Effect of Concentration, Pressure, Temperature and Surface Area

Learning Objectives (OCR Spec 3.2.2)

By the end of this lesson you should be able to:

• Explain the effect of concentration on rate in terms of collision frequency
• Explain the effect of pressure on rate for gas-phase reactions
• Explain the effect of temperature on rate using both collision frequency and the Boltzmann distribution
• Explain the effect of surface area on rate for reactions involving solids
• Describe and interpret experimental results for each factor

Concentration and Rate

Increasing the concentration of a reactant in solution increases the rate of a reaction. The explanation lies in collision theory: doubling the concentration doubles the number of particles per unit volume, so there are twice as many collisions per second. If the proportion of successful collisions remains the same (same temperature, same Ea), the rate should approximately double.

Classic Experiment – Thiosulfate and Acid

Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + SO2(g) + S(s) + H2O(l)

A precipitate of sulfur forms, obscuring a cross drawn on paper under the flask. The time to obscure the cross is used as 1/rate:

[Na2S2O3] / mol dm⁻³	Time to obscure cross / s	1/t / s⁻¹
0.040	200	0.0050
0.080	100	0.0100
0.120	67	0.0149
0.160	50	0.0200
0.200	40	0.0250

A plot of 1/t against [Na2S2O3] is approximately linear, passing through the origin, showing that rate is directly proportional to concentration. (We will meet the formal definition of rate order later.)

Pressure and Gas-Phase Reactions

For gas-phase reactions, pressure plays the same role as concentration. Doubling the pressure doubles the concentration of gas molecules (fewer cm³ per mole) and therefore doubles the collision frequency.

Example: N2(g) + 3H2(g) → 2NH3(g), the Haber process, is carried out at high pressure (~200 atm) partly to speed the reaction up, partly to shift the equilibrium (see Lesson 10).

For reactions between solids and gases (e.g. catalytic cracking) only the gas concentration matters; changing the pressure on a pure solid has no effect on collision frequency.

Temperature and Rate

Temperature is the most dramatic factor because it affects collision frequency and the fraction of effective collisions. As covered in Lesson 7, raising the temperature:

1. Increases the average kinetic energy of particles, so collisions are more energetic.
2. Shifts the Boltzmann distribution to the right, so a greater proportion of molecules have energy ≥ Ea.
3. Increases the frequency of collisions (because particles move faster).

The second effect is by far the most important: a 10 K rise near room temperature typically doubles the rate, even though the collision frequency only rises by about 2%. The dominant contribution comes from the exponential Boltzmann factor.

Exam-style explanation of a temperature rise:

"When temperature increases, the particles have greater kinetic energy. A greater proportion of molecules have energy greater than or equal to the activation energy, so more collisions lead to reaction. Additionally, the frequency of collisions increases because particles move faster. The rate therefore increases."

Surface Area (Heterogeneous Reactions)

When a solid reacts with a liquid or a gas, only the particles at the surface of the solid can collide with the other reactant. Breaking the solid into smaller pieces - or grinding it to a powder - increases the total surface area and therefore the number of sites available for collision. The rate is roughly proportional to the surface area.

Classic Experiment – Marble Chips and Acid

CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)

Large marble chips react slowly; small chips react faster; powder reacts fastest. In each case the same mass of CaCO3 produces the same volume of CO2 - only the rate changes.