Chemistry Chemical Thermodynamics JEE Syllabus

Why does burning coal release heat while melting ice absorbs it? Why do some chemical reactions occur naturally while others require a continuous supply of energy? Chemical Thermodynamics answers these questions by studying the relationship between heat, work, and energy changes during chemical processes.

Before solving thermochemical numericals, it helps to know what Chemical Thermodynamics includes. The chapter explains how energy is transferred between a system and its surroundings, introduces important state functions like internal energy and enthalpy, and develops methods to calculate heat changes in reactions. These ideas also become the foundation for understanding spontaneity, equilibrium, and several advanced topics in physical chemistry.

Fundamentals of Thermodynamic Systems

Chemical thermodynamics begins with a few basic concepts that describe the part of the universe under observation and how it exchanges energy with its surroundings. A clear understanding of these definitions makes it easier to study the laws of thermodynamics and energy calculations.

System, Surroundings, and Thermodynamic Properties

The portion of the universe selected for study is called the system, while everything outside it is known as the surroundings.

Universe = System + Surroundings

A real or imaginary surface separating the system from its surroundings is called the boundary. Depending on the nature of this boundary, matter and energy may or may not pass through it.

Thermodynamic systems are classified into three types:

Open System

• Exchanges both matter and energy with the surroundings.

• Example: Water boiling in an open vessel.

Closed System

• Exchanges only energy but not matter.

• Example: Gas enclosed inside a sealed piston.

Isolated System

• Exchanges neither matter nor energy.

• Example: An ideal thermos flask.

The condition of a system is described by macroscopic properties such as:

• Pressure (P)

• Volume (V)

• Temperature (T)

• Composition

These properties are divided into two categories.

Intensive Properties

• Independent of the amount of substance.

• Examples: Temperature, pressure, density.

Extensive Properties

• Depend on the amount of substance present.

• Examples: Mass, volume, internal energy.

Thermodynamic equilibrium is reached when the system shows no spontaneous change with time. A system may possess:

• Thermal equilibrium

• Mechanical equilibrium

• Chemical equilibrium

Another important classification is based on how properties depend on the process.

State Functions
These depend only on the initial and final states.

Examples:

• Internal energy (U)

• Enthalpy (H)

• Entropy (S)

• Gibbs free energy (G)

Path Functions
These depend on the route by which a change occurs.

Examples:

• Heat (q)

• Work (w)

This distinction is extremely important because many JEE numerical problems are based on identifying state and path functions.

Internal Energy, Heat, Work, and the First Law of Thermodynamics

Energy can neither be created nor destroyed, but it can be transferred from one form to another. The first law of thermodynamics applies this principle to chemical systems and explains how heat and work affect the energy of a reaction.

Internal Energy Changes and Energy Conservation

Internal energy (U) is the total energy possessed by a system due to the motion and interactions of its molecules. It includes translational, rotational, vibrational, electronic, and intermolecular energies.

The absolute value of internal energy cannot be measured directly. Only the change in internal energy can be determined experimentally.

For an ideal gas, internal energy depends only on temperature and is independent of pressure and volume.

The First Law of Thermodynamics is:

ΔU = q + w

where

• ΔU = change in internal energy

• q = heat supplied to the system

• w = work done on the system

In chemistry, the most common form of work is expansion work.

w = -PextΔV

The negative sign indicates that when a gas expands against external pressure, energy leaves the system.

For an infinitesimal change:

δw = -Pext dV

Special cases frequently used in thermodynamics are:

Constant Volume Process

ΔV = 0

Therefore,

w = 0

qv = ΔU

All the supplied heat increases the internal energy.

Constant Pressure Process

qp = ΔH

where H is the enthalpy of the system.

For ideal gases:

ΔU = nCvΔT

where

• n = number of moles

• Cv = molar heat capacity at constant volume

The standard sign convention is:

• q > 0: Heat absorbed

• q < 0: Heat evolved

• w > 0: Work done on the system

• w < 0: Work done by the system

Chemical reactions are also classified according to heat exchange.

Exothermic Reaction

ΔH < 0

Heat is released to the surroundings.

Example:
Combustion of methane.

Endothermic Reaction

ΔH > 0

Heat is absorbed from the surroundings.

Example:
Thermal decomposition of calcium carbonate.

Understanding these sign conventions helps avoid common mistakes in thermodynamic calculations.

Enthalpy and Thermochemical Equations

Many chemical reactions are performed under constant atmospheric pressure. Under these conditions, the heat exchanged during a reaction is represented by the change in enthalpy.

Enthalpy Changes and Thermochemical Calculations

Enthalpy (H) is defined as:

H = U + PV

The enthalpy change during a reaction is:

ΔH = H(products) - H(reactants)

At constant pressure,

qp = ΔH

The standard state of a substance refers to its most stable form at 1 bar pressure and a specified temperature, usually 298 K.

Thermodynamic quantities measured under these conditions are represented by the symbol (°).

Different types of enthalpy changes are included in the chapter.

Standard Enthalpy of Reaction (ΔH°r)

The heat change accompanying a reaction when reactants and products are in their standard states.

Standard Enthalpy of Formation (ΔH°f)

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

For elements in their standard states:

ΔH°f = 0

Standard Enthalpy of Combustion (ΔH°c)

The heat released when one mole of a substance undergoes complete combustion in oxygen.

Standard Enthalpy of Neutralization

The heat evolved when one mole of water is formed during the reaction of an acid and a base.

For strong acid and strong base reactions:

H⁺(aq) + OH⁻(aq) → H₂O(l)

ΔH ≈ -57.1 kJ mol⁻¹

Standard Enthalpy of Atomization

The energy required to convert one mole of a substance into gaseous atoms.

Examples:

Na(s) → Na(g)

½Cl₂(g) → Cl(g)

Phase changes are also associated with enthalpy changes.

Enthalpy of Fusion

• Solid to liquid.

Enthalpy of Vaporization

• Liquid to gas.

Enthalpy of Sublimation

• Solid directly to gas.

These quantities are related by:

ΔHsub = ΔHfus + ΔHvap

Thermochemical equations include both the balanced chemical equation and the corresponding enthalpy change. They indicate whether heat is absorbed or evolved and allow energy calculations based on stoichiometry.

For example:

H₂(g) + ½O₂(g) → H₂O(l)

ΔH = -285.8 kJ mol⁻¹

Such equations are widely used in JEE numerical problems involving energy changes and reaction enthalpies.

Hess's Law and Bond Enthalpy

The enthalpy change of many reactions cannot be measured directly because the reactions may occur very slowly, involve multiple intermediate steps, or be difficult to perform under laboratory conditions. Hess's Law provides a convenient method for calculating such energy changes.

Enthalpy Cycles and Bond Energy Calculations

Hess's Law of Constant Heat Summation states that the total enthalpy change of a reaction depends only on the initial and final states of the system and is independent of the path followed.

This is possible because enthalpy is a state function.

If a reaction takes place through several intermediate steps, then:

ΔH = ΔH₁ + ΔH₂ + ΔH₃ + ...

One of the most important applications of Hess's Law is the calculation of reaction enthalpies using standard enthalpies of formation.

ΔH°reaction = ΣΔH°f(products) − ΣΔH°f(reactants)

Similarly, reaction enthalpy can also be estimated using bond enthalpies.

A bond enthalpy is the average amount of energy required to break one mole of a particular bond in the gaseous state.

General relation:

ΔH = Σ(Bond Enthalpies of Bonds Broken) − Σ(Bond Enthalpies of Bonds Formed)

Breaking chemical bonds requires energy and is therefore an endothermic process, while bond formation releases energy and is exothermic.

The chapter also introduces the idea of average bond enthalpy because the strength of the same type of bond may vary slightly in different molecules.

Hess's Law is also useful for determining quantities such as:

• Lattice enthalpy

• Resonance energy

• Enthalpy of transition

• Enthalpy changes of difficult chemical reactions

Many JEE numerical questions combine Hess's Law with formation enthalpies and bond energies to calculate unknown reaction enthalpies.

Entropy and the Second Law of Thermodynamics

Heat changes alone cannot explain why certain reactions occur naturally. Some processes proceed spontaneously even when they absorb heat. To understand this behavior, the concept of entropy was introduced.

Randomness, Disorder, and Spontaneous Processes

Entropy (S) is a thermodynamic property that measures the degree of randomness or disorder in a system.

A highly ordered system possesses low entropy, while a highly disordered system possesses high entropy.

In general:

Solid < Liquid < Gas

because gaseous particles have much greater freedom of movement.

Entropy usually increases when:

• Temperature increases.

• A solid melts.

• A liquid vaporizes.

• A substance dissolves.

• The number of gaseous molecules increases.

• Different gases mix.

For a reversible process:

ΔS = qrev/T

where:

• qrev = heat absorbed reversibly

• T = absolute temperature

The Second Law of Thermodynamics states that the entropy of the universe tends to increase for a spontaneous process.

ΔSuniverse > 0

Examples of spontaneous processes include:

• Flow of heat from a hot body to a cold body.

• Diffusion of perfume in the air.

• Expansion of a gas into a vacuum.

• Melting of ice above its melting point.

A process may be spontaneous even if it absorbs heat because the increase in entropy can compensate for the energy requirement.

Gibbs Free Energy and Spontaneity

The ideas of enthalpy and entropy are combined into a single function known as Gibbs free energy. This quantity is one of the most useful criteria for predicting whether a chemical reaction can occur on its own.

Free Energy Changes and Chemical Equilibrium

Gibbs free energy is defined as:

G = H − TS

The change in free energy for a process is:

ΔG = ΔH − TΔS

This equation combines the effects of heat change and randomness.

The conditions for spontaneity are:

• ΔG < 0: Spontaneous process

• ΔG > 0: Non-spontaneous process

• ΔG = 0: Equilibrium

The effect of enthalpy and entropy on spontaneity can be summarized as follows:

For non-standard conditions:

ΔG = ΔG° + RT ln Q

where:

• ΔG° = standard free energy change

• R = gas constant

• T = absolute temperature

• Q = reaction quotient

At equilibrium:

ΔG = 0

Therefore,

ΔG° = -RT ln K

where K is the equilibrium constant.

This relation forms an important link between thermodynamics and chemical equilibrium.

A large positive value of K indicates that products are favoured, while a very small value suggests that reactants are favoured.

The Gibbs free energy equation is frequently used to predict the direction and feasibility of chemical reactions.