Classical Thermodynamics

Classical Thermodynamics

• It is a branch of Chemistry focused on the study of energy change in chemical reactions or physical processes (chemical energetics).
• Thermodynamics is derived from the Greek words "therme" (heat) and "dynamics" (flow).
• It studies heat and other energy forms, informing about heat conversion into work, the relation between heat and other energy forms, and process efficiency.

Limitations of Thermodynamics

• Laws apply only to matter in bulk (macroscopic systems), not individual atoms or molecules.
• It predicts process feasibility but doesn't reveal the rate (how fast or slow).
• Focuses solely on the initial and final states, not the process path.

Thermodynamic Terms and Concepts

System

• Specific part of the universe under experimental study, separated by a boundary.

Surroundings

• Part of the universe other than the system, able to exchange energy and matter.

Boundary

• Real or imaginary surface separating the system from surroundings.

System Types

• Homogeneous: System with uniform contents (single phase, e.g., pure solid, gas mixture).
• Heterogeneous: System with multiple phases unevenly distributed (e.g., ice in water).

Types of Systems

• Open System: Exchanges both matter and energy with surroundings (e.g., tea in a cup).
• Closed System: Exchanges only energy, not matter, with surroundings; non-insulated but sealed (e.g., heating iodine in a sealed container).
• Isolated System: Exchanges neither matter nor energy; insulated and sealed (e.g., ice in a thermos flask).

State Functions

• Thermodynamic parameters depending only on initial and final states, independent of path (e.g., internal energy, enthalpy, entropy, Gibbs free energy, pressure, temperature, volume).

Path Functions

• Thermodynamic properties dependent on path (e.g., heat, work).

State Variables

• Intensive properties: State functions independent of the amount of substance (e.g., pressure, temperature, concentration, density, viscosity, surface tension, specific heat).
• Extensive properties: State variables dependent on the amount of substance (e.g., volume, mass, internal energy, enthalpy, entropy, work).

Thermodynamic Process

• Operation by which a thermodynamic system changes from one state to another.

Types of Thermodynamic Processes

• Isothermal: Temperature remains constant (ΔT = 0; e.g., freezing, melting, evaporation).
• Isobaric: Pressure remains constant (ΔP = 0; e.g., heating water at its boiling point).
• Isochoric: Volume remains constant (ΔV = 0; e.g., heating a substance in a closed system).
• Adiabatic: No heat exchange with surroundings (Δq = 0; e.g., process in a thermos bottle).
• Cyclic: System returns to its initial state after several steps (ΔE = 0, ΔH = 0).

Reversible Processes

• Change occurs infinitesimally slowly, and direction is reversible by an infinitesimal change.

Irreversible processes

• Proceeds spontaneously in a single step to equilibrium and cannot be reversed (e.g, combustion of methane).

Thermodynamic Equations

• ΔU = Q + W.
• W = -PΔV.
• H = U + PV.
• G = H - TS.
• A = U - TS.

Heat (Q)

• Defined as energy flowing between system and surroundings due to temperature difference.
• Flows from high to low temperature.
• Heat flow from a system is negative (-ve).
• Heat flow into a system is positive (+ve).
• Measured in Joules (J) or calories (cal).
• 1 kJ = 10^3 J, 1 cal = 4.184 J, 1 kcal = 4.184 kJ, 1 J = 10^7 ergs, 1 liter-atmosphere = 101.33 J.

Work (W)

• Transference of energy occurring when system and surroundings have different pressures.
• For an object displacing distance dx against force F, the work done is equal to Fdx.
• W = Fdx = PAdV = PΔV (pressure x volume change).
• Work done on the system by surroundings: W = +PAV.
• Work done by the system: W = -PAV.

Internal Energy (U)

• Energy associated with a thermodynamic system.
• Total of all possible energy forms within the system.

Properties of Internal Energy

• Extensive property.
• State function
• No change in internal energy for a cyclic process.

Exothermic reactions

• ΔU is negative (-ve).

Endothermic reactions

• ΔU is positive (+ve).

Units for measuring reactions

• Joule or calorie.

Exothermic Reactions

• Reactions accompanied by heat evolution.
• Example: C(s) + O2(g) -> CO2(g) + 393.5 kJ.
• ΔH = Hp - HR, HR > HP, ΔH = -ve.

Endothermic Reactions

• Reactions accompanied by heat absorption.
• Example: N2(g) + O2(g) -> 2NO(g) - 180.5 kJ.
• ΔH = Hp - HR, HP > HR, ΔH = +ve.

Enthalpy (H)

• Defined as heat content of a system at constant pressure.
• Sum of internal energy and pressure-volume product: H = E + PV.

Calculating Enthalpy

• Since E is a state function, we determine change in enthalpy (ΔH), not absolute value. The formula: ΔH = ΔE + Δ(PV).
• ΔH = ΔE + PΔV + VΔP.
• At constant pressure (ΔP = 0): ΔH = ΔE + PΔV.
• Relationship between ΔH and ΔU: ΔH = ΔU + PAV.
• Using Ideal gas equation (PV = nRT):

Isothermal expansion with an ideal gas

• W = -nRT ln(V2/V1).