6.6 Introduction to Enthalpy of Reaction

What is Enthalpy?

Enthalpy (H) is a key concept in thermochemistry that describes the total internal energy of a system, including the energy needed to change its temperature and pressure. In simple terms, enthalpy refers to the heat content of a system. When discussing chemical reactions, we focus on the change in enthalpy (ΔH), which is the difference between the enthalpy of the products and the reactants.

• ΔH > 0 (Positive): The reaction absorbs heat (endothermic).
• ΔH < 0 (Negative): The reaction releases heat (exothermic).

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Enthalpy and Reaction Energy

The enthalpy of reaction (ΔH) describes the energy change during a chemical reaction:

• Exothermic Reactions: Heat is released, resulting in a negative ΔH. Example: The combustion of propane releases heat into the surroundings, often seen as flames.
• Endothermic Reactions: Heat is absorbed, resulting in a positive ΔH. Example: The dissolution of anhydrous copper sulfate in water causes the solution to become colder as it absorbs heat.

Key Examples:

1. Combustion of Propane (Exothermic):

   $$\text{C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(g)}$$
   • ΔH < 0; heat is released, observable as an increase in temperature or a flame.

2. Dissolution of Anhydrous CuSO₄ (Endothermic):

   $$\text{CuSO₄(s) + H₂O(l) → CuSO₄(aq)}$$
   • ΔH > 0; heat is absorbed, observable as a temperature drop.

Positive vs. Negative ΔH

The sign of ΔH provides critical insight into the reaction’s heat flow direction:

• Positive ΔH (Endothermic): Heat flows into the system from the surroundings.
• Negative ΔH (Exothermic): Heat flows out of the system into the surroundings.

Tip: While the temperature change in the surroundings can provide clues about a reaction’s enthalpy change, the total energy exchange depends on the system’s internal energy changes, not just temperature shifts.

Example Reactions:

1. Exothermic Reaction: Combustion of methane (CH₄) to form CO₂ and H₂O. $$\text{CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)}$$
   • ΔH < 0, indicating heat release.
2. Endothermic Reaction: Melting of ice at 0°C. $$\text{H₂O(s) → H₂O(l)}$$
   • ΔH > 0, indicating heat absorption.

Thermodynamic Feasibility of Reactions

While a negative ΔH (exothermic reaction) often suggests a thermodynamically favorable reaction, it’s not the sole determining factor. The Gibbs free energy change (ΔG) and the equilibrium constant (Kc) also play crucial roles.

ΔG = ΔH – TΔS, where:

• ΔS is the change in entropy.
• T is the temperature in Kelvin.

Understanding Enthalpy Changes

Internal Energy (E) and Related Terms

• Energy: The capacity to do work or transfer heat; measured in joules (J).
• Internal Energy (E): Sum of all kinetic and potential energies within a system.
  • ΔE = q + w (where q = heat, w = work)
• Heat (q): Energy transfer due to a temperature difference.
• Work (w): Energy transfer due to a force acting over a distance; for gases, w = -PΔV.

Quick Examples of Energy Transfer

1. Heating Water on a Stove: Heat (q) flows into the pot, increasing water’s temperature.
2. Lifting a Weight: Work (w) is done as a force moves the weight.
3. Chemical Reactions in a Test Tube: Measure ΔH through the heat absorbed or released.
4. Gas Compression: Work done on a gas increases its internal energy.
5. Car Engines: Convert chemical energy from fuel into mechanical energy, with heat as a byproduct.

Practice Examples

1. Combustion of Methane (Exothermic)
Equation:

$$\text{CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)}$$

Description: ΔH is negative, indicating heat release.

2. Dissolution of Ammonium Nitrate (Endothermic)
Equation:

$$\text{NH₄NO₃(s) + H₂O(l) → NH₄NO₃(aq)}$$

Description: ΔH is positive, indicating heat absorption.

Image Courtesy of Mr Lowe

Additional Terms to Remember

• Energy (Joules): The ability to do work.
• Internal Energy (ΔE): Sum of all energies within a system.
• Heat (q): Energy transfer due to temperature difference.
• Work (w): Energy transfer due to force over a distance.

Summary

Enthalpy of reaction (ΔH) is a crucial measure in thermodynamics, indicating whether a reaction absorbs or releases heat. It helps predict reaction feasibility, understand energy transfers, and quantify heat changes. Mastering these concepts provides a strong foundation for tackling energy calculations and real-world applications in thermochemistry.