Thermodynamics Lecture Notes Review

Thermodynamics Notes

Page 2: Spontaneous Physical and Chemical Processes

• Spontaneous Processes:

  • A waterfall runs downhill.
  • A lump of sugar dissolves in a cup of coffee.
  • Ice melts above 0 °C and freezes below 0 °C at 1 atm.
  • Heat flows from hot to cold objects.
  • A gas expands in a vacuum.
  • Iron rusts when exposed to oxygen and water.

• Definitions:

  • Spontaneous: Processes that occur without requiring external work.
  • Nonspontaneous: Processes that do not occur without external influence.

Page 4: Reactions and Enthalpy

• A decrease in enthalpy (ΔH) does not guarantee that a reaction will be spontaneous.
• Example Equations:
  • H^+(aq) + OH^-(aq) \rightarrow H2O(l) \quad ΔH^0 = -56.2 \text{ kJ/mol}
  • H2O(s) \rightarrow H2O(l) \quad ΔH^0 = +6.01 \text{ kJ/mol}
  • NH4NO3(s) \rightarrow NH4^+(aq) + NO3^-(aq) \quad ΔH^0 = +25 \text{ kJ/mol}
  • CH4(g) + 2O2(g) \rightarrow CO2(g) + 2H2O(l) \quad ΔH^0 = -890.4 \text{ kJ/mol}

Page 5: Entropy (S)

• Entropy (S): Measure of randomness or disorder of a system.
  • Transition: ΔS = Sf - Si
  • Key Concept: For a spontaneous process where randomness increases, ΔS > 0.
  • Order: S{solid} < S{liquid} << S_{gas}
  • Example: Transition from solid to liquid increases entropy, hence:
    • H2O(s) \rightarrow H2O(l), \ ΔS > 0

Page 6: Microstates and Entropy

• Formula: W = ext{number of microstates}
• S = k \ln W
• Changes in entropy can be determined by:
  • ΔS = Sf - Si
  • ΔS = k \ln \left( \frac{Wf}{Wi} \right)

Page 9: State Functions

• State Functions: Properties dependent only on the state of a system, not the path taken. Examples include energy, enthalpy, pressure, volume, temperature, and entropy.
• Analogy: Two hikers at the same height have the same potential energy regardless of their paths.

Page 14: Laws of Thermodynamics

• First Law: Energy can be converted but not created/destroyed.
• Second Law: Total entropy of the universe increases in spontaneous processes:
  • ΔS{univ} = ΔS{sys} + ΔS_{surr}
  • For spontaneous processes: ΔS_{univ} > 0
  • For equilibrium processes: ΔS_{univ} = 0

Page 21: Entropy Changes in the System (ΔS_sys)

• if a reaction produces more gas molecules (moles), then ΔS_{rxn} > 0.
• If a reaction consumes gas molecules, then ΔS_{rxn} < 0.
• No net change may result in small ΔS_{rxn}.

Page 27: Third Law of Thermodynamics

• At absolute zero, the entropy of a perfect crystalline substance is zero: S = k \ln W where W = 1, S = 0

Page 28: Gibbs Free Energy

• Gibbs Free Energy equation: ΔG = ΔH{sys} - TΔS{sys}
  • Interpretation:
  • ΔG < 0: Spontaneous in the forward direction.
  • ΔG > 0: Nonspontaneous.
  • ΔG = 0: System is in equilibrium.

Page 56: Example of Coupling Reactions

• Coupling nonspontaneous processes with spontaneous ones (e.g., lifting a weight by letting another weight fall).

Page 58: ATP and Coupled Reactions

• Reaction Example: ATP + H2O + ext{Amino Acids} \rightarrow ADP + H3PO4 + ext{Peptides}
  • Reaction energetics: Determine if K < 1 (nonspontaneous) or K > 1 (spontaneous).

Pathways in biochemical reactions typically involve coupled reactions to drive nonspontaneous processes.