Objective: Describe how the following factors affect the speed of a chemical reaction:
Think of molecules as people at a party. If more people (molecules) are present, the chances of two people bumping into each other increase, so the reaction goes faster. The rate law shows this mathematically:
\$rate = k[A]^m[B]^n\$
• Doubling the concentration of one reactant typically increases the rate by a factor of \$2^m\$ (where \$m\$ is the reaction order with respect to that reactant).
• If the reaction is first‑order in A, doubling [A] doubles the rate.
• For a second‑order reaction in A, doubling [A] quadruples the rate.
Imagine gas molecules as cars on a highway. Higher pressure means cars are packed tighter, so collisions happen more often. For reactions involving gases:
• Increasing pressure increases the concentration of gas molecules (according to the ideal gas law \$PV = nRT\$).
• The rate increases roughly proportionally to the pressure if the reaction is elementary and involves those gases.
• Example: The synthesis of ammonia \$3H2 + N2 \rightarrow 2NH3\$ speeds up when the pressure of \$H2\$ and \$N_2\$ is raised.
Picture a sponge: the more surface area it has, the more water it can absorb. Similarly, a solid reactant with a larger surface area offers more “contact points” for the reaction.
• Crushing a solid into a fine powder increases its surface area and thus the rate.
• Example: The reaction of zinc metal with hydrochloric acid is faster when zinc is powdered than when it is a solid block.
Think of molecules as dancers: at higher temperatures they move faster and collide more energetically. The Arrhenius equation quantifies this:
\$k = A e^{-E_a/(RT)}\$
• Raising the temperature increases the rate constant \$k\$, often dramatically.
• A 10 °C rise can double or triple the rate for many reactions.
• Example: The decomposition of hydrogen peroxide is much faster at 80 °C than at 20 °C.
A catalyst is like a shortcut on a road: it provides an alternative pathway with a lower activation energy \$E_a\$, so more molecules have enough energy to react.
• Catalysts do not appear in the rate law; they only change the rate constant \$k\$.
• Enzymes are biological catalysts that are highly specific and work best at a particular temperature and pH.
• Example: The enzyme catalase breaks down hydrogen peroxide into water and oxygen much faster than the uncatalysed reaction.
| Factor | Effect on Rate | Analogy |
|---|---|---|
| Concentration | Higher → Faster (proportional to concentration powers) | More people at a party → more chances to bump into each other |
| Pressure (gases) | Higher → Faster (due to increased collisions) | Cars on a crowded highway → more collisions |
| Surface Area (solids) | Greater → Faster (more contact points) | Sponge soaking up water → more surface to absorb |
| Temperature | Higher → Faster (exponential increase in \$k\$) | Dancers moving faster → more energetic collisions |
| Catalyst / Enzyme | Faster (lower \$E_a\$) without being consumed | Shortcut on a road → quicker journey |