Unit 9: Advanced Thermodynamics and Electrochemistry

Unit 9: Applications of Thermodynamics

This unit connects energy changes (Thermodynamics) with equilibrium and the ability of chemical reactions to do work (Electrochemistry). It answers the fundamental question: Will a reaction happen without outside intervention?

Entropy and the Second Law

Defining Entropy ($S$)

Entropy ($S$) is often described as a measure of disorder, but in AP Chemistry, it is more accurately defined as the dispersal of energy (or matter) at a specific temperature. The Second Law of Thermodynamics states that the entropy of the universe increases for any spontaneous process.

  • State Function: Like Enthalpy ($ H$), Entropy depends only on the final and initial states, not the path taken.
  • Units: Standard entropy ($S^\circ$) is usually measured in $J/(mol\cdot K)$.
    • Note: $\Delta H$ is usually in $kJ$, while $\Delta S$ is in $J$. This unit mismatch is a major trap in calculations!

Predicting Entropy Changes ($ \Delta S^\circ$)

A positive $\Delta S^\circ$ indicates increased dispersal/disorder. You must be able to predict the sign of $\Delta S$ based on physical changes:

  1. Phase Changes: Solid $\to$ Liquid $\to$ Gas (Entropy increases significantly).
    • $S{solid} < S{liquid} \ll S_{gas}$
  2. Number of Moles of Gas: If a reaction produces more moles of gas than it consumes, entropy increases ($ \Delta S > 0$).
    • Example: $2H2O(l) \to 2H2(g) + O_2(g)$ (0 mol gas $\to$ 3 mol gas; $\Delta S$ is positive).
  3. Dissolution: Dissolving a solid or liquid in a solvent generally increases entropy (particles are more dispersed).
  4. Temperature: Increasing temperature increases molecular motion and entropy.
  5. Complexity: Larger, more complex molecules generally have higher absolute entropy than smaller, simpler ones due to more possible arrangements of atoms (microstates).

Calculating Standard Entropy Change

The calculation is straightforward