Transition State — Definition & Importance

(Video) Transition State vs Intermediate

by Denovo Tutor

In this video, Denovo Tutor clarifies the difference between a transition state and an intermediate.

We've been studying kinetics. At first, it was very math focussed. We investigated reaction rates, the rate law, and determining the rate law from experimental data. But, we're beginning to think a little more broadly about how chemical reactions happen. When we studied collision theory, we learned that molecules need the correct energy and orientation. If they have enough energy but the wrong orientation, nothing happens. If they have the correct orientation but too little energy nothing happens. These are the assumptions that led to the development of the Arrhenius Equation. But, his equation was largely based on experimental data. With the development of quantum mechanics in the 1930s, physical chemists began to build a theoretical framework for how reactions occur. That framework developed into Transition State Theory.

What is a transition state?

A transition state is a momentary, unstable configuration of atoms. On a reaction coordinate diagram, they occur at energy maxima. Do not confuse them with an intermediate. Intermediates exist at energy minimums. A transition state possesses three important qualities:

• contains partial bonds

• has a very short lifetime

• cannot be isolated

In a chemical reaction, the arrangement of atoms that exists during this transition state is known as the activated complex. The molecules have collided. The bonds are partially broken and partially formed. So, the system is incredibly unstable. That's why it's at the highest energy state. At this point, the complex has the potential to either backtrack and re-form the reactants or to proceed and form the products. Because it can backtrack, it's similar to equilibrium. But remember, the complex is neither the reactant nor its product. It actually resembles both to some extent.

Imagine that moment in a relay race where the baton is in passing. Neither runner has complete possession of the baton, but both share it for a just small moment. In this analogy, the baton actually refers to electrons. In the transition state, the activated complex is partially sharing its electrons with the reactants, but also partially with the products--this is where partial bonds are seen.

The presence of these partial bonds causes the complex to be highly unstable, and in turn, highly unhappy. Complexes that are unhappy tend to have extremely short lifespans, measured in femto seconds ($10^{-15}$), and are difficult to isolate.

Furthermore, this point of maximum energy corresponds to the activation energy ($E_{A}$) of a reaction. It's the difference in energy between the reactants and the transition state.

This is an example of a reaction coordinate diagram, a type of energy diagram. The components that you should be familiar with include:

• the x-axis represents the reaction progress in the general direction of reactants (starting materials) to products.

• If you take more advanced chemistry, like organic chemistry, you'll learn that we can use bond length, bond angle, or radius as our reaction coordinates. But, in general chemistry, reaction progress is all you need to know.

• the y-axis records the potential energy and has units of joules.

• the energy curve shows the energy states of the varying components at different points during the reaction.

• transition states are seen at energy maximums.

• reaction intermediates are found at energy lows between the transition states.

• the activation energy $E_{A}$ is the difference in energy between the starting materials and the transition state.

• Gibb's Free Energy ($\Delta$ G) is the difference in energy between the starting material and the products. This is also how we tell if a reaction is exothermic or endothermic.