Class 11 Chemistry Thermodynamics Notes

1. Introduction to Thermodynamics

Thermodynamics is the study of energy changes in physical and chemical processes.
It deals with heat, work, and internal energy of a system.
The system is the part of the universe under study; the surroundings include everything else.
Types of systems:
1. Open system – exchanges both matter and energy with surroundings.
2. Closed system – exchanges only energy, not matter.
3. Isolated system – exchanges neither matter nor energy.

2. Types of Energy

Kinetic Energy (KE): Energy due to motion.
Potential Energy (PE): Energy due to position or configuration.
Internal Energy (U): Sum of all kinetic and potential energies of molecules in a system.

3. First Law of Thermodynamics

Statement: Energy can neither be created nor destroyed; it can only change from one form to another.
Mathematical Form:

$\Delta U = q – w$ ΔU=q−w

where,
$\Delta U$ ΔU = change in internal energy,
$q$ q = heat absorbed by the system,
$w$ w = work done by the system.

Important Points:
- Heat absorbed increases internal energy.
- Work done by the system decreases internal energy.

4. Work and Heat

Work done by a gas:

$w = P \Delta V$ w=PΔV

where $P$ P = pressure, $\Delta V$ ΔV = change in volume.

Heat (q): Energy transferred due to temperature difference.

5. Enthalpy (H)

Definition: Total heat content of a system at constant pressure.

$H = U + PV$ H=U+PV

Change in Enthalpy (ΔH\Delta HΔH):
- Exothermic reaction: $\Delta H < 0$ ΔH<0 (heat released)
- Endothermic reaction: $\Delta H > 0$ ΔH>0 (heat absorbed)
Standard Enthalpy of Formation (ΔHf∘\Delta H_f^\circΔHf∘): Heat change when 1 mole of compound forms from elements in standard states.
Standard Enthalpy of Combustion (ΔHc∘\Delta H_c^\circΔHc∘): Heat released when 1 mole of substance burns completely in oxygen.

6. Hess’s Law of Constant Heat Summation

Statement: Total enthalpy change of a reaction is independent of the pathway and depends only on initial and final states.
Useful in calculating enthalpy changes of reactions that are difficult to measure directly.

7. Second Law of Thermodynamics

Statement: Heat cannot spontaneously flow from a colder body to a hotter body.
Introduces entropy (S) – a measure of disorder or randomness.
Spontaneous process: $\Delta S_{universe} > 0$ ΔSuniverse>0
Non-spontaneous process: $\Delta S_{universe} < 0$ ΔSuniverse<0

8. Gibbs Free Energy (G)

Definition: Predicts spontaneity of a reaction at constant temperature and pressure.

$\Delta G = \Delta H – T\Delta S$ ΔG=ΔH−TΔS

Interpretation:
- $\Delta G < 0$ ΔG<0 → Spontaneous
- $\Delta G > 0$ ΔG>0 → Non-spontaneous
- $\Delta G = 0$ ΔG=0 → Equilibrium

9. Key Points to Remember

Energy is conserved (First Law).
Entropy tends to increase (Second Law).
Exothermic reactions release heat, endothermic reactions absorb heat.
Gibbs free energy combines enthalpy and entropy to predict reaction spontaneity.

5.1 Thermodynamic Terms

System: Part of the universe under study.
Surroundings: Everything outside the system.
Types of Systems:
- Open system: Exchanges energy and matter with surroundings.
- Closed system: Exchanges only energy.
- Isolated system: Exchanges neither energy nor matter.
State Functions: Properties that depend only on the current state of the system (e.g., internal energy (U), enthalpy (H), entropy (S)).
Process Functions: Depend on the path taken (e.g., heat (q), work (w)).

5.2 Applications of Thermodynamics

Explains chemical reactions and physical changes in terms of energy transfer.
Helps in predicting reaction spontaneity.
Determines equilibrium conditions for reactions.
Applications in calorimetry, engines, refrigeration, and industrial processes.

5.3 Measurement of ∆U and ∆H: Calorimetry

Internal Energy Change (∆U): Measured using calorimeters.

$\Delta U = q_v \quad \text{(at constant volume)}$ ΔU=qv(at constant volume)

Enthalpy Change (∆H): Measured at constant pressure.

$\Delta H = q_p$ ΔH=qp

Calorimetry: Technique to measure heat changes during chemical or physical processes.
- Bomb Calorimeter: Used for combustion reactions at constant volume.
- Coffee Cup Calorimeter: Measures reactions at constant pressure.

5.4 Enthalpy Change, ∆rH of a Reaction – Reaction Enthalpy

Reaction Enthalpy (∆rH): Heat change when a reaction occurs at constant pressure.
Exothermic Reaction: ∆rH < 0, heat is released.
Endothermic Reaction: ∆rH > 0, heat is absorbed.
Standard Enthalpy Change (∆rH°): Measured under standard conditions (1 atm, 298 K).

5.5 Enthalpies for Different Types of Reactions

Enthalpy of Formation (∆fH°): Heat change when 1 mole of compound forms from its elements.
Enthalpy of Combustion (∆cH°): Heat released when 1 mole of substance burns completely in oxygen.
Enthalpy of Neutralization (∆neutH): Heat released when an acid reacts with a base to form 1 mole of water.
Enthalpy of Reaction (∆rH°): Heat change for any reaction at standard conditions.

5.6 Spontaneity

A reaction is spontaneous if it occurs naturally without external intervention.
Factors affecting spontaneity:
- Enthalpy change (∆H): Negative favors spontaneity.
- Entropy change (∆S): Positive favors spontaneity.
Second Law of Thermodynamics:
- Entropy of the universe increases for a spontaneous process:
ΔSuniverse>0\Delta S_{universe} > 0ΔSuniverse>0

5.7 Gibbs Energy Change and Equilibrium

Gibbs Free Energy (G): Combines enthalpy and entropy to predict spontaneity.