Collision Theory and Factors Affecting Rate

What Is Rate of Reaction?

The rate of reaction is a measure of how quickly reactants are converted into products. It can be expressed as:

rate = amount of product formed ÷ time

or

rate = amount of reactant used ÷ time

Units depend on what is measured: g/s, cm³/s, or mol/s.

Collision Theory

For a reaction to occur, reactant particles must:

1. Collide with each other

2. With sufficient energy (at least the activation energy, Ea)

Not every collision leads to a reaction. Collisions without enough energy simply result in particles bouncing off each other. Only collisions with energy equal to or greater than the activation energy are successful collisions.

Anything that increases the frequency of successful collisions will increase the rate of reaction.

Factors Affecting Rate

1. Concentration (of Solutions)

Increasing concentration increases the rate.

Why: More particles are present in the same volume. This means particles are closer together and collide more frequently, increasing the number of successful collisions per second.

Example: Reacting magnesium ribbon with hydrochloric acid — using 2 mol/dm³ HCl reacts faster than 0.5 mol/dm³ HCl.

2. Pressure (of Gases)

Increasing pressure increases the rate.

Why: The gas particles are pushed closer together into a smaller volume. This increases the frequency of collisions, leading to more successful collisions per second.

This is essentially the same effect as concentration but for gases.

3. Temperature

Increasing temperature increases the rate.

Why — there are TWO effects:

1. Particles have more kinetic energy, so they move faster and collide more frequently

2. A greater proportion of collisions have energy equal to or exceeding the activation energy, so more collisions are successful

The second effect is more important. A 10°C rise in temperature can roughly double the rate of many reactions.

4. Surface Area (of Solids)

Increasing surface area increases the rate.

Why: More surface is exposed to the other reactant, so more collisions can happen at the same time.

Powdered solid reacts faster than lumps of the same mass because the total surface area is much greater.

Example: Powdered calcium carbonate reacts much faster with HCl than large marble chips.

5. Catalysts

A catalyst increases the rate without being used up. It provides an alternative reaction pathway with a lower activation energy, so a greater proportion of particles have enough energy to react.

The catalyst is not consumed — it can be used again and again.

Required Practical: Investigating Rate of Reaction

Method 1: Measuring Gas Volume

React magnesium with acid and collect the hydrogen gas in a gas syringe.

1. Measure a fixed volume of HCl in a conical flask

2. Add a known length of magnesium ribbon

3. Immediately connect to a gas syringe

4. Record the volume of gas every 10 seconds

5. Repeat with different concentrations of HCl (keeping everything else the same)

Method 2: Disappearing Cross (Sodium Thiosulfate + HCl)

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

The sulfur produced makes the solution go cloudy. Time how long it takes for a cross drawn on paper underneath the flask to disappear.

A shorter time = faster rate.

Method 3: Change in Mass

Place the reaction on a balance. As gas escapes, the mass decreases. Record mass at intervals.

Interpreting Rate Graphs

On a graph of product formed (y-axis) vs time (x-axis):

• A steeper gradient = faster rate
• The line levels off when the reaction is complete (all of the limiting reactant is used up)
• Higher temperature or concentration gives a steeper initial gradient but the same final amount of product (if the same amount of reactant is used)
• Using a catalyst makes the line steeper but the final amount stays the same

Exam Tips

• Always explain rate changes using collision theory: state the effect on collision frequency AND whether collisions are successful (have enough energy)
• For temperature, you must mention BOTH more frequent collisions AND a greater proportion exceeding Ea — just saying "faster" is not enough
• Catalysts lower Ea; they do not increase temperature or give particles more energy
• Control variables: in any rate experiment, change only one factor and keep everything else constant