1. Reaction Profile Diagrams

What Is a Reaction Profile Diagram?

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

It allows us to see:

• the energy of the reactants
• the energy of the products
• the activation energy
• the highest-energy point of the reaction
• whether energy is released or absorbed overall

The vertical axis normally represents:

Energy

The horizontal axis represents:

Progress of reaction or Reaction pathway

The horizontal axis does not normally represent time.

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The Basic Shape of a Reaction Profile

A typical reaction profile starts with the reactants, rises to a peak, and then falls to the products.

The general pattern is:

Reactants

↓

energy increases

↓

highest-energy point

↓

energy decreases

↓

Products

The rise in energy shows that the reaction must overcome an energy barrier before products can form.

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Identifying the Reactants

The reactants are the starting substances in a chemical reaction.

On a standard reaction profile diagram, they are shown on the:

left side

For example:

A + B → C + D

A and B are the reactants.

Their position on the vertical axis tells us their starting energy.

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Identifying the Products

The products are the substances formed by the reaction.

They are normally shown on the:

right side

For:

A + B → C + D

C and D are the products.

The difference between the energy levels of the reactants and products tells us whether energy has been released or absorbed.

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Reading the Axes

Reaction profile diagrams usually contain two axes.

Vertical Axis

Energy

This may also be labelled:

• potential energy
• enthalpy
• energy of the system

Horizontal Axis

Reaction progress

or:

Reaction coordinate

It represents the progress from reactants to products.

It is not usually a direct measurement of:

• time
• distance
• concentration

This is an important distinction.

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The Energy Barrier

Chemical reactions do not usually move directly from reactants to products.

Reactant particles must first reach a higher-energy arrangement.

This is why the reaction profile rises to a peak.

The energy required to reach that peak from the reactants is called the:

activation energy

Symbol:

Eₐ

The activation energy is the minimum energy barrier that must be overcome for the reaction to occur.

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Activation Energy on the Diagram

Activation energy is measured vertically from:

the energy level of the reactants

to:

the highest point of the reaction profile

Therefore:

Activation energy = peak energy − reactant energy

The activation energy is not measured from zero.

It is the difference between the reactant energy and the top of the energy barrier.

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Why Is Activation Energy Needed?

During a chemical reaction:

• existing bonds may need to be weakened or broken
• atoms must rearrange
• new bonds must form

The reactants must therefore reach a temporary high-energy arrangement before they can become products.

This creates the activation-energy barrier.

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The Transition State

The top of the reaction profile represents a very unstable, high-energy arrangement called the:

transition state

It is sometimes also called the:

activated complex

The transition state exists for an extremely short time.

It is not usually isolated as a normal substance.

A useful sequence is:

reactants

↓

activation energy supplied

↓

transition state

↓

products

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Exothermic Reaction Profiles

An exothermic reaction transfers energy from the reacting system to the surroundings.

Examples can include:

• combustion
• many oxidation reactions
• many neutralization reactions

On an exothermic reaction profile:

products are lower in energy than reactants

The system has lost energy overall.

---

Recognizing an Exothermic Diagram

Look at the starting and ending energy levels.

If:

product energy < reactant energy

then the reaction is:

exothermic

The profile generally looks like:

higher reactants

↓

peak

↓

lower products

Energy is released to the surroundings.

---

Overall Energy Change in an Exothermic Reaction

The overall energy change is the difference between the energy of the products and the reactants.

This is often written:

ΔH = Hproducts − Hreactants

For an exothermic reaction:

Hproducts < Hreactants

Therefore:

ΔH < 0

The energy change is negative.

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Example: Exothermic Reaction

Suppose:

Reactant energy = 180 kJ mol⁻¹

Product energy = 120 kJ mol⁻¹

Then:

ΔH = 120 − 180

ΔH = −60 kJ mol⁻¹

The negative value shows that the reaction is:

exothermic

Energy released:

60 kJ mol⁻¹

---

Endothermic Reaction Profiles

An endothermic reaction absorbs energy from the surroundings.

On an endothermic reaction profile:

products are higher in energy than reactants

The reacting system has gained energy overall.

---

Recognizing an Endothermic Diagram

If:

product energy > reactant energy

then the reaction is:

endothermic

The profile generally looks like:

lower reactants

↓

peak

↓

higher products

Energy is absorbed from the surroundings.

---

Overall Energy Change in an Endothermic Reaction

Again:

ΔH = Hproducts − Hreactants

For an endothermic reaction:

Hproducts > Hreactants

Therefore:

ΔH > 0

The energy change is positive.

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Example: Endothermic Reaction

Suppose:

Reactant energy = 90 kJ mol⁻¹

Product energy = 145 kJ mol⁻¹

Then:

ΔH = 145 − 90

ΔH = +55 kJ mol⁻¹

The positive value shows that the reaction is:

endothermic

Energy absorbed:

55 kJ mol⁻¹

---

Exothermic vs Endothermic Profiles

The quickest way to identify the reaction type is to compare the reactant and product energy levels.

---

Interpreting the Energy Change

The vertical distance between the reactants and products represents the overall energy change.

For an exothermic reaction:

ΔH points downward

because the products are lower in energy.

For an endothermic reaction:

ΔH points upward

because the products are higher in energy.

The size of this vertical difference tells us the magnitude of the energy change.

---

Activation Energy Is Different from ΔH

Students often confuse:

activation energy

and:

overall energy change

They are not the same.

Activation Energy

Measured from:

reactants → peak

Overall Energy Change

Measured from:

reactants → products

Therefore:

Eₐ tells us about the energy barrier

while:

ΔH tells us about the overall energy change

---

Example

Suppose a reaction profile has:

Reactants = 100 kJ mol⁻¹

Peak = 250 kJ mol⁻¹

Products = 60 kJ mol⁻¹

Activation Energy

Eₐ = 250 − 100

Eₐ = 150 kJ mol⁻¹

Overall Energy Change

ΔH = 60 − 100

ΔH = −40 kJ mol⁻¹

Therefore the reaction is:

exothermic

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A Useful Reading Method

When you are given a reaction profile, follow these steps.

Step 1

Find the reactants.

Usually on the left.

Step 2

Find the products.

Usually on the right.

Step 3

Compare their energy levels.

Products lower:

exothermic

Products higher:

endothermic

Step 4

Find the peak.

This represents the transition state.

Step 5

Measure from the reactants to the peak.

That is:

activation energy

Step 6

Measure from reactants to products.

That is:

overall energy change

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Comparing Two Reaction Profiles

Reaction profiles can be compared in several ways.

You can compare:

• reactant energy
• product energy
• activation energy
• energy change
• reaction type
• number of energy barriers

A reaction with a higher peak relative to its reactants has a larger activation energy.

---

Example: Comparing Activation Energies

Reaction A:

Reactants = 50 kJ mol⁻¹

Peak = 130 kJ mol⁻¹

Reaction B:

Reactants = 50 kJ mol⁻¹

Peak = 200 kJ mol⁻¹

For Reaction A:

Eₐ = 130 − 50 = 80 kJ mol⁻¹

For Reaction B:

Eₐ = 200 − 50 = 150 kJ mol⁻¹

Therefore:

Reaction B has the larger activation energy.

---

Activation Energy and Reaction Rate

A larger activation-energy barrier generally makes it harder for reacting particles to react at a given temperature.

A lower activation-energy barrier means that a greater proportion of collisions may have enough energy to react.

Therefore, under comparable conditions:

lower Eₐ → reaction can generally proceed more readily

However, reaction rate also depends on factors such as:

• temperature
• concentration
• pressure
• surface area
• catalysts
• molecular orientation

---

Catalysts and Reaction Profiles

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

The catalyst does not change the energy of:

• the reactants
• the products

Therefore, it does not change the overall:

ΔH

It only lowers the energy barrier.

---

Catalyzed vs Uncatalyzed Reaction

On a reaction profile:

uncatalyzed pathway

has a higher peak.

catalyzed pathway

has a lower peak.

But both begin and end at the same energy levels.

Therefore:

Eₐ decreases

but:

ΔH remains unchanged

This is an important feature when comparing reaction profiles.

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Example: Catalyst

Without catalyst:

Reactants = 80 kJ mol⁻¹

Peak = 210 kJ mol⁻¹

Products = 40 kJ mol⁻¹

Activation energy:

210 − 80 = 130 kJ mol⁻¹

With catalyst:

Peak = 145 kJ mol⁻¹

Activation energy:

145 − 80 = 65 kJ mol⁻¹

The catalyst has lowered the activation energy from:

130 kJ mol⁻¹ to 65 kJ mol⁻¹

But:

ΔH = 40 − 80 = −40 kJ mol⁻¹

in both cases.

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Comparing Exothermic Reactions

Two reactions can both be exothermic but have different:

• activation energies
• amounts of energy released

For example:

Reaction A:

ΔH = −20 kJ mol⁻¹

Reaction B:

ΔH = −100 kJ mol⁻¹

Both are exothermic.

However, Reaction B releases more energy per mole according to these profiles.

This does not automatically mean Reaction B is faster.

Reaction rate depends strongly on activation energy and other kinetic factors.

---

Comparing Endothermic Reactions

Two reactions can both be endothermic but absorb different amounts of energy.

For example:

Reaction C:

ΔH = +30 kJ mol⁻¹

Reaction D:

ΔH = +90 kJ mol⁻¹

Reaction D requires a larger net energy input.

Again, this does not automatically tell us which reaction is faster.

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Reverse Reactions

Reaction profiles can also be read backwards.

Suppose:

A → B

is exothermic.

Then:

B → A

must be endothermic.

The energy difference has the same magnitude but opposite sign.

For example:

Forward:

ΔH = −60 kJ mol⁻¹

Reverse:

ΔH = +60 kJ mol⁻¹

---

Activation Energy of the Reverse Reaction

The activation energy for the reverse reaction is measured from:

products → peak

rather than:

reactants → peak

Therefore, the forward and reverse reactions usually have different activation energies.

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Example: Forward and Reverse Reaction

Suppose:

Reactants = 120 kJ mol⁻¹

Peak = 280 kJ mol⁻¹

Products = 70 kJ mol⁻¹

Forward activation energy:

280 − 120 = 160 kJ mol⁻¹

Reverse activation energy:

280 − 70 = 210 kJ mol⁻¹

Forward ΔH:

70 − 120 = −50 kJ mol⁻¹

Reverse ΔH:

+50 kJ mol⁻¹

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More Complex Reaction Profiles

Some reactions occur through several steps.

Their reaction profiles may contain:

• more than one peak
• valleys between the peaks

Each peak represents a transition state.

The valleys can represent intermediates.

A simplified multistep profile might look like:

reactants

↓

first peak

↓

intermediate

↓

second peak

↓

products

The largest relevant energy barrier can strongly influence the overall reaction rate.

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Reaction Intermediate

An intermediate is a species that:

• forms during one reaction step
• is used up in a later step
• does not appear in the overall chemical equation

On an energy profile, an intermediate is often represented by a valley between two peaks.

This is different from a transition state, which occurs at the top of a peak.

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Transition State vs Intermediate

This becomes especially useful when interpreting more complicated reaction mechanisms.

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Reaction Profiles Do Not Show Reaction Time

A common mistake is to assume that the horizontal axis represents time.

It usually does not.

The horizontal axis represents the:

reaction pathway

or:

reaction coordinate

A wider-looking curve therefore does not necessarily mean that the reaction takes longer.

Reaction speed cannot be determined simply from the horizontal width of the diagram.

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Reaction Profiles Do Not Directly Show Temperature

Reaction profiles show relative energy changes.

They do not directly show:

• temperature
• reaction time
• concentration

For example, an exothermic reaction may cause the surroundings to become warmer, but the vertical axis is not a temperature axis.

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Worked Example 1

A reaction profile has:

Reactants = 70 kJ mol⁻¹

Peak = 190 kJ mol⁻¹

Products = 20 kJ mol⁻¹

Activation Energy

Eₐ = 190 − 70

Eₐ = 120 kJ mol⁻¹

Energy Change

ΔH = 20 − 70

ΔH = −50 kJ mol⁻¹

Reaction Type

Because ΔH is negative:

exothermic

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Worked Example 2

Reactants = 40 kJ mol⁻¹

Peak = 160 kJ mol⁻¹

Products = 95 kJ mol⁻¹

Activation energy:

Eₐ = 160 − 40

Eₐ = 120 kJ mol⁻¹

Energy change:

ΔH = 95 − 40

ΔH = +55 kJ mol⁻¹

Therefore:

endothermic

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Worked Example 3: Comparing Profiles

Reaction A:

• Reactants = 60
• Peak = 140
• Products = 20

Reaction B:

• Reactants = 60
• Peak = 200
• Products = 20

Both have:

ΔH = −40 kJ mol⁻¹

But:

Reaction A:

Eₐ = 80 kJ mol⁻¹

Reaction B:

Eₐ = 140 kJ mol⁻¹

Therefore:

• both reactions release the same net amount of energy
• Reaction B has the larger activation-energy barrier

This demonstrates that activation energy and energy change describe different features of a reaction.

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Worked Example 4: Comparing Energy Changes

Reaction X:

Reactants = 100 kJ mol⁻¹

Products = 40 kJ mol⁻¹

Reaction Y:

Reactants = 100 kJ mol⁻¹

Products = 80 kJ mol⁻¹

For X:

ΔH = −60 kJ mol⁻¹

For Y:

ΔH = −20 kJ mol⁻¹

Both reactions are exothermic.

However:

Reaction X releases more energy overall.

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Did You Know?

A reaction can be strongly exothermic and still happen very slowly.

This is because:

thermodynamics tells us about the overall energy change,

while:

kinetics tells us how quickly the reaction occurs.

A reaction may have products that are much lower in energy but still require a very large activation energy to reach them.

This is one reason activation-energy diagrams are so useful: they connect chemical energetics with reaction rates.

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A Quick Visual Checklist

When looking at any reaction profile, ask:

1. Where are the reactants?
2. Where are the products?
3. Which is higher in energy?
4. Where is the highest point?
5. What is the activation energy?
6. What is the overall energy change?
7. Is ΔH positive or negative?
8. Is the reaction exothermic or endothermic?
9. Is there a catalyst shown?
10. Are there multiple peaks suggesting a multistep reaction?

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Connecting the Ideas

A reaction profile can be interpreted as:

reactants

↓

energy barrier

↓

activation energy

↓

transition state

↓

products

Then compare reactant and product energy:

products lower

→ exothermic

→ ΔH negative

or:

products higher

→ endothermic

→ ΔH positive

A catalyst changes:

activation energy

but does not change:

reactant energy, product energy, or ΔH

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Key Terms

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

Reactants – Starting substances in a reaction.

Products – Substances formed during a reaction.

Reaction coordinate – A representation of the progress of a chemical reaction.

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

Transition state – The temporary highest-energy arrangement along a reaction pathway.

Activated complex – Another term often used for the high-energy arrangement at the transition state.

Enthalpy change (ΔH) – The difference between the energy of the products and reactants at constant pressure.

Exothermic reaction – A reaction that transfers energy to the surroundings.

Endothermic reaction – A reaction that absorbs energy from the surroundings.

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

Intermediate – A temporary species formed during one step of a multistep reaction and consumed during another.

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Key Takeaways

• Reaction profile diagrams show how the energy of a reacting system changes during a reaction.
• The reactants are normally shown on the left and the products on the right.
• The vertical axis represents energy.
• The horizontal axis represents reaction progress, not necessarily time.
• The highest point of the curve represents the transition state.
• Activation energy is measured from the reactant energy level to the peak of the profile.
• The overall energy change is:

ΔH = Hproducts − Hreactants

• If the products are lower in energy than the reactants, the reaction is exothermic and ΔH is negative.
• If the products are higher in energy than the reactants, the reaction is endothermic and ΔH is positive.
• Activation energy and overall energy change are not the same quantity.
• A catalyst lowers the activation energy but does not change the overall ΔH.
• Different reaction profiles can be compared by examining their reactant energy, product energy, activation energy, and overall energy change.
• Multistep reactions can contain several peaks and intermediates.
• Reaction profiles connect chemical energetics with the study of reaction rates and mechanisms.

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2. Activation Energy

 

3. Catalysts and Enzymes

What Is a Catalyst?

A catalyst is a substance that increases the rate of a chemical reaction without being consumed overall.

Catalysts work by providing an alternative reaction pathway with a lower activation energy, Eₐ. This allows a greater proportion of reacting particles to overcome the energy barrier and react successfully.

A catalyst may participate in intermediate steps of a reaction, but it is regenerated before the overall reaction is complete.

In simple terms:

Catalyst added → alternative pathway → lower activation energy → faster reaction

Catalysts Lower Activation Energy

Every chemical reaction has an energy barrier called the activation energy. Reacting particles must overcome this barrier before products can form.

A catalyst provides another route from reactants to products that has a lower activation-energy barrier.

Notice something very important in these reaction profiles.

The catalyzed and uncatalyzed pathways:

• start at the same reactant energy
• finish at the same product energy
• have different activation energies

Therefore:

Eₐ(catalyzed) < Eₐ(uncatalyzed)

The catalyst changes the pathway, not the beginning or end of the reaction.

What Does a Catalyst Change?

A catalyst changes the reaction pathway, activation energy and reaction rate.

It does not change:

• the energy of the reactants
• the energy of the products
• the overall enthalpy change, ΔH
• whether the reaction is exothermic or endothermic
• the equilibrium constant or equilibrium position

The catalyzed pathway may also contain additional intermediate steps and transition states.

Worked Example: Activation Energy

Suppose a reaction has:

Reactant energy = 60 kJ mol⁻¹

Uncatalyzed peak = 180 kJ mol⁻¹

Therefore:

Eₐ = 180 − 60

Eₐ = 120 kJ mol⁻¹

Now a catalyst provides a pathway whose highest point is only 110 kJ mol⁻¹.

Eₐ = 110 − 60

Eₐ = 50 kJ mol⁻¹

The catalyst has lowered the activation energy by:

120 − 50 = 70 kJ mol⁻¹

Why Does Lower Activation Energy Increase Reaction Rate?

At any particular temperature, particles do not all have exactly the same kinetic energy.

Some have relatively low energies, while others have enough energy to overcome the activation-energy barrier.

If the activation energy is lowered, more particles have sufficient energy to react successfully.

Therefore:

lower Eₐ → more successful collisions → faster reaction

Importantly, the catalyst has not given the particles extra kinetic energy. It has lowered the energy barrier they must overcome.

Catalysts and Collision Theory

According to collision theory, particles generally need to:

• collide
• collide with sufficient energy
• collide with an appropriate orientation

A catalyst increases reaction rate because its alternative mechanism requires a lower activation energy.

Some catalysts can also help arrange reacting particles in favourable positions.

A Different Reaction Pathway

It is tempting to imagine a catalyst simply pushing down the top of the original energy curve.

In reality, a catalyst can provide an entirely different reaction mechanism.

The catalyzed mechanism may involve:

• different intermediate substances
• different transition states
• additional reaction steps
• several smaller energy barriers rather than one large barrier

This is why some catalyzed reaction profiles contain several smaller peaks.

Why Isn't a Catalyst Consumed?

Consider this simplified mechanism:

A + catalyst → intermediate

Then:

intermediate + B → products + catalyst

The catalyst appears during the reaction, but it is produced again during a later step.

When the steps are combined:

A + B → products

The catalyst cancels from the overall equation.

So saying that a catalyst is "not consumed" does not mean it never participates in the reaction. It means it is regenerated overall.

Homogeneous and Heterogeneous Catalysts

A homogeneous catalyst is in the same phase as the reactants.

For example:

reactants in solution + catalyst in solution

A heterogeneous catalyst is in a different phase from the reactants.

A common industrial situation is:

gaseous reactants + solid catalyst

The reaction then occurs on the surface of the solid catalyst.

The image compares homogeneous catalysis with a heterogeneous reaction occurring on a catalyst surface.

How a Heterogeneous Catalyst Works

A simplified surface-catalyzed reaction involves several stages:

1. Reactant particles approach the catalyst.
2. Reactants attach to the surface.
3. Bonds in the reactants may weaken.
4. The particles react.
5. Products form.
6. Products leave the surface.
7. The catalyst surface becomes available again.

Attachment to the surface is called adsorption.

Release from the surface is called desorption.

Active Sites and Surface Area

Reactions do not necessarily occur equally well everywhere on a catalyst.

Particular locations where catalytic interactions occur are called active sites.

Increasing the surface area of a solid catalyst can expose more active sites.

Therefore:

greater surface area → more exposed active sites → more simultaneous catalytic interactions

This is why industrial catalysts are often porous or finely divided rather than simply being large solid blocks.

Catalysts in Industry

Catalysts are extremely important in industrial chemistry because manufacturers need reactions that are fast enough to be practical while controlling energy costs and unwanted products.

Three important examples are:

• iron-based catalysts in the Haber process
• vanadium(V) oxide in the Contact process
• platinum-group metals in catalytic converters

The Haber Process

The Haber process manufactures ammonia:

N₂ + 3H₂ ⇌ 2NH₃

An iron-based catalyst is used.

Nitrogen molecules contain the extremely strong N≡N triple bond. Breaking and rearranging these bonds involves a substantial activation-energy barrier.

On the catalyst surface, nitrogen and hydrogen can be adsorbed and react through a lower-energy mechanism.

The catalyst therefore allows ammonia to be produced at a much more useful industrial rate.

Catalysts and Equilibrium

The Haber reaction is reversible.

A catalyst speeds up both the forward and reverse reactions.

This means equilibrium is reached more quickly.

However:

a catalyst does not shift the equilibrium position.

It changes how quickly equilibrium is reached, not where the equilibrium eventually lies.

The Contact Process

Another important industrial example occurs during the manufacture of sulfuric acid.

One stage of the Contact process is:

2SO₂ + O₂ ⇌ 2SO₃

A commonly used catalyst is vanadium(V) oxide, V₂O₅.

The catalyst allows sulfur dioxide and oxygen to react at a useful industrial rate. The sulfur trioxide produced is then used in later stages of sulfuric acid manufacture.

Catalytic Converters

Catalysis is also used every day in vehicle exhaust systems.

A catalytic converter contains catalytic materials that commonly include platinum, palladium and rhodium.

Notice the honeycomb structure inside the converter.

The honeycomb provides a very large surface area while still allowing exhaust gases to pass through.

Catalytic converters help convert pollutants such as:

• carbon monoxide
• nitrogen oxides
• unburned hydrocarbons

into less harmful substances.

This is an excellent real-world example of heterogeneous catalysis.

Catalysts and Green Chemistry

Catalysts can make industrial chemistry more sustainable.

Using an effective catalyst may allow:

• lower operating temperatures
• reduced energy consumption
• faster production
• improved selectivity
• fewer unwanted by-products
• less chemical waste

Finding better catalysts is therefore an important area of green chemistry.

Enzymes: Biological Catalysts

Living organisms also depend on catalysts.

Biological catalysts are called enzymes.

Most enzymes are proteins. They allow biochemical reactions to occur rapidly enough to sustain life under relatively mild biological conditions.

Enzymes obey the same fundamental energetic principle as other catalysts:

they lower activation energy.

Why Do Living Organisms Need Enzymes?

Cells cannot simply heat themselves to hundreds of degrees Celsius whenever they need a chemical reaction to happen faster.

Instead, they use enzymes.

Enzymes are involved in processes such as:

• digestion
• cellular respiration
• photosynthesis
• DNA replication
• protein synthesis
• metabolism
• cell repair
• detoxification

Without enzymes, many of these reactions would occur far too slowly to sustain life.

Enzyme, Substrate and Active Site

The reactant acted upon by an enzyme is called the substrate.

An enzyme contains a particular region called its active site.

The substrate interacts with the active site to form a temporary enzyme-substrate complex.

The shape and chemical properties of the active site help determine which substrates can interact effectively with the enzyme.

The Enzyme Reaction Cycle

A simplified enzyme-controlled reaction can be represented as:

enzyme + substrate

↓

enzyme-substrate complex

↓

reaction occurs

↓

products form

↓

products leave

↓

enzyme available again

Symbolically:

E + S ⇌ ES → E + P

where:

E = enzyme
S = substrate
ES = enzyme-substrate complex
P = product

The enzyme can then catalyze another reaction.

Lock-and-Key and Induced Fit

Two models are commonly used to explain enzyme-substrate interactions.

The lock-and-key model treats the active site as having a shape complementary to the substrate.

The induced-fit model recognizes that enzymes are flexible. As a substrate begins to bind, the active site can change shape slightly to interact more effectively with it.

The lock-and-key model is useful for introducing enzyme specificity, but the induced-fit model gives a more realistic picture of many enzyme interactions.

How Do Enzymes Lower Activation Energy?

Enzymes can lower activation energy in several ways.

Depending on the enzyme, the active site may:

• bring substrates close together
• orient substrates correctly
• strain particular chemical bonds
• provide acidic or basic groups
• create a favourable local environment
• stabilize the transition state

The enzyme is therefore doing more than simply "holding" the substrate.

Examples of Enzymes

Catalase

One particularly useful example is catalase.

Hydrogen peroxide can be produced during cellular metabolism. Because it can damage cells, organisms need to break it down efficiently.

Catalase speeds up:

2H₂O₂ → 2H₂O + O₂

The oxygen produced can be observed as bubbles.

Investigating Catalase

Catalase activity can be investigated experimentally using materials such as potato or liver.

Hydrogen peroxide is added to the catalase-containing material, and the oxygen produced can be collected.

The volume of oxygen produced over a measured period can be used to determine reaction rate.

For example:

reaction rate = volume of O₂ produced ÷ time

Raw and Cooked Potato

A simple catalase experiment can also demonstrate enzyme denaturation.

Raw potato contains active catalase.

If potato is heated strongly, the structure of the enzyme can be disrupted.

When hydrogen peroxide is added, untreated potato typically produces more vigorous oxygen bubbling than heavily heated potato because heating can denature the catalase.

Temperature and Enzyme Activity

Temperature has a major effect on enzyme-controlled reactions.

At relatively low temperatures, particles have less kinetic energy and enzyme-substrate collisions occur less frequently.

As temperature rises, collision frequency increases and the reaction generally becomes faster.

Eventually an optimum temperature is reached.

Above the optimum, activity can decrease rapidly because excessive heat disrupts the enzyme's three-dimensional structure.

Importantly, there is no universal optimum temperature for all enzymes. Different enzymes and organisms are adapted to different conditions.

Enzyme Denaturation

An enzyme's function depends on its three-dimensional structure.

High temperatures can disrupt interactions that maintain this structure.

The active site changes shape, so the substrate may no longer bind effectively.

This process is called denaturation.

Therefore:

high temperature → structural change → active site altered → fewer successful interactions → lower enzyme activity

Low temperature is different. It generally slows an enzyme-controlled reaction but does not necessarily denature the enzyme.

Temperature and Catalysts Are Different

Increasing temperature and adding a catalyst can both increase reaction rate, but they work differently.

Increasing temperature → increases particle kinetic energy

Adding a catalyst → lowers activation energy

Heating does not normally lower the activation energy of a particular reaction pathway.

pH and Enzyme Activity

Enzymes can also be strongly affected by pH.

Different enzymes have different optimum pH values.

Digestive enzymes provide good examples.

Pepsin operates effectively under acidic conditions in the stomach.

Salivary amylase operates closer to neutral conditions.

Enzymes operating farther along the digestive system may function best under more alkaline conditions.

Changing pH can affect electrical charges and interactions within an enzyme, which can alter its shape and activity.

Enzymes in Industry

Enzymes are not restricted to living cells. Humans use them extensively in technology and industry.

Applications include:

• biological detergents
• cheese making
• baking
• brewing
• fruit juice production
• lactose-free dairy products
• pharmaceuticals
• biotechnology

One major advantage is specificity. An enzyme can sometimes produce a desired chemical change while minimizing unwanted side reactions.

Lactase and Lactose-Free Milk

Lactase catalyzes the breakdown of lactose.

Lactose is converted into the simpler sugars glucose and galactose.

This reaction can be used commercially to manufacture lactose-free milk.

Immobilized Enzymes

Industrial enzymes can sometimes be attached to or trapped within a solid material. These are called immobilized enzymes.

One method traps enzymes inside small alginate beads.

For example, lactase-containing beads can be packed into a column.

Milk enters the column and flows around the beads.

As it passes through:

lactose → glucose + galactose

The lactase remains in the column while the treated milk flows out.

Advantages of Immobilized Enzymes

Immobilized enzymes can offer several advantages:

• the enzyme can be reused
• enzyme and product are easier to separate
• continuous processing is possible
• product contamination by enzyme is reduced
• the enzyme may be more stable under some conditions

However, immobilization can also require specialized equipment and may sometimes reduce the rate at which substrate reaches the enzyme.

Catalysts Can Lose Activity

Although catalysts are not consumed in the intended overall reaction, they do not necessarily work forever.

Catalysts can become less effective because of:

• contamination
• overheating
• structural changes
• chemical damage
• blockage of active sites

This is called catalyst deactivation.

Catalyst Poisoning

If another substance binds strongly to the active sites of a heterogeneous catalyst, reactants may no longer be able to adsorb there.

This is known as catalyst poisoning.

The sequence is:

impurity attaches → active sites blocked → less adsorption → fewer catalytic reactions → lower reaction rate

This does not contradict the statement that catalysts are not consumed overall. The catalyst has instead become deactivated.

Enzyme Inhibition

Enzymes can also have their activity reduced by substances called inhibitors.

An inhibitor may:

• block the active site
• compete with the substrate
• bind elsewhere on the enzyme
• change enzyme shape
• interfere with the catalytic mechanism

Enzyme inhibition is important in medicine, metabolism and biochemical regulation.

Comparing Industrial Catalysts and Enzymes

Despite their differences, both rely on the same fundamental idea:

alternative reaction pathway → lower activation energy → faster reaction

Worked Example: Catalyst and ΔH

Consider a reaction with:

Reactants = 80 kJ mol⁻¹

Products = 30 kJ mol⁻¹

Uncatalyzed peak = 220 kJ mol⁻¹

Catalyzed peak = 140 kJ mol⁻¹

Without catalyst:

Eₐ = 220 − 80

Eₐ = 140 kJ mol⁻¹

With catalyst:

Eₐ = 140 − 80

Eₐ = 60 kJ mol⁻¹

But the enthalpy change remains:

ΔH = 30 − 80

ΔH = −50 kJ mol⁻¹

Therefore, the catalyst has changed Eₐ, but it has not changed ΔH.

Did You Know?

An iron-based catalyst inside an ammonia plant and an enzyme inside one of your cells may look completely different, but energetically they perform the same fundamental job.

Both provide a reaction mechanism with a lower activation-energy barrier.

That makes catalysis an important link between chemical energetics, reaction kinetics, industrial chemistry and biology.

Common Misconceptions

"A catalyst gives particles more energy."

No. A catalyst lowers the activation-energy barrier.

"A catalyst changes ΔH."

No. The energy difference between reactants and products remains unchanged.

"A catalyst never participates in the reaction."

It may participate in intermediate steps, but it is regenerated overall.

"A catalyst shifts equilibrium."

No. It allows equilibrium to be reached faster.

"A catalyst lasts forever."

Not necessarily. Catalysts can become poisoned or otherwise deactivated.

"Enzymes supply energy to cells."

No. Enzymes lower activation energy.

"All enzymes work best at 37°C."

No. Optimum temperature varies considerably between enzymes and organisms.

"All enzymes work best at pH 7."

No. Different enzymes can have very different optimum pH values.

Key Terms

Catalyst – A substance that increases reaction rate without being consumed overall.

Activation energy, Eₐ – The energy barrier that must be overcome for a reaction to occur.

Alternative pathway – A different reaction route with a lower activation-energy barrier.

Reaction mechanism – The sequence of elementary stages through which a reaction occurs.

Homogeneous catalyst – A catalyst in the same phase as the reactants.

Heterogeneous catalyst – A catalyst in a different phase from the reactants.

Active site – A location where catalytic interactions occur.

Adsorption – Attachment of particles to a surface.

Desorption – Release of particles from a surface.

Enzyme – A biological catalyst.

Substrate – A reactant acted upon by an enzyme.

Enzyme-substrate complex – The temporary combination formed when a substrate interacts with an enzyme.

Induced fit – A model in which substrate binding causes a change in the enzyme's active site.

Denaturation – Loss of an enzyme's functional three-dimensional structure.

Immobilized enzyme – An enzyme attached to or trapped within a solid support.

Catalyst poisoning – Reduction in catalytic activity caused by substances blocking or altering active sites.

Inhibitor – A substance that reduces enzyme activity.

Key Takeaways

• Catalysts increase reaction rate without being consumed overall.
• They provide an alternative reaction pathway with lower activation energy.
• Lower Eₐ allows more reacting particles to overcome the energy barrier.
• Catalysts do not give particles extra kinetic energy.
• Catalysts do not change ΔH.
• Catalysts may participate in intermediate stages but are regenerated.
• Heterogeneous catalysts often operate through adsorption at surface active sites.
• Greater catalyst surface area can expose more active sites.
• Important industrial examples include the Haber process, Contact process and catalytic converters.
• Catalysts speed both forward and reverse reactions, allowing equilibrium to be reached faster without shifting its position.
• Enzymes are biological catalysts.
• Enzymes have active sites that interact with substrates.
• The induced-fit model recognizes that an enzyme can change shape during substrate binding.
• Temperature and pH can strongly affect enzyme activity.
• Excessive heating can cause enzyme denaturation.
• Immobilized enzymes can be retained and reused in industrial processes.
• Catalysts can become deactivated even though they are not consumed by the intended reaction.
• Catalysis connects energetics, kinetics, industry and biology.

4. Exothermic Reaction Profiles

What Is an Exothermic Reaction?

An exothermic reaction is a chemical reaction that transfers energy from the reacting system to the surroundings.

This energy is commonly transferred as heat, although some reactions may also release energy as light or sound.

Common examples include:

• combustion
• many oxidation reactions
• acid-base neutralization
• cellular respiration
• some reactions between metals and acids

The key idea is:

chemical system loses energy → surroundings gain energy

On an energy profile diagram, this means the products are at a lower energy level than the reactants.

Recognizing an Exothermic Reaction Profile

An energy profile diagram usually has two axes.

The vertical axis represents:

Energy or Enthalpy

The horizontal axis represents:

Reaction progress or Reaction coordinate

The horizontal axis does not normally represent time. It simply represents the pathway from reactants to products.

For an exothermic reaction:

reactants start higher → curve rises to a peak → products finish lower

The easiest way to identify an exothermic profile is therefore to compare the beginning and end of the graph.

Products below reactants = exothermic reaction

Why Are the Products Lower in Energy?

Chemical reactions involve both breaking bonds and forming bonds.

Breaking existing chemical bonds requires energy.

Bond breaking → energy absorbed

Forming new chemical bonds releases energy.

Bond making → energy released

In an exothermic reaction, the formation of new bonds releases more energy than was required to break the bonds in the reactants.

For example, imagine:

Energy needed to break bonds = 300 kJ mol⁻¹

Energy released when new bonds form = 450 kJ mol⁻¹

The overall energy change is:

300 − 450 = −150 kJ mol⁻¹

Therefore, the reaction releases 150 kJ mol⁻¹ overall.

The products finish at a lower energy because energy has been transferred from the chemical system to the surroundings.

Enthalpy Change, ΔH

The overall energy change of a reaction is commonly represented by the enthalpy change, ΔH.

It can be calculated using:

ΔH = Hproducts − Hreactants

For an exothermic reaction:

Hproducts < Hreactants

Therefore:

ΔH < 0

The enthalpy change is negative.

The downward difference between the reactant and product energy levels represents the energy released by the reaction.

Worked Example 1: Identifying an Exothermic Reaction

Suppose an energy profile shows:

Reactants = 180 kJ mol⁻¹

Products = 110 kJ mol⁻¹

Because:

110 < 180

the products have less energy than the reactants.

Therefore, the reaction is:

exothermic

The enthalpy change is:

ΔH = 110 − 180

ΔH = −70 kJ mol⁻¹

The reaction releases:

70 kJ mol⁻¹

Notice the difference in wording:

ΔH = −70 kJ mol⁻¹

but:

energy released = 70 kJ mol⁻¹

The negative sign tells us the direction of the energy transfer.

Energy Released

The amount of energy released can be found from the vertical difference between the reactant and product energy levels.

For example:

Reactants = 250 kJ mol⁻¹

Products = 90 kJ mol⁻¹

Energy difference:

250 − 90 = 160 kJ mol⁻¹

Therefore:

160 kJ mol⁻¹ is released

and:

ΔH = −160 kJ mol⁻¹

A larger vertical drop between reactants and products represents a greater energy release.

Activation Energy Is Still Required

A common misconception is that an exothermic reaction should begin automatically because it releases energy.

That is not necessarily true.

Exothermic reactions still normally require an initial amount of energy called the activation energy, Eₐ.

Activation energy is measured from:

reactant energy level → peak of the reaction profile

It should not be measured from the bottom of the graph.

The peak represents the transition state, a very high-energy and unstable arrangement of atoms.

Worked Example 2: Calculating Activation Energy

Suppose:

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
• energy released = 60 kJ mol⁻¹
• ΔH = −60 kJ mol⁻¹
• reaction = exothermic

This illustrates an important point:

activation energy and energy released are not the same thing.

Why Does an Exothermic Reaction Need Activation Energy?

Consider the combustion of a fuel.

Fuel can be surrounded by oxygen without immediately bursting into flames.

An initial energy input may be needed from:

• a spark
• a flame
• electrical energy
• sufficient heating

That initial energy helps reacting particles overcome the activation-energy barrier.

Once the reaction begins, the energy released by the reaction may allow it to continue.

Combustion therefore provides a useful example of a reaction that is strongly exothermic but still requires activation energy.

Activation Energy vs Energy Released

These two quantities describe completely different parts of the profile.

Do not assume that a reaction releasing a large amount of energy must have a small activation energy.

It can have a large energy release and a large activation-energy barrier.

Comparing Exothermic Reactions

Energy profiles can be used to compare different exothermic reactions.

Suppose Reaction A has:

Reactants = 100 kJ mol⁻¹
Products = 40 kJ mol⁻¹

Then:

ΔH = −60 kJ mol⁻¹

Reaction B has:

Reactants = 100 kJ mol⁻¹
Products = −20 kJ mol⁻¹

Then:

ΔH = −120 kJ mol⁻¹

Both reactions are exothermic.

However, Reaction B releases more energy because its products finish much farther below its reactants.

Therefore:

larger downward energy difference → greater energy release

Comparing Activation Energies

Now suppose two exothermic reactions have the same reactant and product energy levels.

Reaction A:

Reactants = 80 kJ mol⁻¹
Peak = 150 kJ mol⁻¹
Products = 30 kJ mol⁻¹

Reaction B:

Reactants = 80 kJ mol⁻¹
Peak = 230 kJ mol⁻¹
Products = 30 kJ mol⁻¹

For Reaction A:

Eₐ = 150 − 80 = 70 kJ mol⁻¹

For Reaction B:

Eₐ = 230 − 80 = 150 kJ mol⁻¹

Both have:

ΔH = 30 − 80 = −50 kJ mol⁻¹

Therefore, the reactions release the same amount of energy, but they have different activation energies.

Under otherwise comparable conditions, Reaction A has the smaller energy barrier.

This shows why energy released and activation energy must be considered separately.

Exothermic vs Endothermic Profiles

It is useful to compare the two major types of energy profile.

For an exothermic reaction:

products below reactants

ΔH is negative

energy released

For an endothermic reaction:

products above reactants

ΔH is positive

energy absorbed

The position of the products relative to the reactants is therefore the quickest way to classify the reaction.

Exothermic Does Not Mean Fast

Another common mistake is to assume:

"A strongly exothermic reaction must be fast."

This is incorrect.

Exothermic describes the overall energy change.

Reaction rate describes how quickly the reaction occurs.

A reaction can release a great deal of energy but still proceed slowly if it has a large activation-energy barrier.

This is an important distinction between thermodynamics and kinetics.

Catalysts and Exothermic Profiles

A catalyst lowers the activation energy by providing an alternative reaction pathway.

However, it does not change:

• the reactant energy
• the product energy
• the amount of energy released
• ΔH

Therefore, on a catalyzed exothermic profile:

same starting level + same finishing level + lower peak

The reaction remains equally exothermic.

Only the activation-energy barrier changes.

Forward and Reverse Reactions

An exothermic reaction can also be considered in reverse.

Suppose:

Reactants = 120 kJ mol⁻¹

Peak = 280 kJ mol⁻¹

Products = 70 kJ mol⁻¹

For the forward reaction:

Eₐ = 280 − 120 = 160 kJ mol⁻¹

and:

ΔH = 70 − 120 = −50 kJ mol⁻¹

The forward reaction is exothermic.

For the reverse reaction, the products become the starting materials:

Eₐ(reverse) = 280 − 70

Eₐ(reverse) = 210 kJ mol⁻¹

and:

ΔH(reverse) = +50 kJ mol⁻¹

So if the forward reaction is exothermic, the reverse reaction is endothermic.

Reading an Exothermic Profile Step by Step

When you encounter an unfamiliar reaction profile, use this sequence.

First, locate the reactants.

Next, locate the products.

Ask:

Are the products below the reactants?

If yes, the reaction is exothermic.

Then determine:

ΔH = Hproducts − Hreactants

Next locate the highest point of the curve.

Determine:

Eₐ = Hpeak − Hreactants

Finally, compare these quantities with other profiles if required.

Worked Example 3

An energy profile has:

Reactants = 150 kJ mol⁻¹

Transition state = 310 kJ mol⁻¹

Products = 50 kJ mol⁻¹

Step 1: Identify the reaction.

Products are lower than reactants.

Exothermic

Step 2: Calculate ΔH.

ΔH = 50 − 150

ΔH = −100 kJ mol⁻¹

Step 3: Determine energy released.

100 kJ mol⁻¹

Step 4: Calculate activation energy.

Eₐ = 310 − 150

Eₐ = 160 kJ mol⁻¹

Worked Example 4: Comparing Two Profiles

Reaction X:

Eₐ = 90 kJ mol⁻¹
ΔH = −40 kJ mol⁻¹

Reaction Y:

Eₐ = 140 kJ mol⁻¹
ΔH = −100 kJ mol⁻¹

Which releases more energy?

Reaction Y, because its ΔH has the larger negative magnitude.

Which has the smaller activation-energy barrier?

Reaction X

Which reaction is exothermic?

Both

This demonstrates why a complete comparison should consider both Eₐ and ΔH.

Did You Know?

An exothermic reaction can sometimes release a large amount of energy but remain effectively dormant until something supplies its activation energy.

This is one reason fuels can be stored in the presence of oxygen without continuously burning. Once sufficient activation energy is supplied, however, combustion can proceed and transfer substantial energy to the surroundings.

Common Mistakes

"Products are lower, so activation energy must be negative."

Incorrect. Activation energy is the upward energy difference from the reactants to the transition-state peak and is normally positive.

"The energy released is measured from the peak to the products."

Incorrect. Overall energy release is determined from the difference between the reactant and product energy levels.

"A large ΔH means a reaction is fast."

Incorrect. ΔH describes overall energy change, not reaction rate.

"Exothermic reactions do not need activation energy."

Incorrect. They normally still need to overcome an initial energy barrier.

"The x-axis represents time."

Usually incorrect. It represents reaction progress or reaction coordinate.

Key Terms

Exothermic reaction – A reaction that transfers energy from the reacting system to the surroundings.

Energy profile – A diagram showing how the energy of a reacting system changes as a reaction progresses.

Reactants – The substances present at the beginning of a reaction.

Products – The substances formed by a reaction.

Enthalpy, H – A measure of the energy content of a system under constant-pressure conditions.

Enthalpy change, ΔH – The difference between product and reactant enthalpy.

Activation energy, Eₐ – The energy barrier that must be overcome for a reaction to occur.

Transition state – The unstable, high-energy arrangement associated with the peak of an energy profile.

Reaction coordinate – A conceptual representation of progress from reactants to products.

Key Takeaways

• An exothermic reaction releases energy to the surroundings.
• On an exothermic energy profile, the products are lower in energy than the reactants.
• Exothermic reactions have negative ΔH.
• ΔH = Hproducts − Hreactants.
• The magnitude of the vertical difference between reactants and products represents the amount of energy released.
• Products are lower in energy because bond formation releases more energy than is required overall to break bonds.
• Exothermic reactions still require activation energy.
• Activation energy is measured from the reactant level to the peak.
• Activation energy is not the same as the energy released.
• Two exothermic reactions can release different amounts of energy.
• Two reactions can release the same amount of energy while having different activation energies.
• A strongly exothermic reaction is not necessarily a fast reaction.
• A catalyst lowers activation energy but does not change ΔH.
• The reverse of an exothermic reaction is endothermic.

5. Endothermic Reaction Profiles

What Is an Endothermic Reaction?

An endothermic reaction is a chemical reaction that absorbs energy from the surroundings.

The reacting system gains energy, usually as heat.

In simple terms:

surroundings lose energy → reacting system gains energy

Because the system has gained energy, the products finish at a higher energy level than the reactants. Endothermic reactions therefore have a positive enthalpy change, ΔH.

Recognizing an Endothermic Reaction Profile

An energy profile diagram normally has:

Vertical axis: Energy, potential energy or enthalpy

Horizontal axis: Reaction progress or reaction coordinate

The horizontal axis represents the progress of the reaction from reactants to products. It does not normally represent time.

For an endothermic reaction:

reactants start lower → curve rises to a peak → products finish higher

The quickest way to identify an endothermic profile is therefore:

Products above reactants = endothermic

Why Do the Products Have More Energy?

Chemical reactions involve both:

• breaking existing bonds
• forming new bonds

Breaking bonds requires energy.

Bond breaking → energy absorbed

Forming bonds releases energy.

Bond making → energy released

In an endothermic reaction, more energy is required to break bonds than is released when new bonds form. The system therefore gains energy overall.

For example:

Energy required to break bonds = 500 kJ mol⁻¹

Energy released when new bonds form = 350 kJ mol⁻¹

Overall energy change:

ΔH = 500 − 350

ΔH = +150 kJ mol⁻¹

The reaction absorbs:

150 kJ mol⁻¹

Enthalpy Change, ΔH

The overall energy change of a reaction is represented by the enthalpy change, ΔH.

The equation is:

ΔH = Hproducts − Hreactants

For an endothermic reaction:

Hproducts > Hreactants

Therefore:

ΔH > 0

The enthalpy change is positive.

For example:

Reactants = 80 kJ mol⁻¹

Products = 140 kJ mol⁻¹

Then:

ΔH = 140 − 80

ΔH = +60 kJ mol⁻¹

The reaction absorbs:

60 kJ mol⁻¹

Energy Absorbed

The amount of energy absorbed is shown by the vertical difference between the reactant and product energy levels.

Suppose:

Reactants = 120 kJ mol⁻¹

Products = 210 kJ mol⁻¹

Energy absorbed:

210 − 120 = 90 kJ mol⁻¹

Therefore:

ΔH = +90 kJ mol⁻¹

and the reaction absorbs:

90 kJ mol⁻¹

Notice the wording:

ΔH = +90 kJ mol⁻¹

but:

energy absorbed = 90 kJ mol⁻¹

The positive sign shows that the reacting system has gained energy.

Activation Energy Is Still Required

Like exothermic reactions, endothermic reactions also have an activation energy, Eₐ.

Activation energy is the minimum energy barrier that reacting particles must overcome before products can form.

On an energy profile, activation energy is measured from:

reactant energy level → highest point of the profile

The highest point represents the transition state or activated complex.

The activation energy should not be measured from the bottom of the graph.

It is always measured relative to the reactants.

Worked Example 1: Activation Energy and ΔH

Suppose an endothermic reaction has:

Reactants = 60 kJ mol⁻¹

Peak = 240 kJ mol⁻¹

Products = 130 kJ mol⁻¹

Activation energy:

Eₐ = 240 − 60

Eₐ = 180 kJ mol⁻¹

Enthalpy change:

ΔH = 130 − 60

ΔH = +70 kJ mol⁻¹

Therefore:

• activation energy = 180 kJ mol⁻¹
• energy absorbed = 70 kJ mol⁻¹
• ΔH = +70 kJ mol⁻¹
• reaction = endothermic

This shows that activation energy and energy absorbed are different quantities.

Activation Energy vs Energy Absorbed

A reaction can have:

• a large activation energy but a small ΔH
• a small activation energy but a large ΔH
• similar ΔH values but very different activation energies

These quantities should always be considered separately.

A Real-World Example: Instant Cold Packs

Some instant cold packs use an endothermic dissolving process.

When the internal barrier is broken, water mixes with a substance such as ammonium nitrate.

The dissolving process absorbs energy from the surroundings.

As thermal energy is removed from the surroundings, the pack feels cold.

This is an important point:

The reaction or process does not create cold.

Instead:

energy is absorbed from the surroundings

→ surroundings lose thermal energy

→ temperature falls

Endothermic Does Not Simply Mean "Cold"

An endothermic reaction is defined by energy transfer, not by temperature alone.

A reaction is endothermic when the reacting system absorbs energy from its surroundings.

The surroundings may become colder, but that temperature change is a consequence of the energy transfer.

So the best definition is:

An endothermic reaction absorbs energy from the surroundings.

Not simply:

"An endothermic reaction gets cold."

Comparing Endothermic and Exothermic Profiles

The position of the products provides the quickest comparison.

For an endothermic reaction:

• products are above reactants
• ΔH is positive
• energy is absorbed

For an exothermic reaction:

• products are below reactants
• ΔH is negative
• energy is released

This difference can be summarized as:

Endothermic: energy goes in

Exothermic: energy goes out

Comparing Different Endothermic Reactions

Suppose Reaction A has:

Reactants = 50 kJ mol⁻¹

Products = 90 kJ mol⁻¹

Then:

ΔH = +40 kJ mol⁻¹

Reaction B has:

Reactants = 50 kJ mol⁻¹

Products = 150 kJ mol⁻¹

Then:

ΔH = +100 kJ mol⁻¹

Both reactions are endothermic.

However, Reaction B absorbs more energy because the vertical difference between reactants and products is greater.

Therefore:

larger upward energy difference → greater energy absorbed

Comparing Activation Energies

Now consider two endothermic reactions with the same ΔH.

Reaction X:

Reactants = 70 kJ mol⁻¹

Peak = 150 kJ mol⁻¹

Products = 110 kJ mol⁻¹

Reaction Y:

Reactants = 70 kJ mol⁻¹

Peak = 230 kJ mol⁻¹

Products = 110 kJ mol⁻¹

For Reaction X:

Eₐ = 150 − 70

Eₐ = 80 kJ mol⁻¹

For Reaction Y:

Eₐ = 230 − 70

Eₐ = 160 kJ mol⁻¹

Both reactions have:

ΔH = 110 − 70

ΔH = +40 kJ mol⁻¹

Therefore:

• both absorb the same amount of energy
• Reaction X has the lower activation-energy barrier
• Reaction Y has the higher activation-energy barrier

This shows that the amount of energy absorbed does not determine the activation energy.

Endothermic Does Not Mean Slow

Another common misconception is:

"Endothermic reactions are always slow."

This is incorrect.

Whether a reaction is endothermic or exothermic tells us about its overall energy change.

Reaction rate depends on factors such as:

• activation energy
• temperature
• concentration
• pressure
• surface area
• catalysts
• reaction mechanism

An endothermic reaction can be relatively fast if its activation-energy barrier is sufficiently low.

Catalysts and Endothermic Profiles

A catalyst lowers the activation energy of an endothermic reaction by providing an alternative pathway.

However, the catalyst does not change:

• the energy of the reactants
• the energy of the products
• ΔH
• the amount of energy absorbed

Therefore:

same reactant level + same product level + lower peak

The reaction remains endothermic.

Only the activation-energy barrier changes.

Worked Example 2: A Catalyzed Endothermic Reaction

Suppose:

Reactants = 100 kJ mol⁻¹

Products = 170 kJ mol⁻¹

Uncatalyzed peak = 300 kJ mol⁻¹

Catalyzed peak = 220 kJ mol⁻¹

Without catalyst:

Eₐ = 300 − 100

Eₐ = 200 kJ mol⁻¹

With catalyst:

Eₐ = 220 − 100

Eₐ = 120 kJ mol⁻¹

But:

ΔH = 170 − 100

ΔH = +70 kJ mol⁻¹

in both cases.

The catalyst lowers Eₐ but does not change the 70 kJ mol⁻¹ absorbed.

Forward and Reverse Reactions

Suppose an endothermic profile has:

Reactants = 90 kJ mol⁻¹

Peak = 250 kJ mol⁻¹

Products = 140 kJ mol⁻¹

For the forward reaction:

ΔH = 140 − 90

ΔH = +50 kJ mol⁻¹

The forward reaction is endothermic.

Forward activation energy:

Eₐ = 250 − 90

Eₐ = 160 kJ mol⁻¹

Now consider the reverse reaction.

Reverse activation energy:

Eₐ(reverse) = 250 − 140

Eₐ(reverse) = 110 kJ mol⁻¹

The reverse enthalpy change is:

ΔH(reverse) = −50 kJ mol⁻¹

Therefore:

forward reaction = endothermic

reverse reaction = exothermic

The energy changes are equal in magnitude but opposite in sign.

Why Is Reverse Activation Energy Smaller Here?

In this example, the products of the forward reaction are already higher in energy than the reactants.

Therefore, when the reaction runs backward, the starting point is closer to the transition-state peak.

This means:

Eₐ(reverse) < Eₐ(forward)

For a simple single-step endothermic profile, this relationship is expected because:

Eₐ(forward) − Eₐ(reverse) = ΔH

Reading an Endothermic Profile Step by Step

When interpreting a reaction profile, follow this sequence.

First, identify the reactant energy level.

Next, identify the product energy level.

Ask:

Are the products above the reactants?

If yes:

The reaction is endothermic.

Then calculate:

ΔH = Hproducts − Hreactants

Next, locate the peak.

Calculate:

Eₐ = Hpeak − Hreactants

Finally, compare ΔH and Eₐ with those of other reactions if required.

Worked Example 3

An energy profile shows:

Reactants = 120 kJ mol⁻¹

Peak = 350 kJ mol⁻¹

Products = 200 kJ mol⁻¹

Step 1: Identify the reaction.

Products are higher than reactants.

Endothermic

Step 2: Calculate ΔH.

ΔH = 200 − 120

ΔH = +80 kJ mol⁻¹

Step 3: Determine the energy absorbed.

80 kJ mol⁻¹

Step 4: Calculate activation energy.

Eₐ = 350 − 120

Eₐ = 230 kJ mol⁻¹

Worked Example 4: Comparing Two Reactions

Reaction A:

Eₐ = 70 kJ mol⁻¹

ΔH = +90 kJ mol⁻¹

Reaction B:

Eₐ = 150 kJ mol⁻¹

ΔH = +40 kJ mol⁻¹

Which reaction absorbs more energy?

Reaction A

because +90 kJ mol⁻¹ is a larger energy gain.

Which has the smaller activation-energy barrier?

Reaction A

Which reactions are endothermic?

Both

This illustrates why a reaction comparison should consider both ΔH and Eₐ.

Other Examples of Endothermic Processes

Common examples include:

• some dissolving processes
• thermal decomposition reactions
• photosynthesis overall
• melting
• evaporation

For example, the thermal decomposition of calcium carbonate requires continued energy input:

CaCO₃ → CaO + CO₂

Energy must be supplied to drive the process.

Photosynthesis also requires energy input, provided by sunlight, to build higher-energy molecules from lower-energy starting materials.

Did You Know?

An instant cold pack feels cold for exactly the opposite energetic reason that a hand warmer feels hot.

A cold pack contains a process that absorbs thermal energy from its surroundings.

A hand warmer contains a process that releases thermal energy to its surroundings.

So the temperature you feel is evidence of energy moving between the chemical system and the surroundings.

Common Misconceptions

"Endothermic means the reaction contains cold energy."

Incorrect. The system absorbs energy from the surroundings.

"Products are higher, so activation energy equals ΔH."

Incorrect. Activation energy is measured from the reactants to the peak. ΔH is measured from reactants to products.

"The energy absorbed is measured from the reactants to the peak."

Incorrect. That is activation energy.

"All endothermic reactions are slow."

Incorrect. Reaction rate and overall energy change are different ideas.

"A catalyst makes an endothermic reaction exothermic."

Incorrect. A catalyst lowers activation energy but does not change ΔH.

"The x-axis shows how long the reaction takes."

Usually incorrect. It represents reaction progress.

Key Terms

Endothermic reaction – A reaction that absorbs energy from the surroundings.

Energy profile – A diagram showing how the energy of a reacting system changes as the reaction progresses.

Reactants – Substances present at the beginning of a reaction.

Products – Substances formed by a reaction.

Enthalpy, H – A measure of the energy content of a system under constant-pressure conditions.

Enthalpy change, ΔH – The difference between product and reactant enthalpy.

Activation energy, Eₐ – The energy barrier that reacting particles must overcome.

Transition state – The high-energy arrangement at the peak of a reaction profile.

Reaction coordinate – A conceptual representation of progress from reactants to products.

Key Takeaways

• An endothermic reaction absorbs energy from the surroundings.
• On an endothermic profile, the products are higher in energy than the reactants.
• Endothermic reactions have positive ΔH.
• ΔH = Hproducts − Hreactants.
• The vertical difference between reactants and products represents the energy absorbed.
• Products have more energy because the reaction absorbs more energy during bond breaking than it releases overall during bond formation.
• Endothermic reactions still require activation energy.
• Activation energy is measured from the reactant level to the peak.
• Activation energy and energy absorbed are different quantities.
• Two endothermic reactions can absorb different amounts of energy.
• Two reactions can absorb the same amount of energy but have different activation energies.
• An endothermic reaction is not necessarily slow.
• A catalyst lowers activation energy but does not change ΔH.
• The reverse of an endothermic reaction is exothermic.
• Cold packs are a familiar real-world example of an endothermic process.