Reaction rates must be controlled in industrial processes.

→ If the rate is too low then the process will not be economically viable; if it is too high there will be a risk of explosion.

Chemists must be able to predict the quantity of heat taken in or given out in an industrial process.

→ Because thermal runaway — like in the cases of the disasters in Bhopal and Seveso — occur when the rate at which a chemical reaction releases energy exceeds the capabilities of the power plant to remove heat. This is dangerous and can lead to explosions.

→ Also because endothermic reactions require heat energy to continue which energy costs the power plant should want to minimise to maximise profits.

Calculating the rate of a chemical reaction at various stages by calculating the change in mass, volume or concentration divided by time is fine for chemical reactions where it is easy to measure such changes but in many cases it can be difficult to do this.

→ An easier method of measuring the rate involves only measuring the time — that is, the time taken to produce a fixed mass, the time taken to produce a fixed volume, the time taken for a reaction to go to completion, or the time taken to produce a colour change. This way, a reaction rate can be calculated using the formula below — i.e. a value that describes a reaction’s rate relative to other reactions which they are timed to the same fixed constant, that can be used to compare them.

$$ r=\dfrac{1}{t} $$

In the above relationship $r$ denotes relative rate and $t$ denotes time, the units of $r$ are the units of time $^{-1}$, e.g. $s^{-1}$.

Collision theory states that before a reaction can take place, the particles of the reactants must successfully collide together — with the sufficient activation energy, and the correct collision geometry.

To change the reaction rate the total number of successful collisions of a reaction, within a given time, must be altered.

The five factors below alter the reaction rate in different ways, all of which can be explained using collision theory, as below.

• In a reaction where one of the reactants is a solute, the reaction can be sped up by increasing the concentration of said reactant, and slowed down by decreasing the concentration of the reactant.

  • If the concentration of a reactant is increased then this means there are more reactant particles moving around within the same volume, increasing the probability of particles colliding, so the particles collide more frequently — i.e. increasing the total number of total collisions within a given time — which increases the number of successful collisions within a given time, which increases the rate of the reaction.
  • If the concentration of a reactant is decreased then this means there are less particles reactant particles moving around within the same volume, decreasing the probability of particles colliding, so the particles collide less frequently — i.e. decreasing the total number of total collisions within a given time — which decreases the number of successful collisions within a given time, which decreases the rate of the reaction.

• In a reaction where one of the reactants is gaseous, increasing the pressure increases the reaction rate and decreasing the pressure decreases the reaction rate.

  • If the pressure of gaseous reactants is increased, there are more reactant particles for a given volume, increasing the probability of particles colliding, so particles will collide more frequently — i.e. increasing the total number of collisions in a given time — which increases the number of successful collisions within a given time, which increases the reaction rate.
  • If the pressure of gaseous reactants is decreased, there are less reactant particles for a given volume, decreasing the probability of particles colliding, so particles will collide less frequently — i.e. decreasing the total number of collisions in a given time — which decreases the number of successful collisions within a given time, which decreases the reaction rate.

• Reactions in which one of the reactants is a solid can be sped up by decreasing the particle size of the solid, and slowed down by increasing the particle size of the solid.

  • Decreasing the particle size of a reactant increases the total surface area presented to other reactants, which means particles collide more frequently — i.e. increases the total number of collisions within a given time. Hence, the total number of successful collisions within a given time is increased, which increases the rate of the reaction.
  • Increasing the particle size of a reactant decreases the total surface area presented to other reactants, which means particles collide less frequently — i.e. decreases the total number of collisions within a given time. Hence, the total number of successful collisions within a given time is decreased, which decreases the rate of the reaction.

• Increasing the temperature of a reaction increases the rate of the reaction, and decreasing the temperature decreases the rate of the reaction.

  • If the temperature of a reaction is increased then the average kinetic energy of the reactant particles is increased.

    → This means the particles move around more quickly and collide more frequently — i.e. the total number of collisions within a given time increases, which increases the number of successful collisions within a given time, which increases the rate of the reaction.

    → It also means that the proportion of particles that have energy greater than (or equal to) the activation energy increases, which increases the probability of each collision that occurs being a successful one, which increases the proportion of successful collisions, increasing the number of successful collisions within a given time, which increases the rate of reaction.

  • If the temperature of a reaction is decreased then the average kinetic energy of the reactant particles is decreased.

    → This means the particles move around less quickly and collide less frequently — i.e. the total number of collisions within a given time decreases, which decreases the number of successful collisions within a given time, which decreases the rate of the reaction.

    → It also means that the proportion of particles that have energy greater than (or equal to) the activation energy decreases, which decreases the probability of each collision that occurs being a successful one, which decreases the proportion of successful collisions, decreasing the number of successful collisions within a given time, which decreases the rate of reaction.

• The collision geometry of particles — the orientation with which they collide — also impacts the rate of the reaction, as some collisions will not produce a successful reaction even if the particles collide with the necessary $E_a$. Favourable collision geometry of reactants results in a successful reaction being being more likely and dis-favourable collision geometry of reactants results in a successful collision being less likely.

  • Collision geometry is more favourable is where the reactive parts of the particles are colliding — the bonds are parallel and hit each other, as in (b). This increases the proportion of collisions that are successful, increasing the number of successful collisions within a given time, which increases the rate of reaction.

  • Collision geometry is less favourable where an unreactive part of a molecule hits a reactive part — where the bonds are perpendicular-ish, as in (a). This decreases the proportion of collisions that are successful, decreasing the number of successful collisions within a given time, which decreases the rate of reaction.

    

    

    

The activation energy is

→ the minimum energy required by colliding particles to form an activated complex.

→ the minimum kinetic energy required by colliding particles before a reaction may occur.

Temperature is a measure of the average kinetic energy of the particles in a substance.

Another way the rate of a reaction can be increased is by adding a catalyst.

A catalyst provides an alternative reaction pathway with a lower activation energy.

This means that the proportion of particles with energy greater than (or equal to) the activation energy increases, which increases the probability of each collision that occurs being a successful one, which increases the proportion of successful collisions, increasing the number of successful collisions within a given time, which increases the rate of reaction.

Potential energy diagrams show the pathway of a reaction — this is different for endothermic and exothermic reactions.

example exothermic potential energy diagram

The $y$-axis has the potential energy of the reactants and products (or their enthalpy — $H$) and the $x$-axis shows the pathway of the reaction.

The peak itself is the where the activated complex is formed.

The activated complex is an unstable arrangement of atoms formed at the maximum of the potential energy barrier, during a reaction.

potential energy diagram with activated complex marked

Exothermic reactions are those where heat energy is given out to the surroundings. In these reactions the reactants have more energy than the products and the enthalpy change is negative.

exothermic potential energy diagram

Endothermic reactions are those where heat energy is taken in from the surroundings. In these reactions the reactants have less energy than the products and the enthalpy change is positive.

endothermic potential energy diagram

example potential energy diagrams for exothermic and endothermic reactions