Reactions occur by molecules colliding. However only a small fraction of these collisions lead to a reaction. For a collision to be effective molecules must collide in the correct orientation and must have sufficient energy. The minimum energy needed is called the activation energy.
- In an exothermic reaction, reacting chemicals lose energy and give it out to the surroundings, so the products have a lower energy than the reactants.
- In an endothermic reaction the enthalpy of the products is more than the enthalpy of the reactants and heat is taken in.
ΔH = Eaf - Eab
Where Eaf is the activation energy of the forward reaction
Where Eab is the activation energy of the reverse reaction
For an exothermic reaction, Eaf is smaller than Eab and ΔH is negative.
For an endothermic reaction, Eaf is bigger than Eab and ΔH is positive
Effect of Concentration (pressure) on reaction rates
If the concentration of the reactant increases, the reaction rate increases. There are more molecules in the same volume so there is a greater chance of collision, and therefore a greater chance that there will be more collisions with energy greater than the activation energy.
For a solid increasing surface area has the same effect
In a gaseous reaction, increasing the pressure is the same as increasing the concentration.
Effect of temperature on reaction rate
If the temperature of a reaction increases, so does the rate. At higher temperatures molecules have a higher kinetic energy and are moving faster, meaning more collisions and more molecules with energy greater than the activation energy.
The areas under the two curves are equal and proportional to the total number of molecules in the sample.
- The curves do not touch the energy axis
- At as higher temperature, T2, the peak moves to the right (higher energy) with a lower height.
- Only the molecules with energy equal to or greater than the activation energy, Ea, will be able to react.
- At the higher temperature, T2, many more molecules have sufficient energy to react and so the rate increases significantly.
- A catalyst lowers the activation energy of a reaction by providing an alternative route for the reaction.
- A catalyst increases the rate of a chemical reaction without being used up. A catalyst does take part in the reaction but is unchanged and can be recovered at the end.
At the same temperature, a greater proportion of the reactant molecules will have sufficient energy to overcome the activation energy for a catalysed reaction.
For a reversible reaction, a catalyst increases the rate of the forward and backward reactions by the same amount, therefore it does not affect the position of equilibrium, but the position of equilibrium is reached quicker.
There are two types of catalyst: heterogeneous and homogeneous.
A heterogeneous catalyst is in a different phase from the reactants, i.e. a different form (solid, liquid, gas.)
A homogeneous catalyst is in the same phase as the reactants. These catalysts take an active part in the reaction rather than being an inactive spectator.
This is the end of the topic!
Drafted by Cherry (Chemistry)