Lesson Notes By Weeks and Term v5 - Grade 11

Lesson Video

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Performance objectives

• Define enthalpy and standard enthalpy change.
• Distinguish between exothermic and endothermic reactions.
• Represent reactions using energy profile diagrams.
• Apply Hess's Law to calculate enthalpy changes.
• Write balanced equations including the enthalpy change (ΔH).

Lesson summary

Chemical reactions are fundamental to life and industry. Understanding the energy changes associated with these reactions is crucial for many applications, from designing more efficient fuel sources to predicting the stability of chemical compounds. In South Africa, this knowledge is particularly relevant in areas such as mining (where reactions release energy), agriculture (where fertilizer production depends on energy input), and addressing air pollution (which often involves understanding the energy of combustion reactions). Understanding energy changes in chemical reactions allows us to control and harness these processes for the benefit of society.

Lesson notes

2.1 Enthalpy (H) Enthalpy (H) is a thermodynamic property of a system that is defined as the sum of the internal energy (U) of the system and the product of its pressure (P) and volume (V): H = U + PV In simpler terms, enthalpy is a measure of the total heat content of a system at constant pressure. We are usually more concerned with the change in enthalpy (ΔH) during a chemical reaction because it represents the heat absorbed or released during the reaction at constant pressure. We cannot measure absolute enthalpy, only changes in enthalpy. 2.2 Standard Enthalpy Change of Reaction (ΔH°) The standard enthalpy change of reaction (ΔH°) is the enthalpy change when a reaction is carried out under standard conditions.

Standard conditions are defined as: Pressure: 101.3 kPa (1 atmosphere)

Temperature: 298 K (25 °C).

Note: even though temperature is defined as 25°C for standard conditions, reactions can happen at any temperature, and the enthalpy change can be calculated at that temperature using more advanced methods than covered in Grade

1

1. Concentration: 1 mol·dm⁻³ for all solutions. The degree symbol (°) indicates standard conditions. 2.3 Exothermic and Endothermic Reactions Chemical reactions involve breaking and forming chemical bonds. Breaking bonds requires energy (endothermic process), while forming bonds releases energy (exothermic process). Whether a reaction is exothermic or endothermic depends on the net energy change: Exothermic Reactions: Reactions that release energy into the surroundings. In these reactions, the energy required to break bonds is less than the energy released when new bonds are formed. The enthalpy change (ΔH) is negative (ΔH 0). Examples include melting ice, boiling water, and the thermal decomposition of calcium carbonate (limestone).

For example: CaCO₃(s) → CaO(s) + CO₂(g) ΔH = +178 kJ/mol (Thermal decomposition of calcium carbonate) This means that for every mole of calcium carbonate that decomposes, 178 kJ of heat energy must be supplied. This reaction is crucial in the cement industry. 2.4 Energy Profile Diagrams Energy profile diagrams visually represent the energy changes that occur during a chemical reaction. They plot the potential energy of the reactants and products against the reaction pathway.

Exothermic Reaction Energy Profile: The products have lower potential energy than the reactants. The difference in potential energy represents the energy released (ΔH is negative). The activation energy (Ea) is the energy required to start the reaction. ``` Potential Energy ^ | Ea | /\ | / \ Reactants --/ \---- Products | -- | |ΔH (negative) --------------------> Reaction Pathway ``` Endothermic Reaction Energy Profile: The products have higher potential energy than the reactants. The difference in potential energy represents the energy absorbed (ΔH is positive). The activation energy (Ea) is the energy required to start the reaction. ``` Potential Energy ^ | Ea | /\ | / \---- Products | / / Reactants --/ / | | |ΔH (positive) --------------------> Reaction Pathway ``` 2.5 Activation Energy (Ea) Activation energy is the minimum amount of energy required for a chemical reaction to occur. It's the energy needed to break the initial bonds in the reactants, allowing the reaction to proceed to form products. Even exothermic reactions require activation energy to get started. Think of it like pushing a rock over a hill; you need to exert some initial energy to get it rolling down the other side. 2.6 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 by different routes, the total enthalpy change will be the same, regardless of the route. This law is based on the fact that enthalpy is a state function – its value depends only on the initial and final states of the system, not on the path taken to get there. Mathematically, Hess's Law can be expressed as: ΔH (overall) = ΔH₁ + ΔH₂ + ΔH₃ + ... Where ΔH₁, ΔH₂, ΔH₃, etc., are the enthalpy changes for individual steps in the reaction pathway.

Example (Applying Hess's Law): Consider the following reaction: C(s) + O₂(g) → CO₂(g) ΔH = ? It's difficult to measure the enthalpy change for this reaction directly because the incomplete combustion of carbon can produce carbon monoxide (CO) as well.

However, we can use the following reactions with known enthalpy changes: C(s) + ½O₂(g) → CO(g) ΔH₁ = -110.5 kJ/mol CO(g) + ½O₂(g) → CO₂(g) ΔH₂ = -283.0 kJ/mol According to Hess's Law, the enthalpy change for the overall reaction is the sum of the enthalpy changes for the individual steps: ΔH = ΔH₁ + ΔH₂ = -110.5 kJ/mol + (-283.0 kJ/mol) = -393.5 kJ/mol Therefore, the enthalpy change for the reaction C(s) + O₂(g) → CO₂(g) is -393.5 kJ/mol.