CBSE Class 11 Chemistry Chapter 6 Thermodynamics Notes

Key Points to Remember in Thermodynamics

• First Law of Thermodynamics (Conservation of Energy): Energy can neither be created nor destroyed; it can only be converted from one form to another. ΔU = Q + W.
• System & Surroundings: Define the boundaries for energy transfer. Systems can be open, closed, or isolated.
• State Functions: Properties that depend only on the initial and final states of the system, not on the path taken (e.g., U, H, S, G). Path functions include Q and W.
• Enthalpy (H): Heat content of a system at constant pressure. ΔH = ΔU + PΔV. For exothermic reactions, ΔH < 0; for endothermic, ΔH > 0.
• Entropy (S): A measure of the disorder or randomness of a system. For a spontaneous process, total entropy (system + surroundings) tends to increase.
• Gibbs Free Energy (G): Determines the spontaneity of a process at constant temperature and pressure. ΔG = ΔH - TΔS. ΔG < 0 for spontaneous, ΔG > 0 for non-spontaneous, ΔG = 0 for equilibrium.
• Standard States: Defined conditions (298 K, 1 atm pressure, 1 M concentration) for comparing thermodynamic properties.
• Hess's Law: The total enthalpy change for a reaction is independent of the pathway between the initial and final states.
• Second Law of Thermodynamics: Entropy of an isolated system always increases during a spontaneous process.
• Third Law of Thermodynamics: The entropy of a perfectly crystalline substance at absolute zero (0 K) is taken to be zero.

Essential Thermodynamic Terms

System

The part of the universe under thermodynamic study, separated from the rest of the universe by a real or imaginary boundary.

Surroundings

Everything in the universe other than the system, which can exchange energy and matter with the system.

Internal Energy (U)

The total energy contained within a thermodynamic system, including kinetic and potential energies of molecules. It is a state function.

Work (W)

Energy transferred between a system and its surroundings due to a force acting through a distance (e.g., expansion/compression work).

Heat (Q)

Energy transferred between a system and its surroundings due to a temperature difference.

Enthalpy of Formation (ΔfH°)

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states.

Bond Enthalpy

The average amount of energy required to break one mole of a particular type of bond in gaseous molecules.

Entropy (S)

A thermodynamic property that is a measure of the disorder or randomness of a system. Units are J/K·mol.

Gibbs Free Energy (G)

A thermodynamic potential that measures the 'useful' or process-initiating work obtainable from an isothermal, isobaric thermodynamic system.

Spontaneous Process

A process that occurs without external intervention once initiated. It proceeds in a definite direction towards equilibrium.

The First Law of Thermodynamics: Energy Conservation

The First Law of Thermodynamics is essentially the law of conservation of energy applied to thermodynamic systems. It states that energy can neither be created nor destroyed in an isolated system; it can only be converted from one form to another. This fundamental principle is expressed mathematically as:

ΔU = Q + W

Where:

• ΔU represents the change in internal energy of the system. Internal energy (U) is a state function, meaning its value depends only on the current state of the system, not on how that state was reached. It includes all forms of energy associated with the molecules of the system (kinetic, potential, electronic, vibrational, etc.).
• Q represents the heat absorbed by the system from the surroundings. By convention, if heat is absorbed by the system (endothermic process), Q is positive (+ve). If heat is released by the system (exothermic process), Q is negative (-ve).
• W represents the work done on the system by the surroundings. The convention for work is crucial: if work is done on the system (e.g., compression), W is positive (+ve). If work is done by the system (e.g., expansion against external pressure), W is negative (-ve). The most common type of work in chemistry is pressure-volume work (P-V work), given by W = -P_extΔV, where P_ext is the external pressure and ΔV is the change in volume. For expansion, ΔV is positive, making W negative. For compression, ΔV is negative, making W positive.

Understanding these sign conventions is vital for solving numerical problems. For example, if a system absorbs 100 J of heat (Q = +100 J) and expands, doing 20 J of work on the surroundings (W = -20 J), then its internal energy change would be ΔU = (+100 J) + (-20 J) = +80 J. This indicates an overall increase in the system's internal energy.

The First Law establishes that the total energy of the universe remains constant. Any energy lost by the system must be gained by the surroundings, and vice versa. This principle underpins all energy transformations in chemical reactions and physical processes, providing a quantitative framework for analyzing energy flow.

Worked Example: Calculating Enthalpy Change

• Example 1: Internal Energy Change Q: A system absorbs 50 kJ of heat from the surroundings and performs 15 kJ of work on the surroundings. Calculate the change in internal energy (ΔU) of the system. A: According to the First Law of Thermodynamics, ΔU = Q + W. Heat absorbed by the system, Q = +50 kJ (positive as heat is absorbed). Work done by the system, W = -15 kJ (negative as work is done by the system). ΔU = (+50 kJ) + (-15 kJ) = 35 kJ. The internal energy of the system increases by 35 kJ.
• Example 2: Enthalpy of Reaction from Enthalpies of Formation Q: Calculate the standard enthalpy of reaction (ΔrH°) for the combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l). Given standard enthalpies of formation (ΔfH°): CH₄(g) = -74.8 kJ/mol CO₂(g) = -393.5 kJ/mol H₂O(l) = -285.8 kJ/mol O₂(g) = 0 kJ/mol (element in its standard state) A: Using the formula ΔrH° = ΣΔfH°(products) - ΣΔfH°(reactants): ΔrH° = [1 × ΔfH°(CO₂) + 2 × ΔfH°(H₂O)] - [1 × ΔfH°(CH₄) + 2 × ΔfH°(O₂)] ΔrH° = [1 × (-393.5 kJ/mol) + 2 × (-285.8 kJ/mol)] - [1 × (-74.8 kJ/mol) + 2 × (0 kJ/mol)] ΔrH° = [-393.5 - 571.6] - [-74.8] ΔrH° = -965.1 + 74.8 = -890.3 kJ/mol The standard enthalpy of combustion for methane is -890.3 kJ/mol, indicating an exothermic reaction.

Comparison of Thermodynamic Processes

Spontaneity and Gibbs Free Energy

Exam Tip: Sign Conventions are Crucial!

One of the most common pitfalls in Thermodynamics problems is getting the sign conventions wrong for heat (Q) and work (W). Always remember:

• Heat (Q): +ve if absorbed by the system (endothermic); -ve if released by the system (exothermic).
• Work (W): +ve if done on the system (e.g., compression); -ve if done by the system (e.g., expansion).

Double-check your calculations, especially when applying the First Law (ΔU = Q + W) or calculating P-V work (W = -P_extΔV). A single sign error can lead to a completely incorrect answer, losing you valuable marks. Pay close attention to keywords like 'absorbs', 'releases', 'expands', 'compresses', 'on the system', 'by the system'.

Practice Questions with Solutions

• Q: Differentiate between a state function and a path function, providing one example of each. A: A state function depends only on the initial and final states of a system (e.g., Internal Energy, Enthalpy), while a path function depends on the path taken (e.g., Heat, Work).
• Q: What are the conditions for a reaction to be spontaneous in terms of Gibbs Free Energy (ΔG)? A: For a reaction to be spontaneous at constant temperature and pressure, the change in Gibbs Free Energy (ΔG) must be negative (ΔG < 0).
• Q: State the First Law of Thermodynamics mathematically and explain each term. A: ΔU = Q + W, where ΔU is the change in internal energy, Q is heat absorbed by the system, and W is work done on the system. (See detailed explanation in prose section).
• Q: For an isolated system, what happens to its total entropy during a spontaneous process? A: For an isolated system, the total entropy always increases during a spontaneous process, according to the Second Law of Thermodynamics.

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