CHAPTER 5 THERMODYNAMICS · 4. Thermodynamics is independent of time factor i. e. it does not deals with the rate of physical changes or chemical reactions. Thermodynamic terms and

Thermodynamics – Prof. G.B.Patil 1

CHAPTER 5

THERMODYNAMICS

Prof. Ganesh B. Patil

Content

Objectives of the chapter

Introduction of thermodynamics

Scope and Limitations

Basic Terms & Definitions

Thermodynamics

System, Surrounding and Boundary

Homogeneous and Heterogeneous system

Isolated, Closed and Open system

Intensive and Extensive properties

Equilibrium and Non-equilibrium states of system

Isothermal, Adiabatic, Isochoric, Isobaric & Cyclic Processes

Reversible and Irreversible Processes

Internal Energy & Enthalpy

Exothermic & Endothermic reactions / processes

First Law of Thermodynamics

Statement & Mathematical statement

Various forms of first law

Significance of Ist law of thermodynamics

Heat of Reaction, Heat of Formation & Heat of Combustion

Hess’s law of constant heat summation

Statement & Explanation

Application of Hess’s law

Second & Third law of thermodynamics

Spontaneous & Non Spontaneous process

Entropy

Concept of Gibbs and Helmholtz free energy

Zeroth law of thermodynamics.


THERMODYNAMICS

Objectives of the chapter

After completion of this lecture, the student should be able to -

1. Heat, Energy and its various forms.

2. System, Surrounding, Boundary and various thermodynamic processes.

3. Ist law of thermodynamics, Its various forms & Significance

4. Internal Energy, Enthalpy and Entropy of a system.

5. Second & Third law of thermodynamics & Hess’s law.

6. Gibbs free energy and its significance

Introduction

Thermodynamics is concerned with the qualitative relationship between heat and

other forms of energy such as mechanical, chemical electrical or radiant energy.

A body is said to possess kinetic energy because of its motion or motion of its parts

i.e. its molecules, atoms and electrons and to posses potential energy by virtue of its position

or the position or configuration of its parts.

It is not possible to know the absolute value of energy of a system and hence, it is

sufficient to record the changes in the energy that occur when a system undergoes some

transformation. Mechanical energy changes are expressed in ergs or joules while heat change

is expressed in calories. Count Rumford in 1798 and James Joules in 1849 showed the

relationship between mechanical work and heat. Today the calories (as defined by U.S.

National Bureau of standards) is equal to 4.1840 X 10 7 ergs or 4.1840 joules, so that work

and heat can be expressed in the some units.

The study of thermodynamics is based on three laws or facts of experience that have

never been proved in a direct way. Various conclusion usually expressed in the form of

mathematical equation however may be deduced from these three principles, and the result

consistently agree with observation consequently, the laws of thermodynamics, from which

these equation are obtained, are accepted as valid for systems involving large numbers of

molecules.


Scope of thermodynamics

Limitations of thermodynamics

1. Thermodynamics is only applicable to the macroscopic systems and not to the

microscope systems.

2. It is concerned with only initial and final states of system.

3. Thermodynamics is independent of the internal structure of atoms and molecules.

4. Thermodynamics is independent of time factor i. e. it does not deals with the rate of

physical changes or chemical reactions.

Thermodynamic terms and basic concepts

Thermodynamics

“Thermodynamics deals with the study of flow of heat or any other form of

energy, into or out of the system during physical or chemical changes.”

System

“A system is that part of the universe which is under thermodynamic study and the

rest of universe is surrounding.”

Surrounding

“The real or imaginary surface separating the system from the surrounding is called as

boundary.”

E.g. A cylinder containing a gas is placed in a thermostatic bath which is maintained at a

constant temperature (25℃) as shown in above figure

In this case the gas constitutes system where as constant temperature bath is the

surrounding the walls of cylinder constitute the boundary.

Universe

“Combination of System and surrounding constitute Universe”

System + Surrounding = Universe

Homogeneous System

“When a system is uniform (consist of only one phase) throughout then it is called as

a homogeneous system.”

A homogeneous system is made up of one phase only.

E.g. - A pure single solid, liquid or gas or mixture of gases.


A Phase

“A phase is defined as a homogeneous, physically distinct portion of a system that is

separated from other portions of the system by bounding surface.”

Heterogeneous System

“When a system consists of two or more phases, then it is called as a heterogeneous

system.”

A heterogeneous system is made up of two or more phases.

E.g. - Ice cube in contact with water or Ice cube in contact with vapour.

Thermodynamics systems

There are three types of thermodynamic system depending on the nature of boundary.

If the boundary is closed or sealed, no matter can pass through it. If the boundary is insulated,

no energy (say heat) can pass through it.

1. Isolated System

“An isolated system is one that can transfer neither energy nor matter to and

from its surrounding.

Thus, when boundary is both sealed and insulated, no interaction is possible with the

surrounding.

E.g. Tea or coffee contained in a thermos flask.

2. Closed System

“A closed system is one which can transfer energy in the form of heat, and

radiation but cannot transfer the matter to and from its surrounding.”

Here, in closed system, boundary is only sealed but not insulated.

E.g. A specific quantity of hot water contained in a sealed tube. Here, water vapors (Matter )

cannot escape from the system (tube) but heat (energy) can transfer to the surrounding

through the walls of the tube.

3. Open System

“An open system is one which can transfer energy as well as matter to and from

its surrounding”.

In open system, boundary remains open i.e. unsealed and un- insulated.

E.g. Hot water contained in a beaker placed on a table.

Here waster vapour (matter) and also heat (energy) can transfer to the surrounding through

the imaginary boundary


.

INTENSIVE AND EXTENSIVE PROPERTIES OF A SYSTEM

The macroscopic or bulk properties of a system area classified into two categories:

1) Intensive Properties

“A property which does not depend upon the quantity of matter present in the

system is called as intensive property”

E.g. Temperature, Pressure, Density, Viscosity, Concentration, Surface tension,

Refractive index, Boiling point, etc.

2) Extensive Properties

“A property that depends upon the quantity of matter present in the system is called

as extensive property”

E.g. Volume, Mass, Area, Enthalpy, Entropy etc.

STATE OF A SYSTEM

“A thermodynamic system is said to be in a certain state when all its properties

are fixed throughout the system”. The fundamental properties that determines the state of a

system these are Pressure (P), Temperature (T), Volume (V), Mass (M) and Composition (C).

Since change in magnitude of the e properties, alters the state of a system these are

called as state variables or state functions of thermodynamic parameters. It is also believed

that a change of a system from initial state to final state will be accompanied by change in the

state variables. It is not necessary to state all the properties (state variable) to define a system

completely.

For a pure gas, its composition is fixed and remaining properties (state variables like

P, T, and V) remains interrelated in the form of an equation, called as equation of state.

For one mole of pure gas, equation of state can be written as;

PV = RT …………………………………….. (1)

Where, R = gas constant.

Out of these three variables (P,V,T), if P and T are used to define the state of a

system, are called as ‘Independent variables’ and the remaining state variables like

volume (V) which depends on the value of P and T, is called as ‘Dependent variable’.


EQUILIBRIUM AND NON-EQUILIBRIUM STATES

1) EQUILIBRIUM STATE

“A system in which the state variables have constant value throughout the

system is said to be in a state of thermodynamic equilibrium”

Thermodynamics is concerned only with equilibrium state.

E.g. Consider, we have a gas filled in a cylinder having frictionless, stationary piston. If the

piston is stationary, the values of pressure, temperature, volume and density remains constant

and the systems is said to be in a state of thermodynamic equilibrium.

2) NON – EQUILIBRIUM STATE

“A system is said to be in a non-equilibrium state when all its variables have

different properties in different parts of that system”

E.g. Consider, a gas confined (filled) in a cylinder having moving piston. When a gas is

compressed rapidly by moving down the piston, system passes through various states in

which its properties (P, V, T and D) does not remains constant. The gas then would be said

to be in non – equilibrium state.

THERMODYNAMIC PROCESSES

A Thermodynamic Process

“When a thermodynamic system changes from one state to another, the

operation is called as process”

These processes involve the change of condition (Temperature, Pressure & Volume).

The various types of thermodynamic processes are as

1) Isothermal process

2) Adiabatic process

3) Isobaric process

4) Isochoric process

5) Cyclic process,

1) Isothermal process

“The process which occurs at constant temperature is called as isothermal

process”

This is often achieved by placing the system in a thermostat (a constant temperature bath). It

means that during this process, heat is allowed to enter or leave the system so that

temperature remains constant.


Examples

1) When a gas is compressed suddenly, some heat is produced but if the compression is slow

and the heat produced is removed at once, so that the temperature remains constant, the

change is isothermal.

2) Similarly when a gas is allowed to expand suddenly work is done by the gas and some heat

is absorbed. If heat is continuously supplied from outside so that the temperature remains

constant, the change is isothermal.

From the above discussion it follows that in an isothermal change the temp is kept constant

by adding heat or taking it away from the substance.

Therefore, for an isothermal process,

T = 0 …………………………………….. (2)

Where T → Temperature.

2) Adiabatic Process

“When a process is carried out under such condition that not exchange of heat takes

place bet n the system and its surroundings, the process is called as adiabatic process”

Adiabatic condition can be approached by carrying the process in an insulated container such

as ‘thermos’ bottle /Flask. High vacuum and highly polished surfaces help to achieve thermal

insulation.

Example

The compression of the mixture of oil vapour and air during the compression stroke of an

internal combustion engine is an adiabatic process and there a rise of temperature.

If the process is exothermic, the heat evolved will remain in the system, and therefore

the temperature of system rises. If on the other hand, an adiabatic process is endothermic, the

heat required to be absorbed is supplied by the system itself and hence the temperature of the

of the system falls.

Therefore, for an adiabatic process

Q = 0 …………………………………….. (3)

Where Q → heat

3) Isobaric process.

“When a process is carried out at a constant pressure then the process is

called as Isobaric process”


Example

Heating of water to its boiling point and its vaporization take place at the same atmospheric

pressure. These changes are, therefore designated as isobaric processes and are said to take

place isobarically.

From the above example, it is clear that in isobaric process volume may change for a isobaric

processes,

P = 0 …………………………………….. (4)

Where, P → Pressure.

4) Isochoric process –

“A process is defined as isochoric if the volume of the system remains constant

during each step of process”

The heating of a substance in a non - expanding chamber is an example of isochoric process.

Example

Let us consider the reaction between hydrogen and oxygen. On passing an electric spark the

gases will react to form liquid water resulting in considerable non- expanding chamber is

used for the reaction, the volume of system will remain constant and there will be

considerable fall in the pressure system.

From the above example, it is clear that in isothermal process, change in pressure may occur

for isochoric process.

V = 0 …………………………………….. (5)

Where V = Volume.

5) Cyclic process

“When a system in a given state goes through a number of different processes

and finally returns to its initial state, the overall process is called as a cyclic process”

For cyclic process,

H = 0 …………………………………….. (6)

Where, H = enthalpy.

DIFFERENTIATE BETWEEN ISOTHERMAL AND ADIABATIC PROCESS.

Sr. No. Isothermal Process Adiabatic Process

1 In an isothermal process the

temperature of the system remains

constant.

In an adiabatic process, the

temperature of the system may change

according to conditions.


2 In an isothermal process, the

system exchange heat with the

surrounding.

In an adiabatic process the system is

completely insulated from the

surrounding and it does not exchange

heat with the surrounding.

3 Here, T = o Here, Q = o

4 E.g. Slow compression of a gas. E.g. Tea or coffee in a thermos/ Flask.

REVERSIBLE AND IRREVERSIBLE PROCESSES

Reversible Process

“A thermodynamic reversible process is one which takes place infinitesimally slowly

and its direction can be reversed at any point by an infinitesimally small change in the

state of system.”

In fact, a reversible process is considered to proceed from initial to final state through a series

of infinitely small intermediate stages. At the initial, final and all intermediate stages, the

system is in equilibrium state.

Irreversible Process

“A process which goes from initial to final state of a system in a single step and

which cannot be carried out in a reverse order is known as an irreversible process.”

Here the system is in equilibrium in initial and final stage but not at points in between

process.

Naturally occurring all processes are irreversible.

Example

Consider a certain quantity of a gas contained in a cylinder having frictionless piston. The

expansion of gas can be carried out by two methods as shown in fig.

1) Let the pressure applied to the piston be P and this is equal to the internal pressure of the

gas since the external and internal pressures are counterbalanced the piston remains stationary

and.

a) Reversible expansion occurs by

decreasing the pressure on the piston by

infinitesimal a mounts.

B) Irreversible expansion occurs by sudden

decrease of pressure from, P to P when the

gas expands rapidly in a single operation


There is no change in volume of the gas. Now, suppose the pressure on the piston is

decreased by an infinitely small amount dp. Thus, the external pressure on the gas being P –

dp, the piston moves up and gas will expand infinitesimally slowly i.e. by thermodynamically

reversible process. (as shown in fig a) At all stages in the expansion of the gas dp being

negligibly small and the gas is maintained in a state of equilibrium throughout the process.

On the other hand, the expansion will be irreversible if the pressure on the gas is

decreased suddenly (as shown in fig b) and the piston will move upwards rapidly in a single

operation. The gas is in equilibrium state in the initial and final stages only. The expansion

of gas in this case occurs by an irreversible manner.

DIFFERENCE BETWEEN REVERSIBLE AND IRREVERSIBLE PROCESS.

Reversible process Irreversible Process

1 It takes place in infinite number of

small steps and it takes infinite time

to occur.

It takes place in finite time in a single

step.

2 It is imaginary. It is real and can be performed

practically.

3 It is equilibrium state at all stages of

operation.

It is equilibrium state only at initial

and final stages of operation.

4 All changes are reversed when the

process is carried out in reversible

direction.

Changes do not return to the initial

state as the process is irreversible.

5 It is extremely slow It proceeds at a measurable sped or

rapid than reversible process.

6 Work done by a reversible process is

greater than the corresponding

irreversible process.

Work done by an irreversible process

is smaller than the corresponding

reversible process.

7 Natural processes are not reversible. All natural processes are irreversible

8 Ex. Reversible expansion of gas. Ex. Irreversible expansion of gas.


INTERNAL ENERGY

A thermodynamic system containing some quantity of matter has a definite amount of

heat within itself. This energy includes not only the kinetic and potential energies but also

the molecular energies, electrical and nuclear energies.

Definition

“The sum of all possible kinds of energies of a system is known as internal energy of

that system.”

Internal Energy (E) = K.E + P.E + R.E + V.E + T.E

Where,

K.E = Kinetic energy

P.E = Potential energy

R.E = Rotational energy

V.E = Vibrational energy

T.E = Translational energy

Since the value of internal energy of a system depends on the mass of the matter

containing in a system, it is classed as an extensive property.

Internal energy is symbolized by E or U. The SI unit for internal energy is Joule (J)

and other unit of energy is the

1 Calorie = 4.184 Joules.

It is not possible to calculate the absolute value of internal energy of a system. In

thermodynamics we are concerned only with the energy changes when a system changes

from one state to another. If E be the difference of energy of the initial state (Einitial) and the

final state (E final), we can write.

E = E initial – E final

E → + Ve if, E final > E initial


As energy of final state is more than that of initial state, energy is absorbed from the

surrounding hence the reaction / process is endothermic.

E → - Ve if, E initial > E final

As energy of initial state is more than that of final state, energy is evolved or given out to the

surrounding hence the reaction/ process is exothermic.

FIRST LAW OF THERMODYNAMICS

The first law of thermodynamics is infact an application of the broad principle known as the

law of conservation of energy to the thermodynamic system.

STATEMENT

“Energy can neither be created nor be destroyed, but one form of energy can be

converted into another.”

ALTERNATIVE STATEMENT

“The total energy of an isolated system remains constant though, it may change

from one form to another” and the total energy of universe is always remains constant.”

When a system changed from initial state to final state, it undergoes a change in internal

energy from E initial to E final

Thus we can write.

E = E final – E initial

Mathematical statement

This energy change is brought about by the evolution or absorption of heat or by work being

done by the system. Because the total energy of system must remain constant, we can write

the mathematical statement of first law as,

E = q - w

Where, q = Amount of heat supplied to the system.


w = Work done by the system.

VARIOUS FORMS OF FIRST LAW OF THERMODYNAMICS

1. The total energy of the universe is always constant.

2. The net energy change for a closed system is equal to the heat absorbed by the system

minus work done by the system.

3. When a quantity of energy of one form disappears, an equal quantity of energy of another

kind makes its appearance.

4. It is impossible for any machine to do work without consuming energy.

SOME SPECIAL CASES OF FIRST LAW (E = Q - W)

Case I

For a cyclic processes involving isothermal expansion of an ideal gas,

i.e. E = 0. Thus, q = w.

Case II

For an adiabatic process, there is no exchange of heat.

i.e. q = 0. Thus, E = -W.

Case III

For an isochoric process, there is no work of expansion.

i.e. W = 0. Thus, E= q.

Case IV

For an isobaric process, there is no change in pressure i.e. p remains constant

As E= q –w, (since w = p v)

E = q – p v;


SIGNIFICANCE OF FIRST LAW OF THERMODYNAMICS

First law establishes an exact relation between heat and work. i.e. a certain quantity of

heat produces a definite amount of work and vice- versa, from first law, it was confirmed that

work can’t appear without disappearance of heat.

ENTHALPY (H)

When we deal with certain process in an open vessel (constant pressure), it becomes essential

to introduce a new function, in place of internal energy and this function is known as

Enthalpy.

Definition

“Enthalpy of a system is defined as the sum of its internal energy and product of its

pressure and volume.”

It is denoted by H, while change in enthalpy as H.

It can be expressed as,

H = E + PV.

Where, H = Enthalpy,

P = Pressure,

V = Volume.

Like internal energy, it is not possible to measure the absolute value of enthalpy. However a

change in enthalpy ( H) accompanying a process can be measured accurately and is given by

the expression.

H = H product - H reactant

THERE ARE TWO CONDITIONS.

1. When the enthalpy of a product is greater than that of reactant (H Product > HReactant )

then energy will be absorbed from the surrounding i.e. the reaction will be

endothermic and H will be positive ( H= + Ve)


2. When the enthalpy of a reactant is greater than that of product (H Reactant > H Product )

then energy will be given out to the surrounding i.e. the reaction will be exothermal

and H will be negative ( H= - Ve)

EXOTHERMIC AND ENDOTHERMIC REACTION

EXOTHERMIC REACTION

Definition

A reaction which proceeds with the evaluation of heat is called as an exothermic

reaction.

Examples

1) 2 Zn (s) + O 2 (g) → 2 ZnO (s) + 693.8 KJ

2) C (s) + O 2 (g) → Co2 (g) + 393.5 KJ

3) C2H2 (g) + 2 H2 (g) → C2H6 (g) + 314 KJ.

In all the above reaction, beat is evolved and therefore these are examples of

exothermic reaction.

In an exothermal reaction, the total internal energy of the reactants is more than that of

product. Hence E is negative for exothermic reaction. Thus for above reactions, E can

be written as.

2 Zn (s) + O2 (g) → 2 ZnO (s) E = - 693.8 kJ

C (s) + O2 (9) → CO2 (g) E = - 393.5 kJ

C2H2 (g) + 2H2 (g) → C2H4 (g) E = - 314 kJ

In exothermic reaction, the total heat content of reactants is more than that of product and

hence change in enthalpy ( H) is also likely to be negative. Thus the above reactions in terms

of H can be written as

2 Zn (s) + O2 (g) → 2 ZnO (s) H = - 693. 8 KJ

C (s) + O2 (9) → CO2 (g) H = - 393.5 kJ

C2H2 (g) + 2H2 (g) → C2H4 (g) H = - 314 .0 kJ

Illustration of exothermic reaction is shown in fig (A)


ENDOTHERMIC REACTION

Definition

A reaction which proceeds with the absorption of heat is called as an endothermic

reaction.

Examples

2 HgO (g) + 180.4 KJ → 2Hg (l) + O2 (g)

H2O (g) + C(s) + 131.2 KJ → CO (g) + H2 (g)

N2 (g) + O2 (g) + 180.5 kJ → 2 NO (g)

In endothermic reaction, the total heat content of reactant is less than that of product and

hence (change in enthalpy) H and (change in internal energy) E is positive. Thus the

above reactions in terms of H can be written as

2 HgO (g) → 2Hg (l) + O2 H = 180.4 kJ

H2O (g) + C(s) → H2O (g) + C(s) H= 131.2 kJ

N2 (g) + O2 (g) → 2 NO (g) H = 180.5 kJ

Melting of solids and vaporization of liquids are endothermic reaction.

Illustration of endothermic reaction is showoin fig (B)

SIGN CONVERSIONS FOR ENERGY.

Energy Term used Sign

Released Exothermic -

Absorbed Endothermic +

Difference between exothermic and endothermic reaction:

Exothermic Reaction Endothermic Reaction.

1

Heat is supplied or given to the surround ding

Heat is absorbed from the surrounding.

2 The internal energy of the reactant

is more than that of the product.

i.e. E is negative.

The internal energy of the reactant is less

than that of the product.

i.e. E is positive.

3 The enthalpy of the reactant is

more than that of the product.

i.e. H is negatve

The enthalpy of the reactant is less than

that of the product

i.e. H is positive.


4 C(s)+O2(g) → CO2 (g) + 393.5 KJ N2 (g)+ O2 (g) + 180.5 kJ → 2 NO (g)

HEAT OF REACTION (ENTHALPY OF REACTION)

Definition

The heat of reaction is simply the amount of heat absorbed or evolved in a reaction

We also know that the amount of heat absorbed or evolved at constant temperature and

pressure is called as enthalpy.

The amount of heat change during a reaction at constant temperature and pressure is called

as enthalpy change.

If value depends upon the number of moles of reactants which have reacted in a given

chemical reaction to give product. Thus, “Heat of reaction may be defined as the amount

of Heat absorbed or evolved in a chemical reaction when the no. of moles of reactants

will change completely into the product.

Example

The heat change for the reaction of one mole of carbon monoxide with half mole of oxygen

to form one mole of carbon dioxide is 284.5 kJ This means that 284.5 KJ of heat is evolved

during the reaction and is the heat of reaction. It can be represented as,

1. CO(g) + ½ O2 (g) → CO2 (g) H – 284.5 KJ

2. C (s) + O2 (g) → CO2 (g) H – 393.5 KJ

It is very important to note that heat of reaction varies with the change in temperature

therefore; we must mention the temperature at which the reaction is taking places. It is also

convenient for comparison to fix up some temperature as standard or reference. Thus here in

thermodynamics the temperature of 298 is (25℃) under a pressure of and atmosphere has

been fixed as the standard state the heat change accompanying a reaction taking place at 29 k

and one atmosphere pressure is called the standard heat change of standard enthalpy change.

Heat of formation:-

The heat of formation of a compound is defined as the change in enthalpy that takes

place when one mole of compound is formed from its element. It is denoted by H F.

e. g. The Heat of formation of ferrous sulphate and acetylene may be expressed as-

Fe (s) + 5 (s) → Fe S (s) HF = 24.0 kcal

2 C (s) + H2 (9) → S2 H2 (9) HF = 53.1415 cal.


Similarly the reaction between gaseous hydrogen and gaseous chloride to form

gaseous hydrogen chloride is represented by the eqn

H2 (9) + Cl 2 (9) → 2Hcl (9)

H = 44.0kcal.

It may be noted in this case that -44.0 kcal is not the heat of formation of

hydrogen chloride because this amount of heat is evolved when two moles of hydrogen

chloride are formed the heat of formation of hydrogen chloride, there fore. would be

- 44.0 = - 22.0kcal

2

Thus the above equation can be written as ½

(H2)9 + ½ (½ 9) - Hcl (9)

HF = - 22.0 kcal

The standard heat of formation of a compound is defined as the change in enthalpy.

This take place when one mole of compound is elements.

All substances being in their standard states ( 298 KJ 250C) & 1atr pressure

Heat of combustion :-

The heat of combustion of a substance is defined as the change in enthalpy of a

system when one mole of the substance is completely burnt in excess of air or oxygen. It is

denoted by HC

For e. g. The compound is burned in the presence of oxygen in a sealed calorimeter to

convent it completely to carbon dioxide and water. the combustion of methane is written as,

CH4 (9) + 2 O2 (9) → CO2 (9) + 2H 2 0 ( )

HC = - 212.8 kcal

Now consider the chemical equation.

( (S) + O2 (9)→ CO2 (9) H = -943 kcal

( (S) + ½ O2 (9) → �� (�) H = -26.0 kcal

It may noted that -94.3 kcal and not – 26.0 kcal is the heat of combustion of

carbon as the combustion is complete only in first reaction. In the second case, oxidation has

converted carbon to carbon monoxide and is by no means complete as carbon monoxide and

is by no means complete as carbon monoxide can be further oxidised to carbon dioxide

It should be noted clearly that the heat of combustion of a substance ( HC) is

always negative of Heat energy is evolved during the process of combustion

i.e. HC = -ve


HESS’S LAW OF CONSTANT HEAT SUMMATION

We have already seen that the heat changes in a chemical reaction are equal to the

difference in internal energy ( E) or heat content ( H) of the reactants and products

depending upon whether reaction is studied at constant pressure or volume.

Since E and H are state functions, they depends only upon the initial and

final states of the system and not on the manner or steps in which the change takes place.

The generalization is known as Hess’s law.

STATEMENT

“The resultant heat change accompanying (associated with / related to) a

chemical reaction depends on the initial and final states and not on the intermediate

states”

Example

1. Benzene can be converted into m-dinitrobenzene in two different ways,

q1 q2

Benzene Nitrobenzene M-Dinitrobenzene

q3

M- Dinitrobenzene

Here q3 = q1 + q2

2. Burning of carbon to CO2

Carbon can be burnt to CO2 directly or it may first be converted to CO, which may farther be

oxidized to CO2 as shown below-

a) ( (S) + O2 (9) → CO2 (9); H = 94.05 Kcal.

b) ( CS) + ½ O2 (9) → CO (9); H1 = -26.42 Kcal.

CO (9) + ½ O2 (9) →CO2 (9); H2 = -67.71 Kcal.

H = H1 + H2 = - 26.42 + -67.71

= - 94. 13 Kcal .

APPLICATIONS OF HESS’S LAW

1) Determination of heat of various reactions.

2) Determination of heat of transition of one allotropic form to another form.

3) Determination of heat of formation of substances.

4) Determination of lattice energy of a crystal.


Lattice energy

“Lattice energy of a crystal may be defined as amount of heat released when the

numbers of ions in a gas combine to yield n mole of a crystal lattice”.

SECOND LAW OF THERMODYNAMICS

SPONTANEOUS (NATURAL) PROCESS

“A process which proceeds of its own accord, without any external assistance is called as

spontaneous process”. The tendency of a process to occur naturally on its own is called as

spontaneity.

Example –

1. Rolling ball

Uphill

Downhill

A ball rolls down hill spontaneously but it can’t roll up-hill unless work is done on it.

2. Gas law

When a vessel containing a gas is connected to another evacuated vessel, the gas

spreads throughout spontaneously unless the pressure is the same in both the vessels. The

reverse process of compressing the gas in to the original vessel cannot occur unless work is

done on it.

NON – SPONTANEOUS (UN-NATURAL) PROCESS

“A process which does not proceed of its own accord is called as non- spontaneous process”

It is denoted symbolically as ‘s’, while change in entropy as S’

Entropy of a system is a state function, since it depends upon initial and final states of

the system.

Thus,

S = S Final –S intial

If Sfinal > S initial, S is +ve.

Heat is absorbed by the system and process is endothermic.


If S final > S intial , S is –ve

Heat is evolved by the system and process of exothermic.

In short, the entropy is a measure of randomness of a system. The entropy of gaseous

state is greater than its liquid and entropy of liquid is greater than the solid.

Thus, entropy or randomness of substances decreases in the following order:

Gas > liquid > solid.

Numerical definition Of Entropy :

We have discussed the physical defn of entropy. But classical thermodynamics does

not require a physical explanation of concept of entropy. All that we need is an operational

definition so that we can calculate the enthalpy change of the system and surrounding.

Entropy :

For many years scientists believed that only exothermic changes resulting in a

lowering of internal energy or enthalpy could occur spontaneously. But melting of ice is an

endothermic process and yet occurs spontaneously. On a warm day. ice melts by itself. The

evaporation of water is another example of a spontaneous endothermic process. Thus arose

the need of inventing another driving Force that affects the spontaneity. This was known as

the entropy change. S

Spontaneity and Randomness. :

Careful examination shows that in each of the process viz, melting of ice and

evaporation 0f water there is an increase in randomness or disorder of the system. The water

molecules in ice are arranged in a highly organized crystal pattern which permits little

movement. As the ice melts, the water molecules become disorganized and can move more

freely the movement of molecules becomes free still when the water evaporates into space as

now they can room about throughout the entire atmosphere. In other words, we can say that

the randomness of the water molecules increases as ice melts into or water evaporates into

space.

Increase in randomness favors a spontaneous change. :

E.g. : Suit of playing cards arranged numerically highly organized.

When some tossed into air, collect them and restack them →highly disordered.


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CHAPTER 5 THERMODYNAMICS · 4. Thermodynamics is independent of time factor i. e. it does not deals with the rate of physical changes or chemical reactions. Thermodynamic terms and