Thermodynamics Revision Notes

Overview

• Thermodynamics: Study of heat
• Aim: Quick revision of theory and formulae

Key Terminologies

• System: Part of the universe under observation
• Surroundings: Everything outside the system
• Boundary: Separates system from surroundings; can be real or imaginary

Types of Systems

• Open System: Exchanges both energy and matter with surroundings (e.g., human body)
• Closed System: Exchanges only energy with surroundings
• Isolated System: Exchanges neither energy nor matter (ideal systems)

State Variables

• Describes the condition of a system: pressure, volume, temperature, number of moles

Properties of Systems

• Extensive Properties: Depend on size (e.g., total energy, volume, mass)
• Intensive Properties: Independent of size (e.g., pressure, density)

Thermodynamic Processes

1. Isothermal: Temperature remains constant (ΔT = 0)
2. Isobaric: Pressure remains constant (ΔP = 0)
3. Isochoric: Volume remains constant (ΔV = 0)
4. Adiabatic: No heat transfer (Q = 0)
5. Cyclic: System returns to initial state

Energy Concepts

• Heat (Q): Energy transfer due to temperature differences; flows from higher to lower temperatures
• Work (W): Energy transfer via mechanical means; can be compression or expansion
  • Formula: W = -P_external dV (integration of PV graph)
  • Sign Convention:
    • Heat supplied to system: Positive
    • Heat removed from system: Negative
    • Work done on the system: Positive
    • Work done by the system: Negative

Internal Energy (U)

• Sum of kinetic and potential energy of internal components
• For ideal gases: U depends only on temperature
• First Law of Thermodynamics: ΔU = Q + W

Enthalpy (H)

• Defined as H = U + PV
• Heat exchanged at constant pressure: ΔH = Q_P
• Heat exchanged at constant volume: ΔU = Q_V

Heat Capacity

• Heat required to raise the temperature by 1 K
• Formula: C = Q / ΔT
• Molar heat capacities:
  • C_P = Q_P / (N ΔT)
  • C_V = Q_V / (N ΔT)

Ideal Gas Relationships

• For monoatomic gases:
  • C_V = (3/2)R, C_P = (5/2)R
• For diatomic gases:
  • C_V = (5/2)R, C_P = (7/2)R
• Relationship: C_P = C_V + R

Important Formulas

• Isothermal Process: W = nRT ln(V_1/V_2)
• Isobaric Process: W = -PΔV
• Isochoric Process: W = 0; Q_V = ΔU = N C_V ΔT
• Adiabatic Process: PV^γ = constant

Entropy (S)

• Measure of randomness; ΔS = Q_reversible / T
• Second Law of Thermodynamics: Entropy of isolated systems tends to increase
• Spontaneous processes have ΔS_total > 0
• For reversible processes: ΔS_universe = 0

Gibbs Free Energy (G)

• ΔG = -T ΔS_total
• At constant temperature and pressure: ΔG = ΔH - T ΔS

Types of Enthalpy

• Enthalpy of Formation: Energy change when 1 mole of compound forms from elements
• Enthalpy of Combustion: Energy change during combustion
• Enthalpy of Solution: Energy change when dissolving 1 mole in excess solvent
• Enthalpy of Hydration: Energy change when adding water of crystallization
• Enthalpy of Neutralization: Energy change when acid and base react
• Bond Dissociation Energy: Energy needed to break a bond

Hess's Law

• Based on the fact that enthalpy is a state function
• ΔH = ΔH1 + ΔH2 + ΔH3 for combined reactions

Ideal Processes

• Free Expansion: No work done, no heat exchange, ΔU = 0, but ΔS > 0
• Polytropic Process: PV^n = constant (n is a real number)

Conclusion

• Revision of thermodynamics covered theory and formulae
• Refer to PDF in the description for detailed notes
• Next session: Question-solving in thermodynamics