Rate of Reaction
Quick Notes
- The rate of a reaction tells us how fast reactants turn into products.
- Usually given as the change in concentration of a reactant or product per unit time.
- For a reaction to happen, particles must collide with enough energy (≥ activation energy) and in the correct orientation.
- Effective collisions lead to a reaction; non-effective collisions don’t.
- Increasing concentration or pressure leads to more frequent collisions, so reactions happen faster.
- We can use data from experiments to calculate rate using:
Rate = Δ[concentration] ÷ Δt
Full Notes
Collision theory and activation energy has been outlined with more background theory and detail here
This page is just what you need to know for CIE A-level Chemistry :)
Rate of Reaction and Collision Theory
The rate of a reaction is the change in concentration of a reactant or product per unit time.
Let’s consider a general reaction:
R → P
Where [R] and [P] represent the concentrations of reactant and product respectively.
Average Rate:
Over a time interval Δt:
Rate = – Δ[R]/Δt = Δ[P]/Δt
The negative sign indicates that the concentration of the reactant is decreasing with time.
Collision Theory
For a reaction to occur:
- Reactant particles must collide.
- They must collide with enough energy (called the activation energy, Ea).
- The collision must also be in the correct orientation (this is covered more in later topics).
If particles collide without enough energy, they bounce off each other and remain unchanged – this is called a non-effective collision.
If they collide with sufficient energy, they may react – this is an effective collision.
The rate of reaction depends on how many effective collisions occur per second.
Effect of Concentration and Pressure on Rate
Increasing concentration (in solutions) means more particles per unit volume, so collisions happen more often. This increases the rate of reaction.
Increasing pressure (in gases) has the same effect — it pushes particles closer together, so they collide more frequently.
More frequent collisions = more effective collisions per second = faster reaction.
Increasing the concentration increases the rate as the number of collisions per second increases. However, the proportion of collisions that are successful or ‘effective’ remains the same. Only changing temperature or activation energy (using a catalyst) can change the proportion of collisions that are successful.
Calculating Rate of Reaction from Data
You can calculate rate from experimental data such as:
- Loss of mass (e.g. CO2 escaping)
- Volume of gas produced
- Change in concentration (e.g. colour change or pH)
Rate = Δ[concentration] ÷ time
This gives the average rate over a time interval. To find the initial rate, use data from the first few seconds of the reaction.
In an experiment, the concentration of H2O2 drops from 0.50 mol dm−3 to 0.30 mol dm−3 in 20 seconds.
Rate = (0.50 – 0.30) ÷ 20
= 0.20 ÷ 20
= 0.010 mol dm−3 s−1
Summary
- The rate of a reaction is the change in concentration of a reactant or product per unit time.
- Reactions occur when particles collide with enough energy (≥ activation energy) and in the correct orientation.
- Effective collisions cause reaction; non-effective collisions do not.
- Increasing concentration or pressure increases frequency of collisions, speeding up the reaction.
- Rate can be measured by mass loss, gas volume, or concentration change.
- Average rate = Δ[concentration] ÷ Δt; initial rate from the first few seconds.