Mastering Activation Energy in Chemical Reactions

Dive into the world of activation energy and discover its crucial role in chemical reactions. Learn how it influences reaction rates, the impact of catalysts, and its applications in various fields.

In this lesson, we will learn:

• To use the potential energy curve to describe reactant to product transition in chemical reactions.
• The difference between a transition state and an intermediate in chemical reactions.
• The definition of successful collision and the conditions for it to occur in a chemical reaction.

• In the 'Introduction to kinetics' (C12.1.1) lesson we introduced activation energy. It is the minimum energy that reactant molecules need to successfully collide and start changing from reactants to products. It is an obstacle that must be overcome for any reaction and rate of reaction to be measured.
  Here, we will talk about what happens to molecules when that obstacle is overcome.
  
• The chemical reaction actually occurs at the peak of the curve marked E_a on an energy diagram – when the reactant molecules have enough potential energy to form an activated complex. This can either be a specific intermediate or just a transition state.
• The transition state is a temporary high energy state seen during the breaking of reactant bonds and forming of product bonds – these are extremely short lived and can't be isolated as chemical substances on their own.
• An intermediate is an unstable short-lived state the reactants take when reactant bonds have broken, but product bonds have not formed. Not all reactions have intermediates. Unlike transition states, intermediates can sometimes be isolated as they have a specific lifetime, though it may be extremely short, before the products are formed. An intermediate is shown by a small 'well' cut into the peak of the potential energy curve.
• Potential energy diagrams during reactions can be explained in the following way. Events marked as points on the potential energy curve are highlighted in red:
• Moving molecules have energy in the form of motion (AKA kinetic energy), so they will collide with each other. Remember that molecules at room temperature and pressure generally collide a few billion times per second!
• Before they collide, when they get near each other repulsion between electrons in the separate molecules will start happening (in a similar way to when London forces start to form). This reduces the energy of motion in the molecules; this reduced 'kinetic energy' is stored as potential energy in the electrons. This is where an increase in potential energy begins in reactant molecules.
• The faster the motion of the molecules originally was, the greater the increase in potential energy this will cause. If this rise in potential energy does not match the activation energy, nothing will happen. The molecules will eventually repel back, the excited electrons will repel each other again and send the unchanged reactant molecules back in motion in opposing directions.
• If the molecules do reach a sufficiently high potential energy then the molecules, still close together, can form an activated complex or transition state. This is where the curve flattens out at the top of the potential energy 'hill'.
• Once this occurs, the chemical bonds get rearranged. The reactant bonds are broken and the bonds of the product made, in a relatively short time period.
• The newly made products begin repelling each other by their electrons, moving away from each other. This reduces potential energy in the electrons, which is now being used again as kinetic energy. This is where the curve begins dropping again, as PE decreases and kinetic energy increases.
• As the molecules get further away, the remaining potential energy from the electron repulsion is 'converted' to kinetic energy until the products reach their 'ground state', their lowest energy state available, like in the reactants before the molecules collided.
• Molecules that go through the reactant to product process described above have undergone a successful collision. There are two conditions for a successful collision. A molecule must have:
• Sufficient energy – the activation energy required as explained above.
• Correct orientation – the molecules have to be arranged in space in an ideal way that the reactant bonds will break and product bonds form when colliding. If they are not, the activation energy for the reaction will be significantly higher. Activation energy on PE curves is shown assuming there is perfect orientation – its lowest possible value.
• In the lesson on the Boltzmann Distribution (C12.1.5) we saw the distribution of molecules and the kinetic energy they possess. The activation energy barrier is directly related to the rate of reaction through this – the lower the activation energy, the more likely any given molecule will possess enough kinetic energy that can be stored as potential energy when molecules approach and collide. Therefore, the lower the activation energy 'hill', the faster the reaction.