R2.2 How fast? The rate of chemical changeIB Chemistry SL: Revision notes
Section 1
Measuring rate
The rate of reaction is the change in concentration of a reactant or product per unit time, in mol dm⁻³ s⁻¹. In practice we follow any property that changes with concentration: volume of gas (gas syringe), mass loss as gas escapes, colour (colorimeter), conductivity, or the time for a fixed amount of precipitate to form.
The mean rate over an interval is Δ(quantity) ÷ Δt. When a fixed amount of product is formed each time, rate ∝ 1/time. Rate is highest at the start and falls as reactants are used up.
Always check the time unit: a rate in cm³ s⁻¹ is 60 times bigger as a number when written in cm³ min⁻¹.
Section 2
Collision theory
Particles react only when they collide with energy at least equal to the activation energy and with the correct orientation (collision geometry): the reacting parts of the particles must meet. Most collisions are unsuccessful.
The average kinetic energy of particles is proportional to the temperature in kelvin. A rise from 293 K to 303 K increases average kinetic energy by only about 3%, so collision frequency hardly changes.
Section 3
Factors affecting rate
- Concentration (solutions) and pressure (gases): more particles per unit volume → more frequent collisions.
- Surface area of a solid: powder exposes more particles → more frequent collisions.
- Temperature: particles move faster (more collisions) and, far more importantly, a much larger proportion of collisions have E ≥ Ea.
- Catalyst: lowers Ea so a larger proportion of collisions succeed.
Increasing the volume of a solution at the same concentration does not change the rate.
Saying higher temperature works mainly because of 'more collisions'. The key reason is the larger proportion of collisions with E ≥ Ea.
Section 4
Maxwell–Boltzmann distributions
A Maxwell–Boltzmann distribution plots the number of particles against their kinetic energy. It starts at the origin, rises to a peak (the most probable energy), then falls with a long tail that never touches the energy axis. The area under the curve is the total number of particles.
Higher temperature: the peak is lower and moves to higher energy; the area is unchanged; much more area lies beyond Ea.
Lower Ea (catalyst): the curve is unchanged; the Ea line moves left, so more area lies beyond it.
Section 5
Catalysts and energy profiles
A catalyst increases the rate by providing an alternative reaction pathway with a lower Ea; it is not used up. On an energy profile, reactants and products are at the same levels with or without a catalyst, so ΔH is unchanged; only the maximum is lower.
Exothermic: products below reactants. Endothermic: products above reactants. For an exothermic reaction, Ea(reverse) = Ea(forward) + |ΔH|; for an endothermic reaction, Ea(reverse) = Ea(forward) − ΔH. A catalyst lowers the forward and reverse Ea by the same amount.
Must know
- Rate = change in concentration per unit time (mol dm⁻³ s⁻¹).
- Successful collision: E ≥ Ea and correct orientation.
- Average kinetic energy ∝ temperature in kelvin.
- Temperature raises the proportion with E ≥ Ea; a catalyst lowers Ea.
- Catalysts do not change ΔH and lower both forward and reverse Ea.
That's the notes covered.
Carry on to the next subtopic.