What Is Activation Energy?

The activation energy, written as Eₐ, is the minimum energy barrier that reacting particles must overcome for a chemical reaction to occur.

Even when a reaction releases energy overall, the reactants usually need an initial input of energy before they can begin changing into products.

A useful way to think about activation energy is as an energy hill.

The reacting particles must have enough energy to get over the hill before products can form.

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Why Do Reactions Need Activation Energy?

During a chemical reaction, atoms must rearrange.

This often requires:

• existing bonds to stretch
• some bonds to break
• atoms to move into new positions
• new bonds to form

Breaking or weakening existing bonds requires energy.

Therefore, reacting particles cannot usually move directly from reactants to products without first reaching a higher-energy state.

This creates an activation-energy barrier.

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A Simple Example

Consider:

H₂ + Cl₂ → 2HCl

The reaction forms hydrogen chloride.

However, the H–H and Cl–Cl bonds do not simply rearrange without an initial energy input.

The reactant molecules must first collide with enough energy to begin breaking or weakening their existing bonds.

Once that barrier is overcome, new H–Cl bonds can form.

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Collision Theory

Activation energy is closely connected to collision theory.

For particles to react successfully:

1. they must collide
2. they must collide with sufficient energy
3. they must collide with a suitable orientation

A collision that leads to a reaction is called a:

successful collision

or:

effective collision

A collision with energy below the activation energy will normally not produce products.

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Successful and Unsuccessful Collisions

Imagine two reactant particles collide.

Collision A

Collision energy:

40 kJ mol⁻¹

Activation energy:

80 kJ mol⁻¹

Because:

40 < 80

the particles do not have enough energy.

The collision is unsuccessful.

Collision B

Collision energy:

100 kJ mol⁻¹

Activation energy:

80 kJ mol⁻¹

Because:

100 > 80

the particles have enough energy to overcome the energy barrier.

If their orientation is also suitable, the collision may produce a reaction.

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Activation Energy on a Reaction Profile

A reaction profile shows how the energy of the reacting system changes as the reaction proceeds.

The activation energy is the vertical energy difference between:

the reactants

and:

the highest point of the reaction pathway

The highest point represents the transition state.

Therefore:

Eₐ = energy of transition state − energy of reactants

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The Transition State

The transition state is a temporary, unstable, high-energy arrangement of atoms.

At this point:

• old bonds may be partly broken
• new bonds may be partly formed
• atoms are arranged differently from both reactants and products

The transition state exists for an extremely short time.

On a reaction profile, it is located at the:

top of the energy peak

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How to Identify Activation Energy

When looking at a reaction profile:

Step 1

Find the energy level of the reactants.

Step 2

Find the highest point on the curve.

Step 3

Measure the vertical energy difference.

That difference is the activation energy.

Do not measure activation energy from zero.

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Example 1

Suppose:

Reactants = 70 kJ mol⁻¹

Peak = 190 kJ mol⁻¹

Then:

Eₐ = 190 − 70

Eₐ = 120 kJ mol⁻¹

The activation energy is:

120 kJ mol⁻¹

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Exothermic Reactions Still Need Activation Energy

An exothermic reaction releases energy overall.

Its products are lower in energy than its reactants.

However, this does not mean that no activation energy is required.

The reaction may still need an initial input of energy to overcome the energy barrier.

For example, fuels such as methane can release large amounts of energy when burned.

Yet methane does not normally burst into flame at room temperature.

An ignition source may be needed to provide enough energy to begin the reaction.

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Combustion and Activation Energy

Consider combustion.

A fuel may be mixed with oxygen, but little or no visible reaction occurs until:

• a spark is provided
• a flame is introduced
• sufficient heating occurs

The spark does not provide all the energy released by the reaction.

It provides enough energy for some particles to overcome the initial activation barrier.

Once the reaction starts, released energy can help activate additional particles.

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Endothermic Reactions Also Have Activation Energy

Endothermic reactions absorb energy overall.

Their products are higher in energy than their reactants.

They also require activation energy.

The important difference is:

activation energy describes the barrier to starting the reaction.

ΔH describes the overall energy difference between products and reactants.

These are different quantities.

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Activation Energy vs Enthalpy Change

This distinction is extremely important.

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Example 2

A reaction has:

Reactants = 100 kJ mol⁻¹

Peak = 260 kJ mol⁻¹

Products = 40 kJ mol⁻¹

Activation energy:

Eₐ = 260 − 100

Eₐ = 160 kJ mol⁻¹

Enthalpy change:

ΔH = 40 − 100

ΔH = −60 kJ mol⁻¹

Therefore:

• activation energy = 160 kJ mol⁻¹
• reaction is exothermic
• energy released overall = 60 kJ mol⁻¹

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Comparing Reactions with Different Activation Energies

Suppose two reactions occur under similar conditions.

Reaction A

Activation energy:

40 kJ mol⁻¹

Reaction B

Activation energy:

120 kJ mol⁻¹

Reaction B has the larger energy barrier.

At the same temperature, fewer particles are likely to have enough energy to overcome the barrier in Reaction B.

Therefore Reaction B would generally be expected to proceed more slowly, assuming other factors are comparable.

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Low Activation Energy

A reaction with a low activation energy has a relatively small energy barrier.

This means that a larger proportion of collisions may have enough energy to react.

Therefore:

lower Eₐ → more successful collisions → faster reaction

under otherwise similar conditions.

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High Activation Energy

A reaction with a high activation energy has a larger energy barrier.

Fewer particles will have sufficient energy to react at the same temperature.

Therefore:

higher Eₐ → fewer successful collisions → slower reaction

under otherwise similar conditions.

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Activation Energy and Reaction Rate

Reaction rate depends on how many successful collisions occur each second.

A reaction with lower activation energy generally has more particles capable of reacting.

A reaction with higher activation energy has fewer particles above the required energy threshold.

So activation energy strongly influences reaction rate.

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Particle Energy Distribution

Not all particles in a sample have the same kinetic energy.

Some move slowly.

Some move faster.

At any temperature, there is a distribution of particle energies.

Only particles with energy equal to or greater than the activation energy have enough energy to overcome the energy barrier.

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The Activation Energy Threshold

On an energy distribution graph, the activation energy can be shown as a vertical line.

Particles to the right of this line have:

E ≥ Eₐ

These particles have enough energy for a potentially successful collision.

Particles to the left have:

E < Eₐ

These collisions generally do not have enough energy to react.

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Temperature and Activation Energy

Increasing temperature does not usually change the activation energy of a reaction.

Instead, increasing temperature changes the energy distribution of the particles.

At higher temperature:

• particles have greater average kinetic energy
• collisions occur more energetically
• a larger fraction of particles has E ≥ Eₐ
• more collisions are successful

Therefore the reaction becomes faster.

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Important Distinction

A common mistake is to say:

Heating lowers activation energy.

This is usually incorrect.

Heating normally:

increases particle energy

while:

Eₐ remains essentially the same

A catalyst, on the other hand, can reduce the activation energy.

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Effect of a Catalyst

A catalyst increases reaction rate by providing an alternative reaction pathway with a lower activation energy.

The catalyst does not supply the reaction with extra energy.

Instead, it creates a route with a smaller energy barrier.

Therefore:

catalyst → lower Eₐ → more successful collisions → faster reaction

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Catalyzed and Uncatalyzed Profiles

A reaction profile may show two pathways.

Uncatalyzed Pathway

Higher peak

→ larger activation energy

Catalyzed Pathway

Lower peak

→ smaller activation energy

The reactants and products remain at the same energy levels.

Therefore, the catalyst does not change the overall enthalpy change, ΔH.

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Example 3: Catalyst

Without catalyst:

Reactants = 60 kJ mol⁻¹

Peak = 180 kJ mol⁻¹

Therefore:

Eₐ = 120 kJ mol⁻¹

With catalyst:

Peak = 110 kJ mol⁻¹

Therefore:

Eₐ = 50 kJ mol⁻¹

The catalyst lowers the activation energy by:

120 − 50 = 70 kJ mol⁻¹

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Why Does Lowering Eₐ Increase Rate?

At a particular temperature, particles have many different kinetic energies.

Suppose only 5% have enough energy to overcome the original activation energy.

After adding a catalyst, the activation energy is lower.

Perhaps now 30% of the particles have enough energy.

There are therefore far more successful collisions each second.

The reaction rate increases.

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Catalyst and Particle Distribution

A catalyst does not need to change the particle energy distribution.

Instead, it changes the position of the activation-energy threshold.

The threshold moves to a lower energy.

This means a larger fraction of the existing particles already has enough energy to react.

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Reaction Rate Is Not the Same as Energy Release

A reaction that releases a large amount of energy is not necessarily fast.

For example, Reaction A could have:

ΔH = −300 kJ mol⁻¹

but:

Eₐ = 200 kJ mol⁻¹

Reaction B might have:

ΔH = −30 kJ mol⁻¹

but:

Eₐ = 40 kJ mol⁻¹

Reaction A releases much more energy overall, but Reaction B may proceed faster because its activation-energy barrier is much lower.

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Thermodynamics vs Kinetics

This leads to an important distinction.

Thermodynamics asks:

• Is the reaction energetically favourable?
• How much energy is released or absorbed?

Kinetics asks:

• How fast does the reaction happen?
• What is the activation energy?
• How frequently do successful collisions occur?

Activation energy is mainly a kinetic concept.

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Why Some Reactions Appear Not to Happen

Some reactions are energetically possible but happen extremely slowly.

This may be because their activation energy is very high.

A large activation barrier prevents most particles from reaching the transition state under normal conditions.

This explains why substances can sometimes remain unchanged for long periods even though lower-energy products are possible.

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Diamond and Graphite

Diamond provides a useful example.

Under ordinary conditions, graphite is thermodynamically more stable than diamond.

Yet diamonds do not rapidly turn into graphite at room temperature.

Why?

The transformation has a very large activation-energy barrier.

Therefore, the reaction rate is extremely slow.

This shows why knowing only the overall energy change is not enough to predict how quickly a process occurs.

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Biological Reactions

Living cells perform thousands of chemical reactions.

Many would occur far too slowly without catalysts.

Biological catalysts are called:

enzymes

Enzymes lower activation energies for specific reactions.

This allows reactions to occur rapidly at the relatively mild temperatures found inside living organisms.

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Enzymes and Activation Energy

The basic idea is:

uncatalyzed biological reaction

→ high activation energy

→ slow

With an enzyme:

alternative pathway

→ lower activation energy

→ faster reaction

The enzyme itself is not consumed overall.

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Reverse Activation Energy

Reaction profiles can also be read in reverse.

For a forward reaction:

Eₐ(forward) = peak − reactants

For the reverse reaction:

Eₐ(reverse) = peak − products

These values are often different.

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Example 4: Reverse Reaction

Reactants:

120 kJ mol⁻¹

Peak:

250 kJ mol⁻¹

Products:

70 kJ mol⁻¹

Forward activation energy:

250 − 120 = 130 kJ mol⁻¹

Reverse activation energy:

250 − 70 = 180 kJ mol⁻¹

Therefore:

Eₐ(reverse) > Eₐ(forward)

in this example.

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Activation Energy and Exothermic Reactions

For an exothermic forward reaction:

products are lower than reactants

Therefore, the reverse reaction usually has a larger activation energy than the forward reaction when both travel through the same transition state.

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Activation Energy and Endothermic Reactions

For an endothermic forward reaction:

products are higher than reactants

Therefore, the reverse reaction may have a lower activation energy than the forward reaction when they share the same transition state.

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Multistep Reactions

Some reactions happen in several stages.

Their energy profiles contain:

• several peaks
• one or more intermediates

Each step has its own activation energy.

A simplified sequence might be:

Reactants

↓

Transition State 1

↓

Intermediate

↓

Transition State 2

↓

Products

The slowest step is often associated with the largest relevant activation barrier and may control the overall reaction rate.

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Activation Energy in Different Reactions

Different reactions can have very different activation energies because:

• different bonds must be broken
• different particles are involved
• different molecular arrangements are required
• different reaction pathways exist

There is therefore no single activation energy that applies to all chemical reactions.

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Units of Activation Energy

Activation energy is usually expressed in:

kJ mol⁻¹

For example:

Eₐ = 75 kJ mol⁻¹

This means that the energy barrier corresponds to 75 kilojoules for one mole of reacting events.

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Worked Example 1

A reaction profile shows:

Reactants = 50 kJ mol⁻¹

Peak = 170 kJ mol⁻¹

Products = 90 kJ mol⁻¹

Calculate the activation energy.

Eₐ = 170 − 50

Eₐ = 120 kJ mol⁻¹

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Worked Example 2

Reaction A has:

Eₐ = 45 kJ mol⁻¹

Reaction B has:

Eₐ = 105 kJ mol⁻¹

Both reactions are performed at the same temperature and under similar conditions.

Which reaction would generally be expected to occur faster?

Reaction A

because its activation-energy barrier is lower.

A greater proportion of particle collisions can overcome the barrier.

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Worked Example 3

A catalyst changes a reaction's activation energy from:

150 kJ mol⁻¹

to:

85 kJ mol⁻¹

Decrease:

150 − 85 = 65 kJ mol⁻¹

The catalyst lowers the activation energy by:

65 kJ mol⁻¹

This increases the number of successful collisions.

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Worked Example 4

A student says:

"This reaction is very exothermic, so it must happen quickly."

This conclusion is not necessarily correct.

A reaction can:

• release a large amount of energy
• still have a very high activation energy
• therefore proceed slowly

To predict reaction rate, we need kinetic information such as activation energy.

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Worked Example 5

At a particular temperature:

Reaction X has Eₐ = 60 kJ mol⁻¹

Reaction Y has Eₐ = 160 kJ mol⁻¹

Which reaction has the larger energy barrier?

Reaction Y

Which reaction would generally have a larger fraction of successful collisions?

Reaction X

Which is likely faster if other factors are similar?

Reaction X

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Did You Know?

A spark from a lighter contains only a small amount of energy compared with the total energy released when fuel burns.

Its role is not to supply all of the reaction's energy.

Instead, the spark provides enough initial energy for some particles to overcome the activation barrier.

Once combustion begins, the reaction itself releases energy that can help additional reactant particles overcome the barrier.

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Common Misconceptions

Misconception 1

Activation energy is the total energy released by a reaction.

Incorrect.

Activation energy is the barrier that must be overcome.

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Misconception 2

Exothermic reactions do not need activation energy.

Incorrect.

Many exothermic reactions still require a substantial initial energy input.

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Misconception 3

Heating lowers activation energy.

Usually incorrect.

Heating increases the energies of particles.

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Misconception 4

A catalyst gives particles more energy.

Incorrect.

A catalyst provides an alternative pathway with a lower activation energy.

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Misconception 5

Lower activation energy changes ΔH.

Incorrect.

A catalyst can lower activation energy while leaving reactant energy, product energy, and ΔH unchanged.

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A Quick Activation-Energy Checklist

When studying a reaction, ask:

1. What are the reactants?
2. Where is the transition state?
3. What is the energy difference from reactants to the peak?
4. What is the activation energy?
5. Is another reaction's activation energy larger or smaller?
6. How would this affect the number of successful collisions?
7. Is a catalyst present?
8. Does the catalyst lower the energy barrier?
9. Has ΔH changed?
10. How does the activation energy affect reaction rate?

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Connecting the Ideas

The relationship can be summarized as:

reactant particles collide

↓

particles need sufficient energy

↓

collision energy must reach Eₐ

↓

transition state forms

↓

bonds rearrange

↓

products form

Therefore:

higher Eₐ

→ fewer particles can overcome the barrier

→ fewer successful collisions

→ slower reaction

while:

lower Eₐ

→ more particles can overcome the barrier

→ more successful collisions

→ faster reaction

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Key Terms

Activation energy (Eₐ) – The minimum energy barrier that reacting particles must overcome for a reaction to occur.

Collision theory – A model explaining reactions in terms of particle collisions.

Successful collision – A collision with sufficient energy and suitable orientation to lead to reaction.

Transition state – A temporary, unstable, high-energy arrangement during a reaction.

Reaction profile – A diagram showing how energy changes as a reaction proceeds.

Reaction rate – The speed at which reactants are converted into products.

Catalyst – A substance that increases reaction rate by providing an alternative pathway with lower activation energy.

Enzyme – A biological catalyst.

Kinetics – The study of reaction rates and the factors affecting them.

Thermodynamics – The study of energy changes and energetic feasibility.

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Key Takeaways

• Activation energy, Eₐ, is the minimum energy barrier that must be overcome for a reaction to occur.
• Reactions need activation energy because atoms and bonds must rearrange before products can form.
• On a reaction profile, activation energy is measured from the reactant energy level to the peak.
• The peak represents the transition state.
• Activation energy and overall enthalpy change are different quantities.
• A reaction can be highly exothermic and still have a large activation energy.
• According to collision theory, reacting particles must collide with enough energy and suitable orientation.
• At a given temperature, only some particles have enough energy to overcome Eₐ.
• A larger activation energy generally means fewer successful collisions and a slower reaction, assuming comparable conditions.
• A smaller activation energy generally allows more successful collisions and a faster reaction.
• Increasing temperature usually does not lower Eₐ; it increases the fraction of particles with enough energy to overcome it.
• A catalyst lowers activation energy by providing an alternative pathway.
• A catalyst does not change the overall energy difference between reactants and products.
• Enzymes are biological catalysts that allow many reactions to proceed rapidly at normal biological temperatures.
• Understanding activation energy connects energy profiles, collision theory, catalysts, temperature, and reaction rate.