Chemical kinetics is the study of reaction rates and the molecular mechanisms by which reactants convert into products. At the microscopic level, chemical reactions occur via particle collisions. Collision theory provides the framework for understanding how and why chemical reactions happen at different speeds.
🔑 Key Principle
For a chemical reaction to occur, reactant particles must collide. However, not all collisions result in a reaction. A collision will only lead to a reaction (making it a successful collision) if the particles collide with sufficient energy (greater than or equal to the activation energy) and in the correct orientation.
The minimum energy that colliding reactant particles must possess for a collision to result in a chemical reaction by overcoming the barrier to break existing bonds.
The change in concentration of a reactant or a product per unit time. The standard unit is \( \text{mol dm}^{-3}\text{s}^{-1} \).
The total number of collisions occurring between reactant particles per unit time in a given volume of the reaction mixture.
The Energy Profile Diagram
A reaction profile diagram illustrates the energy changes during a chemical reaction. The curve starts at the energy level of the reactants, climbs to a peak representing the transition state (where bonds are being broken and made), and ends at the energy level of the products.
The activation energy (\( E_a \)) is represented as the energy difference between the reactants and the highest point on the curve. This is the energy barrier that must be overcome.
Factors Affecting the Rate of Reaction
Collision theory explains how several key factors alter the rate of a chemical reaction by changing either the frequency of collisions or the proportion of successful collisions.
1. Concentration (Solutions) and Pressure (Gases)
When the concentration of a reactant in solution is increased, or the pressure of a gaseous reactant mixture is increased, there are more particles in a given volume.
- This causes the particles to be closer together on average.
- Consequently, the collision frequency (number of collisions per second) increases.
- Since collision frequency increases, the rate of reaction increases.
2. Surface Area (Solids)
If a solid reactant is broken into smaller pieces (or a powder), its total surface area increases. This exposes more reactant particles to the surrounding fluid reactants.
- With more particles exposed at the surface, more collisions can occur simultaneously.
- The collision frequency increases.
- The rate of reaction increases.
3. Temperature
Increasing the temperature of a reaction mixture increases the average kinetic energy of the particles. This has two separate effects, one of which is far more important:
- The particles move faster, which increases the frequency of collisions slightly.
- More importantly, a much greater proportion of particles possess energy equal to or greater than the activation energy (\( E \ge E_a \)).
- This results in a dramatic increase in the frequency of successful collisions, leading to a large increase in reaction rate.
4. Adding a Catalyst
A catalyst increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy (\( E_a \)). This is explained in detail in subsequent lessons.
When writing exam answers, avoid vague statements like "there are more collisions". You must specify "more collisions per unit time" or "increased frequency of collisions". Additionally, for temperature changes, you must distinguish between the overall collision rate (which increases slightly) and the number of successful collisions per unit time (which increases dramatically because more particles have energy \( \ge E_a \)).
Worked Examples
Solution:
1. As the reaction proceeds, acid particles (hydrochloric acid) and solid calcium carbonate are consumed to form products.
2. Therefore, the concentration of the acid decreases (fewer acid particles per unit volume).
3. This leads to a decrease in the collision frequency between the acid particles and the remaining calcium carbonate.
4. Since there are fewer collisions per unit time, the rate of reaction decreases.
Solution:
1. Increase the total frequency of collisions: This is achieved by increasing concentration, increasing gas pressure, or increasing the surface area of a solid reactant.
2. Increase the proportion of collisions that are successful (have \( E \ge E_a \)): This is achieved by increasing the temperature (giving particles more energy) or by adding a catalyst (lowering the activation energy barrier).
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