversible Reactions?

Some reactions are very straightforward. The reactants are brought together, the reaction starts and the products are formed. That’s where it ends, it’s a one-way street.

We could represent these reactions in this general equation:

The equation for a straightforward chemical reaction showing reactants and products
A non- reversible reaction

Some reactions don’t work like that, they don’t go in one direction only. In these reactions, molecules of products can react in the opposite direction to produce some molecules of the original reactants again. These are reversible reactions.

We can represent reversible reactions like this:

Reversible reaction equation showing reversible symbol
A reversible reaction

Examples of Reversible Reactions

There are a huge number of reversible reactions, so we are going to look at some of the most common example you will come across in your studies.

Please bear in mind that you may be given any example in an exam question, you are not limited to these examples. And be assured that you will be able to cope with any reversible reaction or equilibrium question so long as you learn the basics that we will cover.

Reaction Producing Ammonia

The first example we will consider is the reaction between nitrogen and hydrogen to produce ammonia. You probably recognise this reaction as the Haber process, the industrial process used to manufacture ammonia.

In the Haber process certain conditions are used to make the reaction proceed in a useful way. We will look at those details in another article, and focus on the reversible nature of the reaction here.

Let’s consider what happens when we start the reaction with the reactants: hydrogen and nitrogen, both in the gas state. Under suitable conditions the reaction proceeds with some nitrogen and hydrogen molecules reacting to become ammonia molecules.

As soon as ammonia is produced, some of the ammonia molecules will react in the opposite direction, the reverse reaction. Naturally, the products of the reverse reaction are hydrogen and nitrogen.

The equation for this reversible reaction:

haber process reversible reaction

By the way, you will have noticed that the ammonia produced is gaseous. Yet, you may be familiar with the bottle of “ammonia” on the shelf in your college lab. To help you to understand this apparent conflict there is an article that clarifies the difference between ammonia, ammonium and ammonia solution.

Reaction Producing an Ester

The reaction between an alcohol and a carboxylic acid produces an ester and water. This is a general reaction type, as it is possible to produce different esters by starting with a different alcohol and/or a different carboxylic acid. (These reactions require a catalyst, but we don’t need to consider this here).

The reaction is known as esterification.

The particular example we will look at is the reaction between ethanol and ethanoic acid, which produces ethyl ethanoate and water.

the equation between ethanol and ethyl ethanoate can be shown as this:

Reversible reaction between ethanl and ethyl ethanoate producing ethyl ethanoate and water
The reaction is usually carried out in aqueous conditions

Equilibria vs Reversible Reactions

So these above reactions are all equilibria, right?


It’s a common misconception that equilibria are types of reaction. That a reaction can be an equilibrium. But equilibrium is not a noun, it is a condition or state.

These reversible reactions can be at equilibrium, but aren’t necessarily at equilibrium.

So, exactly what do we mean by equilibrium?

The state of equilibrium is reached when there is no further change to the concentration of reactants or products.

That means that the forward reaction and the backward/reverse reaction are happening at the same rate as each other. This is why we often refer to the state as dynamic equilibrium. Both reactions do continue at equilibrium, but there is no overall change to concentrations.

A prerequisite for equilibrium is that the reaction must be carried out in a closed system. A closed system just means that substances can’t be lost from the reaction mixture, and nothing can be added either.

What Happens if the Equilibrium is Disturbed?

We’ve said that once equilibrium is reached there is no further change to the concentration, no further change in the “equilibrium position”. And that is true if we don’t disturb the conditions.

So, what happens if we do disturb the conditions? And what conditions are we talking about?

Changing Conditions at Equilibrium

The position of equilibrium can be changed if we change the temperature or pressure of a reaction, or change the concentration of any of the reactants or products.

That may sound awkward. Once an equilibrium state is established how do you predict what will happen when a condition is altered?

Luckily for us, there is a principle that was developed by a French chemist Henry Louis Le Chatelier. The principle can be applied when we change any of these conditions, to help us to deduce how the equilibrium position will change. You have almost certainly know the name of Le Chatelier Principle, and now we will get to know and use it.

Le Chatelier’s Principle

The principle states:

If a system at equilibrium is disturbed, the equilibrium moves in the direction that tends to reduce the disturbance.

By system he means the reaction mixture in that closed system we mentioned earlier in this article.

So, if we make a change to the conditions, the equilibrium position moves towards the right hand side (RHS) or the left hand side (LHS) of the equation – whichever direction reduces the disturbance that was made.

Let’s consider each of those conditions, starting with temperature.

Changing the Temperature at Equilibrium

According to Le Chatelier’s principle, if the temperature of a system at equilibrium is changed the position of equilibrium will move in the direction to reduce that change.

What does that mean? What happens if we increase the temperature of a system at equilibrium? Which way does the equilibrium position move?

It’s not as simple as being able to say that the equilibrium moves more to the RHS or the LHS. Because it’s not necessarily the case that the forward reaction is always exothermic (or always endothermic). Why not? Because a reversible reaction is reversible, there is no reason why we couldn’t swap the LHS and RHS of the written equation. It would still make as much sense.

We just tend to write the substances we started with on the LHS, but there’s no reason why we couldn’t have started with the ‘products’. It’s important to realise that an equilibrium can be approached from either direction.

In our earlier example of producing an ester, we started with ethanol and ethanoic acid. We could equally start another experiment of the same reaction using ethyl ethanoate and water – these would initially produce the alcohol and acid and gradually an equilibrium would be established, an equilibrium equivalent to the first experiment. If we did these two experiments under the same conditions we would have two identical mixtures once their equilibria had established.

Anyway, back to raising the temperature of our system at equilibrium. The equilibrium will move to reduce that disturbance, and to reduce an increase in temperature the reaction must move in the endothermic direction, the direction that uses some of that additional heat energy.

It follows that, if the temperature of our system at equilibrium is lowered, then the equilibrium moved more in the direction of the exothermic reaction.

If we raise the temperature of a system at equilibrium, the equilibrium position will move in the endothermic direction.

If we lower the temperature of a system at equilibrium, the equilibrium position will move in the exothermic direction.

Changing the Concentration of a Substance in a System at Equilibrium

What happens if we change the concentration of one of the substances in a reversible reaction at equilibrium? Again we need to consider Le Chatelier’s Principle. The equilibrium will move in the direction that reduces this disturbance.

So, if we increase the concentration of a “reactant”, a substance on the LHS of our equation, the equilibrium position will move to the right. And if we increase the concentration of a “product”, a substance on the RHS the equilibrium position will move towards the LHS.

It’s also possible to reduce the concentration of substance. For example, in our example of producing an ester we could add a little sodium hydroxide. This would neutralise some of the ethanoic acid hence reducing its’ concentration.

If we reduce the concentration of a “reactant” (on the LHS of the equation) the equilibrium will move towards the left. And if we reduce the concentration of a “product” (on the RHS of the equation) the equilibrium will move towards the right.

Changing the Pressure of a Equilibrium

If there are gaseous substances in the system then we may effect the equilibrium position by changing the pressure of the system.

To work out how the equilibrium will move you must count the number of moles of gases on each side of our equation. For example, in our equation for the production of ammonia, there are 4 moles of gas on the LHS (one mole of nitrogen and three moles of hydrogen). And there are two moles of gas on the RHS.

If we increase the pressure of the system, Le Chatelier’s Principle tells us that the position of equilibrium moves to reduce that disturbance. That would mean reducing the number of moles of gas (which reduces the pressure). In this case, that means the equilibrium would move towards the RHS.

If we increase the pressure of the system, the equilibrium position moves towards the side with the fewest moles of gases.

If we decrease the pressure of the system, the equilibrium position moves towards the side with the greater number of moles.

How about if there are the same number of moles of gas on each side? There would be no change to the equilibrium position because neither direction would reduce the disturbance of changing the pressure. The same is true if there are no gases in the reaction mixture.

Quantifying Equilibria

We will look more closely at the equilibrium constant in the next article.

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