Kinetics is the study of how fast reactions occur.
How quickly a reaction happens is called its reaction rate.
Reactions occur when reactant particles collide together with enough energy (activation energy), creating a successful collision.
Increasing the frequency of successful collisions increases the rate of reaction.
Increasing the temperature, concentration and pressure of reacting particles increases the rate of a reaction.
The study of kinetics (in chemistry) is concerned with the movement of particles and how quickly reactions happen. Altering certain conditions can change the speed of reactions. The speed at which a reaction occurs is called the rate of reaction.
Reactions occur between different particles that are constantly moving. All particles are in constant motion (even in solids, particles are constantly vibrating – albeit with limited motion). Therefore, particles are constantly ‘bumping’ into each other. This ‘bumping’ we call collisions.
If two particles collide into each other hard enough, a reaction takes place – this is called a successful collision. How ‘hard’ a collision is depends upon how much energy each particle has when they collide. The minimum amount of energy that results in a successful collision is called activation energy (Ea).
The majority of collisions between particles are not successful, as the particles colliding do not collide with enough energy to react (activation energy).
Activation energy is fixed for a particular reaction and cannot be changed (unless a catalyst is used, see below). So, when we want to increase the rate of a reaction, we have to give the reacting particles more energy. This means that when they collide, they have the required activation energy. This is why reactions (usually) happen faster at higher temperatures.
Effect of temperature on reaction rate:
Imagine the reaction A + B → C. If we carry out this experiment at 25°C and at 75°C, at which temperature would we expect the reaction to happen fastest?
At 75°C, because if the temperature is higher particles A + B would have more energy. When particles collide with more energy, more successful collisions happen (collisions with the required activation energy). In other words, more reactions occur between particles. It is important to note that at both 25°C and 75°C successful collisions are occurring, it is just that at 75°C more successful collisions occur in a given amount of time – so the reaction is happening faster.
Effect of concentration on reaction rate:
The rate of reaction is a measure of how many successful collisions are happening in a given amount of time (frequency of successful collisions). Increasing the reaction rate means we have to increase the number of successful collisions. By increasing the total number of collisions that are occurring in a given amount of time, we are also increasing the number of successful collisions that are occurring. It is just a game of odds!
If the concentration of colliding particles is increased, there are more particles in the same amount of space. This means there will be more collisions – if you increase the number of people in a room and they are moving around, there will be more collisions because there are more people to collide with.
By increasing the concentration of reacting particles, the rate of a reaction is increased. The reacting particles are not given more energy, but there is simply an increase in the chances of a successful collision.
Effect of pressure on a reaction rate:
In a gaseous system, increasing the pressure acting on the reactants effectively forces them closer together. This gives the same effect on the rate of reaction as increasing the concentration of reactants.
Remember – concentration is just a measure of how much of substance you have in a given volume.
By forcing particles closer together, there will be more particles in the same volume as before the pressure was increased. Therefore, the number of the particles in that volume increases.
A higher number of particles colliding gives a greater chance for successful collisions. More successful collisions in a given amount of time increases the rate of the reaction.
Effect of a catalyst on a reaction rate:
The activation energy of a particular reaction route is fixed – it does not change. If we want particles with lower energies to successfully collide, the only thing we can do is find another way for the reaction to happen. Rather like going for a hike – if you don’t want to climb over a mountain to get to the other side then you need to find another way, as you cannot change the mountain.
A catalyst is a species that can provide an alternative pathway for a reaction to happen. To use the above example, imagine the catalyst as being a buggy that can drive you around the mountain, getting you to the same destination but saving you a lot of energy. A catalyst remains unchanged at the end of a reaction (it can actually change during in the reaction but it will reform at the end, so it is not considered a reactant).
By providing a lower energy alternative pathway, the catalyst lowers the activation energy required for a reaction to occur. This means particles’ collisions that happen at a lower energy now have sufficient energy to react, and thus there are more successful collisions.
A catalyst does not alter the energy of the particles colliding, but simply enables collisions that happen at lower energy to become successful. These collisions were still occurring before a catalyst was used, but whereas before they were not successful, now they are!
The rate of reaction is how many successful collisions are happening in a given amount of time. If we lower the activation energy, we increase the number of successful collisions occurring – so a catalyst increases the rate of a reaction.