HSSLIVE Plus One Chemistry Chapter 6: Thermodynamics Notes – Hsslive | Hsslive.co.in | Hss Live

Energy transformations in chemical reactions take center stage in this chapter. Students learn about enthalpy, entropy, and Gibbs free energy, which determine whether reactions occur spontaneously. The laws of thermodynamics are presented with their profound implications for natural processes. This chapter helps students understand why certain reactions occur while others don’t, connecting chemistry with energy concepts from physics and providing insight into industrial processes and biological systems.

Chapter 6: Thermodynamics

Thermodynamics is the branch of chemistry that deals with energy changes accompanying physical and chemical processes. It helps us determine whether a process is spontaneous and how much energy is transferred during the process.

Basic Concepts

Several fundamental concepts form the foundation of thermodynamics:

A system is the part of the universe under investigation. Everything else constitutes the surroundings. Together, the system and surroundings make up the universe. Systems can be classified as:

Open system: Can exchange both matter and energy with surroundings
Closed system: Can exchange energy but not matter with surroundings
Isolated system: Cannot exchange either matter or energy with surroundings

Energy is the capacity to do work or transfer heat. In chemistry, we’re primarily concerned with two types:

Kinetic energy: Energy of motion, directly proportional to mass and velocity squared (KE = ½mv²)
Potential energy: Energy due to position or composition, dependent on intermolecular forces and chemical bonds

Work is energy transferred when a force moves an object through a distance. In chemistry, work often involves volume changes against external pressure. For a gas expanding against constant external pressure, work is calculated as w = -PΔV. The negative sign indicates that work is done by the system on the surroundings.

Heat is energy transferred due to a temperature difference. Heat flows spontaneously from a hotter body to a cooler one. Unlike temperature (which is an intensive property), heat depends on the amount of substance and is an extensive property.

First Law of Thermodynamics

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, only converted from one form to another. Mathematically, it’s expressed as:

ΔU = q + w

Where:

ΔU is the change in internal energy of the system
q is the heat transferred to the system (positive if absorbed, negative if released)
w is the work done on the system (positive if done on the system, negative if done by the system)

The internal energy (U) represents the total energy contained within a system, including both kinetic and potential energy of all particles. It is a state function, meaning it depends only on the current state of the system, not on how it reached that state.

Enthalpy

Enthalpy (H) is defined as the sum of internal energy and the product of pressure and volume:

H = U + PV

The change in enthalpy (ΔH) for a process is:

ΔH = ΔU + Δ(PV)

For a reaction occurring at constant pressure (which is common in open containers), the change in enthalpy equals the heat transferred:

ΔH = qp

This makes enthalpy particularly useful for describing heat changes in chemical reactions. Reactions are classified as:

Exothermic: Release heat to surroundings (ΔH < 0)
Endothermic: Absorb heat from surroundings (ΔH > 0)

Important enthalpies in chemistry include:

Enthalpy of formation (ΔHf): Heat change when one mole of a compound forms from its elements in their standard states
Enthalpy of combustion (ΔHc): Heat released when one mole of a substance completely burns in oxygen
Enthalpy of neutralization: Heat released when one mole of an acid neutralizes one mole of a base
Enthalpy of solution: Heat change when one mole of a substance dissolves in a specified amount of solvent

Hess’s Law states that the enthalpy change for a reaction is the same whether it occurs in one step or multiple steps. This allows us to calculate enthalpy changes for reactions that are difficult to measure directly by breaking them into simpler reactions with known enthalpy values.

Types of Processes

Thermodynamic processes can be classified based on which variables remain constant:

Isothermal processes occur at constant temperature. During these processes, heat is exchanged with surroundings to maintain temperature. For an ideal gas, PV = constant (Boyle’s Law).

Adiabatic processes occur without heat exchange with surroundings (q = 0). For these processes, ΔU = w. For an ideal gas, PVᵞ = constant, where γ is the ratio of heat capacities.

Isobaric processes occur at constant pressure. For these processes, work is simply w = -PΔV, and ΔH = qp.

Isochoric processes occur at constant volume. Since no expansion or compression work is done (ΔV = 0, so w = 0), ΔU = qv.

Spontaneity and Second Law

The First Law tells us about energy conservation but doesn’t explain why some processes occur spontaneously while others don’t. The Second Law of Thermodynamics addresses this through the concept of entropy.

Entropy (S) is a measure of disorder or randomness in a system. The Second Law states that the total entropy of an isolated system always increases for a spontaneous process: ΔSuniverse = ΔSsystem + ΔSsurroundings > 0.

Factors that generally increase entropy include:

Phase changes from solid to liquid to gas
Temperature increase
Volume increase (for gases)
Increasing number of particles (e.g., during dissociation)
Mixing of different substances

The Third Law of Thermodynamics states that the entropy of a perfect crystalline substance at absolute zero (0 K) is zero. This provides a reference point for calculating absolute entropies.

Gibbs Free Energy

To determine spontaneity under common conditions (constant temperature and pressure), we use Gibbs free energy (G), defined as:

G = H – TS

The change in Gibbs free energy for a process is:

ΔG = ΔH – TΔS

Where:

ΔG < 0: Spontaneous process
ΔG = 0: System at equilibrium
ΔG > 0: Non-spontaneous process (reverse process is spontaneous)

This equation shows that both enthalpy and entropy influence spontaneity. A process tends to be spontaneous if it releases heat (ΔH < 0) and/or increases entropy (ΔS > 0). At higher temperatures, the TΔS term becomes more significant, making entropy change more important in determining spontaneity.

The standard Gibbs free energy change (ΔG°) is related to the equilibrium constant (K) by the equation:

ΔG° = -RT ln K

This relationship allows us to calculate equilibrium constants from thermodynamic data and vice versa.

Complete Chapter-wise Hsslive Plus One Chemistry Notes

Our HSSLive Plus One Chemistry Notes cover all chapters with key focus areas to help you organize your study effectively: