CBSE Class 11 Chemistry Revision Notes Chapter 5 Thermodynamics

Thermodynamics studies energy changes in chemical reactions and physical processes. In CBSE Class 11 Chemistry Chapter 5, students learn how heat, work, internal energy, enthalpy, entropy and Gibbs free energy help explain whether a process can occur.

Chemical reactions involve energy changes. Some reactions release heat, some absorb heat, and some can perform work. Thermodynamics gives a framework to measure these changes and connect them with the direction of a process.

These CBSE class 11 chemistry revision notes chapter 5 follow the chapter in a quick revision format. The chapter begins with system and surroundings, then moves to internal energy, heat, work, first law of thermodynamics, enthalpy, calorimetry, Hess’s law, entropy, Gibbs free energy and equilibrium constant.

Key Takeaways

• Thermodynamics: Study of energy changes in physical and chemical processes.
• System and surroundings: The part under study is the system, and everything else is the surroundings.
• First law: Energy can be transferred as heat or work, but total energy remains conserved.
• Enthalpy: Heat change at constant pressure is measured as enthalpy change.
• Gibbs free energy: It helps predict whether a process is spontaneous.

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Thermodynamics Class 11 Chemistry Notes: Chapter Overview

Thermodynamics explains how energy changes take place during reactions and physical processes. It deals with measurable properties such as pressure, volume, temperature, heat and work.

Class 11 Chemistry Chapter 5 Notes help students understand how energy conservation, heat transfer and entropy decide the feasibility of a process. The chapter also explains why some reactions occur naturally while others need external energy.

What Is Thermodynamics?

Thermodynamics is the branch of Chemistry that studies energy changes during chemical reactions and physical processes. It focuses on macroscopic properties of matter, such as pressure, volume and temperature.

Thermodynamics does not explain the speed of a reaction. It tells whether a process is energetically possible and how much energy change is involved.

System and Surroundings

System and surroundings are the basic terms used in thermodynamics.

A system is the part of the universe selected for study. The surroundings include everything outside the system that can exchange energy or matter with it.

Universe = System + Surroundings

Types of Thermodynamic Systems

Thermodynamic systems are classified based on exchange of matter and energy.

Open System

An open system can exchange both matter and energy with the surroundings. An open beaker containing a reaction mixture is an example.

Closed System

A closed system can exchange energy but does not exchange matter. A closed conducting vessel is an example.

Isolated System

An isolated system exchanges neither matter nor energy with the surroundings. A thermos flask is a common example.

State of a System

The state of a system is described by measurable properties such as pressure, volume, temperature and amount of substance.

These properties define the condition of the system at a given moment.

When these properties are fixed, the state of the system is defined.

State Functions

State functions are properties whose values depend only on the present state of the system. They do not depend on the path used to reach that state.

Heat and work are path functions because their values depend on how the process is carried out.

Internal Energy

Internal energy is the total energy stored in a system. It includes all forms of molecular energy, such as translational, rotational, vibrational, electronic and chemical energy.

It is represented by U.

A system’s absolute internal energy cannot be measured directly, but the change in internal energy can be measured.

ΔU = U₂ - U₁

Heat

Heat is energy transferred between system and surroundings because of temperature difference.

It is represented by q.

Heat is measured in joule or kilojoule.

Work

Work is another mode of energy transfer. In Chemistry, pressure-volume work is commonly studied.

When a gas expands or compresses, work is done.

w = -pΔV

First Law of Thermodynamics

First law of thermodynamics states that energy can neither be created nor destroyed. It can only be transferred from one form to another.

Mathematically:

ΔU = q + w

This law is also called the law of conservation of energy.

First Law for Different Conditions

Thermodynamic Processes

Thermodynamic processes are changes that take a system from one state to another.

Isothermal Process

In an isothermal process, temperature remains constant. For an ideal gas, internal energy depends only on temperature, so ΔU is zero.

Adiabatic Process

In an adiabatic process, no heat is exchanged between the system and surroundings.

Isobaric Process

In an isobaric process, pressure remains constant.

Isochoric Process

In an isochoric process, volume remains constant. Since ΔV = 0, no pressure-volume work is done.

Enthalpy

Enthalpy is a thermodynamic state function represented by H.

H = U + pV

At constant pressure, heat absorbed or released is equal to enthalpy change.

ΔH = qₚ

Relation Between ΔH and ΔU

For reactions involving gases:

ΔH = ΔU + ΔnᵍRT

The difference between ΔH and ΔU is more important when gases are involved.

Extensive and Intensive Properties

Extensive and intensive properties help describe thermodynamic systems.

If a system is divided into two equal parts, extensive properties are divided, but intensive properties remain the same.

Heat Capacity

Heat capacity is the amount of heat required to raise the temperature of a system by one degree Celsius or one kelvin.

q = CΔT

Molar Heat Capacity

Molar heat capacity is the heat required to raise the temperature of one mole of a substance by one degree.

Specific Heat Capacity

Specific heat capacity is the heat required to raise the temperature of unit mass of a substance by one degree.

Relationship Between Cp and Cv

For one mole of an ideal gas:

Cp - Cv = R

This relation is useful for ideal gases.

Calorimetry

Calorimetry is the experimental method used to measure heat changes in physical and chemical processes.

The instrument used is called a calorimeter.

Bomb Calorimeter

A bomb calorimeter measures heat change at constant volume. It is commonly used for combustion reactions.

Coffee Cup Calorimeter

A coffee cup calorimeter measures heat change at constant pressure. It is commonly used for reactions in solution.

Enthalpy Change

Enthalpy change is the heat absorbed or released during a process at constant pressure.

Standard Enthalpy Change

Standard enthalpy change is the enthalpy change measured under standard conditions.

It is represented by ΔH°.

Hess’s Law

Hess’s law states that the total enthalpy change for a reaction remains the same whether the reaction takes place in one step or several steps.

This is because enthalpy is a state function.

Hess’s Law Example

If a reaction can be written as the sum of two or more reactions, the enthalpy change of the main reaction equals the sum of the enthalpy changes of those reactions.

ΔH = ΔH₁ + ΔH₂ + ΔH₃

Bond Enthalpy and Reaction Enthalpy

Bond enthalpy is the energy required to break one mole of bonds in gaseous molecules.

For a reaction:

ΔH = Bond enthalpy of bonds broken - Bond enthalpy of bonds formed

Spontaneous Process

A spontaneous process is a process that can take place on its own under given conditions.

A spontaneous process may be slow or fast. Thermodynamics only tells whether the process is possible, not the rate of the process.

Non Spontaneous Process

A non spontaneous process does not occur on its own under given conditions. It needs continuous external help.

Entropy

Entropy is a thermodynamic state function that measures the randomness or disorder of a system. It is represented by S.

Higher disorder means higher entropy.

Entropy helps explain the direction of natural processes.

Entropy Change

Entropy change is represented as ΔS.

For a reversible process:

ΔS = qrev / T

A process is favoured when total entropy of the universe increases.

Gibbs Free Energy

Gibbs free energy is a thermodynamic function that helps predict spontaneity at constant temperature and pressure.

It is represented by G.

ΔG = ΔH - TΔS

Gibbs Free Energy and Spontaneity

Gibbs free energy connects enthalpy and entropy in one equation.

Effect of Temperature on Spontaneity

The spontaneity of a process depends on ΔH, ΔS and temperature.

Gibbs Free Energy and Equilibrium Constant

Gibbs free energy is related to equilibrium constant.

ΔG° = -RT ln K

When ΔG° is negative, K is greater than 1 and products are favoured.

Thermodynamics Class 11 Chemistry Chapter 5 CBSE Notes: Quick Revision Tables

Important Formulae

System Types Summary

State and Path Functions

Thermodynamic Processes

Key Terms from CBSE Class 11 Chemistry Revision Notes Chapter 5

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