Laws of Thermodynamics

Introduction

Conservation of Energy:
- Energy cannot be created or destroyed, only transformed.
- Examples of energy forms: potential, kinetic, heat.
- Generally holds true in chemistry despite exceptions at the quantum level.
- Preferred direction in energy transformation.

Entropy:
- Difficult to grasp but can be described as disorder.
- Entropy must always increase in a closed system and its surroundings.
- Example: A messy bedroom tends to stay messy over time.
- Entropy measures energy dispersion in a system.
- Analogies:
  - Ionic solid vs. liquid states (solid = more ordered, liquid = more disordered).
  - Describing solid requires more information (geometry, intermolecular distances); liquid requires less (volume, shape).
Heat Transfer:
- Heat flows from hot to cold to increase disorder (entropy).

States that a perfectly crystalline solid at absolute zero has zero entropy.
Entropy is measured in joules per Kelvin.
It's a measure of how energy is distributed, not energy itself.

Relates to spontaneity of processes.
Change in Gibbs Free Energy Equation:
- Includes changes in enthalpy (ΔH), entropy (ΔS), and temperature (T).
- If ΔG < 0, process is spontaneous; if ΔG > 0, it's non-spontaneous.

Both favorable: ΔH negative, ΔS positive ➔ ΔG negative (spontaneous).
Both unfavorable: ΔH positive, ΔS negative ➔ ΔG positive (non-spontaneous).
One favorable requires calculation:
- Endothermic (ΔH positive) but entropically favorable can be spontaneous at high T.
- Exothermic (ΔH negative) but entropically unfavorable can be spontaneous at low T.

Incorrect conclusions suggest order can't happen spontaneously.
Example of soap:
- Soap forms micelles, with polar heads facing water and nonpolar tails trapping dirt.
- Micelles are water-soluble due to polar heads, showcasing how ordered structures can form spontaneously.