Chapter 5 Thermodynamics Class11

Thermodynamics 

By Aarish Sir

1. Introduction

• Thermodynamics is the branch of chemistry that deals with the study of energy changes, especially heat and work during chemical reactions and physical processes.

• It tells us whether a process is possible or not, but not how fast (that is studied in kinetics).

• Energy is conserved, but it can be converted from one form to another.

2. Important Terms

1. System and Surroundings

   • System: Part of the universe under study (e.g., reaction mixture).

   • Surroundings: Everything outside the system.

   • Universe = System + Surroundings.

   Types of systems:

   • Open system: Exchange of matter and energy (e.g., open beaker of water).

   • Closed system: Exchange of only energy, not matter (e.g., closed vessel).

   • Isolated system: No exchange of matter or energy (e.g., thermos flask).

2. State of a System

   • Described by properties like temperature, pressure, volume, composition, etc.

   • State function: Property depending only on state, not path (e.g., ΔE, ΔH, P, V, T).

   • Path function: Depends on path taken (e.g., heat, work).

3. Work (w)

   • Work done by/on the system due to volume change.

   • w = –PΔV (P = external pressure).

4. Heat (q)

   • Energy transferred due to temperature difference.

5. Internal Energy (U)

   • Total energy of a system (kinetic + potential).

   • Change in internal energy (ΔU) = q + w.

3. First Law of Thermodynamics

• Energy can neither be created nor destroyed, only transformed.

• Mathematically:
  ΔU = q + w

• If system absorbs heat (q > 0), or work is done on system (w > 0) → ΔU increases.

• If system loses heat (q < 0), or does work on surroundings (w < 0) → ΔU decreases.

4. Enthalpy (H)

• Heat change at constant pressure.

• H = U + PV

• Change in enthalpy (ΔH) = q (at constant pressure).

Interpretation:

• ΔH > 0 → Endothermic (absorbs heat).

• ΔH < 0 → Exothermic (releases heat).

Examples:

• Combustion of fuel → exothermic.

• Photosynthesis → endothermic.

5. Types of Processes

1. Isothermal process

   • Temperature constant (ΔT = 0).

   • ΔU = 0, so q = –w.

2. Adiabatic process

   • No heat exchange (q = 0).

   • ΔU = w.

3. Isobaric process

   • Pressure constant.

   • q = ΔH.

4. Isochoric process

   • Volume constant (ΔV = 0, so w = 0).

   • q = ΔU.

6. Heat Capacity

• Heat capacity (C): Amount of heat required to raise temperature by 1 K.

  • C = q/ΔT.

• Molar heat capacity: Heat capacity per mole.

• Specific heat capacity: Heat capacity per unit mass.

• Relation: Cp – Cv = R (for ideal gases).

  • Cp = heat capacity at constant pressure.

  • Cv = heat capacity at constant volume.

7. Enthalpy Changes in Reactions

1. Enthalpy of Reaction (ΔHrxn): Heat change during reaction.

2. Enthalpy of Combustion (ΔHc): Heat change when 1 mole of substance completely burns in oxygen.

   • Example: CH₄ + 2O₂ → CO₂ + 2H₂O (ΔHc < 0).

3. Enthalpy of Formation (ΔHf): Heat change when 1 mole of compound forms from its elements.

   • Example: H₂ + ½O₂ → H₂O (l).

4. Enthalpy of Neutralisation (ΔHn): Heat change when acid reacts with base to form 1 mole water.

   • Example: HCl + NaOH → NaCl + H₂O (ΔHn ≈ –57 kJ).

5. Enthalpy of Atomisation (ΔHat): Heat change to convert 1 mole of element into atoms.

6. Enthalpy of Bond Dissociation (ΔHd): Heat required to break 1 mole of bonds.

7. Enthalpy of Sublimation (ΔHs): Heat change when solid directly changes into gas.

8. Enthalpy of Vaporisation (ΔHv): Heat required to convert liquid into vapour.

9. Enthalpy of Fusion (ΔHf): Heat required to convert solid into liquid.

8. Hess’s Law of Constant Heat Summation

• Statement: The total enthalpy change of a reaction is the same, whether it takes place in one step or many steps.

• Important because enthalpy is a state function.

Example:
C(graphite) + O₂ → CO₂ (ΔH = –393.5 kJ)
C(graphite) + ½O₂ → CO (ΔH = –110.5 kJ)
CO + ½O₂ → CO₂ (ΔH = –283 kJ)

Adding 2nd + 3rd gives the 1st reaction.
ΔH is same.

9. Spontaneity and Second Law of Thermodynamics

1. Spontaneous Process: Occurs on its own without external help.

   • Example: Heat flows from hot to cold, ice melts at room temp.

2. Entropy (S): Measure of randomness/disorder.

   • ΔS > 0 → randomness increases.

   • ΔS < 0 → randomness decreases.

3. Second Law: For a spontaneous process, total entropy of system + surroundings increases.

10. Gibbs Free Energy (G)

• G = H – TS

• ΔG = ΔH – TΔS

Criteria for spontaneity:

• ΔG < 0 → Process spontaneous.

• ΔG > 0 → Non-spontaneous.

• ΔG = 0 → System in equilibrium.

11. Important Summary Table

Quantity	Symbol	Condition	Spontaneity
ΔH < 0, ΔS > 0	–	Always spontaneous	Exothermic + more disorder
ΔH > 0, ΔS < 0	–	Never spontaneous	Endothermic + less disorder
ΔH < 0, ΔS < 0	Low T	Spontaneous
ΔH > 0, ΔS > 0	High T	Spontaneous

12. Applications of Thermodynamics

• Explains energy changes in chemical reactions.

• Helps in designing engines, refrigerators, and power plants.

• Predicts feasibility of reactions (via ΔG).

• Important in understanding biological processes like respiration, photosynthesis.