Chemical Thermodynamics

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Chemical Thermodynamics 

1. Basic Concepts

System: Part of universe chosen for study.
Surroundings: Everything else outside the system.
Boundary: Real/imaginary surface separating system and surroundings.

Types of Systems:
- Open System: Exchange of matter + energy.
- Closed System: Exchange of energy only, not matter.
- Isolated System: No exchange of matter or energy.

2. State & Path Functions

State functions: Depend only on initial & final states (P, V, T, ΔU, ΔH, ΔG, S).
Path functions: Depend on path (work, heat).

3. Internal Energy (U)

Total energy of molecules.
First Law of Thermodynamics:
ΔU = q + w
where q = heat absorbed, w = work done on system.

4. Work in Expansion/Compression

Constant pressure:
w = -Pext(Vf - Vi)

Reversible expansion of ideal gas:
w = -nRT ln(Vf/Vi)

5. Enthalpy (H)

Definition: H = U + PV
Change in enthalpy: ΔH = ΔU + PΔV
At constant pressure: ΔH = qp

6. Relation between ΔH and ΔU

For reactions involving gases:
ΔH = ΔU + Δn(gas)RT
where Δn(gas) = n(products) – n(reactants) (gaseous moles only).

7. Heat Capacity

C = dq/dT
At constant volume: Cv = (dq/dT)V
At constant pressure: Cp = (dq/dT)P
Mayer’s relation: Cp – Cv = R

8. Hess’s Law

Enthalpy change is independent of path, depends only on initial and final states.

9. Entropy (S)

Measure of disorder/randomness.
Change in entropy: ΔS = qrev/T

10. Gibbs Free Energy (G)

Definition: G = H – TS
Change in free energy: ΔG = ΔH – TΔS

Spontaneity conditions:
- ΔG < 0 → spontaneous
- ΔG > 0 → non-spontaneous
- ΔG = 0 → equilibrium

11. Laws of Thermodynamics

Zeroth Law:
If two systems are in thermal equilibrium with a third, they are in thermal equilibrium with each other. Basis of temperature measurement.

First Law (Law of Conservation of Energy):
ΔU = q + w

Second Law:
ΔSuniverse = ΔSsystem + ΔSsurroundings > 0
Heat cannot flow spontaneously from colder body to hotter body.

Third Law:
Entropy of a pure crystalline substance at absolute zero (0 K) is zero (S = 0 at T = 0 K).

12. Standard Enthalpies

ΔHf°: Enthalpy change when 1 mole compound is formed from elements in standard states.
ΔHc°: Enthalpy change when 1 mole of substance is completely burnt in O2.
ΔHr°: Enthalpy change for a reaction under standard conditions.

13. Important Derivations

1. First Law: ΔU = q + w
2. Work of isothermal reversible expansion of ideal gas:
   w = -nRT ln(Vf/Vi)
3. Relation between ΔH and ΔU:
   ΔH = ΔU + ΔngasRT
4. Mayer’s relation:
   Cp – Cv = R
5. Gibbs free energy:
   ΔG = ΔH – TΔS

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