⚡ Collision theory says that the more molecules you have, the more collisions happen per second.
If the concentration of a reactant increases from \$[A]\$ to \$2[A]\$, the number of collisions roughly doubles, so the rate increases by a factor of 2 (for a first‑order reaction).
Analogy: Think of people in a crowded room. The more people (molecules) there are, the more chances they bump into each other.
Exam tip: Remember that the rate law is \$rate = k[A]^m[B]^n\$. Changing concentration changes the rate by the power of that concentration in the law.
🔥 Raising the pressure of a gas compresses the molecules into a smaller volume, increasing the number of collisions per unit volume.
For a reaction involving gases, the rate often increases proportionally to the pressure (for a first‑order gas reaction).
Analogy: Imagine a crowded subway car. The more people (higher pressure), the faster they bump into each other.
Exam tip: Use the ideal gas law \$PV = nRT\$ to relate pressure changes to concentration changes for gases.
🧩 When a solid reactant is broken into smaller pieces, its surface area increases, giving more sites for collisions.
The rate increases because more molecules can collide with the solid at once.
Analogy: Think of a pizza. A whole pizza (small surface area) is slower to eat than many thin slices (large surface area).
Exam tip: For heterogeneous reactions, rate ∝ surface area. Remember to mention “surface area” in your answer.
🌡️ Increasing temperature gives molecules more kinetic energy, so they collide faster and with more energy.
The Arrhenius equation \$k = A e^{-E_a/RT}\$ shows that the rate constant \$k\$ rises exponentially with temperature.
Analogy: Like a hot summer day, people move faster and bump into each other more often.
Exam tip: Write that the rate increases with temperature and explain it using the Arrhenius equation or collision theory.
🏃♂️ A catalyst provides an alternative reaction pathway with a lower activation energy \$E_a\$.
It does not change the equilibrium but speeds up the approach to equilibrium.
Enzymes are biological catalysts that bind reactants (substrates) in a “lock‑and‑key” pocket, further lowering \$E_a\$.
Analogy: Think of a traffic cop (catalyst) who directs cars (molecules) through a shortcut, reducing the time to reach the destination.
Exam tip: Mention that catalysts lower \$Ea\$, increase \$k\$ (but not \$K{eq}\$), and that enzymes are specific and can be inhibited.
| Factor | Effect on Rate | Key Point |
|---|---|---|
| Concentration | Increases (proportional to concentration power in rate law) | \$rate = k[A]^m[B]^n\$ |
| Pressure (gases) | Increases (∝ pressure for first‑order) | Use \$PV = nRT\$ to relate pressure to concentration |
| Surface area (solids) | Increases with more surface area | Rate ∝ surface area for heterogeneous reactions |
| Temperature | Exponential increase (Arrhenius) | \$k = A e^{-E_a/RT}\$ |
| Catalyst / Enzyme | Increases rate by lowering \$E_a\$ | Does not change equilibrium constant |