MYP Integrated Sciences Topic Reversible Reactions…

What Is a Reversible Reaction, and How Is It Different from an Irreversible Reaction?

In chemistry, reactions can be:

Irreversible (one way), or
Reversible (two-way).

Irreversible Reactions

An irreversible reaction is a reaction that effectively goes in one direction only:

Reactants → Products

Once the reactants have formed products, they do not easily change back into the original substances.
These reactions often release a lot of energy and produce very stable products.

Example

Combustion of methane

$$\mathrm{CH}_4(\mathrm{~g})+2 \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}$$

Methane burns in oxygen to form carbon dioxide and water.
Under normal conditions, CO₂ and H₂O do not react together to reform methane and oxygen.
The reaction is practically one-way → irreversible.

Burning magnesium

$$2 \mathrm{Mg}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{MgO}(\mathrm{~s})$$

Magnesium burns in air with a bright white flame to form magnesium oxide.
Magnesium oxide is very stable; it does not spontaneously decompose back into Mg and O₂.
The reaction is effectively irreversible in normal conditions.

Reversible Reactions

A reversible reaction can go in both directions:

Reactants ⇌ Products

The products can react again to re-form the original reactants.
Both the forward and reverse reactions can occur in the same system, often at the same time.

Example

Formation of hydrogen iodide

$$\mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{HI}(\mathrm{~g})$$

Forward reaction: H₂ and I₂ combine to form HI.
Reverse reaction: HI decomposes back into H₂ and I₂.

If this reaction is carried out in a closed container, both reactions occur, and over time the system reaches a state where:

H₂ and I₂ are still reacting to form HI.
HI is still decomposing back into H₂ and I₂.

But the concentrations of all three gases become constant. This state is called dynamic equilibrium.

Key Differences

Irreversible reactions
1. Proceed in one direction only.
2. Products are usually very stable and do not easily revert to reactants.
3. Often release a large amount of energy and go to completion.
Reversible reactions
1. Can proceed in both forward and reverse directions.
2. Products can react to reform reactants.
3. Can reach a state of dynamic equilibrium in a closed system, where
  1. the rate of the forward reaction = the rate of the reverse reaction,
  2. and the concentrations of reactants and products remain constant (but not necessarily equal).

Hint

Reversible reactions are written with a double arrow, e.g. ⇌.
Irreversible reactions are written with a single arrow, e.g. →.

How Do We Represent Reversible Reactions with Symbols and Dynamic Models?

Symbolic Representation

A general reversible reaction is written as: $$\mathrm{A}+\mathrm{B} \rightleftharpoons \mathrm{C}+\mathrm{D}$$

Forward reaction: $$\mathrm{A}+\mathrm{B} \to \mathrm{C}+\mathrm{D}$$
Reverse reaction: $$\mathrm{C}+\mathrm{D} \to \mathrm{A}+\mathrm{DB}$$

The double arrow (⇌) tells us the reaction can go both ways.

Dynamic Models – Forward and Reverse Processes

We can represent the dynamic nature of reversible reactions with arrows and rates.

At the start (only reactants present):
1. Forward reaction rate is high (long arrow).
2. Reverse reaction rate is low (short arrow or almost zero), because there are almost no products yet.
As time passes:
1. Reactants are used up → forward rate decreases.
2. Products are formed → reverse rate increases.
At dynamic equilibrium:
1. The rates of the forward and reverse reactions become equal.
2. Arrows can be drawn the same length to show this.
3. The system looks unchanged at the macroscopic level, but particles are still reacting both ways.

Hint

Key point: At equilibrium, rates are equal, but concentrations are constant (not necessarily equal).

Example

The Haber Process (Ammonia Synthesis)

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Forward: Nitrogen + hydrogen → ammonia.
Reverse: Ammonia → nitrogen + hydrogen.
In a closed system, an equilibrium mixture is formed containing N₂, H₂, and NH₃.
At equilibrium:
- Ammonia is being formed and decomposed at the same rate.
- The mixture of gases remains constant in composition.

N₂O₄ / NO₂ Colour Change

$$\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$$

N₂O₄ is colourless.
NO₂ is brown.
At equilibrium in a closed tube:
- Both gases are present.
- The overall colour depends on the relative amounts of N₂O₄ and NO₂.
Changing conditions (like temperature) can shift the balance and make the mixture more brown (more NO₂) or more colourless (more N₂O₄).
This is a nice visual model for equilibrium shifts.

Reversible Reactions in Everyday Systems - Hydrated salts (e.g. cobalt(II) chloride)

Hydrated cobalt(II) chloride: CoCl₂·6H₂O (pink).
Anhydrous cobalt(II) chloride: CoCl₂ (blue).
Dehydration (heating): $$\mathrm{CoCl}_2 \cdot 6 \mathrm{H}_2 \mathrm{O}(\text { pink }) \rightleftharpoons \mathrm{CoCl}_2(\text { blue })+6 \mathrm{H}_2 \mathrm{O}(\mathrm{~g})$$
Heating drives the reaction to the right (removing water).
Adding water drives the reaction back to the left (re-forming the hydrated salt).
These reversible colour changes are used in humidity indicators and desiccant (drying) packs.

How Do Changes in Conditions Affect the Position of Equilibrium?

For a reversible reaction in a closed system, the mixture of reactants and products at equilibrium is called the equilibrium mixture.
The position of equilibrium tells us whether the mixture contains mostly reactants, mostly products, or significant amounts of both.

Le Chatelier’s Principle

Le Chatelier’s Principle states:

If a system at equilibrium is disturbed by a change in conditions, the system will adjust to oppose the changeand establish a new equilibrium.

We’ll focus on:

Concentration
Temperature
(and briefly) Pressure and catalysts

Changing Concentration

For a general reaction:

$$\mathrm{A}+\mathrm{B} \rightleftharpoons \mathrm{C}+\mathrm{D}$$

Increase [A] or [B] (add reactant)
1. The system responds by using up extra reactant.
2. Equilibrium shifts to the right, making more products C and D.
Decrease [C] or [D] (remove product)
1. The system responds by making more product.
2. Again, equilibrium shifts to the right.
Increase [C] or [D] (add product)
1. The system responds by using up extra product.
2. Equilibrium shifts to the left, making more A and B.
Decrease [A] or [B] (remove reactant)
1. The system responds by reforming reactants.
2. Equilibrium shifts to the left.

Example

Haber process

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Increasing [N₂] or [H₂] → shifts equilibrium to the right, producing more NH₃.
Removing NH₃ as it forms also shifts equilibrium to the right, favouring ammonia production.

Changing Temperature

Temperature changes affect which direction is favoured, depending on whether the forward reaction is exothermic or endothermic.
1. Exothermic reaction (releases heat): $$\text { Reactants ⇌ Products + heat }$$
2. Endothermic reaction (absorbs heat): $$\text { Reactants + heat ⇌ Products }$$
Add heat (increase temperature):
1. The system will try to use up the extra heat.
2. Equilibrium shifts in the endothermic direction.
Remove heat (decrease temperature):
1. The system will try to produce heat.
2. Equilibrium shifts in the exothermic direction.

Example

Haber process (exothermic forward reaction)

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})+\text { heat }$$

Increasing temperature → shifts equilibrium to the left, reducing ammonia yield.
Decreasing temperature → shifts equilibrium to the right, increasing ammonia yield (but slows the rate).

This is why industrial conditions are a compromise: moderately high temperature for a reasonable rate, but not so high that the yield is too low.

Pressure (for Gaseous Reactions)

For reactions involving gases, changing the pressure (by changing volume) can affect the equilibrium position.

Le Chatelier’s Principle:

Increasing pressure → equilibrium shifts to the side with fewer gas molecules.
Decreasing pressure → equilibrium shifts to the side with more gas molecules.

Example

Haber process

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Left side: 1 + 3 = 4 moles of gas.
Right side: 2 moles of gas.

Thus:

Increasing pressure → shifts equilibrium to the right (towards fewer moles of gas) → more ammonia.
Decreasing pressure → shifts equilibrium to the left.

Catalysts

A catalyst:

Speeds up both the forward and reverse reactions equally.
Helps the system reach equilibrium faster.
Does not change the position of equilibrium or the final equilibrium composition.

Active recall

In your own words, what is the difference between a reversible and an irreversible reaction?
Write the reversible reaction for the Haber process with the correct symbol.
What does “dynamic equilibrium” mean in this context?
For the reaction $$\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$$ predict what happens to the colour of the gas mixture if:
- more NO₂ is added,
- the mixture is cooled.
In the Haber process, why are high pressure and moderate temperature used in industry?

What Is a Reversible Reaction, and How Is It Different from an Irreversible Reaction?

In chemistry, reactions can be:

Irreversible (one way), or
Reversible (two-way).

Irreversible Reactions

An irreversible reaction is a reaction that effectively goes in one direction only:

Reactants → Products

Once the reactants have formed products, they do not easily change back into the original substances.
These reactions often release a lot of energy and produce very stable products.

Example

Combustion of methane

$$\mathrm{CH}_4(\mathrm{~g})+2 \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})+2 \mathrm{H}_2 \mathrm{O}$$

Methane burns in oxygen to form carbon dioxide and water.
Under normal conditions, CO₂ and H₂O do not react together to reform methane and oxygen.
The reaction is practically one-way → irreversible.

Burning magnesium

$$2 \mathrm{Mg}(\mathrm{~s})+\mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{MgO}(\mathrm{~s})$$

Magnesium burns in air with a bright white flame to form magnesium oxide.
Magnesium oxide is very stable; it does not spontaneously decompose back into Mg and O₂.
The reaction is effectively irreversible in normal conditions.

Reversible Reactions

A reversible reaction can go in both directions:

Reactants ⇌ Products

The products can react again to re-form the original reactants.
Both the forward and reverse reactions can occur in the same system, often at the same time.

Example

Formation of hydrogen iodide

$$\mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{HI}(\mathrm{~g})$$

Forward reaction: H₂ and I₂ combine to form HI.
Reverse reaction: HI decomposes back into H₂ and I₂.

If this reaction is carried out in a closed container, both reactions occur, and over time the system reaches a state where:

H₂ and I₂ are still reacting to form HI.
HI is still decomposing back into H₂ and I₂.

But the concentrations of all three gases become constant. This state is called dynamic equilibrium.

Key Differences

Irreversible reactions
1. Proceed in one direction only.
2. Products are usually very stable and do not easily revert to reactants.
3. Often release a large amount of energy and go to completion.
Reversible reactions
1. Can proceed in both forward and reverse directions.
2. Products can react to reform reactants.
3. Can reach a state of dynamic equilibrium in a closed system, where
  1. the rate of the forward reaction = the rate of the reverse reaction,
  2. and the concentrations of reactants and products remain constant (but not necessarily equal).

Hint

Reversible reactions are written with a double arrow, e.g. ⇌.
Irreversible reactions are written with a single arrow, e.g. →.

How Do We Represent Reversible Reactions with Symbols and Dynamic Models?

Symbolic Representation

A general reversible reaction is written as: $$\mathrm{A}+\mathrm{B} \rightleftharpoons \mathrm{C}+\mathrm{D}$$

Forward reaction: $$\mathrm{A}+\mathrm{B} \to \mathrm{C}+\mathrm{D}$$
Reverse reaction: $$\mathrm{C}+\mathrm{D} \to \mathrm{A}+\mathrm{DB}$$

The double arrow (⇌) tells us the reaction can go both ways.

Dynamic Models – Forward and Reverse Processes

We can represent the dynamic nature of reversible reactions with arrows and rates.

At the start (only reactants present):
1. Forward reaction rate is high (long arrow).
2. Reverse reaction rate is low (short arrow or almost zero), because there are almost no products yet.
As time passes:
1. Reactants are used up → forward rate decreases.
2. Products are formed → reverse rate increases.
At dynamic equilibrium:
1. The rates of the forward and reverse reactions become equal.
2. Arrows can be drawn the same length to show this.
3. The system looks unchanged at the macroscopic level, but particles are still reacting both ways.

Hint

Key point: At equilibrium, rates are equal, but concentrations are constant (not necessarily equal).

Example

The Haber Process (Ammonia Synthesis)

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Forward: Nitrogen + hydrogen → ammonia.
Reverse: Ammonia → nitrogen + hydrogen.
In a closed system, an equilibrium mixture is formed containing N₂, H₂, and NH₃.
At equilibrium:
- Ammonia is being formed and decomposed at the same rate.
- The mixture of gases remains constant in composition.

N₂O₄ / NO₂ Colour Change

$$\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$$

N₂O₄ is colourless.
NO₂ is brown.
At equilibrium in a closed tube:
- Both gases are present.
- The overall colour depends on the relative amounts of N₂O₄ and NO₂.
Changing conditions (like temperature) can shift the balance and make the mixture more brown (more NO₂) or more colourless (more N₂O₄).
This is a nice visual model for equilibrium shifts.

Reversible Reactions in Everyday Systems - Hydrated salts (e.g. cobalt(II) chloride)

Hydrated cobalt(II) chloride: CoCl₂·6H₂O (pink).
Anhydrous cobalt(II) chloride: CoCl₂ (blue).
Dehydration (heating): $$\mathrm{CoCl}_2 \cdot 6 \mathrm{H}_2 \mathrm{O}(\text { pink }) \rightleftharpoons \mathrm{CoCl}_2(\text { blue })+6 \mathrm{H}_2 \mathrm{O}(\mathrm{~g})$$
Heating drives the reaction to the right (removing water).
Adding water drives the reaction back to the left (re-forming the hydrated salt).
These reversible colour changes are used in humidity indicators and desiccant (drying) packs.

How Do Changes in Conditions Affect the Position of Equilibrium?

For a reversible reaction in a closed system, the mixture of reactants and products at equilibrium is called the equilibrium mixture.
The position of equilibrium tells us whether the mixture contains mostly reactants, mostly products, or significant amounts of both.

Le Chatelier’s Principle

Le Chatelier’s Principle states:

If a system at equilibrium is disturbed by a change in conditions, the system will adjust to oppose the changeand establish a new equilibrium.

We’ll focus on:

Concentration
Temperature
(and briefly) Pressure and catalysts

Changing Concentration

For a general reaction:

$$\mathrm{A}+\mathrm{B} \rightleftharpoons \mathrm{C}+\mathrm{D}$$

Increase [A] or [B] (add reactant)
1. The system responds by using up extra reactant.
2. Equilibrium shifts to the right, making more products C and D.
Decrease [C] or [D] (remove product)
1. The system responds by making more product.
2. Again, equilibrium shifts to the right.
Increase [C] or [D] (add product)
1. The system responds by using up extra product.
2. Equilibrium shifts to the left, making more A and B.
Decrease [A] or [B] (remove reactant)
1. The system responds by reforming reactants.
2. Equilibrium shifts to the left.

Example

Haber process

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Increasing [N₂] or [H₂] → shifts equilibrium to the right, producing more NH₃.
Removing NH₃ as it forms also shifts equilibrium to the right, favouring ammonia production.

Changing Temperature

Temperature changes affect which direction is favoured, depending on whether the forward reaction is exothermic or endothermic.
1. Exothermic reaction (releases heat): $$\text { Reactants ⇌ Products + heat }$$
2. Endothermic reaction (absorbs heat): $$\text { Reactants + heat ⇌ Products }$$
Add heat (increase temperature):
1. The system will try to use up the extra heat.
2. Equilibrium shifts in the endothermic direction.
Remove heat (decrease temperature):
1. The system will try to produce heat.
2. Equilibrium shifts in the exothermic direction.

Example

Haber process (exothermic forward reaction)

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})+\text { heat }$$

Increasing temperature → shifts equilibrium to the left, reducing ammonia yield.
Decreasing temperature → shifts equilibrium to the right, increasing ammonia yield (but slows the rate).

This is why industrial conditions are a compromise: moderately high temperature for a reasonable rate, but not so high that the yield is too low.

Pressure (for Gaseous Reactions)

For reactions involving gases, changing the pressure (by changing volume) can affect the equilibrium position.

Le Chatelier’s Principle:

Increasing pressure → equilibrium shifts to the side with fewer gas molecules.
Decreasing pressure → equilibrium shifts to the side with more gas molecules.

Example

Haber process

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Left side: 1 + 3 = 4 moles of gas.
Right side: 2 moles of gas.

Thus:

Increasing pressure → shifts equilibrium to the right (towards fewer moles of gas) → more ammonia.
Decreasing pressure → shifts equilibrium to the left.

Catalysts

A catalyst:

Speeds up both the forward and reverse reactions equally.
Helps the system reach equilibrium faster.
Does not change the position of equilibrium or the final equilibrium composition.

Active recall

In your own words, what is the difference between a reversible and an irreversible reaction?
Write the reversible reaction for the Haber process with the correct symbol.
What does “dynamic equilibrium” mean in this context?
For the reaction $$\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$$ predict what happens to the colour of the gas mixture if:
- more NO₂ is added,
- the mixture is cooled.
In the Haber process, why are high pressure and moderate temperature used in industry?

MYP Integrated Sciences Reversible Reactions Notes

MYP Integrated Sciences Reversible Reactions Notes

What Is a Reversible Reaction, and How Is It Different from an Irreversible Reaction?

Irreversible Reactions

Combustion of methane

Burning magnesium

Reversible Reactions

Formation of hydrogen iodide

Key Differences

How Do We Represent Reversible Reactions with Symbols and Dynamic Models?

Symbolic Representation

Dynamic Models – Forward and Reverse Processes

The Haber Process (Ammonia Synthesis)

N₂O₄ / NO₂ Colour Change

Reversible Reactions in Everyday Systems - Hydrated salts (e.g. cobalt(II) chloride)

How Do Changes in Conditions Affect the Position of Equilibrium?

Le Chatelier’s Principle

Changing Concentration

Changing Temperature

Pressure (for Gaseous Reactions)

Catalysts

What Is a Reversible Reaction, and How Is It Different from an Irreversible Reaction?

Irreversible Reactions

Combustion of methane

Burning magnesium

Reversible Reactions

Formation of hydrogen iodide

Key Differences

How Do We Represent Reversible Reactions with Symbols and Dynamic Models?

Symbolic Representation

Dynamic Models – Forward and Reverse Processes

The Haber Process (Ammonia Synthesis)

N₂O₄ / NO₂ Colour Change

Reversible Reactions in Everyday Systems - Hydrated salts (e.g. cobalt(II) chloride)

How Do Changes in Conditions Affect the Position of Equilibrium?

Le Chatelier’s Principle

Changing Concentration

Changing Temperature

Pressure (for Gaseous Reactions)

Catalysts

MYP Integrated Sciences Reversible Reactions Notes

Science skills

Scale and structure

Classification and organisation

Materials and bonding

Chemical reactions

Energy

Forces and motion

Thermal physics

Electricity and magnetism

Waves

Atoms and space

Genetics and evolution

Body systems

Ecosystems and environment

Human impact and sustainability

Biotechnology and the future

MYP Integrated Sciences Reversible Reactions Notes

Science skills

Scale and structure

Classification and organisation

Materials and bonding

Chemical reactions

Energy

Forces and motion

Thermal physics

Electricity and magnetism

Waves

Atoms and space

Genetics and evolution

Body systems

Ecosystems and environment

Human impact and sustainability

Biotechnology and the future

What Is a Reversible Reaction, and How Is It Different from an Irreversible Reaction?

Irreversible Reactions

Combustion of methane

Burning magnesium

Reversible Reactions

Formation of hydrogen iodide

Key Differences

How Do We Represent Reversible Reactions with Symbols and Dynamic Models?

Symbolic Representation

Dynamic Models – Forward and Reverse Processes

The Haber Process (Ammonia Synthesis)

N₂O₄ / NO₂ Colour Change

Reversible Reactions in Everyday Systems - Hydrated salts (e.g. cobalt(II) chloride)

How Do Changes in Conditions Affect the Position of Equilibrium?

Le Chatelier’s Principle

Changing Concentration

Changing Temperature

Pressure (for Gaseous Reactions)

Catalysts

What Is a Reversible Reaction, and How Is It Different from an Irreversible Reaction?

Irreversible Reactions

Combustion of methane

Burning magnesium

Reversible Reactions

Formation of hydrogen iodide

Key Differences

How Do We Represent Reversible Reactions with Symbols and Dynamic Models?

Symbolic Representation

Dynamic Models – Forward and Reverse Processes

The Haber Process (Ammonia Synthesis)

N₂O₄ / NO₂ Colour Change

Reversible Reactions in Everyday Systems - Hydrated salts (e.g. cobalt(II) chloride)

How Do Changes in Conditions Affect the Position of Equilibrium?

Le Chatelier’s Principle

Changing Concentration

Changing Temperature

Pressure (for Gaseous Reactions)

Catalysts