Class 11 Chemistry – Thermodynamics

Chapter Notes (Part 1)

Topics Covered

• Introduction to Thermodynamics
• System and Surroundings
• Boundary
• Types of Systems
• State of a System
• State Functions
• Internal Energy
• Work and Heat (Basic Concepts)

---

Chapter 5: Thermodynamics

What is Thermodynamics?

Thermodynamics is the branch of chemistry that studies the relationship between heat, work, and energy during physical and chemical changes.

It helps us understand:

• How energy changes during a reaction
• Whether a reaction absorbs or releases heat
• Whether a reaction can occur naturally
• How much energy is involved in a process

Examples

• Burning LPG releases heat.
• A battery converts chemical energy into electrical energy.
• An engine converts chemical energy into mechanical energy.

Important Point

Thermodynamics studies only the initial and final states of a system. It does not explain how fast a reaction occurs.

---

Important Terms

1. System

A system is the part of the universe chosen for study.

Examples

• Water inside a beaker
• Gas inside a cylinder
• Chemicals reacting in a flask

---

2. Surroundings

Everything outside the system is called the surroundings.

Example:

• If water in a beaker is the system, then the beaker, air, and room are the surroundings.

---

3. Universe

Universe = System + Surroundings

This is one of the most important definitions in thermodynamics.

---

Boundary

The surface that separates the system from its surroundings is called the boundary.

It may be:

• Real (glass wall of a flask)
• Imaginary

The boundary controls the exchange of matter and energy.

---

Types of Systems

Thermodynamic systems are classified according to the exchange of matter and energy.

System	Matter Exchange	Energy Exchange	Example
Open System	Yes	Yes	Open beaker
Closed System	No	Yes	Closed steel vessel
Isolated System	No	No	Thermos flask

---

1. Open System

Definition

An open system can exchange both matter and energy with the surroundings.

Example

• Tea in an open cup
• Water boiling in an open pan

Memory Trick

Open = Everything can move.

---

2. Closed System

Definition

A closed system can exchange energy only, but not matter.

Example

• Water in a sealed pressure cooker
• Gas inside a closed cylinder

Memory Trick

Closed = Heat can move, matter cannot.

---

3. Isolated System

Definition

An isolated system exchanges neither matter nor energy.

Example

• Thermos flask (approximately)
• Perfect insulated container (ideal)

Memory Trick

Isolated = Nothing enters or leaves.

---

Comparison of Systems

Property	Open	Closed	Isolated
Matter transfer	✓	✗	✗
Energy transfer	✓	✓	✗
Example	Open beaker	Closed bottle	Thermos flask

---

State of a System

The condition of a system at any moment is called its state.

The state is described using measurable quantities such as:

• Pressure (P)
• Volume (V)
• Temperature (T)
• Amount of substance (n)

These quantities are called state variables.

---

State Variables

State variables describe the present condition of a system.

Examples:

• Pressure
• Temperature
• Volume
• Number of moles

Changing any one of these changes the state of the system.

---

State Function

A state function depends only on the initial and final states, not on the path followed.

Examples

• Internal Energy (U)
• Pressure (P)
• Volume (V)
• Temperature (T)
• Enthalpy (H)

Example

If you travel from Delhi to Mumbai:

• by train,
• by bus,
• by plane,

your final location is the same. Similarly, a state function depends only on the initial and final states, not on the route taken.

---

Internal Energy (U)

Internal energy is the total energy present inside a system.

It includes:

• Kinetic energy of particles
• Potential energy of particles
• Chemical energy
• Electrical energy
• Other microscopic forms of energy

Symbol:
U

Unit:
Joule (J) or kJ

---

Internal Energy Changes When

1. Heat is supplied or removed.
2. Work is done on or by the system.
3. Matter enters or leaves the system.

---

Change in Internal Energy

The change in internal energy is represented as:

ΔU = U₂ − U₁

where:

• U₁ = Initial internal energy
• U₂ = Final internal energy

---

Why is Internal Energy a State Function?

Suppose water is heated:

Method 1:

• Stirring with paddles

Method 2:

• Electric heater

If both methods raise the water to the same final temperature, the change in internal energy is the same.

Therefore,

Internal energy depends only on the initial and final states, not on the method used.

---

Work (w)

Work is done whenever energy is transferred by applying force.

In thermodynamics:

• Compression → Work is done on the system.
• Expansion → Work is done by the system.

IUPAC Sign Convention

Process	Sign of Work
Work done on system	Positive (+)
Work done by system	Negative (−)

---

Heat (q)

Heat is the energy transferred because of a temperature difference.

Sign Convention

Process	Sign
Heat absorbed by system	Positive (+)
Heat released by system	Negative (−)

---

Important Sign Convention Table

Quantity	Positive When	Negative When
Heat (q)	System absorbs heat	System releases heat
Work (w)	Work done on system	Work done by system

---

Quick Revision Box

Learn These Definitions

System: Part of the universe under study.

Surroundings: Everything outside the system.

Boundary: Surface separating system and surroundings.

Open System: Exchanges matter and energy.

Closed System: Exchanges only energy.

Isolated System: Exchanges neither matter nor energy.

State Function: Depends only on initial and final states.

Internal Energy: Total energy possessed by a system.

Heat: Energy transferred because of temperature difference.

Work: Energy transferred by force.

---

Important Exam Questions

1. Define system and surroundings with examples.

2. Differentiate between open, closed, and isolated systems.

3. What is a state function? Give four examples.

4. Explain internal energy. Why is it called a state function?

---

Chapter Notes (Part 3)

Topics Covered

• Enthalpy (H)
• Enthalpy Change (ΔH)
• Relation between ΔH and ΔU
• Heat Capacity
• Molar Heat Capacity
• Specific Heat Capacity
• Relation between Cp and Cv
• Calorimetry

(These notes are rewritten in simple original language for Class 11 students.)

---

Enthalpy (H)

Why do we need Enthalpy?

Most chemical reactions are performed at constant atmospheric pressure.

At constant volume:$$\Delta U=q_v$$ΔU=qv​

But in laboratories, reactions usually occur at constant pressure. Therefore, a new thermodynamic function called enthalpy is used.

---

Definition of Enthalpy

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

It is represented by:$$\boxed{H}$$H​

Mathematically:$$\boxed{H=U+PV}$$H=U+PV​

Where:

• H = Enthalpy
• U = Internal energy
• P = Pressure
• V = Volume

---

Change in Enthalpy (ΔH)

The change in enthalpy is:$$\boxed{\Delta H=H_2-H_1}$$ΔH=H2​−H1​​

At constant pressure:$$\boxed{\Delta H=q_p}$$ΔH=qp​​

This means:

Heat absorbed or released at constant pressure is equal to enthalpy change.

---

Types of Enthalpy Change

1. Exothermic Reaction

A reaction that releases heat is called an exothermic reaction.

Example:$$C+O_2 \rightarrow CO_2+\text{heat}$$C+O2​→CO2​+heat

For exothermic reactions:$$\boxed{\Delta H<0}$$ΔH<0​

Heat leaves the system.

---

2. Endothermic Reaction

A reaction that absorbs heat is called an endothermic reaction.

Example:$$CaCO_3+\text{heat}\rightarrow CaO+CO_2$$CaCO3​+heat→CaO+CO2​

For endothermic reactions:$$\boxed{\Delta H>0}$$ΔH>0​

Heat enters the system.

---

Relationship Between ΔH and ΔU

We know:$$H=U+PV$$H=U+PV

Change in enthalpy:$$\Delta H=\Delta U+\Delta(PV)$$ΔH=ΔU+Δ(PV)

At constant pressure:$$\boxed{\Delta H=\Delta U+P\Delta V}$$ΔH=ΔU+PΔV​

---

For Gaseous Reactions

Using ideal gas equation:$$PV=nRT$$PV=nRT

The relationship becomes:$$\boxed{\Delta H=\Delta U+\Delta n_gRT}$$ΔH=ΔU+Δng​RT​

Where:

• Δng = moles of gaseous products − moles of gaseous reactants
• R = Gas constant
• T = Temperature

---

Important Cases

Case 1: No change in number of gas molecules

If:$$\Delta n_g=0$$Δng​=0

Then:$$\Delta H=\Delta U$$ΔH=ΔU

Example:$$H_2+Cl_2\rightarrow2HCl$$H2​+Cl2​→2HCl

---

Case 2: Solids and Liquids

For reactions involving only solids and liquids:$$\Delta H\approx\Delta U$$ΔH≈ΔU

because volume change is very small.

---

Heat Capacity

Definition

Heat capacity is the amount of heat required to increase the temperature of a system by 1 K (or 1°C).

Symbol:$$C$$C

Formula:$$\boxed{q=C\Delta T}$$q=CΔT​

Where:

• q = Heat supplied
• C = Heat capacity
• ΔT = Temperature change

---

Factors Affecting Heat Capacity

Heat capacity depends on:

1. Amount of substance
2. Nature of substance
3. Composition of substance

Example:

Water has high heat capacity, so it requires more heat to increase its temperature.

---

Types of Heat Capacity

1. Molar Heat Capacity

The heat required to raise the temperature of one mole of a substance by 1 K.

Formula:$$\boxed{C_m=\frac{C}{n}}$$Cm​=nC​​

Unit:$$J\,mol^{-1}K^{-1}$$Jmol−1K−1

---

2. Specific Heat Capacity

The heat required to raise the temperature of one gram of a substance by 1 K.

Formula:$$\boxed{q=mc\Delta T}$$q=mcΔT​

Where:

• m = mass
• c = specific heat capacity
• ΔT = temperature change

---

Heat Capacity at Constant Volume and Pressure

There are two types:

1. Heat Capacity at Constant Volume (Cv)

At constant volume:$$q_v=C_v\Delta T$$qv​=Cv​ΔT

Since no expansion occurs:$$\Delta U=q_v$$ΔU=qv​

---

2. Heat Capacity at Constant Pressure (Cp)

At constant pressure:$$q_p=C_p\Delta T$$qp​=Cp​ΔT

Since:$$\Delta H=q_p$$ΔH=qp​

---

Relation Between Cp and Cv

For an ideal gas:$$\boxed{C_p-C_v=R}$$Cp​−Cv​=R​

Where:

• Cp = Heat capacity at constant pressure
• Cv = Heat capacity at constant volume
• R = Gas constant

---

Calorimetry

Definition

Calorimetry is the experimental technique used to measure heat changes during chemical reactions.

The instrument used is called a calorimeter.

---

Types of Calorimetry

1. Constant Volume Calorimetry

Used to measure:$$\Delta U$$ΔU

Instrument:

Bomb Calorimeter

Working:

• Reaction occurs inside a strong steel container.
• Container is surrounded by water.
• Heat released increases water temperature.
• Temperature change is used to calculate heat.

Because volume remains constant:$$\boxed{\Delta U=q_v}$$ΔU=qv​​

---

2. Constant Pressure Calorimetry

Used to measure:$$\Delta H$$ΔH

At constant pressure:$$\boxed{\Delta H=q_p}$$ΔH=qp​​

Used for reactions occurring in open containers.

---

Important Differences

ΔU	ΔH
Change in internal energy	Change in enthalpy
Measured at constant volume	Measured at constant pressure
ΔU = qv	ΔH = qp
Used in bomb calorimeter	Used in pressure calorimeter

---

Quick Revision Formula Sheet

First Law:

$$\Delta U=q+w$$ΔU=q+w

Enthalpy:

$$H=U+PV$$H=U+PV

Enthalpy Change:

$$\Delta H=\Delta U+P\Delta V$$ΔH=ΔU+PΔV

Gas Reaction:

$$\Delta H=\Delta U+\Delta n_gRT$$ΔH=ΔU+Δng​RT

Heat Capacity:

$$q=C\Delta T$$q=CΔT

Specific Heat:

$$q=mc\Delta T$$q=mcΔT

Ideal Gas:

$$C_p-C_v=R$$Cp​−Cv​=R

---

Important Exam Questions

1. Define enthalpy and write its mathematical expression.

2. Differentiate between ΔH and ΔU.

3. Derive the relationship between ΔH and ΔU.

4. What is calorimetry? Explain bomb calorimeter.

5. Explain Cp and Cv relationship.

Chapter Notes (Part 4)

Topics Covered

• Reaction Enthalpy (ΔrH)
• Standard Enthalpy of Reaction
• Enthalpy of Phase Transformations
• Standard Enthalpy of Formation
• Thermochemical Equations
• Hess’s Law of Constant Heat Summation

Reaction Enthalpy (ΔrH)

Definition

The heat change that occurs during a chemical reaction at constant pressure is called reaction enthalpy.

It is represented as:$$\boxed{\Delta_rH}$$Δr​H​

For a reaction:$$\text{Reactants} \rightarrow \text{Products}$$Reactants→Products

Reaction enthalpy is:$$\boxed{\Delta_rH = \text{Enthalpy of products} – \text{Enthalpy of reactants}}$$Δr​H=Enthalpy of products−Enthalpy of reactants​

or$$\boxed{\Delta_rH=\sum H_{products}-\sum H_{reactants}}$$Δr​H=∑Hproducts​−∑Hreactants​​

---

Meaning of Reaction Enthalpy

If:

$$\Delta_rH<0$$Δr​H<0

Reaction is exothermic.

• Heat is released.
• Products have lower enthalpy than reactants.

Example:$$CH_4+2O_2\rightarrow CO_2+2H_2O$$CH4​+2O2​→CO2​+2H2​O

---

If:

$$\Delta_rH>0$$Δr​H>0

Reaction is endothermic.

• Heat is absorbed.
• Products have higher enthalpy.

Example:$$CaCO_3\rightarrow CaO+CO_2$$CaCO3​→CaO+CO2​

---

Standard Enthalpy of Reaction (ΔrH°)

The enthalpy change when a reaction occurs with all substances in their standard states is called standard enthalpy of reaction.

Symbol:$$\boxed{\Delta_rH^\circ}$$Δr​H∘​

---

Standard State

The standard state of a substance is its pure form at:

• Pressure = 1 bar
• Specified temperature (usually 298 K)

Examples:

• Pure liquid water at 298 K and 1 bar
• Solid iron at 1 bar pressure

---

Enthalpy Changes During Phase Transformation

A phase change involves energy change.

Examples:

• Melting
• Boiling
• Sublimation

---

1. Enthalpy of Fusion (ΔfusH°)

Definition

The heat required to convert one mole of solid into liquid at constant temperature and standard pressure is called standard enthalpy of fusion.

Example:$$H_2O(s)\rightarrow H_2O(l)$$H2​O(s)→H2​O(l) $$\Delta_{fus}H^\circ=+6.00\,kJ\,mol^{-1}$$Δfus​H∘=+6.00kJmol−1

Important Point

Melting is an endothermic process.

Therefore:$$\Delta_{fus}H^\circ>0$$Δfus​H∘>0

---

2. Enthalpy of Vaporisation (ΔvapH°)

Definition

The heat required to convert one mole of liquid into vapour at constant temperature and pressure is called standard enthalpy of vaporisation.

Example:$$H_2O(l)\rightarrow H_2O(g)$$H2​O(l)→H2​O(g) $$\Delta_{vap}H^\circ=+40.79\,kJ\,mol^{-1}$$Δvap​H∘=+40.79kJmol−1

---

3. Enthalpy of Sublimation (ΔsubH°)

Definition

The enthalpy change when one mole of a solid changes directly into vapour is called standard enthalpy of sublimation.

Example:$$CO_2(s)\rightarrow CO_2(g)$$CO2​(s)→CO2​(g)

---

Factors Affecting Phase Change Enthalpy

The amount of heat required depends on:

1. Strength of intermolecular forces
2. Nature of substance

Stronger intermolecular forces require more energy.

Example:

• Water has strong hydrogen bonding.
• More heat is needed to vaporise water compared with many organic liquids.

---

Standard Enthalpy of Formation (ΔfH°)

Definition

The enthalpy change when one mole of a compound is formed from its elements in their most stable forms under standard conditions is called standard enthalpy of formation.

Symbol:$$\boxed{\Delta_fH^\circ}$$Δf​H∘​

---

Examples

Formation of Water:

$$H_2(g)+\frac12O_2(g)\rightarrow H_2O(l)$$H2​(g)+21​O2​(g)→H2​O(l) $$\Delta_fH^\circ=-285.8\,kJ\,mol^{-1}$$Δf​H∘=−285.8kJmol−1

---

Formation of Methane:

$$C(graphite)+2H_2(g)\rightarrow CH_4(g)$$C(graphite)+2H2​(g)→CH4​(g) $$\Delta_fH^\circ=-74.81\,kJ\,mol^{-1}$$Δf​H∘=−74.81kJmol−1

---

Important Rule

The standard enthalpy of formation of an element in its most stable state is:$$\boxed{0}$$0​

Examples:

• H₂(g) = 0
• O₂(g) = 0
• C(graphite) = 0

---

Calculation of Reaction Enthalpy Using Formation Enthalpies

Formula:$$\boxed{\Delta_rH^\circ= \sum \Delta_fH^\circ(products) – \sum \Delta_fH^\circ(reactants)}$$Δr​H∘=∑Δf​H∘(products)−∑Δf​H∘(reactants)​

Steps:

1. Write balanced equation.
2. Multiply each substance by its coefficient.
3. Add product enthalpies.
4. Add reactant enthalpies.
5. Subtract.

---

Thermochemical Equation

A balanced chemical equation along with its enthalpy change is called a thermochemical equation.

Example:$$C_2H_5OH(l)+3O_2(g) \rightarrow 2CO_2(g)+3H_2O(l)$$C2​H5​OH(l)+3O2​(g)→2CO2​(g)+3H2​O(l) $$\Delta_rH^\circ=-1367\,kJ\,mol^{-1}$$Δr​H∘=−1367kJmol−1

The physical states of substances must be mentioned.

---

Rules for Thermochemical Equations

Rule 1

Coefficients represent number of moles, not molecules.

---

Rule 2

If the equation is multiplied by a number, ΔH is also multiplied by that number.

Example:

If:$$\Delta H=-100\,kJ$$ΔH=−100kJ

For half reaction:$$\Delta H=-50\,kJ$$ΔH=−50kJ

---

Rule 3

If a reaction is reversed, the sign of ΔH changes.

Example:

Forward:$$N_2+3H_2\rightarrow2NH_3$$N2​+3H2​→2NH3​ $$\Delta H=-91.8\,kJ$$ΔH=−91.8kJ

Reverse:$$2NH_3\rightarrow N_2+3H_2$$2NH3​→N2​+3H2​ $$\Delta H=+91.8\,kJ$$ΔH=+91.8kJ

---

Hess’s Law of Constant Heat Summation

Statement

If a reaction occurs in several steps, the total enthalpy change is equal to the sum of enthalpy changes of all steps.$$\boxed{\Delta H_{overall}=\Delta H_1+\Delta H_2+\Delta H_3+…}$$ΔHoverall​=ΔH1​+ΔH2​+ΔH3​+…​

---

Why Hess’s Law Works?

Because enthalpy is a state function.

It depends only on:

• Initial state
• Final state

It does not depend on the path followed.

---

Example of Hess’s Law

Formation of CO₂:

Direct method:$$C+O_2\rightarrow CO_2$$C+O2​→CO2​

or through two steps:

Step 1:$$C+ \frac12O_2\rightarrow CO$$C+21​O2​→CO

Step 2:$$CO+\frac12O_2\rightarrow CO_2$$CO+21​O2​→CO2​

Total:$$\Delta H=\Delta H_1+\Delta H_2$$ΔH=ΔH1​+ΔH2​

---

Applications of Hess’s Law

It is used to calculate:

• Enthalpy of formation
• Enthalpy of combustion
• Bond energies
• Heat changes that cannot be measured directly

---

Quick Revision Table

Quantity	Meaning
ΔrH°	Standard reaction enthalpy
ΔfH°	Formation enthalpy
ΔfusH°	Melting heat
ΔvapH°	Vaporisation heat
ΔsubH°	Sublimation heat

---

Important Exam Questions

1. Define standard enthalpy of formation.

2. What is a thermochemical equation?

3. State Hess’s law of constant heat summation.

4. Why is enthalpy a state function?

5. Explain the difference between exothermic and endothermic reactions.

Chapter Notes (Part 5 – Final Part)

Topics Covered

• Entropy (ΔS)
• Second Law of Thermodynamics
• Spontaneous and Non-spontaneous Processes
• Gibbs Free Energy (ΔG)
• Relation Between ΔG and Spontaneity
• Complete Chapter Revision Sheet

Entropy (S)

Definition

Entropy is a measure of the randomness or disorder of a system.

Symbol:$$\boxed{S}$$S​

A system with greater disorder has higher entropy.

---

Understanding Entropy

Particles in matter can be arranged differently:

Solid

• Particles are tightly arranged.
• Movement is limited.
• Disorder is low.

Therefore:$$S_{solid}<S_{liquid}$$Ssolid​<Sliquid​

---

Liquid

• Particles have more freedom of movement.

---

Gas

• Particles move freely.
• Disorder is maximum.

Therefore:$$\boxed{S_{solid}<S_{liquid}<S_{gas}}$$Ssolid​<Sliquid​<Sgas​​

---

Change in Entropy (ΔS)

The change in entropy is:$$\boxed{\Delta S=S_{final}-S_{initial}}$$ΔS=Sfinal​−Sinitial​​

---

Cases

1. Increase in Entropy

$$\Delta S>0$$ΔS>0

Example:$$H_2O(l)\rightarrow H_2O(g)$$H2​O(l)→H2​O(g)

Liquid changes into gas, so disorder increases.

---

2. Decrease in Entropy

$$\Delta S<0$$ΔS<0

Example:$$H_2O(g)\rightarrow H_2O(l)$$H2​O(g)→H2​O(l)

Gas changes into liquid, so disorder decreases.

---

Factors Affecting Entropy

1. Physical State

Entropy order:$$Gas > Liquid > Solid$$Gas>Liquid>Solid

---

2. Temperature

Higher temperature increases molecular movement.

Therefore:

Higher temperature → Higher entropy

---

3. Number of Gas Molecules

More gas molecules generally mean more randomness.

Example:$$N_2+3H_2\rightarrow2NH_3$$N2​+3H2​→2NH3​

Gas molecules decrease, so entropy decreases.

---

4. Molecular Complexity

Larger molecules generally have more possible arrangements and higher entropy.

---

Second Law of Thermodynamics

Statement

A spontaneous process is accompanied by an increase in the total entropy of the universe.

Mathematically:$$\boxed{\Delta S_{universe}>0}$$ΔSuniverse​>0​

For an equilibrium process:$$\Delta S_{universe}=0$$ΔSuniverse​=0

---

Spontaneous Process

Definition

A process that occurs naturally under given conditions is called a spontaneous process.

Examples:

• Flow of heat from hot object to cold object
• Expansion of gas
• Burning of fuel

---

Important Point

A spontaneous process:

• Does not always occur quickly.
• Depends on thermodynamic conditions.

Example:

Conversion of diamond into graphite is thermodynamically possible but occurs extremely slowly.

---

Non-spontaneous Process

A process that does not occur naturally and requires continuous energy supply is called a non-spontaneous process.

Examples:

• Electrolysis of water
• Formation of glucose by photosynthesis

---

Factors Affecting Spontaneity

Spontaneity depends on:

1. Enthalpy change (ΔH)
2. Entropy change (ΔS)
3. Temperature

---

Gibbs Free Energy (G)

Definition

Gibbs free energy is the energy available in a system that can be used to perform useful work.

Symbol:$$\boxed{G}$$G​

---

Gibbs Energy Change (ΔG)

The equation is:$$\boxed{\Delta G=\Delta H-T\Delta S}$$ΔG=ΔH−TΔS​

Where:

• ΔG = Gibbs energy change
• ΔH = Enthalpy change
• T = Temperature in Kelvin
• ΔS = Entropy change

---

Significance of ΔG

Case 1: ΔG Negative

$$\boxed{\Delta G<0}$$ΔG<0​

Reaction is spontaneous.

Example:

Heat-releasing reactions.

---

Case 2: ΔG Positive

$$\boxed{\Delta G>0}$$ΔG>0​

Reaction is non-spontaneous.

Energy must be supplied.

---

Case 3: ΔG = 0

System is at equilibrium.

No net change occurs.

---

Relationship Between ΔH, ΔS and ΔG

ΔH	ΔS	Spontaneity
Negative	Positive	Always spontaneous
Positive	Negative	Never spontaneous
Negative	Negative	Depends on temperature
Positive	Positive	Depends on temperature

---

Temperature Effect on Spontaneity

Case 1

$$\Delta H<0,\Delta S>0$$ΔH<0,ΔS>0

Both terms favour spontaneity.

Reaction is spontaneous at all temperatures.

---

Case 2

$$\Delta H>0,\Delta S<0$$ΔH>0,ΔS<0

Both terms oppose spontaneity.

Reaction is never spontaneous.

---

Case 3

$$\Delta H<0,\Delta S<0$$ΔH<0,ΔS<0

Spontaneous at low temperature.

---

Case 4

$$\Delta H>0,\Delta S>0$$ΔH>0,ΔS>0

Spontaneous at high temperature.

---

Relation Between Gibbs Energy and Equilibrium Constant

For a reaction:$$aA+bB\rightarrow cC+dD$$aA+bB→cC+dD

The relationship is:$$\boxed{\Delta G^\circ=-RT\ln K}$$ΔG∘=−RTlnK​

Where:

• R = Gas constant
• T = Temperature
• K = Equilibrium constant

---

Meaning of Equilibrium Constant

If:

$$K>1$$K>1

Products are favoured.$$\Delta G^\circ<0$$ΔG∘<0

Reaction is spontaneous.

---

If:

$$K<1$$K<1

Reactants are favoured.$$\Delta G^\circ>0$$ΔG∘>0

Reaction is non-spontaneous.

---

Complete Thermodynamics Formula Sheet

1. First Law

$$\boxed{\Delta U=q+w}$$ΔU=q+w​

---

2. Work

$$\boxed{w=-P_{ext}\Delta V}$$w=−Pext​ΔV​

---

3. Enthalpy

$$\boxed{H=U+PV}$$H=U+PV​

---

4. Enthalpy Change

$$\boxed{\Delta H=\Delta U+P\Delta V}$$ΔH=ΔU+PΔV​

---

5. Gas Reaction

$$\boxed{\Delta H=\Delta U+\Delta n_gRT}$$ΔH=ΔU+Δng​RT​

---

6. Heat Capacity

$$\boxed{q=C\Delta T}$$q=CΔT​

---

7. Specific Heat

$$\boxed{q=mc\Delta T}$$q=mcΔT​

---

8. Heat Capacity Relation

$$\boxed{C_p-C_v=R}$$Cp​−Cv​=R​

---

9. Entropy Change

$$\boxed{\Delta S=S_2-S_1}$$ΔS=S2​−S1​​

---

10. Gibbs Energy

$$\boxed{\Delta G=\Delta H-T\Delta S}$$ΔG=ΔH−TΔS​

---

11. Gibbs-Energy and Equilibrium

$$\boxed{\Delta G^\circ=-RT\ln K}$$ΔG∘=−RTlnK​

---

Chapter One-Page Revision

Thermodynamics studies:

Heat + Work + Energy changes

System:

Part being studied

Surroundings:

Everything outside the system

First Law:

Energy is conserved$$\Delta U=q+w$$ΔU=q+w

Enthalpy:

Heat change at constant pressure$$\Delta H=q_p$$ΔH=qp​

Entropy:

Measure of disorder

Spontaneous Process:

Natural process

Gibbs Energy:

Predicts spontaneity$$\Delta G<0$$ΔG<0

means spontaneous.

---

Most Important Board Exam Questions

1. State and explain the first law of thermodynamics.

2. Derive the relationship between ΔH and ΔU.

3. Explain Hess’s law with an example.

4. Define entropy and explain factors affecting it.

5. What is Gibbs free energy? Explain its importance.

6. Explain conditions for spontaneity using ΔG.

7. Differentiate between:

• Open, closed and isolated systems
• Exothermic and endothermic reactions
• ΔH and ΔU
• Reversible and irreversible processes