Topic 9.4 Notes – Thermodynamic and Kinetic Control

Thermodynamic Favorability vs Reaction Rate

Two separate ideas control what you observe in a reaction.

Thermodynamics

Thermodynamics is about energy differences between reactants and products.

The key quantity is Gibbs free energy:

• $\Delta G < 0$ → reaction is thermodynamically favored (spontaneous).
• $\Delta G > 0$ → reaction is not favored.
• $\Delta G = 0$ → system is at equilibrium.

What ΔG tells you:

• Whether products are favored at equilibrium.
• The relative stability of products vs reactants.

What ΔG does not tell you:

• How fast the reaction occurs.
• Whether you’ll see it happen during class.

Students mix this up all the time. “Spontaneous” means energetically favorable, not fast.

Kinetics

Kinetics is about rate, meaning how quickly concentrations change.

A rate law looks like:

$$\text{rate} = k[A]^m[B]^n$$

• $k$ = rate constant (depends on temperature and activation energy)
• Exponents $m$, $n$ come from experiment
• Larger rate → reaction happens faster

Rate depends on:

• Activation energy
• Temperature
• Concentration
• Presence of a catalyst
• Reaction mechanism

So thermodynamics tells you where the reaction wants to go.
Kinetics tells you whether it gets there in your lifetime.

Activation Energy and the Energy Barrier

Even if products are lower in energy, particles must overcome an activation energy $E_a$ to react.

Activation energy is:

• The minimum energy needed to form the activated complex
• The height of the energy barrier

Here’s what that looks like on a reaction coordinate diagram.

Reaction coordinate diagram for an exergonic reaction

On this diagram:

• ΔG = vertical difference between reactants and products.
• E_a = height from reactants up to the peak of the curve.

The figure also shows two pathways. The higher peak is the uncatalyzed reaction, and the lower peak is the catalyzed reaction.

Notice something important:
ΔG depends only on the starting and ending points.
Ea depends on the pathway.

If $E_a$ is very large:

• Very few collisions have enough energy.
• Rate is extremely small.
• Reaction may be unobservable.

This explains why many thermodynamically favored processes barely occur.

What Kinetic Control Means

A reaction is under kinetic control when:

• $\Delta G < 0$ (products favored)
• But the reaction does not occur at a measurable rate
• Because the activation energy is very high

This does not mean:

• The system is at equilibrium.
• The reaction is nonspontaneous.
• The reaction is impossible.

It means the reaction is so slow that you don’t detect meaningful change.

AP-style questions often describe a reaction that is “thermodynamically favorable but no observable reaction occurs.” The correct reasoning is high activation energy leading to kinetic control.

How Catalysts Remove Kinetic Control

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

It changes:

• The mechanism
• The rate constant $k$
• The speed of the reaction

It does not change:

• ΔG
• ΔH
• The equilibrium constant $K$
• The energies of reactants or products

On a free energy diagram, the catalyst lowers the height of the activation barrier without changing the overall energy difference between reactants and products.

Catalyzed vs. uncatalyzed free energy diagram

The peak for the catalyzed pathway is lower, but the starting and ending energy levels are the same. That is why ΔG°_rxn is unchanged.

Lower $E_a$ means:

• More particles can overcome the barrier.
• Reaction proceeds at a measurable rate.
• The reaction is no longer under kinetic control.

This is why adding a catalyst can suddenly make a previously “dead” reaction proceed rapidly.

Thermodynamic Control vs Kinetic Control

Here’s how to keep them straight:

In this topic, the most tested idea is simple:
A reaction can be spontaneous and still not happen at a noticeable rate.