Unit 1: Chemical Thermodynamics I

Course Code: CHM-DSC-251

Paper Name: Physical Chemistry - II (Chemical Thermodynamics & Equilibrium)

1. Thermodynamic Variables and Zeroth Law
2. The First Law of Thermodynamics
3. Expansion of Gases: Reversible and Irreversible
4. Thermochemistry: Heats of Reaction
5. Bond Energy and Resonance Energy
6. Temperature Dependence (Kirchhoff’s Equation)
7. Exam Focus: Tips and FAQs

1. Thermodynamic Variables and Zeroth Law

Thermodynamics deals with the study of energy transformations in macroscopic systems. Understanding the nature of variables is fundamental to defining the state of a system.

Intensive and Extensive Variables

Variable Type	Definition	Examples
Intensive	Properties that are independent of the amount of substance present.	Temperature, Pressure, Density, Viscosity, Molar Volume.
Extensive	Properties that depend on the mass or size of the system.	Mass, Volume, Internal Energy, Enthalpy, Entropy.

State and Path Functions

A State Function depends only on the current state of the system and is independent of the path taken to reach that state (e.g., U, H, P, V, T). A Path Function depends on the specific route taken (e.g., Heat, Work).

Zeroth Law of Thermodynamics

If two systems are each in thermal equilibrium with a third system, then they are in thermal equilibrium with each other.

This law provides the basis for the measurement of temperature.

2. The First Law of Thermodynamics

The First Law is essentially the Law of Conservation of Energy applied to thermodynamic systems.

Fundamental Concepts

Internal Energy (U): The total energy contained within a system. It is a state function.
Heat (q): Energy transferred due to a temperature difference.
Work (w): Energy transfer through mechanical means (e.g., expansion against pressure).

Mathematical Statement

ΔU = q + w

For an infinitesimal change: dU = dq + dw.

Enthalpy (H)

Enthalpy is the total heat content of a system at constant pressure.

H = U + PV

At constant pressure, the change in enthalpy is equal to the heat absorbed: ΔH = qp.

3. Expansion of Gases: Reversible and Irreversible

The work done and heat exchanged depend heavily on the conditions of expansion (Isothermal vs. Adiabatic) and the nature of the process.

Isothermal Expansion (Ideal Gas)

In an isothermal process, temperature remains constant (ΔT = 0), so ΔU = 0 for an ideal gas.

Reversible Work: w = -nRT ln(V2/V1) or w = -2.303 nRT log(V2/V1).
Irreversible Work: w = -Pext (V2 - V1).

Adiabatic Expansion

In an adiabatic process, no heat enters or leaves the system (q = 0).

ΔU = w

For an ideal gas, the relation is P V^γ = constant, where γ (gamma) is the ratio of heat capacities (Cp/Cv).

4. Thermochemistry: Heats of Reaction

Thermochemistry studies the heat changes accompanying chemical reactions.

Standard States

The Standard State of a substance is its most stable form at 1 bar pressure and a specified temperature (usually 298 K).

Key Enthalpy Changes

Enthalpy of Formation (ΔHf): Heat change when 1 mole of a compound is formed from its elements in their standard states.
Enthalpy of Combustion (ΔHc): Heat released when 1 mole of a substance is completely burnt in oxygen.

Note: By convention, the standard enthalpy of formation for elements in their most stable state (e.g., O2 gas, C graphite) is zero.

5. Bond Energy and Resonance Energy

Thermochemical data allows us to calculate the strength of chemical bonds.

Bond Dissociation Energy vs. Bond Energy

Bond Dissociation Energy: The energy required to break a specific single bond in a gaseous molecule.
Average Bond Energy: The average value of dissociation energies for the same type of bond in different compounds.

Resonance Energy

Resonance energy is the difference between the experimental enthalpy of formation and the value calculated based on a localized structure (like the Kekulé structure for benzene).

Resonance Energy = ΔH (experimental) - ΔH (calculated)

6. Temperature Dependence (Kirchhoff’s Equation)

The heat of a reaction changes with temperature if the heat capacities of the reactants and products are different.

Kirchhoff’s Equations

1. For Enthalpy Change (constant pressure):

ΔH2 - ΔH1 = ΔCp (T2 - T1)

2. For Internal Energy Change (constant volume):

ΔU2 - ΔU1 = ΔCv (T2 - T1)

Where ΔCp = Σ Cp (products) - Σ Cp (reactants).

7. Exam Focus: Tips and FAQs

Exam Tips:

Always check units! R (gas constant) can be 8.314 J/mol·K or 0.0821 L·atm/mol·K. Use the one that matches your pressure/volume units.
In Isothermal expansion of an ideal gas, ΔU and ΔH are always 0 because they only depend on temperature.
Remember the sign convention: Work done by the system is negative (-), work done on the system is positive (+).

Common Mistakes

Confusing state functions (U, H) with path functions (q, w).
Using Celsius instead of Kelvin in thermodynamic equations. Always use T(K) = T(°C) + 273.15.
Forgetting that ΔHf for pure elements is zero when calculating ΔH of a reaction.

Frequently Asked Questions

Q: What is the significance of the Zeroth Law?
A: It defines the concept of temperature and allows for the creation of thermometers.

Q: Why is Work (w) a path function?
A: Because the amount of work done depends on how the volume changes (e.g., in one step vs. infinite reversible steps), not just the start and end points.

Q: Define an Intensive property with an example.
A: An intensive property is independent of the size of the system. For example, the density of water is the same whether you have a cup or a bucket of it.