Rate of Reaction

Collision theory. For two particles to react, three things must happen:

1. They must COLLIDE.
2. They must collide with enough energy (at least the activation energy, $E_a$).
3. They must be ORIENTED correctly (for many reactions).

Effective collision = a collision that meets all three conditions and leads to a reaction.

Increasing rate means increasing the rate of effective collisions. Each of the five factors below works through this lens.

Factor	Change	Effect on rate	Why — in terms of collisions
Concentration	Increase	Faster	More particles per unit volume → more frequent collisions
Pressure (gases)	Increase	Faster	Gas compressed → more particles per unit volume → more frequent collisions
Temperature	Increase	Faster	Particles move faster → more frequent collisions and more collisions with energy ≥ $E_a$
Surface area (solids)	Increase (smaller pieces)	Faster	More particles exposed → more frequent collisions
Catalyst	Add	Faster	Alternative pathway with lower $E_a$ → a greater fraction of collisions are effective

Worked qualitative. Why doesn't every collision produce a reaction? Most collisions are too gentle (below $E_a$) or have the wrong orientation. Only a small fraction of collisions are effective.

Concentration (solutions). Higher concentration = more particles per dm³.

• More particles in a given volume.
• More likely they'll collide.
• More effective collisions per second → FASTER rate.

Worked qualitative. Doubling the concentration of HCl in a reaction with marble chips roughly DOUBLES the rate (for first-order reactions). The reaction lasts about half the time.

Pressure (gases). Higher pressure = compressing the gas = more particles per volume.

• Same effect as concentration: more collisions per second → faster.

Worked qualitative. Increasing pressure on a gas-phase reaction (e.g. Haber process N₂ + H₂ → NH₃) increases the rate of forward reaction.

Pressure on liquid/solid reactions. Negligible effect — particles already close together.

Two effects of higher temperature.

1. Particles move FASTER → collide MORE OFTEN.
2. Particles have MORE energy on average → MORE collisions exceed the activation energy.

The second effect is much bigger. A 10°C rise typically DOUBLES the rate.

Worked qualitative. Why does food spoil more slowly in a fridge? Lower temp → enzymes (catalysts) and microbial reactions are slowed down → food lasts longer.

Why heating works two ways:

• Counts collisions per second (slight increase).
• Counts the FRACTION of collisions with enough energy (big increase).

Cambridge tip. When asked "why does temperature increase rate?", give BOTH reasons: more frequent collisions AND a higher fraction of collisions with sufficient energy.

Surface area effect. When a solid reacts with a liquid or gas, the reaction can only happen at the SURFACE.

• A large lump has small surface-area-to-volume ratio.
• Powder has huge SA:V — much more surface exposed.
• Powder reacts MUCH faster than a lump of the same mass.

Worked qualitative. Why does a flour mill have explosion safety rules? Powdered flour has enormous surface area. Mixed with air, the slightest spark can trigger rapid combustion of the entire dust cloud — flour-dust explosions are real industrial hazards.

Practical experiment. Marble chips in HCl: vary chip size (small → medium → large). Plot CO₂ volume vs time. Smallest chips give the steepest initial slope.

Cambridge tip. Always describe surface area in terms of SOLIDS reacting with gases or liquids. Two liquids or two gases don't have "surface area" in this sense (they mix freely).

Catalyst. A substance that increases the rate of a reaction without being chemically used up itself. It can be reused indefinitely.

Mechanism. Catalysts provide an ALTERNATIVE PATHWAY with a LOWER activation energy.

• More particles have enough energy to overcome $E_a$.
• More effective collisions.
• Faster rate.

Catalyst does NOT change $\Delta H$. Same products, same energy difference between reactants and products. Just a different (faster) route.

Industrial examples.

• Iron in the Haber process ($\text{N}_{2} + 3\text{H}_{2} \to 2\text{NH}_{3}$).
• Vanadium(V) oxide in the Contact process (manufacture of sulfuric acid).
• Nickel in the hydrogenation of vegetable oils to make margarine.
• Platinum, rhodium, palladium in catalytic converters (cars).
• Enzymes in biology — biological catalysts that speed up reactions in living cells.

Worked qualitative. A factory using a catalyst can run at LOWER temperatures while keeping the same rate. Saves energy → cheaper to operate → less environmental impact.

Typical graph. Plot the volume of gas produced (or mass lost, etc.) on the y-axis vs time on the x-axis.

• Curve rises steeply at first (highest rate when reactants most concentrated).
• Slope decreases as reactants are used up.
• Flattens out when reaction completes (reactants exhausted).

Reading the rate. Rate at any point = gradient of the curve. Use a tangent at that point.

Comparing conditions. Two reactions starting from the same amount but at different temperatures (say) give two curves. The hotter one rises faster — steeper initial slope. Both reach the same FINAL volume (if the reaction goes to completion) — same yield, just different rate.

Worked. Volume of CO₂ from $1\,\text{g}$ marble + 50 cm³ HCl after 1 minute: 50 cm³. After 2 minutes: 80 cm³. After 3 minutes: 95 cm³. After 5 minutes: 100 cm³ (constant).

• Initial rate $\approx 50\,\text{cm}^{3}/\text{min}$.
• Average rate over 5 min: $100/5 = 20\,\text{cm}^{3}/\text{min}$.
• Reaction effectively complete at 5 min.