8.6 Catalyst

This topic explains what a catalyst is, how it works, and why it is important in chemistry and industry.

Definition

A catalyst is a substance that accelerates a chemical reaction but remains chemically unchanged at the end of the reaction. The phenomenon is called catalysis.

Simple explanation: A catalyst speeds up a reaction without being used up itself.

Main Idea

Many industrial reactions are carried out at high temperatures over a period of time to maximize the amount of product that can be synthesized.

However, high-temperature reactions cause safety problems, and many chemicals are not stable at high temperatures.

So it would be useful to use another method to increase the speed of chemical reactions.

Another way to increase the reaction rate is to change the mechanism in a way that lowers the activation energy. This can be done by adding a catalyst.

Simple explanation: Instead of using high temperatures (which can be dangerous), we can use a catalyst to speed up reactions safely.

Core Concept

How a Catalyst Works

• A catalyst provides a new mechanism for the reaction with lower activation energy.
• Thus, a catalyst increases the rate of a reaction by lowering its activation energy.
• The catalyst does not affect the overall thermodynamics or enthalpy of the reaction.

Simple explanation: A catalyst makes it easier for the reaction to happen by lowering the energy barrier.

Key Points About Catalysts

Catalysts in Living Organisms

In the bodies of living organisms, enzymes catalyze chemical reactions.

Simple explanation: Enzymes are biological catalysts — they speed up reactions in your body.

Detailed Explanation

Why Catalysts Are Needed

Many industrial reactions are carried out at high temperatures over long periods to maximize product yield.

• Safety problems — high temperatures can be dangerous.
• Stability problems — many chemicals are not stable at high temperatures.

Solution: Use a catalyst to lower the activation energy, so the reaction can proceed faster at lower temperatures.

Simple explanation: Catalysts allow reactions to happen faster without needing extreme conditions.

How Catalysts Lower Activation Energy

• A catalyst provides an alternative reaction pathway.
• This alternative pathway has a lower activation energy than the uncatalyzed reaction.
• With a lower activation energy, more particles have enough energy to react.
• So the reaction proceeds faster.

Simple explanation: The catalyst creates a shortcut over the energy hill.

Catalysts Do Not Affect ΔH

• Does not affect the overall thermodynamics or enthalpy of the reaction.
• The ΔH (energy difference between reactants and products) remains the same.
• Only the activation energy is lowered.

Simple explanation: A catalyst changes how fast you get there, not where you end up.

Enzymes — Biological Catalysts

• In living organisms, enzymes catalyze chemical reactions.
• Enzymes are proteins that speed up biochemical reactions.
• They work in a similar way to industrial catalysts — by lowering activation energy.

• Amylase — breaks down starch into sugars.
• Protease — breaks down proteins into amino acids.
• Lipase — breaks down fats into fatty acids and glycerol.

Simple explanation: Enzymes are nature's catalysts — they keep your body's chemistry running smoothly.

Effect of Catalyst on Activation Energy (Figure 8.3)

• Reaction without catalyst — higher activation energy.
• Reaction with catalyst — lower activation energy.
• The products end up at the same energy level in both cases (ΔH is unchanged).

Simple explanation: The catalyst lowers the hill, but the starting and ending points stay the same.

Intermediate Level

How Catalysts Work — In Detail

• Reactant molecules approach the catalyst surface (for heterogeneous catalysts).
• Bonds in reactant molecules weaken as they interact with the catalyst.
• New bonds begin to form with lower energy requirement.
• Products are released from the catalyst.
• The catalyst is regenerated and ready for the next reaction.

Simple explanation: The catalyst provides a surface or environment where the reaction can happen more easily.

Types of Catalysts

Effect of Catalyst on Reaction Rate

• Without catalyst: High activation energy → fewer particles have enough energy → slow reaction.
• With catalyst: Lower activation energy → more particles have enough energy → fast reaction.

Simple explanation: Lower barrier = more particles can get over = faster reaction.

Catalysts and Enthalpy (ΔH)

• The enthalpy change (ΔH) of the reaction.
• The position of equilibrium.
• The overall thermodynamics.

It only lowers the activation energy and speeds up the reaction.

Simple explanation: A catalyst doesn't change the "before" and "after" — it only changes how fast you get there.

Industrial Importance of Catalysts

• Reduce energy costs (lower temperatures needed).
• Increase production rate.
• Improve safety (lower temperatures and pressures).
• Enable specific reactions (selectivity).

Examples from the book:

Important Points

• - Catalyst: substance that accelerates a chemical reaction but remains chemically unchanged at the end.
• - Catalysis: the phenomenon of speeding up a reaction with a catalyst.
• - How it works: provides a new mechanism with lower activation energy.
• - Effect on rate: increases the rate of reaction.
• - Effect on ΔH: no effect — ΔH remains the same.
• - Effect on equilibrium: no effect — only speeds up reaching equilibrium.
• - Enzymes: biological catalysts in living organisms.
• - Industrial importance: reduces energy costs, increases production, improves safety.
• - Examples:
• - Iron (Haber process)
• - Platinum, palladium (catalytic converters)
• - Nickel (margarine manufacture)
• - Platinum (Contact process for sulphuric acid)

Remember

• - Catalyst = speeds up reaction, not used up.
• - Lowers activation energy → faster reaction.
• - Does not affect ΔH.
• - Does not affect equilibrium position.
• - Enzymes = biological catalysts.
• - Examples:
• - Fe → Haber process
• - Pt, Pd → catalytic converters
• - Ni → margarine
• - Pt → Contact process

Common Mistakes

• 1. Thinking a catalyst is used up in the reaction. It remains chemically unchanged at the end.
• 2. Thinking a catalyst changes ΔH. It does not — ΔH stays the same.
• 3. Thinking a catalyst changes the equilibrium position. It only speeds up reaching equilibrium.
• 4. Thinking a catalyst increases activation energy. It lowers activation energy.
• 5. Confusing catalysts with enzymes. Enzymes are biological catalysts; catalysts can be inorganic or organic.
• 6. Thinking catalysts are only used in industry. They are also used in living organisms (enzymes) and in everyday products (catalytic converters).

Quick Revision

• Catalyst: a substance that accelerates a chemical reaction but remains chemically unchanged at the end of the reaction. The phenomenon is called catalysis.
• How it works: the catalyst provides a new mechanism for the reaction with lower activation energy.
• Effect: increases the rate of reaction by lowering activation energy.
• No effect on: overall thermodynamics or enthalpy of the reaction.
• Enzymes: biological catalysts in living organisms.
• Industrial importance: many industrial reactions use catalysts to reduce the need for high temperatures (which cause safety problems and are unsuitable for unstable chemicals).
• Examples:
• Iron — Haber process for ammonia synthesis
• Platinum and palladium — catalytic converters
• Nickel — manufacture of margarine
• Platinum — Contact process for sulphuric acid

60-Second Revision

• Catalyst = speeds up reaction; not used up.
• Lowers activation energy.
• Does not change ΔH.
• Enzymes = biological catalysts.
• Examples: Fe (Haber), Pt/Pd (converters), Ni (margarine).
• Warning: a catalyst does not change the enthalpy or equilibrium position — it only speeds up the reaction.