Learning Objectives
- Understand thermodynamic systems, surroundings, and state functions
- Learn the first law of thermodynamics and enthalpy
- Study Hess's law and standard enthalpies of reactions
- Understand the second and third laws of thermodynamics
- Apply Gibbs free energy to predict spontaneity of reactions
Key Concepts
Thermodynamic System and State Functions
System: Part of universe under study. Surroundings: Everything else.
Types: Open (exchanges energy and matter), Closed (exchanges only energy), Isolated (no exchange).
State functions: Depend only on state, not path (U, H, S, G, P, V, T).
Path functions: Depend on the process (q, w).
Extensive properties: Depend on amount (U, H, S, V, mass). Intensive: Independent of amount (T, P, density, molar quantities).
First Law of Thermodynamics
ΔU = q + w (IUPAC convention: q = heat absorbed, w = work done on system).
For expansion work: w = -P_ext ΔV (work done by system against external pressure).
For reversible isothermal expansion of ideal gas: w = -nRT ln(V₂/V₁) = -2.303 nRT log(V₂/V₁).
Enthalpy (H)
H = U + PV. At constant pressure: ΔH = q_p (heat at constant pressure).
At constant volume: ΔU = q_v. Relation: ΔH = ΔU + Δn_g RT (Δn_g = moles of gaseous products - gaseous reactants).
Standard Enthalpies
Standard enthalpy of formation (ΔfH°): Enthalpy change when 1 mole of compound is formed from elements in standard states. ΔfH° of elements = 0.
ΔrH° = ΣΔfH°(products) - ΣΔfH°(reactants)
Standard enthalpy of combustion: Enthalpy when 1 mole burns completely in O₂.
Bond enthalpy: ΔrH° = Σ(bond enthalpies of reactants) - Σ(bond enthalpies of products).
Hess's Law
Total enthalpy change is the same regardless of the route taken (since H is a state function).
Used to calculate enthalpies that cannot be directly measured.
Entropy (S)
Measure of disorder or randomness. ΔS = q_rev/T.
Entropy increases with: temperature increase, phase change (solid → liquid → gas), increase in number of gaseous moles, dissolution.
Second Law: For a spontaneous process, the total entropy of the universe increases. ΔS_universe = ΔS_system + ΔS_surroundings > 0.
Third Law: Entropy of a perfect crystal at absolute zero (0 K) is zero.
Gibbs Free Energy (G)
G = H - TS. At constant T and P: ΔG = ΔH - TΔS
Spontaneity criteria:
- ΔG < 0: Spontaneous (exergonic)
- ΔG = 0: Equilibrium
- ΔG > 0: Non-spontaneous (endergonic)
Relation to equilibrium: ΔG° = -RT ln K = -2.303 RT log K.
ΔG° = ΣΔfG°(products) - ΣΔfG°(reactants)
Predicting Spontaneity from ΔH and ΔS
- ΔH < 0, ΔS > 0: Always spontaneous
- ΔH > 0, ΔS < 0: Never spontaneous
- ΔH < 0, ΔS < 0: Spontaneous at low T
- ΔH > 0, ΔS > 0: Spontaneous at high T (when TΔS > ΔH)
Summary
Thermodynamics deals with energy changes in chemical reactions. The first law relates internal energy, heat, and work. Enthalpy changes at constant pressure are measured calorimetrically. Hess's law allows indirect calculation of enthalpies. Entropy measures disorder and increases in spontaneous processes. Gibbs free energy combines enthalpy and entropy to predict spontaneity: ΔG < 0 for spontaneous reactions.
Important Terms
- Enthalpy (H): Heat content at constant pressure, H = U + PV
- Entropy (S): Measure of randomness or disorder
- Gibbs Free Energy (G): Maximum useful work at constant T, P
- Hess's Law: ΔH is path-independent (state function)
- Exothermic: ΔH < 0 (heat released)
- Endothermic: ΔH > 0 (heat absorbed)
Quick Revision
- ΔU = q + w; ΔH = ΔU + Δn_gRT
- ΔrH° = ΣΔfH°(products) - ΣΔfH°(reactants)
- ΔG = ΔH - TΔS; ΔG < 0 → spontaneous
- ΔG° = -2.303 RT log K
- ΔS_universe > 0 for spontaneous process
- S = 0 at 0 K for perfect crystal (Third Law)