NCERT Class 11 Chemistry Chapter 6 Notes Thermodynamics - Download PDF

Thermodynamics is one of the important branches of chemistry that deals with the change in energy in any Physical and Chemical process. It examines how energy moves between different systems and their surroundings, with an emphasis on important ideas like work and heat, along with the basic rules that control these changes. Thermodynamics explains the reason for the occurrence of any reaction, transfer of energy, etc. Students of NCERT Class 11 Chemistry must grasp thermodynamics. It helps us understand how chemical processes function, how materials change states, and how reactions balance. It also allows us to determine whether a reaction can occur on its own by calculating the amount of energy required.

This Story also Contains

1. NCERT Class 11 Chapter 5 Notes
2. System and Surroundings
3. Types of the system:
4. The State of the System
5. Applications
6. Measurement of ∆U and ∆H: Calorimetry
7. Hess’s Law
8. Born-Haber Cycle
9. Gibb’s Energy and Spontaneity
10. Entropy and Second Law of Thermodynamics
11. Absolute Entropy and Third Law of Thermodynamics
12. NCERT Class 11 Notes Chapter-Wise
13. Subject Wise NCERT Exemplar Solutions
14. Subject Wise NCERT Solutions

Thermodynamics is something we deal with every day. Consider how automobile engines convert fuel into action or how refrigerators keep food fresh. These concepts are vital for using energy effectively. Our bodies use the same ideas- metabolism turns chemical energy from food into the energy needed to live. These thermodynamics notes is according to Class 11 CBSE Chemistry Syllabus gives students valuable knowledge about how energy changes and how to use that understanding in real-world situations.

Also, students can refer,

NCERT Class 11 Chapter 5 Notes

System and Surroundings

A system can be defined as where the observations were made, it is the part of the universe.

Surrounding can be defined as the part other than the system is, surrounding.

The system and surroundings can be differentiated by a wall called a boundary.

Types of the system:

The system can be further divided into three categories such as-

1. Open system- Boundary is imaginary in this system and exchange of mass or matter and energy both can be possible between system and surroundings. Example: Open beaker.

2. Closed system-No exchange of matter occurs but the exchange of energy can be possible in between the system and surroundings. Example: closed container of steel or copper.

3. Isolated system- No exchange of matter or energy can be possible. Example: Closed container like a thermos flask.

The State of the System

In general concern when we have to make any calculations we require measurable quantities or properties.

Those quantities are temperature, pressure, the volume combines to form the composition of the system. But in order to deal with these quantities, we must have average conclusions of such measurements so we introduce a term called state functions or state variables. These quantities like pressure(p), volume(V), temperature(T), and amount(n).such variables like p, V, T, n are stated as state variables or state functions.

Internal energy: The combined energy or sum of energies such as mechanical, electrical or chemical are so-called internal energy of the system. This can be denoted in thermodynamics by the letter ‘U’. This can be affected by a change in any of the following:

• Heat goes into and out of the system

• Work done; it can be done on the system or by the system.

• Matter; Exchange of matter( entering or leaving)

General case: This can b shown using an equation.

∆U=q+w

Here the change in internal energy is the sum of heat passing in and out and the work done. It is independent of the change and also depends on the initial and final state of the system. The equation that we get for change in internal energy is the mathematical form of the First law of thermodynamics.

Adiabatic system: It is the system in which the system and surroundings are separated by a wall, where no transfer or exchange of matter or energy occurs.

Applications

Work: Here we take an example of a cylinder or piston, which is fitted perfectly and is filled with a frictionless container with one mole of gas.

Let us first discuss the variables to calculate the work done.

The pressure of the gas=p

The volume of the gas=V

The external pressure of the gas=pext

Area of cross section=A

The piston is moved to a certain distance, taken as length=l

Change in volume∆V=A×l=(V_f-V_l)

So here, the pressure inside the piston becomes

p=FA

Case1: Expansion process

In the case of this process, work is done by the system, by expanding the volume of the system. When the work is done by the system then the negative sign convention will be used in it. so the value of work done will be

W=-p∆V

Now consider the two cases of expansion of ideal gas.

1. Work is to be done in isothermal and reversible condition, where the expansion of ideal gas occur.

Reversible work is taken from initial volume to final volume.

Wrev=-ViVf_pextdV=-V_iV_f(p_in±dP)dV

The value of (dP)dV is so small, so we can write it as;

Wrev=-ViVfp_indV

The cylinder is filled with an ideal gas, so consider the ideal gas equation.

PV=nRT

P=nRTV

At constant temperature conditions;

Wrev=-ViVfnRTdVV

=-2.303nRTlogVfVi

B. Isothermal condition, where expansion of ideal gas occur.

• For irreversible change:

q=-W=pext(Vf-Vi)

• For reversible change

q=-W=2.303nRTlogVfVi

• For adiabatic change

∆U=W_ad

Enthalpy

Enthalpy can be denoted by the letter ‘H’.

Enthalpy is defined as the sum of pressure volume and internal energy.

Change in enthalpy can be defined as the heat is evolved or absorbed at constant pressure conditions in a given system.

For the Exothermic process; Heat is given out of the system, whereas for the Endothermic process heat is given to the system through surroundings.

Relationship between Change in enthalpy and change in internal energy of the system

p∆V=∆n_gRT

∆H=∆U+∆n_gRT

1. Extensive property:

Extensive property can be defined as that property in which the value depends on the quantity or the size that is present in the system.

Examples of extensive properties are mass, enthalpy, volume, etc.

2. Intensive property:

Intensive property can be defined as that property in which the value does not depend on the quantity or the size that is present in the system.

Examples of intensive properties are temperature, density, etc.

3. Heat capacity:

Heat capacity may be defined increase in temperature is directly proportional to heat transferred.

4. Relationship between the value of C_p and C_v.

Constant volume value of heat capacity; C_v.

Constant pressure value of heat capacity; C_p.

C_p-C_v=R

Measurement of ∆U and ∆H: Calorimetry

Measurement of the value of change in internal energy and change in enthalpy with Bomb calorimeter.

Enthalpy changes occur during the process of phase transformation.

• Enthalpy of Fusion: Enthalpy of fusion can be defined as the conversion of one mole of solid is accompanied to change it in liquid, condition of fusion is governed at its melting point

• Enthalpy of Vaporisation: Enthalpy of vaporization can be defined as the conversion of one mole of solid is accompanied to change it in liquid, the condition of fusion is governed at its boiling point.

• Enthalpy of Sublimation: Enthalpy of sublimation is defined as the change in when one mole of solid is converted to gaseous state at its melting point. But temperature is below at its melting range.

• Enthalpy of Combustion: Enthalpy of fusion is defined as the change in when one mole of substance burn in excess amount of air.

• Standard enthalpy of formation: It is defined as in standard condition of temperature at 273K and pressure at 1 atm formation of one mole of substance occur with the help of their constituent elements.

• Thermochemical equations: It is defined as in chemical reaction when the products and reactants are present with their physical value and value of∆rH present such reaction equation is known as Thermochemical equation.

Hess’s Law

It is defined as Hess’s law for constant heat summation.

In general, to solve the questions we can write the equation as follows:

• On combining the two reactions we get the final result as shown above.

• The constant heat summation law suggest that whether the heat is absorbed or evolved in one or many steps are involved, the total amount of heat is same.

Born-Haber Cycle

Spontaneity

Spontaneous process: It is the process that has been taken by itself or tendency to occur on its own.

The process is not instantaneous in nature. Speed can vary from slow to fast.

Examples of such process: Common salt dissolves in water on its own.

Entropy

Entropy is defined as the degree of disorders and randomness in a system.

The entropy of the gaseous system is found to be more than in solid state.

Change in entropy∆S=q_revT=∆HT;

Change is entropy can be denoted as ∆S, and is zero for equilibrium condition.

Gibb’s Energy and Spontaneity

It is defined as mathematically ∆G=∆H-T∆S

The above reaction is at constant temperature and pressure.

1. The value of ∆G is negative or less than zero, the process would be spontaneous.

2. The value of ∆G is positive or more than zero, the process would be non-spontaneous.

Condition of equilibrium, when all reactants and products are in standard condition, value of ∆G will be:

0=∆_rG-RTlnK

∆rG=∆_rH-T∆_rS

=-RTlnK

Entropy and Second Law of Thermodynamics

We know that for an isolated system the change in energy remains constant. Therefore, increase in entropy in such systems is the natural direction of a spontaneous change. This, in fact is the second law of thermodynamics. Like first law of thermodynamics, second law can also be stated in several ways. The second law of thermodynamics explains why spontaneous exothermic reactions are so common. In exothermic reactions heat released by the reaction increases the disorder of the surroundings and overall entropy change is positive which makes the reaction spontaneous.

Absolute Entropy and Third Law of Thermodynamics

The entropy of any pure crystalline substance approaches zero as the temperature approaches absolute zero. This is called third law of thermodynamics. This is so because there is perfect order in a crystal at absolute zero. The statement is confined to pure crystalline solids because theoretical arguments and practical evidences have shown that entropy of solutions and super cooled liquids is not zero at 0 K.

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