AP Chemistry 5.7: Catalysis

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

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

• Because the activation-energy barrier is lower, a larger fraction of collisions can successfully produce products at the same temperature.
• Therefore:
• lower Eₐ → larger fraction of successful collisions → faster reaction

A catalyst does not increase reaction rate by giving particles more kinetic energy.

It does not increase the temperature.

It does not change the overall energy difference between reactants and products.

It does not make an endothermic reaction become exothermic.

• Instead, it changes the pathway used to get from reactants to products.

Suppose the uncatalyzed reaction has a very high energy peak.

• A catalyst may create a different mechanism with lower peaks.
• The starting reactant energy stays the same.
• The final product energy stays the same.
• Therefore the overall energy change:
• ΔE or ΔH
• stays the same.
• What changes is:
• activation energy
• The catalyzed pathway has a smaller maximum barrier.

A catalyst lowers activation barriers for the pathway in both directions.

• Therefore it can speed up both forward and reverse reactions.
• This becomes especially important when equilibrium is studied later.

A catalyst helps the system reach equilibrium faster, but it does not determine which side is favored at equilibrium.

A catalyst may appear as a reactant in an early elementary step but then reappear as a product in a later step.

• Example:
• Step 1:
• A + X → AX
• Step 2:
• AX + B → AB + X
• X is consumed in Step 1 but regenerated in Step 2.
• Overall:
• A + B → AB
• X cancels from the overall equation.
• Therefore X is the catalyst.

This distinction is important enough to repeat.

An intermediate is:

• formed in one step → consumed later

A catalyst is:

• consumed in one step → regenerated later

Neither normally appears in the final overall equation.

In homogeneous catalysis, the catalyst and reactants are in the same phase.

• For example, a dissolved catalyst participating in a reaction between dissolved reactants is homogeneous.
• Because everything is in the same phase, catalyst and reactant particles can interact throughout the mixture.

In heterogeneous catalysis, the catalyst is in a different phase from the reactants.

• A common situation involves gaseous reactants interacting with a solid catalyst surface.
• A simplified pathway is:
• Reactant particles approach the catalyst surface.
• They attach to, or adsorb onto, the surface.
• Surface interactions can weaken existing bonds or place particles in favorable orientations.
• The reaction occurs on the surface.
• Products detach from the surface.

Because only catalyst surface sites are directly available, surface area can matter greatly for heterogeneous catalysts.

Enzymes are biological catalysts.

• They increase reaction rates in living systems by providing reaction pathways with lower activation-energy barriers.

An enzyme usually interacts with particular reactant molecules, often called substrates, at an active site.

• The important AP Chemistry concept is still the same:
• catalysis lowers activation energy without changing the overall energy difference between reactants and products.

At a fixed temperature, adding a catalyst does not change the Maxwell–Boltzmann kinetic-energy distribution of the molecules.

• Instead, the catalyst lowers the activation-energy threshold.
• Imagine the same energy-distribution curve, but move the Eₐ line to the left.
• Now a larger area of the curve lies beyond Eₐ.
• Therefore more collisions can successfully react.

Compare this with increasing temperature:

• Increasing temperature → changes the distribution
• Adding catalyst → lowers the Eₐ threshold
• Students often mix these up.

A catalyst provides a new mechanism with a lower activation-energy barrier.

• According to the Arrhenius relationship:
• k = Ae^(−Eₐ/RT)
• a lower Eₐ generally means a larger rate constant for the catalyzed pathway.
• That produces a faster reaction.

Do not say a catalyst is permanently used up.

Do not say a catalyst changes ΔH or ΔE.

Do not say a catalyst increases particle kinetic energy.

Do not say a catalyst shifts the energy levels of reactants and products.

Instead:

a catalyst changes the mechanism and lowers the activation-energy pathway.

A catalyst is a shortcut over a smaller energy hill.

• Same starting point.
• Same ending point.
• Different pathway.
• Lower Eₐ.
• Faster reaction.