KJ

Daley_Student_Abbrev_Unit_II_2025_Chem152

Free Energy & Thermodynamics

  • Chapter 18 focuses on the principles of thermodynamics, particularly randomness, free energy, and equilibrium.

Key Thermodynamic Laws

First Law of Thermodynamics

  • Energy cannot be created or destroyed; it can only change forms.

  • Energy can transition from potential energy to kinetic energy.

Second Law of Thermodynamics

  • The entropy of the universe is always increasing.

  • Entropy (S) represents disorder; a solid with greater entropy has higher disorder.

Third Law of Thermodynamics

  • Entropy approaches zero as the temperature approaches absolute zero (0 K).

  • Absolute entropy for a perfect crystal at 0 K is 0 J/mol K.

  • All substances not at absolute zero have positive entropies due to particle dispersion.

Free Energy (G)

  • G, or Gibbs Free Energy, indicates the spontaneity of a process.

  • The equation: ΔG = ΔH - TΔS

  • A spontaneous process has ΔG < 0; non-spontaneous has ΔG > 0.

Spontaneous vs Nonspontaneous Processes

Spontaneous Processes

  • Occur naturally without external assistance (energy).

  • Does not indicate the reaction rate.

Nonspontaneous Processes

  • Cannot occur naturally; require continuous external energy input.

  • Examples include recharging batteries or endothermic reactions.

  • Can be coupled with spontaneous reactions to become spontaneous.

Entropy (S)

  • Entropy indicates the unavailability of energy for work, measured in J mol -1 K -1.

  • Increases with larger atomic mass and increased molecular complexity.

  • Dissolving substances increases entropy because the particles are more dispersed.

Entropy Change (ΔS)

  • Measures how energy disperses at a given temperature.

  • Positive ΔS indicates increased dispersal (favors spontaneity); negative ΔS indicates decreased dispersal (disfavors spontaneity).

  • More complex or larger systems have higher entropy.

Requirements for Spontaneous Reactions

  • Total entropy change of the universe must be positive: ΔSuniv = ΔSsys + ΔSsurr > 0.

  • If ΔSsys decreases, ΔSsurr must increase more than ΔSsys decreases to maintain positivity.

Heat Exchange and Surroundings (ΔS surr)

  • In exothermic reactions (ΔH sys < 0), heat is released to surroundings, increasing their entropy.

  • In endothermic reactions (ΔH sys > 0), heat is absorbed from surroundings, reducing their entropy.

  • The effect on surroundings depends on their initial temperature.

Gibbs Free Energy and Temperature Relation

  • ΔG = ΔH - TΔS; critical for determining spontaneity.

  • The temperature can determine the sign of ΔG.

  • Identifying where reactions switch from spontaneous to non-spontaneous based on ΔH and ΔS conditions.

Equilibrium State

  • Equilibrium occurs when the forward and reverse processes happen at the same rate.

  • At equilibrium, ΔG = 0.

  • All reactions move naturally toward equilibrium.

Phase Changes

  • During phase changes, temperature remains constant, and systems are at equilibrium (ΔG = 0).

  • Latent heat is the energy absorbed or released without temperature change during phase transitions.

Calculating Changes in Free Energy

  • The Gibbs free energy change is also calculated from stoichiometric changes in reactions:

    • ΔG° rxn = ∑ ΔG° products - ∑ ΔG° reactants.

  • Summary of various reactions specifying their spontaneity, heat of fusion, and vaporization properties.

Non-standard Conditions of Concentration and Pressure

  • A non-spontaneous reaction can become spontaneous by adjusting concentrations and/or pressures.

  • Pure solids and liquids are excluded from the Q equation; concentration changes are only for gases and aqueous solutions.