9.4 Thermodynamic and Kinetic Control

This section examines the relationship between thermodynamic favorability and kinetics, specifically the shortcomings of Gibbs Free Energy in predicting reaction behavior. Despite being spontaneous, many reactions do not occur at an observable rate because they are under kinetic control. In this guide, we will explore what kinetic control means and why it occurs.

Before getting into the thermodynamics of kinetic control, we need to briefly review topics from Unit 5. If you feel confident with kinetics, feel free to move past this section.

Recall that kinetics is the study of the rate of a reaction—essentially, how fast a reaction occurs. We measure the rate of a reaction by measuring the change in the concentration of reactants over time: R = Δ[A]/Δt or R = -d[A]/dt (for those familiar with calculus). Higher R values indicate a quicker loss of reactants and formation of products.

Rates can also be described using rate laws, where initial reactant concentration is directly tied to the rate of a reaction as follows:

where [A]^n, [B]^m, etc., are various reactants raised to their reaction order (how much of an impact their concentration has on the rate).

Take a look at the following rate law as an example:

rate = k[A]

These laws help us describe how quickly a reaction will occur, with a higher rate implying a faster reaction overall.

Most important to our study of kinetic control is understanding activation energy. Due to the kinetic molecular theory, chemical reactions occur when molecules hit each other at the right angle and speed/energy. This activation energy is the energy required for a chemical reaction to actually occur. The higher the activation energy, the harder it is for the reaction to occur at an observable rate.

The following diagram visually shows the concept of activation energy: 

A common misconception when looking at thermodynamic favorability is that a thermodynamically favorable reaction occurs quickly. Many reactions that are spontaneous occur incredibly slowly. A good example of this concept that can be applied to the real world is the conversion from diamonds into graphite (represented as Cdiamond(s) → Cgraphite(s)).

For this reaction, ΔG° = -3 kJ. This tells us that the reaction for graphite formation occurs spontaneously and does not require the input of any external energy to occur. However, take a look at the nearest diamond (because people have those around the house…right?). Is it suddenly morphing into graphite? No! (I hope not. Fiveable is not responsible for diamond graphitification…).

The conversion of diamond into graphite is incredibly slow, as in thousands of years slow. We say that this reaction is in kinetic control because it is driven by the slowness of the reaction. These types of reactions are often slow because they have a high activation energy. 

A spontaneous process may take either the thermodynamically controlled or the kinetic controlled pathway. A kinetically controlled path like the one above is driven by a high activation energy. A thermodynamically controlled reaction is driven by the difference in free energy between the products and reactants, the type of reaction we saw in the last section. 

As we mentioned before, the primary reason for a reaction to be under kinetic control is because of a high activation energy. Because of this, even if the reaction is thermodynamically favorable, it may not continue at a measurable rate. There is a way around this, however, and that is through the use of a catalyst! Catalysts change the mechanism behind a reaction in order to decrease the activation energy and make reactions quicker. By reducing the activation energy, the reaction can proceed at a measurable rate.

For example, the decomposition of hydrogen peroxide usually occurs at an unmeasurable rate. However, when iodide ions are used as a catalyst, it creates the famous "Elephant’s Toothpaste" reaction, which you may be familiar with. 

This example shows us how catalysts can transform a reaction from one at an immeasurable rate to one at a measurable rate. The concept of catalysis can be applied to kinetic control by noticing that catalysts can help get a reaction out of kinetic control by lowering the activation energy. For example, if there was a catalyst for the reaction: Cdiamond(s) → Cgraphite(s), the reaction would be able to proceed at a measurable rate. However, without one, it cannot because the rate of the reaction is incredibly slow.