Catalysts play a major role in chemical kinetics. They are used across industry, biology, and even everyday processes because they make reactions happen more quickly and efficiently. In IB Chemistry, the effect of a catalyst on activation energy is a frequently tested idea in Topic 6. This article explains how catalysts work, why activation energy matters, and how reaction pathways change when a catalyst is present.
What Is Activation Energy?
Activation energy (Ea) is the minimum amount of energy particles need to collide successfully and react.
Even if molecules collide, the reaction only occurs if they collide with energy equal to or greater than the activation energy. This is why heating a reaction speeds it up—particles gain more kinetic energy, increasing the chance of successful collisions.
Activation energy is represented on an energy profile diagram as the “hump” that reactants must climb before forming products.
What Does a Catalyst Do?
A catalyst lowers the activation energy of a reaction by providing an alternative reaction pathway.
Key points:
- Catalysts speed up the rate of reaction
- Catalysts do not get used up
- Catalysts do not change ΔH (enthalpy change)
- Catalysts increase the number of successful collisions
Because more particles now have enough energy to react, the reaction rate increases significantly.
How Catalysts Lower Activation Energy
Catalysts affect reaction pathways in the following ways:
1. They provide an alternative mechanism
Catalysts introduce steps with lower energy requirements.
For example, heterogeneous catalysts may adsorb reactants onto a surface, weakening bonds and reducing the energy needed to break them.
2. They stabilize transition states
By lowering the energy of the activated complex (transition state), catalysts reduce the energy barrier.
3. They improve molecular orientation
Catalysts bring reactants into better alignment, increasing effective collision frequency.
The combination of these effects results in a lower activation energy and a higher reaction rate.
Energy Profile Diagrams With and Without a Catalyst
In an energy diagram:
- The uncatalyzed pathway shows a high activation energy peak.
- The catalyzed pathway shows a lower peak.
- The overall enthalpy change (ΔH) remains unchanged.
This diagram visually demonstrates why catalysts increase the rate without altering the final energy difference between reactants and products.
Biological Catalysts: Enzymes
Enzymes are natural catalysts found in living organisms.
They:
- Bind to substrates at active sites
- Lower activation energy dramatically
- Work under mild temperatures and pH conditions
- Are highly specific to particular reactions
Enzymes illustrate how powerful catalytic effects can be, enabling life’s essential biochemical processes.
Importance of Catalysts in Industry
Catalysts are essential for:
- Haber process (NH₃ production)
- Contact process (H₂SO₄ production)
- Polymer production
- Hydrocarbon cracking
- Environmental catalysis (catalytic converters)
Without catalysts, these processes would require extreme temperatures or pressures, making them inefficient or economically unfeasible.
Why Catalysts Don’t Affect Equilibrium Position
A common misconception is that catalysts shift equilibrium.
They do not.
Catalysts speed up both forward and reverse reactions equally, helping equilibrium reach its position faster but not changing that position.
Equilibrium depends on ΔG, ΔH, and temperature—not on activation energy.
FAQs
Does lowering activation energy always make a reaction faster?
Yes. With lower activation energy, more particles meet the required threshold for reaction, so the reaction rate increases.
Can a catalyst make a reaction spontaneous?
No. Catalysts do not change ΔH or ΔG. They make reactions faster but cannot make non-spontaneous reactions occur on their own.
Do catalysts always stay unchanged?
Catalysts are not consumed, but they can become poisoned or deactivated, especially in industrial environments.
Conclusion
Catalysts lower the activation energy of a reaction by offering an alternative pathway, stabilizing transition states, and improving molecular orientation. This significantly increases the reaction rate without changing the overall enthalpy or equilibrium position. Understanding this mechanism is essential for success in IB Chemistry and for applying kinetic principles in both laboratory and industrial settings.
