Rate of Reaction

Rate of reaction is the change in concentration of a reactant or product per unit time (units: mol dm⁻³ s⁻¹).

Collision theory: reactions happen when particles collide. But not every collision causes a reaction. A collision is effective (leads to reaction) only if:

1. the colliding particles have at least the activation energy ($E_a$), and
2. they collide with the correct orientation.

A non-effective collision has too little energy or the wrong orientation, so no reaction occurs.

Frequency of collisions is how often collisions happen per second. Anything that increases the frequency of effective collisions increases the rate.

Increasing concentration of a reactant in solution:

• packs more particles into the same volume,
• so collisions become more frequent,
• giving a higher frequency of effective collisions → faster rate.

Increasing pressure of a gas does the same thing: it forces the gas molecules closer together (effectively raising their concentration), increasing the collision frequency and the rate.

Key phrasing. Always link the change to the frequency of effective collisions — not just 'more collisions'. The molecules' energy is unchanged by concentration/pressure; only how often they collide changes.

Rate is found from how a measurable quantity changes with time — e.g. gas volume, mass loss, colour/absorbance, or concentration (from titration).

$\text{rate} = \frac{\text{change in concentration}}{\text{time}}$

On a concentration–time graph:

• the rate at any instant = the gradient of the tangent at that point;
• the reaction is fastest at the start (gradient steepest, highest concentration) and slows as reactants are used up;
• the graph levels off when the reaction stops (a reactant runs out).

Worked. If [reactant] falls from 0.80 to 0.50 mol dm⁻³ in 60 s, the average rate is: $\text{rate} = \frac{0.80 - 0.50}{60} = 5.0 \times 10^{-3}\,\text{mol dm}^{-3}\,\text{s}^{-1}$