Summary — What This Chapter Covers
- Collision Theory: reactions only happen when particles collide with enough energy (≥ activation energy) AND the correct orientation.
- Rates of Reaction: how to measure and calculate rate from graphs (mass loss, gas volume, tangents).
- Maxwell-Boltzmann Distributions: a graph showing the spread of particle energies, and how temperature reshapes it.
- Catalysts: speed up reactions by lowering activation energy via an alternative pathway, without being used up.
- Dynamic Equilibrium: in a closed system, reversible reactions reach a state where forward rate = backward rate.
- Le Chatelier's Principle: a system at equilibrium shifts to counteract any change made to it (concentration, pressure, temperature).
- Industrial Compromises: real factories (like the Haber Process) balance rate vs. yield vs. cost vs. safety.
1. Collision Theory
Why do some collisions "work" and others don't?
Imagine two people trying to shake hands while running past each other in a crowded street. If they're moving too slowly, or their hands miss each other completely, nothing happens — they just bump and carry on. But if they're moving fast enough and their hands meet just right, a proper handshake happens.
That's exactly what's going on at the molecular level. When reactant particles collide, most of these collisions are unsuccessful (ineffective) — the particles just bounce off each other without reacting. Only a small fraction are successful (effective) collisions, where a chemical reaction actually takes place.
For a collision to be effective, two conditions must both be met:
- The particles must collide with the correct orientation (the reactive parts of the molecules need to actually meet).
- The particles must collide with enough energy — at least equal to the activation energy (Ea) of the reaction.
H—H H H
+ → \ / → H—Cl + H—Cl (SUCCESSFUL — right
Cl—Cl Cl-Cl orientation + enough energy)
H—H
↓ → bounce off, NO REACTION
Cl—Cl (wrong orientation or not enough energy)
Think of it like this
Having enough energy but the wrong angle = still no reaction. Having the perfect angle but not enough energy = still no reaction. You need
BOTH at the same time. This is the single most commonly missed point in exam answers — students often only mention energy and forget orientation.
Collision Frequency — what makes collisions happen more often?
Collision frequency is simply how many collisions happen per unit time. It can be increased by:
- Increasing concentration — more particles squeezed into the same volume means they bump into each other more often.
- Increasing pressure (for gases) — same number of particles in a smaller volume = more frequent collisions.
- Increasing temperature — particles move faster (more kinetic energy), so they collide more often and harder.
- Increasing surface area
Common Mistake
successful
proportion
a
| Reaction Type |
Energy Comparison |
a
|
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