3.1.5 Kinetics

With collision theory, ask two things about a collision: did the particles meet with at least the activation energy, and was the orientation suitable where that matters? The rate depends on how often those successful collisions occur.

Enough energy means collision energy at least equal to Eₐ. Collision frequency matters too, but a collision below the activation threshold does not produce products.

On a Maxwell–Boltzmann diagram, draw Eₐ as a vertical threshold. Molecules to its right have sufficient energy; molecules to the left do not.

If you are explaining a rate change, use Eₐ as the threshold: ask what fraction of particles has E ≥ Eₐ, then connect that fraction to the number of successful collisions.

I’d learn it as the minimum energy that colliding particles must have for a reaction to occur. Do not define it as the energy “released” or as the total energy of the reactants.

For the definition, keep both ideas: minimum energy and colliding particles. That wording stops activation energy drifting into a vague idea of “energy needed by the reaction”.

For an unsuccessful collision, check two possible reasons: the particles may not have enough energy or, where orientation matters, they may meet the wrong way. At a given temperature only part of the energy distribution lies at E ≥ Eₐ.

There are two filters on a collision: enough energy and a suitable orientation. If either fails, the particles separate without reacting, even though a collision occurred.

With a Maxwell–Boltzmann curve, treat it as a distribution of molecular energies at one temperature. The total area represents the number of molecules, and the peak tells you the most probable energy.

The curve starts at the origin and approaches the axis at high energy without touching it. Molecules have a range of energies; the peak is not a special energy shared by most molecules.

When temperature rises, the Maxwell–Boltzmann curve becomes lower and broader and its peak moves to higher energy. The total area stays the same if the number of molecules is unchanged.

Look especially at the area with E ≥ Eₐ: that fraction increases substantially at higher temperature, which is why the rate effect is much larger than the modest shift in average energy might suggest.

Rate is change in concentration of a reactant or product per unit time, with sign convention chosen so reaction rate is reported positively. A graph gradient can therefore be used to estimate rate.

On a concentration–time graph, you can obtain a mean rate from Δconcentration/Δtime or an instantaneous rate from the gradient of a tangent. Keep the units consistent with the axes.

When you explain why heating speeds a reaction up, look at the area beyond Eₐ. The distribution becomes broader and lower and its peak moves to higher energy, so a much larger fraction of molecules now has E ≥ Eₐ.

Higher temperature makes particles move faster and collide more frequently, but the larger effect is the increased fraction with E ≥ Eₐ. That is the part worth emphasising in a full explanation.

Shade or compare the area to the right of Eₐ. Even a modest shift in the distribution can greatly increase that area, which is why rate can rise sharply with temperature. The total area under the curve is unchanged: you have redistributed particle energies, not created more particles.

Increasing concentration puts more reacting particles into the same volume, so collision frequency rises. If the energy distribution and reaction pathway are otherwise unchanged, successful collisions occur more often per second.

Finish the explanation at successful collisions. “More particles” alone does not yet explain why the measured reaction rate increases.

Increasing the pressure of reacting gases pushes the particles into a smaller volume, increasing collision frequency. That gives more successful collisions per second and a faster reaction.

Keep this pressure explanation for gases. For solutions, concentration is the appropriate way to describe how many reacting particles occupy a given volume.

Higher concentration, or higher pressure for gases, puts reacting particles closer together. Collision frequency rises, so successful collisions occur more often per second.

For gases, increasing pressure by decreasing the volume raises the number of particles per unit volume. Do not explain that effect by saying the molecules move faster; at unchanged temperature their energy distribution has not shifted.

Keep two catalyst ideas together: it provides an alternative pathway with a lower activation energy, and it is regenerated overall. On a Maxwell–Boltzmann distribution the temperature curve stays put; the Eₐ threshold moves to a lower energy.

A catalyst participates in the mechanism but is regenerated, so there is no net consumption. “Unchanged at the end” refers to its overall amount/composition, not to doing nothing during the reaction.

If you add a catalyst, do not move the Maxwell–Boltzmann curve. You lower the Eₐ threshold for the alternative pathway, which increases the fraction of collisions energetic enough to react at the same temperature.

On a Maxwell–Boltzmann diagram, the temperature curve stays where it is when a catalyst is added; the Eₐ threshold moves to a lower value for the catalysed route. That is the visual link to a larger reacting fraction.

Keep the Maxwell–Boltzmann curve unchanged when a catalyst is added at the same temperature. Draw the catalysed Eₐ threshold further to the left, so a larger area under the existing curve has E ≥ Eₐ.

That extra area represents a larger fraction of molecules energetic enough to react by the catalysed pathway. The catalyst changes the activation-energy threshold, not the molecular-energy distribution.