Class 11 Chemistry Notes on Chapter 6 – Thermodynamics

Chapter 6 – Thermodynamics

Thermodynamics is a cornerstone of physical chemistry, exploring energy transformations within systems and the universe. For students diving into class 11 chemistry notes on [Chapter 6 – Thermodynamics], mastering the fundamentals here paves the way for understanding advanced topics later. This guide unpacks key concepts with clarity, including definitions, laws, and applications.

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Outline

1. Introduction
2. Terminology Involved
   • Definitions of System, Surroundings, and Universe
   • Types of Systems
   • State Functions vs. Path Functions
3. Internal Energy
   • Definition and Concept
   • Factors Affecting Internal Energy
   • Mathematical Representation
4. 1st Law of Thermodynamics
   • Statement of the Law
   • Mathematical Expression
   • Applications in Chemistry
5. Enthalpy
   • Enthalpy Definition
   • Enthalpy Change and Its Importance
   • Standard Enthalpy of Formation
6. Heat Capacity, Specific Heat Capacity & Molar Heat Capacity
   • Definitions and Differences
   • Importance in Thermodynamics
   • Calculations and Examples
7. Relation Between Cp & Cv
   • Derivation and Explanation
   • Applications in Chemistry
8. Calorimeter
   • Types of Calorimeters
   • Principle of Calorimetry
   • Applications in Chemistry
9. Enthalpies of Reaction
   • Types of Reaction Enthalpies
   • Exothermic vs. Endothermic Reactions
   • Practical Applications
10. Hess’s Law of Heat Summation
    • Statement and Explanation
    • Mathematical Proof
    • Practical Applications
11. Bond Enthalpy
    • Definition and Significance
    • Bond Dissociation Energy
    • Calculations and Applications
12. Entropy
    • Understanding Disorder
    • Factors Affecting Entropy
    • Entropy Changes in Reactions
13. Second Law of Thermodynamics
    • Spontaneity and the Law
    • Entropy and Heat Flow
    • Applications in Real-Life Systems
14. Gibbs Free Energy
    • Definition and Formula
    • Determining Spontaneity
    • Applications in Reactions
15. Third Law of Thermodynamics
    • Statement and Explanation
    • Absolute Zero and Entropy
    • Implications for Chemistry
16. Conclusion
17. FAQs

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Introduction

Thermodynamics examines how energy moves through systems, whether it's the heat from a cup of coffee or the work done in a chemical reaction. In class 11 chemistry notes on [Chapter 6 – Thermodynamics], we aim to simplify this dense topic, breaking it into digestible parts for students. Ready to decode the universe's energy puzzle?

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Terminology Involved

Definitions of System, Surroundings, and Universe

In thermodynamics:

• System: The part under observation.
• Surroundings: Everything outside the system.
• Universe: The sum of the system and surroundings.

Types of Systems

1. Open Systems: Exchange matter and energy (e.g., boiling water in an open pot).
2. Closed Systems: Exchange energy but not matter (e.g., a pressure cooker).
3. Isolated Systems: Exchange neither matter nor energy (e.g., a thermos flask).

State Functions vs. Path Functions

State functions depend only on the initial and final states (e.g., enthalpy, entropy), while path functions depend on the process path (e.g., heat, work).

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Internal Energy

Definition and Concept

Internal energy ($U$) is the total energy contained within a system, including kinetic and potential energies of molecules.

Factors Affecting Internal Energy

1. Temperature: Higher temperature increases $U$.
2. Pressure and Volume: Affects potential energy in gases.
3. Chemical Composition: Bond energy contributes to $U$.

Mathematical Representation

The change in internal energy ($\Delta U$) is:

$$\Delta U = q + w$$

Where:

• $q$ = Heat exchanged
• $w$ = Work done on or by the system

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1st Law of Thermodynamics

Statement of the Law

Energy cannot be created or destroyed but can only transform from one form to another.

Mathematical Expression

$$\Delta U = q + w$$

Where:

• $q$ is positive when heat is absorbed.
• $w$ is positive when work is done on the system.

Applications in Chemistry

• Explains heat exchange in reactions.
• Used in calculating bond energies and reaction enthalpies.

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Enthalpy

Enthalpy Definition

Enthalpy ($H$) is the heat content of a system at constant pressure:

$$H = U + PV$$

Enthalpy Change and Its Importance

$$\Delta H = H_{\text{final}} - H_{\text{initial}}$$

A positive $\Delta H$ indicates an endothermic process, while a negative $\Delta H$ signifies an exothermic process.

Standard Enthalpy of Formation

The heat change when one mole of a compound is formed from its elements under standard conditions (298 K, 1 atm).

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Heat Capacity, Specific Heat Capacity & Molar Heat Capacity

Definitions and Differences

• Heat Capacity ($C$): The amount of heat required to raise the temperature of a system by 1 degree Celsius.
• Specific Heat Capacity ($C_s$): Heat required to raise the temperature of 1 gram of a substance by 1 degree Celsius.
• Molar Heat Capacity ($C_m$): Heat required to raise the temperature of 1 mole of a substance by 1 degree Celsius.

Importance in Thermodynamics

These properties allow scientists to quantify how different substances react to temperature changes. Water, with its high specific heat capacity, for instance, stabilizes Earth's climate by absorbing and releasing heat slowly.

Calculations and Examples

The heat ($q$) absorbed or released can be calculated using:

$$q = mc\Delta T$$

Where:

• $m$ = Mass of the substance
• $c$ = Specific heat capacity
• $\Delta T$ = Temperature change

For example, to heat 10 g of water ($c = 4.18 \, \text{J/g°C}$) by 25°C:

$$q = 10 \times 4.18 \times 25 = 1045 \, \text{J}$$

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Relation Between Cp & Cv

Derivation and Explanation

For an ideal gas, the relation between molar heat capacities at constant pressure ($C_p$) and constant volume ($C_v$) is:

$$C_p - C_v = R$$

Where $R$ is the gas constant.

Applications in Chemistry

This relationship helps in calculating the thermodynamic properties of gases, like adiabatic processes where no heat is exchanged with the surroundings.

Key Example

For diatomic gases:

$$C_v = \frac{5}{2}R, \, C_p = \frac{7}{2}R$$

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Calorimeter

Types of Calorimeters

1. Bomb Calorimeter: Measures heat of combustion at constant volume.
2. Coffee Cup Calorimeter: Measures heat of reaction at constant pressure.

Principle of Calorimetry

Heat lost by one part of the system equals heat gained by another, maintaining energy conservation:

$$q_{\text{lost}} + q_{\text{gained}} = 0$$

Applications in Chemistry

Calorimeters are indispensable in determining enthalpies of reactions, specific heat capacities, and the calorific values of fuels.

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Enthalpies of Reaction

Types of Reaction Enthalpies

1. Heat of Combustion: Heat released during complete combustion.
2. Heat of Formation: Heat change when a compound forms from its elements.
3. Heat of Neutralization: Heat released during acid-base reactions.

Exothermic vs. Endothermic Reactions

• Exothermic: Releases heat ($\Delta H < 0$).
• Endothermic: Absorbs heat ($\Delta H > 0$).

Practical Applications

• Understanding reaction energies in fuel combustion.
• Designing industrial chemical processes.

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Hess's Law of Heat Summation

Statement and Explanation

Hess’s law states:

$$\Delta H_{\text{reaction}} = \sum \Delta H_{\text{products}} - \sum \Delta H_{\text{reactants}}$$

This implies that the total enthalpy change is independent of the reaction pathway.

Mathematical Proof

By combining two or more thermochemical equations, the overall enthalpy change equals the sum of individual enthalpy changes.

Practical Applications

• Calculating enthalpy changes for reactions that are challenging to measure directly.
• Predicting energies for multi-step processes in industrial settings.

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Bond Enthalpy

Definition and Significance

Bond enthalpy is the energy required to break one mole of a specific bond in a gaseous molecule. It’s an average value for bonds in polyatomic molecules.

Bond Dissociation Energy

This refers to the energy needed to break a specific bond in a molecule, influencing reaction energetics.

Calculations and Applications

For a reaction:

$$\Delta H = \text{Sum of bond enthalpies (reactants)} - \text{Sum of bond enthalpies (products)}$$

Example: In $H_2 + Cl_2 \rightarrow 2HCl$, Bond energies: $H-H = 436 \, \text{kJ/mol}, \, Cl-Cl = 242 \, \text{kJ/mol}, \, H-Cl = 431 \, \text{kJ/mol}$.

$$\Delta H = [436 + 242] - [2 \times 431] = -184 \, \text{kJ/mol}$$

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Entropy

Understanding Disorder

Entropy ($S$) quantifies the randomness or disorder in a system. Greater disorder means higher entropy.

Factors Affecting Entropy

1. State of Matter: Gases have higher entropy than liquids or solids.
2. Temperature: Higher temperature increases molecular motion, raising entropy.
3. Mixing: Mixing substances usually increases entropy.

Entropy Changes in Reactions

Reactions with an increase in gas moles often show positive $\Delta S$, indicating higher disorder.

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Second Law of Thermodynamics

Spontaneity and the Law

The Second Law of Thermodynamics states that in any spontaneous process, the total entropy of the universe increases:

$$\Delta S_{\text{universe}} = \Delta S_{\text{system}} + \Delta S_{\text{surroundings}} > 0$$

This law explains why certain processes occur naturally, like heat flowing from a hot object to a cold one, but not the reverse without external work.

Entropy and Heat Flow

The transfer of heat ($q$) affects entropy ($S$) according to:

$$\Delta S = \frac{q_{\text{rev}}}{T}$$

Where $q_{\text{rev}}$ is the heat exchanged in a reversible process and $T$ is the temperature.

Applications in Real-Life Systems

• Heat engines and refrigerators rely on this principle.
• Understanding energy losses in processes helps improve efficiency.

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Gibbs Free Energy

Definition and Formula

Gibbs free energy ($G$) determines the spontaneity of a process. Its formula is:

$$G = H - TS$$

Where:

• $H$ = Enthalpy
• $T$ = Temperature in Kelvin
• $S$ = Entropy

Determining Spontaneity

1. $\Delta G < 0$: Process is spontaneous.
2. $\Delta G = 0$: System is at equilibrium.
3. $\Delta G > 0$: Process is non-spontaneous.

Applications in Reactions

• Predicting feasibility of chemical reactions.
• Designing processes in industries like pharmaceuticals and energy.

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Third Law of Thermodynamics

Statement and Explanation

The Third Law of Thermodynamics states that the entropy of a perfect crystal at absolute zero (0 K) is zero:

$$S = 0 \text{ at } T = 0 \, \text{K}$$

Absolute Zero and Entropy

At absolute zero, molecular motion ceases, leading to no disorder. This serves as the baseline for measuring absolute entropy values.

Implications for Chemistry

• Enables calculation of standard entropies for substances.
• Helps in understanding the thermodynamics of reactions at very low temperatures.

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Conclusion

Thermodynamics in class 11 chemistry notes on [Chapter 6 – Thermodynamics] provides a gateway to understanding how energy governs the physical and chemical transformations around us. From the elegance of the First Law to the practicality of Gibbs free energy, this chapter equips students with tools to predict and analyze the world at both molecular and macroscopic levels.

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FAQs

1. What is the significance of the First Law of Thermodynamics?

The First Law establishes that energy is conserved in a system. It explains energy transformations in chemical reactions, enabling calculations of work and heat.

2. How do enthalpy and internal energy differ?

Internal energy ($U$) includes all forms of energy in a system, while enthalpy ($H$) specifically accounts for energy at constant pressure, adding the $PV$ work term.

3. Why is entropy important in chemistry?

Entropy measures disorder in a system, helping to predict the spontaneity of reactions and the feasibility of processes based on the Second Law of Thermodynamics.

4. What is the practical use of Hess’s Law?

Hess’s Law simplifies the calculation of enthalpy changes for reactions where direct measurement is impractical, such as multi-step processes.

5. How does Gibbs free energy predict reaction spontaneity?

By considering enthalpy, entropy, and temperature, Gibbs free energy ($G$) determines if a reaction will occur naturally ($\Delta G < 0$) or require external input.

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