What Is a Reversible Reaction?

Many chemical reactions appear to travel in only one direction.

Reactants combine or change to form products:

reactants → products

For example, when a fuel burns in oxygen, new substances are produced. Under ordinary conditions, those products do not simply react to recreate the original fuel and oxygen.

However, some chemical reactions can occur in both directions.

These are called reversible reactions.

In a reversible reaction:

• reactants form products
• the products can react to form the original reactants again

We can represent this as:

reactants ⇌ products

The special double arrow ⇌ shows that the reaction can proceed in both directions.

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Forward and Reverse Reactions

A reversible reaction actually involves two reactions.

Consider:

A + B ⇌ C + D

The forward reaction is:

A + B → C + D

The reactants A and B form products C and D.

The reverse reaction is:

C + D → A + B

The products of the forward reaction react to recreate the original substances.

Therefore, the words reactant and product depend partly on which direction we are considering.

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The Reversible Reaction Symbol

Reversible reactions are represented using:

⇌

rather than the ordinary reaction arrow:

→

For example:

A + B → C

means that the reaction is being represented as proceeding from A and B toward C.

By contrast:

A + B ⇌ C

means that:

A + B can form C

and:

C can form A + B

The double arrow is therefore an important part of a reversible chemical equation.

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Reversible vs Irreversible Reactions

An irreversible reaction is represented as proceeding essentially in one direction under the conditions being considered.

reactants → products

A reversible reaction can proceed in both directions.

reactants ⇌ products

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Why Can Some Reactions Reverse?

For a reaction to reverse, the products must be capable of reacting to recreate the original reactants.

Imagine:

A + B ⇌ C + D

Initially, there may be mostly A and B.

Their particles collide and react:

A + B → C + D

As C and D accumulate, there are more opportunities for C and D particles to collide.

Some of these collisions can cause the reverse reaction:

C + D → A + B

Eventually, both directions may be occurring at the same time.

This leads to the important idea of dynamic equilibrium, which is closely connected to reversible reactions.

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A Simple Particle Model

Imagine a closed container containing particles A and B.

At first, there are many A and B particles, so the forward reaction occurs frequently:

A + B → C

As C is produced, its concentration increases.

Now C particles are available to undergo the reverse reaction:

C → A + B

Over time:

• the forward reaction continues
• the reverse reaction begins
• both reactions can occur simultaneously

This does not mean that individual particles constantly move backward and forward in a predictable pattern.

Instead, some particle collisions produce products while other collisions produce reactants.

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Example: Hydrated Copper(II) Sulfate

A useful classroom example involves hydrated copper(II) sulfate.

Hydrated copper(II) sulfate is blue.

When heated, it loses water and forms white anhydrous copper(II) sulfate.

This can be represented as:

hydrated copper(II) sulfate ⇌ anhydrous copper(II) sulfate + water

Heating favours the removal of water.

Adding water to the anhydrous copper(II) sulfate can recreate the blue hydrated substance.

This provides visible evidence that the chemical change can proceed in opposite directions under suitable conditions.

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Forward Direction: Removing Water

When hydrated copper(II) sulfate is heated:

hydrated copper(II) sulfate → anhydrous copper(II) sulfate + water

The colour changes approximately:

blue → white

Water is removed from the hydrated crystals.

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Reverse Direction: Adding Water

If water is added to anhydrous copper(II) sulfate:

anhydrous copper(II) sulfate + water → hydrated copper(II) sulfate

The colour changes:

white → blue

Together, the two reactions can be represented using the reversible reaction symbol.

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Example: Ammonium Chloride

Another example involves ammonium chloride.

When heated strongly, ammonium chloride vapour can dissociate into ammonia and hydrogen chloride:

NH₄Cl(g) ⇌ NH₃(g) + HCl(g)

Under suitable conditions, ammonia and hydrogen chloride can combine again to form ammonium chloride.

This demonstrates that decomposition reactions are not necessarily irreversible.

Whether a reaction is reversible depends on the chemistry of the substances and the conditions.

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Example: Formation of Hydrogen Iodide

Hydrogen gas and iodine vapour can react to form hydrogen iodide:

H₂(g) + I₂(g) ⇌ 2HI(g)

The reverse reaction can also occur:

2HI(g) ⇌ H₂(g) + I₂(g)

In a closed system under suitable conditions, both reactions can occur simultaneously.

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Example: The Haber Process

One of the most important industrial reversible reactions is the production of ammonia.

Nitrogen reacts with hydrogen:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

The forward reaction produces ammonia.

The reverse reaction breaks ammonia back into nitrogen and hydrogen.

This reversible reaction is extremely important because ammonia is used to manufacture many nitrogen-containing products, especially fertilizers.

Industrial chemists carefully control reaction conditions to obtain a useful rate and yield.

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Reversible Reactions and Closed Systems

Reversible reactions are particularly important in closed systems.

A closed system allows energy to be transferred but prevents the reacting substances from freely entering or leaving.

Imagine:

A + B ⇌ C + D

If C and D remain inside the container, they can react to recreate A and B.

But suppose C escapes continuously.

There may then be much less C available for the reverse reaction.

Keeping all reacting substances together allows the forward and reverse reactions to continue and makes it possible for dynamic equilibrium to become established.

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Reversible Reactions and Dynamic Equilibrium

A reversible reaction in a closed system may eventually reach dynamic equilibrium.

At dynamic equilibrium:

rate of forward reaction = rate of reverse reaction

This does not mean the reactions have stopped.

Both reactions continue.

Imagine that every second:

• 50 forward reactions occur
• 50 reverse reactions occur

There is no overall change in the amounts of reactants and products because they are being produced and consumed at equal rates.

The system looks unchanged at the macroscopic level, even though reactions are still occurring at the particle level.

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Equilibrium Does Not Mean Equal Amounts

A very important point is that equilibrium does not necessarily mean:

amount of reactants = amount of products

Suppose:

A ⇌ B

At equilibrium, there might be:

• 80% A
• 20% B

Or there might be:

• 30% A
• 70% B

Both could represent equilibrium.

What must be equal is the rate of the forward reaction and the rate of the reverse reaction, not the concentrations of the substances.

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How Equilibrium Develops

Consider a reversible reaction that begins with only reactants:

A + B ⇌ C + D

At the beginning

There are many reactant particles and almost no products.

Therefore:

• forward reaction is relatively fast
• reverse reaction is very slow or initially absent

As the reaction continues

Reactants are consumed and products accumulate.

Therefore:

• forward reaction tends to slow
• reverse reaction becomes faster

At equilibrium

Eventually:

forward rate = reverse rate

The concentrations then remain constant as long as the conditions remain unchanged.

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Reversible Reactions Can Respond to Conditions

Changing the conditions of a reversible reaction can change the relative amounts of reactants and products.

Important conditions include:

• temperature
• pressure for reactions involving gases
• concentration

For example, changing the temperature may favour one direction more than the other.

Changing the concentration of one substance can also affect the balance between the forward and reverse reactions.

This behaviour is extremely important in industrial chemistry because manufacturers often want to maximize the production of a particular product.

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Everyday Example: Carbon Dioxide in Fizzy Drinks

Carbon dioxide dissolved in a carbonated drink provides a useful example of processes that can occur in opposite directions.

In a sealed bottle, carbon dioxide exists both in the gas above the liquid and dissolved in the liquid.

A simplified representation is:

CO₂(g) ⇌ CO₂(aq)

where:

• (g) means gas
• (aq) means dissolved in water

When the bottle is opened, carbon dioxide can escape from the system. This changes the conditions and dissolved carbon dioxide leaves the liquid, producing bubbles.

This is why fizzy drinks gradually become flat after being left open.

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Reversible Physical Changes

The idea of processes occurring in opposite directions is also familiar from physical changes.

For example:

liquid water ⇌ water vapour

In a closed container:

• some liquid molecules evaporate
• some gas molecules condense

This is a physical equilibrium, rather than a chemical reaction, because no new chemical substances are formed.

However, it provides a useful model for understanding how two opposing processes can occur simultaneously.

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Reversible and Irreversible Examples

Some familiar processes can be compared.

Often treated as irreversible under ordinary conditions

• combustion of fuels
• many reactions that produce gases which escape
• many precipitation reactions when considered under fixed classroom conditions
• many reactions in which products are continually removed

Reversible under suitable conditions

• hydration and dehydration of some salts
• formation and decomposition of hydrogen iodide
• ammonia formation in the Haber process
• some dissociation reactions

It is important to include the phrase under suitable conditions. Whether a reaction behaves reversibly can depend strongly on the conditions.

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Worked Example 1: Identifying the Reaction Type

Consider:

A + B ⇌ C

Is this reaction reversible?

Step 1: Examine the arrow

The equation contains:

⇌

Step 2: Interpret the symbol

This means:

A + B → C

can occur, and:

C → A + B

can also occur.

Answer

Yes. The equation represents a reversible reaction.

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Worked Example 2: Writing Forward and Reverse Reactions

Consider:

N₂ + 3H₂ ⇌ 2NH₃

Forward reaction

N₂ + 3H₂ → 2NH₃

Nitrogen and hydrogen form ammonia.

Reverse reaction

2NH₃ → N₂ + 3H₂

Ammonia forms nitrogen and hydrogen.

The double arrow combines these two processes into one reversible equation.

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Worked Example 3: Interpreting Equilibrium

Suppose the following reaction reaches equilibrium:

A ⇌ B

Measurements show:

[A] = 0.80 mol/dm³

[B] = 0.20 mol/dm³

A student says:

"The system cannot be at equilibrium because there is more A than B."

This reasoning is incorrect.

Equilibrium does not require equal concentrations.

The system can be at equilibrium if:

rate of A → B = rate of B → A

The amounts of A and B can be different.

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How to Recognize a Reversible Reaction

When examining a chemical equation, look first at the reaction arrow.

Single arrow

→

Usually indicates that the reaction is being represented as proceeding in one direction.

Double reversible arrow

⇌

Indicates that both forward and reverse reactions can occur.

Then ask:

• Can the products react to recreate the reactants?
• Are the substances kept together in a closed system?
• Do the conditions allow both reactions to occur?
• Could the system establish dynamic equilibrium?

These questions help explain the behaviour rather than simply identifying the symbol.

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Common Mistakes

Mistake 1: Using → for a reversible reaction

If both directions are important, use:

⇌

rather than:

→

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Mistake 2: Thinking the reverse reaction is a different, unrelated reaction

The reverse reaction uses the products of the forward reaction to recreate the original reactants.

For:

A + B ⇌ C + D

Forward:

A + B → C + D

Reverse:

C + D → A + B

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Mistake 3: Thinking equilibrium means the reaction stops

At dynamic equilibrium, reactions continue in both directions.

The system appears unchanged because the rates are equal.

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Mistake 4: Thinking equilibrium means equal concentrations

At equilibrium:

forward rate = reverse rate

It does not necessarily mean:

reactant concentration = product concentration

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Mistake 5: Thinking all chemical reactions are reversible

Many reactions do not readily reverse under the same conditions.

For example, the products of combustion do not normally reform the original fuel simply by remaining together.

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Mistake 6: Ignoring the importance of a closed system

If substances continuously escape from a reaction mixture, the conditions required for dynamic equilibrium may not be maintained.

Closed systems are therefore important when studying reversible reactions and equilibrium.

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Key Terms

Reversible reaction — A chemical reaction that can proceed in both the forward and reverse directions.

Irreversible reaction — A reaction treated as proceeding essentially in one direction under the conditions considered.

Forward reaction — The reaction in which the substances written on the left side of a reversible equation form those on the right.

Reverse reaction — The reaction in which the products of the forward reaction recreate the original reactants.

Dynamic equilibrium — A state in a closed system in which the forward and reverse reactions continue at equal rates.

Closed system — A system in which matter does not freely enter or leave.

Reaction rate — How quickly reactants are converted into products.

Equilibrium — A state in which opposing processes occur at equal rates, producing no overall change in macroscopic composition.

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Check Your Understanding

1. Define a reversible reaction.

2. What symbol is used to represent a reversible chemical reaction?

3. Explain the difference between:

→

and:

⇌

4. For the reaction:

A + B ⇌ C + D

write the forward reaction and the reverse reaction separately.

5. Explain why products must remain available for a reverse reaction to occur.

6. Is the following equation reversible or irreversible as written?

X + Y → Z

Explain your answer.

7. Consider:

H₂ + I₂ ⇌ 2HI

What substances react in the forward direction?

8. What substances react in the reverse direction?

9. Explain why a closed system is important when establishing dynamic equilibrium.

10. At equilibrium, a system contains 70% reactants and 30% products. Can the system still be at equilibrium? Explain.

11. A student says, "At equilibrium, the forward and reverse reactions have stopped." Explain why this statement is incorrect.

12. Hydrated copper(II) sulfate can form anhydrous copper(II) sulfate and water. Explain how this example demonstrates reversibility.

13. Explain what happens to the number of forward and reverse reactions per second when dynamic equilibrium is reached.

14. Why might opening a closed container affect a reversible system involving gases?

15. Explain the difference between a reversible reaction and an irreversible reaction.

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Key Takeaways

• A reversible reaction can proceed in both directions.
• The products of the forward reaction can react to recreate the original reactants.
• Reversible reactions are represented using the symbol ⇌.
• The forward reaction proceeds from reactants toward products.
• The reverse reaction proceeds from products toward reactants.
• Not all reactions are readily reversible under the same conditions.
• Reversibility depends on the substances involved and the reaction conditions.
• Reversible reactions are especially important in closed systems.
• A reversible reaction can eventually establish dynamic equilibrium.
• At dynamic equilibrium, the forward and reverse reactions continue.
• The rates of the forward and reverse reactions are equal at equilibrium.
• Equilibrium does not mean that the concentrations of reactants and products are equal.
• Reversible reactions are important in laboratory chemistry, environmental systems, and industrial processes such as ammonia production.

Did You Know?

The arrows used in chemical equations carry important scientific information. A single arrow → describes a reaction in the direction shown, while the equilibrium arrow ⇌ tells a chemist that both forward and reverse reactions are important. That tiny change in notation can completely change how we analyze a chemical system.