Collision theory is a mathematical model for predicting the speeds of chemical reactions, especially in glasses. The collision hypothesis assumes that in order for a reaction to occur, the reacting components (atoms or molecules) must come together or collide with one another. However, not all collisions result in chemical change. Only if the species brought together have a minimum value of internal energy equal to the reaction’s activation energy would a collision be effective in creating chemical change.
Moreover, the colliding components must be orientated in a way that allows for the essential atom and electron rearrangement. As a result, the pace of a chemical reaction is equal to the rate of effective collisions, according to the collision hypothesis. Because atomic or molecular collision frequencies can only be estimated with some precision for glasses (using the kinetic theory), the collision theory’s use is confined to gas-phase processes.
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Collision Theory of Chemical Reactions
“The molecules of reactants are supposed to be hard spheres, and reactions are thought to occur only when these spheres (molecules) clash with each other,” according to the collision hypothesis. As a result, it was necessary to measure the number of collisions that occurred in order to produce products in order to have a full picture of the reaction, and so the term collision frequency was coined.
The amount of collisions per second per unit volume of the reacting substances is known as collision frequency. Z is a common abbreviation for it.
Let’s consider the following bimolecular elementary reaction:
P + Q → Product
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Now as per the collision theory, the rate of the above reaction can be given by:
Rate=ZPQ ρ e−Ea/RT
Where:
- ZPQ = collision frequency of reactants P and Q
- Ea = Activation Energy
- R = Universal Gas Constant
- T = Temperature in absolute scale
- ρ = is the steric factor
The activation energy is another factor that has a substantial impact on the speeds of chemical processes (Ea). Arrhenius coined the phrase activation energy, which refers to the minimal amount of energy required for reactants to generate a product during a chemical reaction.
All molecules with an energy higher than or equal to the chemical potential will collide to produce products, according to the Arrhenius Equation. However, this was not the case for all reactions. In reactions involving complicated compounds, there was a large level of variation.
Some molecules with enough energy (activation energy) to generate the product did not collide. Only a handful of them were able to create successful collisions, which resulted in the production of products. The scientists discovered that the reaction is governed by more than just the kinetic energy of the molecules.
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They came to the conclusion that only molecules with the right orientation and threshold energy (activation energy) will produce products during the collision. To account for the effective collisions, they proposed a probability factor P.
Activation Energy and Temperature
Two billiard balls clash and bounce off of one other. When two molecules, A and B, come into touch, this is also the most likely outcome: they bounce off each other, fully unmodified and undamaged. A and B must collide with enough energy to break chemical bonds for a collision to be successful in causing a chemical reaction. This is because chemical bonds in the reagents are broken and new bonds in the products are generated in any chemical reaction.
As a result, in order to properly commence a reaction, the reactants must be travelling quickly enough to impact with enough force to break bonds. The activation energy is the minimal energy with which molecules must move in order for a collision to occur in a chemical reaction.
Arrhenius Equation
The activation energy is the minimal amount of energy required to generate a product when reactants collide. Because the energy required to generate a product is given by a collision of a component of the mixture with another reactant molecule, the kinetic energy of reactant molecules is crucial in a reaction. (A collision of the reactant molecule with the reaction vessel wall or with molecules of an inert contaminant can generate activation energy in single-reactant reactions.)
The reaction will take a long time if the activation energy is considerably higher than the average kinetic energy of the molecules: Only a small number of fast-moving molecules will be able to respond. If the activation energy is substantially lower than the average kinetic energy of the molecules, the proportion of molecules with the required kinetic energy will be high; most molecule collisions will result in a reaction, and the reaction will happen quickly.
Conclusion
In this article, we studied in detail the concept of collision theory. The concept of collision theory is in itself a vast subject that covers collision theory of bimolecular reaction, etc. Collision theory describes how chemical reactions take place in part. To react, chemical compounds must collide with one another, but not every collision results in a reaction.
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