Thermodynamics Class 11 Notes CBSE Chemistry Chapter 6 (Free PDF Download)

Section – A (1 Mark Questions)

1. What do you mean by system? 

Ans. A system in thermodynamics refers to that part of the universe in which observations are made.

2. What do you mean by surroundings?

Ans. The rest of the universe which might be in a position to exchange energy and matter with the system

is called its surroundings.

3. Explain first law of thermodynamics.

Ans. The first law of thermodynamics states that ‘the energy of an isolated system is constant’.

4. ‘Coffee held in a cup’ is what type of system?

Ans. Coffee held in a cup is an open system because it can exchange matter (water vapors) and energy

(heat) with the surroundings.

5. List one example of isolated system.

Ans. Coffee held in a thermos flask is an isolated system because it can neither exchange energy

nor matter with the surroundings.

6. If work is done by the system, what will be the effect on internal enery of the system.

Ans. The internal energy of the system will decrease if work is done by the system.

7. In thermodynamics we discuss about open and close sytsem. So, animals and plants belongs to?

Ans. Animals and plants belongs to Open system.

8. Explain state of thermodynamic system?

Ans. The state of thermodynamic system may be defined by specifying values of state variables like

temperature, pressure, volume.

9. What do you mean by Enthalpy?

Ans. Enthalpy is defined as total heat content of the system.

10. How do you define enthalpy mathematically?

Ans. Mathematically.

H = U + pV where U is internal energy.z

Section – B (2 Marks Questions)

11. Internal energy is state function while work done is not. Why?

Ans. The change in internal energy during a process depends only upon the initial and final state of the

system.

Therefore, it is a state function. But the work is related the path followed. Therefore, it is not a state function.

12. Predict the sign of $\Delta\text{S}$ for the following reaction

$\text{CaCO}_3(\text{s})\overset{\Delta}{\rightarrow}\text{CaO}(\text{s})+\text{CO}_2(\text{g})$

Ans.  $\Delta$ S is positive because on the product side we have gaseous CO2. 

13. What do you mean by spontaneous process.

Ans. A spontaneous process is an irreversible process and may only be reversed by some external agency.

14. What do you mean by non-spontaneous process?

Ans. A process is said to be non-spontaneous if it does not occur of its own under given condition and

occur only when an external force is continuously applied.

15. The standard heat of formation of Fe2O3(s) is 824.2kJ mol–1. Calculate heat change for the

reaction.

4Fe(s) + 3O2(g) → 2Fe2O3(s)

Ans.

$\Delta\text{H}=\sum\Delta\text{H}^{\text{O}}_{\text{f}}(\text{products})-\sum\Delta\text{H}^{\text{O}}_{\text{f}}(\text{reactants})$

$=[2\times\Delta\text{H}^{\text{O}}_{\text{f}}\text{Fe}_2\text{O}_3(\text{s})]-[4\Delta\text{H}^{\text{O}}_{\text{f}}\text{Fe}(\text{s})+3\Delta\text{H}^{\text{O}}_{\text{f}}\text{O}_2(\text{g})]$

= 2 (-824.2 kJ) - [4 x 0 + 3 x 0]

= - 1648.4 kJ

16. Explain Heat capacity and what are the units of heat capacity?

Ans. The heat capacity for one mole of the substance is the quantity of heat needed to raise the

temperature of one mole by one degree Celsius. 

Its unit is : J/K

17. Explain specific heat.

Ans. Specific heat/specific heat capacity is the quantity of heat required to raise the temperature of one unit

mass of a substance by one degree Celsius (or one Kelvin).

18. For solids or liquids, there is no significant difference between $\Delta\text{H}$ and $\Delta\text{H}$ .

Ans. The difference between $\Delta\text{H}$ and $\Delta\text{U}$ is not usually significant for systems

consisting of only solids and/or liquids because they do not suffer any significant volume changes upon

heating.

19. Define extensive and intensive properties?

Ans. Extensive property is a property whose value depends on the quantity or size of matter present in the

system.

Intensive property is a property which do not depend upon the quantity or size of matter present.

20. C(graphite) + O2(g) → CO2(g) ΔH= - 393.5 kJ mol–1 . This is combustion reaction of C and CO2 is forming. Write the $\Delta\text{H}^{\text{O}}$ values of the two preocesses.

Ans. (i) The standard enthalpy of formation of CO2 is –393.5 kJ per mole of CO2.

That is $\Delta\text{H}^{\text{O}}(\text{CO}_2,\text{g})$ = –393.5 kJ mol–1

(ii) The standard enthalpy of combustion of carbon is 393.5 kJ per mole of carbon i.e $\Delta\text{H}^{\text{O}}$ comb(c)

= –393.5 kJ mol–1

PDF Summary - Class 11 Chemistry Thermodynamics Notes (Chapter 10)

Thermodynamics:

The study of the flow of mass, heat and energy is the study of thermodynamics.

Thermodynamics terminology:

• System:

A notable part of the universe that is kept under observation is known as the system.

• Surrounding:

The remaining part of the universe except for the system which isn’t kept under observation is known as surroundings.

In general, it can be stated as;

Universe = System + Surrounding

• Types of the system:

a) Open system – 

The system where the flow of both, mass and heat energy takes place.

Example: Human body.

b) Closed system –

The system where the flow of heat energy takes place but has constant mass.

Example: Pressure cooker.

c) Isolated system – 

The system where none of the flow takes place.

Example: Thermos flask.

State of the system:

The state of the system can be defined and changed with respect to the changes in state variables i.e., P, V, T and n. These variables define the conditions of the system and change in any one of them, will change the state of the system. 

Properties of the system:

• Intensive properties – 

Properties depending upon concentration and are independent of mass or the total number of particles in the system. They are pressure, refractive index, density, etc.

• Extensive properties – 

Properties depending upon the mass or the total number of particles in the system. They are volume, total energy, etc.

State and path function:

• State function –

The function will be independent of the path followed but will depend upon the initial and final states while bringing up the changes in the system.

Example: internal energy, enthalpy, etc.

• Path function –

The function will depend upon the path followed while bringing up the changes in the system. 

Example: work, heat, etc.

Thermodynamic equilibrium:

The system remains in equilibrium when the state variables do not change and the below three types of equilibrium are satisfied.

• Mechanical equilibrium –

The absence of mechanical motion, constant pressure and volume bring up the mechanical equilibrium.

• Thermal equilibrium – 

The constant heat and temperature with respect to time bring up thermal   equilibrium.

• Chemical equilibrium –

The rate of forward reaction equal to the rate of backward reaction brings up the chemical equilibrium. 

Internal energy:

The sum total of the components of the energy influenced by the internal factors of the system is known as internal energy; often denoted by U or E.

The system under observation acts as an ideal gas system that depends only upon kinetic energy and hence, is the function of temperature as $U\propto T$. Thus, the internal energy is a state function.

Modes of energy transport:

• Heat –

The energy transferred due to temperature differences within the system and surroundings is known as heat (Q). When the system is heated, the kinetic energy of the molecules is being increased which then increases the internal energy. 

• Work –

The energy spent to overcome the external forces acting upon the system is known as work (W). When a system expands, the internal energy is reduced. Whereas, on the contraction of the system the internal energy is increased.   

The first law of thermodynamics:

The first law of thermodynamics states that energy can neither be created nor destroyed.

\[\Delta U=Q+W\] 

The sign conventions are given as;

Work done by the system = - W

Work done on the system = + W

Heat flows into the system = + Q

Heat flows out of the system = - Q

Reversibility:

The process can change its direction by very small i.e., infinitesimal change in the system or surrounding; retracking its original path reaching the same initial state. In a process to follow reversibility, there must not be any dissipative forces and the system must be in Quasi-Static State.

• Quasi-static state –

Here, the system seems to be static at all time intervals but not actually in reality. The motion is so slow that the system seems to be in equilibrium with the surroundings.

Expansion work:

The work done due to changes in the volume of the system is known as expansion work. Note that, let it be expansion or compression, we take external pressure as the driving force. 

Mathematically, it can be represented as;

\[W=-\int{{{P}_{ex}}dV}\] 

For reversible processes, external pressure is considered equal to the pressure of the gas. Thus,

\[W=-\int{{{P}_{gas}}dV}\] 

When a P – V graph is drawn, work done is represented as the area covered under it as shown;

Expansion Work

Sign conventions:

• W –

Positive if the volume of the system is decreasing and negative when the volume of the system is increasing.

• $\Delta U$ - 

When the temperature of the system or product pressure or volume is reducing, it is negative; else is positive.

• Q –

This needs to be determined by the first law of thermodynamics.

Cyclic process:

A process that comes back to its original and initial state is known as a cyclic process. A closed graph determines this process and here, $\Delta U=0$ and ${{Q}_{net}}=-{{W}_{net}}$.

Enthalpy:

A thermodynamic state function is defined as the sum of energy stored in the system and the energy used in doing work. Mathematically, can be represented as;

\[\Delta H=U+PV\] 

• At constant P, $\Delta H={{Q}_{P}}$.

• At constant V, $\Delta U={{Q}_{V}}$.

Molar heat capacity:

• At constant Pressure –

The amount of heat needed to raise the temperature of one mole of gas by a degree at constant pressure. It can be stated as;

\[{{C}_{P}}=\frac{{{Q}_{P}}}{n\Delta T}\] 

• At constant Volume –

The amount of heat needed to raise the temperature of one mole of gas by a degree at constant volume. It can be stated as;

\[{{C}_{V}}=\frac{{{Q}_{V}}}{n\Delta T}\] 

We can now say that, $\Delta H=n{{C}_{P}}\Delta T$ and $\Delta U=n{{C}_{V}}\Delta T$ .

Types of thermodynamic processes:

• Isothermal process –

The constant temperature process is known as the isothermal process. Here, $\Delta U=0$ and $\Delta H=0$ .

\[W=-2.303nRT\log \frac{{{V}_{2}}}{{{V}_{1}}}=-2.303nRT\log \frac{{{P}_{1}}}{{{P}_{2}}}\] 

\[Q=2.303nRT\log \frac{{{V}_{2}}}{{{V}_{1}}}=2.303nRT\log \frac{{{P}_{1}}}{{{P}_{2}}}\] 

• Adiabatic process –

When the heat exchanged with the surrounding is zero, such a process is known as adiabatic process. Here,

\[T{{V}^{\gamma -1}}=C,{{T}^{\gamma }}{{P}^{1-\gamma }}=C,P{{V}^{\gamma }}=C\] 

where, C is constant.

\[Q=0\Rightarrow W=\Delta U\] 

Now,

\[\Delta U=n{{C}_{V}}\Delta T=\frac{\left( {{P}_{2}}{{V}_{2}}-{{P}_{1}}{{V}_{1}} \right)}{\left( \gamma -1 \right)}=\frac{\left( nR\Delta T \right)}{\left( \gamma -1 \right)}\]  and

\[\Delta H=n{{C}_{P}}\Delta T\] 

• Isochoric process –

Constant volume process is known as isochoric process. Here, W = 0, $\Delta H=n{{C}_{P}}\Delta T$ and $\Delta U=n{{C}_{V}}\Delta T={{Q}_{V}}$.

• Isobaric process –

Constant pressure process is known as isobaric process. Here, $W=-P\Delta V=-nR\Delta T$ , $\Delta H=n{{C}_{P}}\Delta T={{Q}_{P}}$ and $\Delta U=n{{C}_{V}}\Delta T$.

Graphs:

Graph of Thermodynamic processes

Note that, the P – V graphs of the isothermal and adiabatic processes are similar but the one for adiabatic is steeper than that of isothermal.

• Irreversible process – 

Work done is given as $W=-\int{{{P}_{ex}}dV}$ in the irreversible process. Here, we cannot say external pressure will be equal to that of the pressure of the gas.

• Free expansion – 

In free expansion, the external pressure of the gas is zero i.e., the gas expanding against the vacuum will have work as zero. Thus, no heat will be supplied to the process showing no changes in the temperature. Hence, it is an isothermal and adiabatic process. 

• Polytropic process – 

A generalised form of any thermodynamic process can be represented as $P{{V}^{n}}$ = constant. 

For the isothermal process, n = 1.

For adiabatic process, $n=\gamma $ .

Thermochemical equation:

A chemical equation giving you all the information like phases of reactants and products in the reaction along with energy changes associated with the same is known as a thermochemical equation.

Types of reaction:

• Endothermic reaction – 

The chemical reactions that absorb energy are known as endothermic reactions. Here, $\Delta H=+ve$ .

• Exothermic reaction –

The chemical reactions that release energy are known as exothermic reactions. Here, $\Delta H=-ve$ .

For any chemical reaction, 

\[\text{ }\!\!\Delta\!\!\text{ }{{\text{H}}_{\text{Reaction}}}\text{= }\!\!\Delta\!\!\text{ }{{\text{H}}_{\text{Products}}}\text{- }\!\!\Delta\!\!\text{ }{{\text{H}}_{\text{Reactants}}}\]

This change in enthalpy occurs due to making and breaking of bonds.

Hess law of constant heat summation:

For a reaction that takes place in a stepwise manner, the net change in enthalpy can be calculated as the enthalpy changes in each step. The governing law is known as the Hess law of constant heat summation.

Enthalpy of reactions:

• Enthalpy of bond dissociation –

The energy needed to break the bonds of one-mole molecules is known as the enthalpy of bond dissociation. It is defined per mole of the molecule.

• Enthalpy of combustion –

The heat released or absorbed when a mole of a substance undergoes combustion in presence of oxygen is known as enthalpy of combustion.

• Enthalpy of formation –

The heat released or absorbed when a mole of a compound is formed from its constituent elements under their standard elemental forms is known as enthalpy of formation.

• Enthalpy of atomization –

The energy required to convert any substance to gaseous atoms is known as the enthalpy of atomization. It is defined per mole of the gaseous atoms.

• Enthalpy of sublimation –

The heat required to change a mole of a substance from solid-state to its gaseous state at STP is known as enthalpy of sublimation.

• Enthalpy of phase transition –

The phase transition from one phase to another release or absorbs a particular standard enthalpy which is known as enthalpy of phase transition. 

• Enthalpy of ionization –

The amount of energy an isolated gaseous atom will take to lose an electron in its ground state is known as the enthalpy of ionization.

• Enthalpy of the solution –

The heat released or absorbed when a mole of a compound is dissolved in excess of a solvent (mostly, water) is known as enthalpy of solution.

• Enthalpy of dilution –

The enthalpy change associated with the dilution process of a component in a solution at constant pressure is known as enthalpy of dilution. It is defined as energy per unit mass or amount of substance.

The second law of thermodynamics:

The state of entropy of the entire universe, as an isolated system will always increase over time, is the standard statement of the second law of thermodynamics.

• Need –

The first law of thermodynamics states the conversion of energy in a process but does not explain the feasibility of the same. This point gave rise to the need for the second law of thermodynamics.

Types of processes:

• Spontaneous process –

The spontaneous process has the tendency to take place naturally and no external work is needed to carry out the same.

• Non-spontaneous process –

The non-spontaneous process is driven by external work and cannot be performed naturally.

Entropy:

The measure of randomness or disorder in the process of a body is known as its entropy. It is a state function and is represented as S.

The spontaneous process is the process in which the total randomness of the universe tends to increase. Thus,

\[\Delta S=\frac{{{Q}_{rev}}}{T}\] 

For spontaneous change, $\Delta {{S}_{Total}}=\Delta {{S}_{System}}+\Delta {{S}_{Surrounding}}>0$ .

For reversible processes where the entropy of the universe remains constant, $\Delta {{S}_{Total}}=0$.

Entropy changes in thermodynamic processes:

The entropy changes in any thermodynamic process can be mathematically represented as;

\[\Delta S=n{{C}_{V}}\ln \frac{{{T}_{2}}}{{{T}_{1}}}+nR\ln \frac{{{V}_{2}}}{{{V}_{1}}}\] 

• Isothermal process –

\[\Delta S=nR\ln \frac{{{V}_{2}}}{{{V}_{1}}}\] 

• Isochoric process –

\[\Delta S=n{{C}_{V}}\ln \frac{{{T}_{2}}}{{{T}_{1}}}\] 

• Isobaric process –

\[\Delta S=n{{C}_{P}}\ln \frac{{{T}_{2}}}{{{T}_{1}}}\] 

• Adiabatic process –

\[\Delta S=0\] 

Gibbs free energy:

This gives us the most convenient parameter to judge the spontaneity of the process from the perspective of the system. At constant temperature it can be represented as;

\[\Delta {{G}_{sys}}=\Delta H-T\Delta {{S}_{sys}}\] 

At constant temperature and pressure, $\Delta G=-T\Delta {{S}_{Total}}$ .

For the process to be spontaneous, $\Delta G<0$.

Third law of thermodynamics:

The entropy of the system will approach a constant value as its temperature approaches absolute zero is the empirical statement of the third law of thermodynamics. 

Class 11 Chemistry Chapter 6 Thermodynamics Notes

The following topics are included in Chemistry ch 6 class 11 notes:

1. Types of Thermodynamic Equilibriums

The three major types of thermodynamic equilibrium are – (a) mechanical equilibrium, (b) thermal equilibrium, and (d) chemical equilibrium.

2. Energy Transfer Modes

The two modes of internal energy transfer are through heat and work. Heat is the transference of energy owing to differences in temperature. On the other hand, work amounts to the energy that is used for overcoming an external force.

3. First Law of Thermodynamics

The First Law of Thermodynamics is well explained in Notes of Chemistry class 11 chapter 6. The law states that energy cannot be created or destroyed. It can only be transferred from one form to another.

4. Concept of Enthalpy

The quantity of heat content in a given system having pressure kept at constant is termed as enthalpy. The change in enthalpy corresponds to the heat that is released or absorbed in a specific reaction when pressure remains constant.

5. Thermodynamic Processes

There are four types of thermodynamic processes discussed in notes of thermodynamics Chemistry class 11. Those are:

• Isothermal Process – Temperature remains the same all through the process.

• Adiabatic Process – Heat exchanged with surroundings amounts to zero.

• Isochoric Process – Volume remains same all through the process.

• Isobaric Process – Pressure remains the same all through the process.

6. Concept of Entropy

The concept of entropy is illustrated in chapter 6 class 11 Chemistry notes. It is the measurement of thermal energy present in a system with respect to per unit temperature, which does not result in quantifiable work. Hence, it is used to measure molecular disorder.

7. Hess Law of Constant Heat Summation

Class 11 Chemistry chapter 6 thermodynamics notes include Hess's law. The law states that notwithstanding the number of steps or stages in a reaction, the total change of enthalpy for a particular reaction amounts to an aggregate of all changes. Such an aspect leads to state function.

8. Born Haber Cycle

Born Haber cycle is the process of creation of ionic compounds from the elements in a series of steps. It is indicative of the factors that are favourable for the formation of ionic bonding.

The topics included in class 11 Chemistry chapter 6 thermodynamics notes can be complex. It is quite vital for examination as well. A student must have a clearing understanding of all the concepts while preparing. If you are facing any difficulty within this, feel free to reach out to us or avail our online classes. Download the Vedantu app today to get started.