Lesson Notes By Weeks and Term v5 - Grade 12

Lesson Video

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Performance objectives

• Define and explain the collision theory.
• Identify factors that influence reaction rate.
• Explain activation energy and the activated complex.
• Use the Maxwell-Boltzmann distribution curve to explain temperature's effect.
• Define a catalyst and explain its function.

Lesson summary

This week, we delve deeper into the fascinating world of chemical reaction rates. Understanding how quickly reactions occur is crucial, not just in a laboratory setting, but also in everyday life and various industries vital to South Africa's economy. For instance, knowing how to control the rate of corrosion of pipes is critical for maintaining our water infrastructure, and understanding reaction rates in fertilizer production helps ensure efficient and cost-effective farming practices. In mining, controlling the rate of reactions during ore processing is paramount for maximizing yields and minimizing environmental impact.

Lesson notes

2.1 Collision Theory The collision theory states that for a chemical reaction to occur, reactant particles (atoms, ions, or molecules) must collide with each other.

However, not all collisions result in a reaction. For a collision to be successful, two conditions must be met: Sufficient Energy: The colliding particles must possess enough kinetic energy to overcome the activation energy barrier (explained below).

Correct Orientation: The particles must collide with the proper orientation, allowing the reactive parts of the molecules to come into contact and form new bonds. Imagine trying to build a Lego castle. You need the right pieces (correct orientation) and enough force (sufficient energy) to snap them together. A gentle push with the wrong side won't work! 2.2 Activation Energy (Ea) and Activated Complex Activation Energy (Ea) is the minimum amount of energy required for a reaction to occur. It's the energy needed to break the existing bonds in the reactants so that new bonds can form in the products. Think of it as the energy needed to push a rock over a hill. Once it's over, it rolls down on its own. The Activated Complex (also called the transition state) is an unstable, high-energy arrangement of atoms that exists momentarily during a reaction. It represents the point where bonds are breaking and forming simultaneously. It's like the rock precariously balanced at the top of the hill.

Potential Energy Diagrams: These diagrams illustrate the energy changes during a reaction. The y-axis represents potential energy, and the x-axis represents the reaction progress. For an exothermic reaction, the products have lower potential energy than the reactants. For an endothermic reaction, the products have higher potential energy. The activation energy is the difference in potential energy between the reactants and the activated complex. 2.3 Factors Affecting Reaction Rate Several factors influence how quickly a reaction proceeds: Concentration: Increasing the concentration of reactants increases the number of particles per unit volume, leading to more frequent collisions and a higher chance of successful collisions. Imagine a crowded taxi rank. The more people crammed into a small space, the more likely they are to bump into each other.

Example:* Burning wood chips vs. a log. Finer wood chips (higher surface area, addressed below, but related to effective concentration of combustible material) ignite faster because the available combustible material is concentrated in smaller particles.

Pressure (for gases): Increasing the pressure of gaseous reactants forces the molecules closer together, effectively increasing their concentration. This leads to more frequent collisions and a faster reaction rate. This principle is utilized in many industrial processes, especially in the petrochemical industry.

Example:* The Haber-Bosch process for producing ammonia (used in fertilizers) uses high pressure to increase the rate of reaction between nitrogen and hydrogen gases.

Temperature: Increasing the temperature increases the average kinetic energy of the particles. This means more particles will have enough energy to overcome the activation energy barrier, resulting in more successful collisions. Higher temperatures also increase the frequency of collisions.

Example:* Food spoils faster at room temperature than in a refrigerator because the higher temperature increases the rate of the reactions that cause spoilage.

Surface Area: Increasing the surface area of a solid reactant increases the area available for collisions. This leads to more frequent collisions and a faster reaction rate.

Example:* A sugar cube dissolves slower than granulated sugar because the granulated sugar has a larger surface area exposed to the solvent (water). Mining companies crush rocks into smaller pieces to increase the surface area for leaching processes to extract valuable minerals like gold.

Catalyst: A catalyst is a substance that speeds up a reaction without being consumed in the reaction itself. It provides an alternative reaction pathway with a lower activation energy. This means that more particles will have enough energy to react at a given temperature.

Example:* Enzymes in our bodies are biological catalysts that speed up biochemical reactions necessary for life. Catalytic converters in cars use catalysts like platinum and palladium to speed up the conversion of harmful pollutants in exhaust gases into less harmful substances. 2.4 Maxwell-Boltzmann Distribution Curve This curve shows the distribution of kinetic energies among the molecules of a gas at a given temperature. The x-axis represents kinetic energy, and the y-axis represents the number of molecules with that energy. The area under the curve represents the total number of molecules. The peak of the curve represents the most probable kinetic energy. As temperature increases, the curve shifts to the right and flattens, indicating that more molecules have higher kinetic energies.