Lesson Notes By Weeks and Term v5 - Grade 11

Lesson Video

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Performance objectives

• Define and differentiate between endothermic and exothermic reactions.
• Explain enthalpy change (ΔH).
• Interpret and draw potential energy diagrams.
• Apply Hess's Law to calculate enthalpy changes.
• Use standard enthalpy of formation data to calculate reaction enthalpy.

Lesson summary

Chemical reactions are not just about mixing substances and observing new products. They are also inherently linked to energy changes. Understanding these energy changes is crucial for a variety of applications, from optimizing industrial processes to understanding biological systems. In South Africa, energy production from coal, the design of more efficient solar panels, and even cooking our food are all underpinned by the principles of energy and chemical change. Without grasping these fundamentals, we cannot effectively tackle challenges related to energy security, environmental sustainability, and technological advancement within our country.

Lesson notes

2.1 Exothermic and Endothermic Reactions: Chemical reactions always involve the breaking and forming of chemical bonds. Breaking bonds requires energy (endothermic process), while forming bonds releases energy (exothermic process). Whether a reaction is ultimately exothermic or endothermic depends on the overall energy balance.

Exothermic Reactions: These reactions release energy into the surroundings, usually in the form of heat. This means the products have less chemical potential energy than the reactants. The temperature of the surroundings increases. Think about burning wood in a braai – the heat released is a clear sign of an exothermic reaction. The formation of water from hydrogen and oxygen is also a highly exothermic reaction.

Endothermic Reactions: These reactions absorb energy from the surroundings. The reactants have less chemical potential energy than the products. The temperature of the surroundings decreases. An example relevant to South Africa is the extraction of metals like gold from their ores, which often requires heating the ore with other substances. Photosynthesis, the process by which plants convert carbon dioxide and water into glucose and oxygen, is another vital endothermic reaction. 2.2 Enthalpy Change (ΔH): Enthalpy (H) is a thermodynamic quantity that represents the total heat content of a system at constant pressure. We're usually more interested in the change in enthalpy (ΔH), which is the heat absorbed or released during a chemical reaction at constant pressure. ΔH = H products - H reactants For an exothermic reaction, ΔH is negative (ΔH 0) because the enthalpy of the products is greater than the enthalpy of the reactants. The energy absorbed is indicated by the positive sign. 2.3 Thermochemical Equations: A thermochemical equation is a balanced chemical equation that includes the enthalpy change (ΔH) for the reaction. It's crucial to include the physical states (solid (s), liquid (l), gas (g), aqueous (aq)) of the reactants and products because these states affect the enthalpy.

Example: ``` CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) ΔH = -890 kJ/mol ``` This equation tells us that when 1 mole of methane gas reacts with 2 moles of oxygen gas to produce 1 mole of carbon dioxide gas and 2 moles of water vapor, 890 kJ of energy is released (exothermic reaction). 2.4 Hess's Law: Hess's Law states that the enthalpy change for a reaction is independent of the pathway taken. In other words, if a reaction can occur in one step or several steps, the total enthalpy change is the same. This is incredibly useful because it allows us to calculate enthalpy changes for reactions that are difficult or impossible to measure directly.

Example: Calculate the enthalpy change for the reaction: C(s) + O₂(g) → CO(g) (This reaction is hard to measure directly as CO₂ is often formed).

Given: C(s) + O₂(g) → CO₂(g) ΔH₁ = -393.5 kJ/mol CO(g) + ½O₂(g) → CO₂(g) ΔH₂ = -283.0 kJ/mol To get the desired reaction, we need to reverse reaction (2) and add it to reaction (1): C(s) + O₂(g) → CO₂(g) ΔH₁ = -393.5 kJ/mol CO₂(g) → CO(g) + ½O₂(g) ΔH₂' = +283.0 kJ/mol (

Note: Reversing the reaction changes the sign of ΔH)

Adding the reactions: C(s) + O₂(g) + CO₂(g) → CO₂(g) + CO(g) + ½O₂(g)

Simplifying: C(s) + ½O₂(g) → CO(g) ΔH = ΔH₁ + ΔH₂' = -393.5 kJ/mol + 283.0 kJ/mol = -110.5 kJ/mol 2.5 Standard Enthalpy of Formation (ΔH f o ): The standard enthalpy of formation (ΔH f o ) is the enthalpy change when one mole of a compound is formed from its elements in their standard states (298 K and 1 atm). The standard enthalpy of formation of an element in its standard state is, by definition, zero. Standard enthalpy of formation data is readily available in tables. We can use standard enthalpies of formation to calculate the enthalpy change for any reaction: ΔH reaction = Σ ΔH f o (products) - Σ ΔH f o (reactants)

Example: Calculate the enthalpy change for the combustion of ethanol: C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l)

Given: ΔH f o [C₂H₅OH(l)] = -277.7 kJ/mol ΔH f o [CO₂(g)] = -393.5 kJ/mol ΔH f o [H₂O(l)] = -285.8 kJ/mol ΔH f o [O₂(g)] = 0 kJ/mol (element in its standard state) ΔH reaction = [2 ΔH f o (CO₂(g)) + 3 ΔH f o (H₂O(l))] - [ΔH f o (C₂H₅OH(l)) + 3 * ΔH f o (O₂(g))] ΔH reaction = [2 (-393.5 kJ/mol) + 3 (-285.8 kJ/mol)] - [-277.7 kJ/mol + 3 * (0 kJ/mol)] ΔH reaction = [-787 kJ/mol - 857.4 kJ/mol] - [-277.7 kJ/mol] ΔH reaction = -1644.4 kJ/mol + 277.7 kJ/mol ΔH reaction = -1366.7 kJ/mol 2.6 Potential Energy Diagrams: Potential energy diagrams visually represent the energy changes that occur during a chemical reaction. The y-axis represents potential energy. The x-axis represents the reaction progress (or reaction coordinate).

Key Features: Reactants: The starting point on the diagram.

Products: The ending point on the diagram.

Activation Energy (E a ): The energy barrier that must be overcome for the reaction to occur.