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D3.5 · Explain reaction-rate factors with collision theory and energy diagrams

Learn to explain reaction-rate factors with collision theory and energy diagrams through clear examples and targeted practice.

Ontario Grade 12 Chemistry

Energy Changes and Rates of Reaction

Reaction-rate factors explained with collision theory and energy diagrams

A crushed antacid tablet usually reacts with water faster than the same tablet left whole. Bubbles may appear sooner, even though both tablets contain the same substance. At the particle level, a reaction depends on particles meeting in ways that can lead to change. This lesson uses that idea to explain reaction rate and energy diagrams.

What you will learn

  • Describe reaction rate as a change in concentration over time.
  • Use collision theory to explain how concentration, gas pressure, surface area, temperature, and catalysts affect reaction rate.
  • Read an energy diagram and identify activation energy and the energy difference between reactants and products.
  • Explain how a catalyst lowers the energy barrier without changing the reactants or products.

1. From observable change to collision theory

A chemical reaction changes reactants into products. Concentration describes the amount of a substance in a given volume. Reaction rate describes how quickly a measurable amount changes. For a reactant that is used up, average rate can be described as its decrease in concentration divided by the time taken. A possible unit is extmolL−1exts−1 ext{mol L}^{-1} ext{s}^{-1}.
Collision theory is a particle model used to explain reaction rate. Reacting particles must collide, but not every collision leads to reaction. An effective collision is one that leads to reaction. For a collision to be effective, the particles need enough energy and must meet in a suitable way. Activation energy, written EextaE_{ ext{a}}, is the minimum energy needed for particles to react when they collide.
More effective collisions per second usually means a faster reaction. A condition can increase how often particles collide, increase the share of collisions with enough energy, or provide a different reaction path with a lower energy barrier. These explanations are related, but they are not identical.
Average rate can be represented as the change in concentration divided by the time interval. For a reactant being used up, describe the change as a decrease.
  • Reaction rate describes change over time.
  • Effective collisions have enough energy and a suitable meeting.
  • Activation energy is the energy barrier particles must overcome.

2. Factors that change reaction rate

In a solution, increasing the concentration of a reacting substance puts more of its particles in the same volume. The particles are more likely to meet. Collisions can therefore occur more often, and effective collisions can happen more often. This does not mean that every collision reacts.
For gases at the same volume, increasing pressure packs particles closer together. They collide more often, so the reaction can become faster. In this comparison, the key change is how closely gas particles are packed. Pressure is not the same as temperature.
When a solid reacts with a gas or a solution, surface area can affect the rate. Surface area is the amount of the solid’s outer area exposed to the other reactant. Breaking a solid into smaller pieces exposes more area. More particles can meet the solid at its surface. This helps explain why a crushed tablet can react faster than an equal amount left whole.
Increasing temperature gives particles greater average kinetic energy, which is energy of motion. They move faster and collide more often. More importantly, a greater share of collisions has energy equal to or greater than the activation energy. The number of effective collisions per second can therefore increase. An explanation based only on faster particle motion is incomplete.
A catalyst is a substance that increases reaction rate without being used up overall. It provides a different reaction path with a lower activation energy. At the same temperature, more collisions can meet the energy requirement. A catalyst does not make particles hotter, and it does not change which reactants and products the overall reaction has.
  • Higher solution concentration and higher gas pressure can increase collision frequency.
  • Greater solid surface area provides more places for particles to meet.
  • Higher temperature increases particle motion and the share of collisions with enough energy.
  • A catalyst provides a reaction path with lower activation energy.

3. Reading energy diagrams

An energy diagram shows how the energy of a reacting system changes as a reaction proceeds. The horizontal axis represents reaction progress, from reactants toward products. The vertical axis represents energy. A curve rises to a peak before reaching the products. The peak represents a high-energy stage; it is not a product to collect.
Activation energy is the energy difference between the reactants and the peak. It is not the same as the overall energy difference between reactants and products. To find the overall difference, compare the energy levels of the products and reactants. A reaction can have a large activation energy even if its overall energy difference is small.
A catalyst’s energy diagram has a lower peak than the uncatalysed route. The reactant and product energy levels remain the same. The lower peak means a smaller activation energy. With a lower barrier, more collisions can have enough energy to react. This connects the energy diagram to collision theory.
Eexta=Eextpeak−EextreactantsE_{ ext{a}}=E_{ ext{peak}}-E_{ ext{reactants}}
  • Activation energy is measured from the reactant energy level to the peak.
  • The overall energy difference compares product and reactant energy levels.
  • A catalyst lowers the activation-energy barrier but leaves reactant and product energy levels unchanged.

4. Applying the model

To explain a rate comparison, name the changed condition first. Then describe what changes for the particles and connect it to effective collisions per second. For example: the solid is crushed, so more surface is exposed. More particles can collide with the solid’s surface, increasing the number of effective collisions per second. This explanation gives a reason for the faster rate.
When using an energy diagram, compare the reactant energy with the peak to find activation energy. Subtract the reactant energy from the peak energy. Use consistent units and report the answer in those units. Keep sensible significant digits, based on the precision of the supplied values. Do not use the product level to find activation energy.
A rate comparison does not provide the detailed sequence of particle changes in a reaction. For this lesson, use the course-level particle model: identify how the condition changes collisions or the energy barrier.
  • Link the changed condition to particle behaviour and then to effective collisions.
  • Find activation energy from the reactant level and the peak.
  • Keep activation energy distinct from the overall energy difference.

Worked example

Finding activation energy from an energy diagram

An energy diagram gives the reactants an energy of 35extkJmol−135 ext{ kJ mol}^{-1} and the peak an energy of 92extkJmol−192 ext{ kJ mol}^{-1}. Find the activation energy. A catalyst provides a path with a peak at 70extkJmol−170 ext{ kJ mol}^{-1} while the reactant energy stays the same. Find the activation energy for that path, then explain why it would be expected to be faster at the same temperature.
  1. Identify the energy levels
    Activation energy is the difference between the peak and the reactant energy. Use the reactant level as the starting point for both paths.
    Eexta=Eextpeak−EextreactantsE_{ ext{a}}=E_{ ext{peak}}-E_{ ext{reactants}}
  2. Calculate the uncatalysed barrier
    Subtract the reactant energy from the peak energy. Because both values use the same units, the difference is in extkJmol−1 ext{kJ mol}^{-1}.
    92extkJmol−1−35extkJmol−1=57extkJmol−192 ext{ kJ mol}^{-1}-35 ext{ kJ mol}^{-1}=57 ext{ kJ mol}^{-1}
  3. Calculate the catalysed barrier
    The catalysed path starts at the same reactant energy but has a lower peak. Subtract the reactant energy from that peak.
    70extkJmol−1−35extkJmol−1=35extkJmol−170 ext{ kJ mol}^{-1}-35 ext{ kJ mol}^{-1}=35 ext{ kJ mol}^{-1}
  4. Connect the result to reaction rate
    At the same temperature, a smaller activation energy means a greater share of collisions can meet the energy requirement. The catalysed path is expected to produce more effective collisions per second and a faster reaction.
Answer: The uncatalysed activation energy is 57extkJmol−157 ext{ kJ mol}^{-1}. The catalysed activation energy is 35extkJmol−135 ext{ kJ mol}^{-1}.
Check: Both calculations compare the peak with the reactant level. The catalyst lowers the barrier by 22extkJmol−122 ext{ kJ mol}^{-1}; it does not change the reactant energy.

Common mistakes and how to avoid them

Saying that every collision causes a reaction.
Correction: Only collisions with enough energy and a suitable meeting are effective.
Explaining higher temperature only by saying particles collide more often.
Correction: Also explain that a greater share of collisions has enough energy to overcome the activation-energy barrier.
Measuring activation energy from the product level.
Correction: Measure from the reactant energy level to the peak.
Saying a catalyst changes the energy of the products.
Correction: A catalyst lowers the activation-energy barrier for a different path. The reactant and product energy levels remain the same.

Lesson summary

  • Reaction rate is a change in concentration over time.
  • Collision theory explains rate through effective collisions: collisions with enough energy and a suitable meeting.
  • Concentration, gas pressure, and solid surface area can increase how often particles meet.
  • Higher temperature increases particle motion and the share of collisions with enough energy to react.
  • A catalyst provides a path with lower activation energy.
  • On an energy diagram, activation energy is the energy difference between reactants and the peak.

Check your understanding

Question 1

Why can increasing the concentration of a reactant in solution increase reaction rate?
  1. There are more reactant particles in the same volume, so collisions can occur more often.
  2. The activation energy automatically becomes zero.
  3. The products gain extra atoms from the solution.
  4. correctIndex: 0,
Show answer and explanation
There are more reactant particles in the same volume, so collisions can occur more often.
More reactant particles in the same volume can increase collision frequency and therefore the number of effective collisions per second.

Question 2

A catalyst changes an energy diagram. Which change is expected?
  1. The reactant and product energy levels both shift, but the peak stays fixed.
  2. The peak is lower relative to the reactants, while reactant and product levels stay the same.
  3. The product energy level must become equal to the reactant energy level.
  4. correctIndex: 1,
Show answer and explanation
The peak is lower relative to the reactants, while reactant and product levels stay the same.
A catalyst provides a route with lower activation energy without changing the energy levels of the overall reactants and products.

Question 3

A solid reactant is crushed into smaller pieces. What particle-level explanation best fits the faster reaction?
  1. More of the solid’s surface is exposed for particles of the other reactant to collide with.
  2. The solid particles acquire a higher temperature automatically.
  3. The reaction’s products become reactants.
  4. correctIndex: 0,
Show answer and explanation
More of the solid’s surface is exposed for particles of the other reactant to collide with.
Crushing increases exposed surface area, allowing more collisions at the solid’s surface.

Key terms

Activation energy
The minimum energy needed for colliding particles to react.
Collision theory
A particle model that explains reaction rate through collisions, energy, and whether collisions can lead to reaction.
Effective collision
A collision with enough energy and a suitable meeting that leads to reaction.
Kinetic energy
Energy of motion.
Reaction rate
How quickly a measurable amount, such as concentration, changes over time.

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