Lesson Notes By Weeks and Term v5 - Grade 12

Lesson Video

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Performance objectives

• Define reaction rate and calculate it from data.
• List the factors that affect reaction rates.
• Explain these factors using collision theory.
• Define what a catalyst is and explain how it works.
• Apply these concepts to solve problems.

Lesson summary

Chemical kinetics, or the study of reaction rates, is a fundamental concept in chemistry with wide-ranging implications. Understanding how quickly or slowly a reaction proceeds allows us to control and optimize chemical processes. From the efficient production of fertilizers to the preservation of food and the development of new pharmaceuticals, the principles of reaction rates are essential. In South Africa, this knowledge is particularly important in industries like mining (optimizing ore extraction), agriculture (improving fertilizer use and crop yield), and manufacturing (controlling the speed and efficiency of chemical processes).

Lesson notes

2.1 Definition of Reaction Rate: Reaction rate is defined as the change in concentration of reactants or products per unit time. It essentially describes how quickly a chemical reaction is occurring. The concentration is typically measured in moles per cubic decimeter (mol/dm³) or molarity (M), and time is often measured in seconds (s), minutes (min), or hours (h).

Therefore, the units of reaction rate are typically mol/dm³/s, mol/dm³/min, or mol/dm³/h. Mathematically, the average reaction rate can be expressed as: Rate = - Δ[Reactant] / Δt Rate = + Δ[Product] / Δt The negative sign in front of the reactant's change in concentration indicates that the reactant is being consumed, and its concentration is decreasing over time. The positive sign in front of the product's change in concentration indicates that the product is being formed, and its concentration is increasing over time. 2.2 Factors Affecting Reaction Rate: Several factors can influence the speed at which a chemical reaction proceeds.

These include: Nature of Reactants: Some substances are inherently more reactive than others. For example, reactions involving ionic compounds in solution often occur much faster than reactions involving covalent compounds. This is because ionic compounds readily dissociate into ions, which can quickly react with other ions in solution.

Concentration of Reactants: Increasing the concentration of reactants generally increases the reaction rate. This is because there are more reactant molecules present, leading to a higher frequency of collisions between them. According to collision theory, a greater number of effective collisions per unit time results in a faster reaction rate.

Pressure (for gaseous reactants): For reactions involving gases, increasing the pressure is analogous to increasing the concentration. Higher pressure means more gas molecules are confined to a smaller volume, leading to more frequent collisions and a faster reaction rate.

Temperature: Increasing the temperature almost always increases the reaction rate. This is because higher temperatures provide more kinetic energy to the reactant molecules. This increased kinetic energy results in more frequent and more energetic collisions. The higher the energy, the higher the proportion of collisions exceeding the activation energy. Surface Area (for heterogeneous reactions): For reactions involving solids, increasing the surface area of the solid reactant increases the reaction rate. A greater surface area means more reactant molecules are exposed and available for collisions with other reactants. Think of trying to burn a log of wood versus wood shavings; the shavings catch fire much faster because they have a much larger surface area exposed to the oxygen in the air.

Catalysts: A catalyst is a substance that speeds up a chemical reaction without being consumed in the reaction. Catalysts work by providing an alternative reaction pathway with a lower activation energy. This allows more reactant molecules to overcome the energy barrier and form products. Catalysts are incredibly important in many industrial processes, allowing reactions to occur at faster rates and lower temperatures, thus reducing energy consumption and production costs. 2.3 Collision Theory: Collision theory is a fundamental concept in understanding reaction rates. It states that for a reaction to occur, reactant molecules must collide with each other.

However, not all collisions result in a reaction. For a collision to be effective, two conditions must be met: Sufficient Energy: The colliding molecules must have enough kinetic energy to overcome the activation energy barrier. Activation energy is the minimum energy required for a reaction to occur.

Correct Orientation: The colliding molecules must be oriented in the correct way for the reaction to take place. For example, if two molecules are reacting at specific sites, those sites must collide with each other. Increasing the temperature increases the kinetic energy of the molecules, leading to more collisions with sufficient energy to overcome the activation energy. Increasing the concentration increases the frequency of collisions. Catalysts lower the activation energy, making it easier for molecules to react upon collision. 2.4 Catalysts in Detail: Catalysts provide an alternate reaction pathway with a lower activation energy. This can involve the formation of intermediate species that are more easily converted to products. Catalysts can be homogeneous (present in the same phase as the reactants) or heterogeneous (present in a different phase).

Homogeneous Catalysis: An example is the acid catalysis of ester hydrolysis. The acid (e.g., H+) is in the same aqueous phase as the ester.

Heterogeneous Catalysis: An example is the use of a platinum catalyst in the catalytic converter of a car. The platinum (solid) catalyzes the oxidation of carbon monoxide and hydrocarbons (gases) in the exhaust fumes.