When talking about collisions, the image of billiard balls hitting each other on the table may come to mind. However, molecules, although this theory assumes that their shapes are spherical, do not behave in a similar way. Molecular collisions differ in several ways, both spatial and energetic.
This theory, although it can be mathematically a bit complex and show considerable variations with respect to the experimental results, offers a picture of visible interpretations and without abstract aspects.
Conditions for reactions to occur
According to collision theory, there are three conditions for reactions to occur:
- Chemical species (ions, molecules, radicals, etc.) must collide efficiently.
- They must carry enough energy to break their bonds.
- Also, the targeting of the impact has to be very specific.
Molecular collisions share something in common with macroscopic ones: they occur predominantly between two species or bodies. Two billiard balls collide with each other, as well as a soccer ball against a wall, or two projectiles in midair. That is, collisions as far as chemistry and their reactions are concerned tend to be of the bimolecular type.
Molecules are not still, but move and rotate through the space around them. In doing so, they are assumed to draw a kind of circle called a cross section, over which there is a probability that another molecule will collide. Also, the theory considers that the molecules are spherical to simplify the mathematical models.
Two hypothetically spherical molecules can collide without any problem, even when there is no chemical reaction. But it does not happen in the same way when dealing with three, four or more molecules.
The more species that must collide to create a product, the more unlikely the phenomenon becomes. This is explained visually by trying to make three balls or projectiles collide with each other at the same time. Therefore, bimolecular collisions are by far the most common.
The collision theory is only valid for gaseous systems or phases. This is because gases show a behavior that can be well described by their kinetics.
Molecules can collide slowly or very quickly. This depends on how great its energy is, which in turn varies significantly with temperature. The stronger the collision, the probability of a chemical reaction will increase, since this energy will be able to break the necessary bonds to form new ones.
This energy is known as the activation energy, E A , and is characteristic of all chemical reactions. When the temperature is increased, the average of the molecules is able to equal or exceed E A , so the number of effective collisions and, therefore, the products formed, increase.
In the presence of a catalyst, E A decreases, since it provides surfaces and electronic means that benefit collisions. The result: the reaction rate increases, without the need to increase the temperature or add other reagents.
Chemical species do collide to react, this theory predicting how fast their reactions will be. However, experiments have shown that the more complex the reagent structures, the greater the deviations or differences between the theoretical and experimental speeds.
This is because the molecules are far from being spherical, but they have all kinds of geometries spatially. This is where the steric factor, ρ , comes in, with which it is sought to correct the reaction rates so that the predictions of the theory better agree with the experimental results.
Examples of reactions
The following reaction:
N 2 O + NO → N 2 + NO 2
It is in common use to explain what effective targeting means in molecular collisions.
The molecule N 2 O, dinitrogen oxide, will not react with NO, nitric oxide, unless during the collision the oxygen atom (red circle) of N 2 O collides directly with the nitrogen atom (blue circle) of NO . Only in this way will products N 2 and NO 2 originate ; otherwise the N 2 O and NO molecules will bounce back unreacted.
This is an example of a bimolecular collision for a bimolecular elemental reaction. Even if the reactants have enough energy to break the bonds, if the orientation is not correct, there will be no chemical reaction.
Molecular collisions can also intervene in a unimolecular elemental reaction, even when only one species undergoes the transformation or the breaking of its bonds.
Consider for example the isomerization of cyclobutane to give a mixture of butenes. As the temperature increases, the cyclobutane molecules will vibrate at higher frequencies and collide with increasing force. The products, however, demonstrate that two cyclobutane molecules do not react with each other because they would otherwise produce a compound with eight carbons.
In the middle there may be impurities (green circle), which also collide with cyclobutane, specifically on any of its CC bonds. These can be noble gases, or little reactive molecules such as nitrogen.
There will come a time when the impurity will collide with enough energy to break one of the CC bonds in the cyclobutane. And then, its molecule will seek to rearrange itself and will give rise to a butene, recognizable by its double bond and its straight chain structure.