ncert ch12 physics class 11

CHAPTER TWELVE

THERMODYNAMICS

12.1 INTRODUCTION

In previous chapter we have studied thermal properties ofmatter. In this chapter we shall study laws that governthermal energy. We shall study the processes where work isconverted into heat and vice versa. In winter, when we rubour palms together, we feel warmer; here work done in rubbingproduces the ‘heat’. Conversely, in a steam engine, the ‘heat’of the steam is used to do useful work in moving the pistons,which in turn rotate the wheels of the train.

In physics, we need to define the notions of heat,temperature, work, etc. more carefully. Historically, it took along time to arrive at the proper concept of ‘heat’. Before themodern picture, heat was regarded as a fine invisible fluidfilling in the pores of a substance. On contact between a hotbody and a cold body, the fluid (called caloric) flowed fromthe colder to the hotter body ! This is similar to what happenswhen a horizontal pipe connects two tanks containing waterup to different heights. The flow continues until the levels ofwater in the two tanks are the same. Likewise, in the ‘caloric’picture of heat, heat flows until the ‘caloric levels’ (i.e., thetemperatures) equalise.

In time, the picture of heat as a fluid was discarded infavour of the modern concept of heat as a form of energy. Animportant experiment in this connection was due to BenjaminThomson (also known as Count Rumford) in 1798. Heobserved that boring of a brass cannon generated a lot ofheat, indeed enough to boil water. More significantly, theamount of heat produced depended on the work done (by thehorses employed for turning the drill) but not on thesharpness of the drill. In the caloric picture, a sharper drillwould scoop out more heat fluid from the pores; but thiswas not observed. A most natural explanation of theobservations was that heat was a form of energy and theexperiment demonstrated conversion of energy from one formto another–from work to heat.

12.1 Introduction

12.2 Thermal equilibrium

12.3 Zeroth law ofThermodynamics

12.4 Heat, internal energy andwork

12.5 First law ofthermodynamics

12.6 Specific heat capacity

12.7 Thermodynamic statevariables and equation ofstate

12.8 Thermodynamic processes

12.9 Heat engines

12.10 Refrigerators and heatpumps

12.11 Second law ofthermodynamics

12.12 Reversible and irreversibleprocesses

12.13 Carnot engine

SummaryPoints to ponderExercises

Thermodynamics is the branch of physics thatdeals with the concepts of heat and temperatureand the inter-conversion of heat and other formsof energy. Thermodynamics is a macroscopicscience. It deals with bulk systems and does notgo into the molecular constitution of matter. Infact, its concepts and laws were formulated in thenineteenth century before the molecular pictureof matter was firmly established. Thermodynamicdescription involves relatively few macroscopicvariables of the system, which are suggested bycommon sense and can be usually measureddirectly. A microscopic description of a gas, forexample, would involve specifying the co-ordinatesand velocities of the huge number of moleculesconstituting the gas. The description in kinetictheory of gases is not so detailed but it does involvemolecular distribution of velocities.Thermodynamic description of a gas, on the otherhand, avoids the molecular description altogether.Instead, the state of a gas in thermodynamics isspecified by macroscopic variables such aspressure, volume, temperature, mass andcomposition that are felt by our sense perceptionsand are measurable*.

The distinction between mechanics andthermodynamics is worth bearing in mind. Inmechanics, our interest is in the motion of particlesor bodies under the action of forces and torques.Thermodynamics is not concerned with themotion of the system as a whole. It is concernedwith the internal macroscopic state of the body.When a bullet is fired from a gun, what changesis the mechanical state of the bullet (its kineticenergy, in particular), not its temperature. Whenthe bullet pierces a wood and stops, the kineticenergy of the bullet gets converted into heat,changing the temperature of the bullet and thesurrounding layers of wood. Temperature isrelated to the energy of the internal (disordered)motion of the bullet, not to the motion of the bulletas a whole.

12.2 THERMAL EQUILIBRIUMEquilibrium in mechanics means that the netexternal force and torque on a system are zero.The term ‘equilibrium’ in thermodynamics appears

* Thermodynamics may also involve other variables that are not so obvious to our senses e.g. entropy, enthalpy,etc., and they are all macroscopic variables.

** Both the variables need not change. It depends on the constraints. For instance, if the gases are in containersof fixed volume, only the pressures of the gases would change to achieve thermal equilibrium.

in a different context : we say the state of a systemis an equilibrium state if the macroscopicvariables that characterise the system do notchange in time. For example, a gas inside a closedrigid container, completely insulated from itssurroundings, with fixed values of pressure,volume, temperature, mass and composition thatdo not change with time, is in a state ofthermodynamic equilibrium.

In general, whether or not a system is in a stateof equilibrium depends on the surroundings andthe nature of the wall that separates the systemfrom the surroundings. Consider two gases A andB occupying two different containers. We knowexperimentally that pressure and volume of agiven mass of gas can be chosen to be its twoindependent variables. Let the pressure andvolume of the gases be (P

A, V

A) and (P

B, V

B)

respectively. Suppose first that the two systemsare put in proximity but are separated by anadiabatic wall – an insulating wall (can bemovable) that does not allow flow of energy (heat)from one to another. The systems are insulatedfrom the rest of the surroundings also by similaradiabatic walls. The situation is shownschematically in Fig. 12.1 (a). In this case, it isfound that any possible pair of values (P

A, V

A) will

be in equilibrium with any possible pair of values(P

B, V

B). Next, suppose that the adiabatic wall is

replaced by a diathermic wall – a conducting wallthat allows energy flow (heat) from one to another.It is then found that the macroscopic variables ofthe systems A and B change spontaneously untilboth the systems attain equilibrium states. Afterthat there is no change in their states. Thesituation is shown in Fig. 12.1(b). The pressureand volume variables of the two gases change to(P

B′ , V

B′ ) and (P

A′ , V

A′ ) such that the new states

of A and B are in equilibrium with each other**.There is no more energy flow from one to another.We then say that the system A is in thermalequilibrium with the system B.

What characterises the situation of thermalequilibrium between two systems ? You can guessthe answer from your experience. In thermalequilibrium, the temperatures of the two systems

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are equal. We shall see how does one arrive atthe concept of temperature in thermodynamics?The Zeroth law of thermodynamics provides theclue.

12.3 ZEROTH LAW OF THERMODYNAMICS

Imagine two systems A and B, separated by anadiabatic wall, while each is in contact with a thirdsystem C, via a conducting wall [Fig. 12.2(a)]. Thestates of the systems (i.e., their macroscopicvariables) will change until both A and B come tothermal equilibrium with C. After this is achieved,suppose that the adiabatic wall between A and Bis replaced by a conducting wall and C is insulatedfrom A and B by an adiabatic wall [Fig.12.2(b)]. Itis found that the states of A and B change nofurther i.e. they are found to be in thermalequilibrium with each other. This observationforms the basis of the Zeroth Law ofThermodynamics, which states that ‘twosystems in thermal equilibrium with a thirdsystem separately are in thermal equilibriumwith each other’. R.H. Fowler formulated thislaw in 1931 long after the first and second Lawsof thermodynamics were stated and so numbered.

Fig. 12.1 (a) Systems A and B (two gases) separatedby an adiabatic wall – an insulating wallthat does not allow flow of heat. (b) Thesame systems A and B separated by adiathermic wall – a conducting wall thatallows heat to flow from one to another. Inthis case, thermal equilibrium is attainedin due course.

The Zeroth Law clearly suggests that when twosystems A and B, are in thermal equilibrium,there must be a physical quantity that has thesame value for both. This thermodynamicvariable whose value is equal for two systems inthermal equilibrium is called temperature (T ).Thus, if A and B are separately in equilibriumwith C, TA = TC and TB = TC. This implies thatTA = TB i.e. the systems A and B are also inthermal equilibrium.

We have arrived at the concept of temperatureformally via the Zeroth Law. The next questionis : how to assign numerical values totemperatures of different bodies ? In other words,how do we construct a scale of temperature ?Thermometry deals with this basic question towhich we turn in the next section.

Fig. 12.2 (a) Systems A and B are separated by anadiabatic wall, while each is in contactwith a third system C via a conductingwall. (b) The adiabatic wall between Aand B is replaced by a conducting wall,while C is insulated from A and B by anadiabatic wall.

12.4 HEAT, INTERNAL ENERGY AND WORK

The Zeroth Law of Thermodynamics led us tothe concept of temperature that agrees with ourcommonsense notion. Temperature is a marker

(a)

(b)

(a)

(b)

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of the ‘hotness’ of a body. It determines thedirection of flow of heat when two bodies areplaced in thermal contact. Heat flows from thebody at a higher temperature to the one at lowertemperature. The flow stops when thetemperatures equalise; the two bodies are thenin thermal equilibrium. We saw in some detailhow to construct temperature scales to assigntemperatures to different bodies. We nowdescribe the concepts of heat and other relevantquantities like internal energy and work.

The concept of internal energy of a system isnot difficult to understand. We know that everybulk system consists of a large number ofmolecules. Internal energy is simply the sum ofthe kinetic energies and potential energies ofthese molecules. We remarked earlier that inthermodynamics, the kinetic energy of thesystem, as a whole, is not relevant. Internalenergy is thus, the sum of molecular kinetic andpotential energies in the frame of referencerelative to which the centre of mass of the systemis at rest. Thus, it includes only the (disordered)energy associated with the random motion ofmolecules of the system. We denote the internalenergy of a system by U.

Though we have invoked the molecularpicture to understand the meaning of internalenergy, as far as thermodynamics is concerned,U is simply a macroscopic variable of the system.The important thing about internal energy isthat it depends only on the state of the system,not on how that state was achieved. Internalenergy U of a system is an example of athermodynamic ‘state variable’ – its valuedepends only on the given state of the system,not on history i.e. not on the ‘path’ taken to arriveat that state. Thus, the internal energy of a givenmass of gas depends on its state described byspecific values of pressure, volume andtemperature. It does not depend on how thisstate of the gas came about. Pressure, volume,temperature, and internal energy arethermodynamic state variables of the system(gas) (see section 12.7). If we neglect the smallintermolecular forces in a gas, the internalenergy of a gas is just the sum of kinetic energiesassociated with various random motions of itsmolecules. We will see in the next chapter thatin a gas this motion is not only translational(i.e. motion from one point to another in the

volume of the container); it also includesrotational and vibrational motion of themolecules (Fig. 12.3).

What are the ways of changing internalenergy of a system ? Consider again, forsimplicity, the system to be a certain mass ofgas contained in a cylinder with a movablepiston as shown in Fig. 12.4. Experience showsthere are two ways of changing the state of thegas (and hence its internal energy). One way isto put the cylinder in contact with a body at ahigher temperature than that of the gas. Thetemperature difference will cause a flow ofenergy (heat) from the hotter body to the gas,thus increasing the internal energy of the gas.The other way is to push the piston down i.e. todo work on the system, which again results inincreasing the internal energy of the gas. Ofcourse, both these things could happen in thereverse direction. With surroundings at a lowertemperature, heat would flow from the gas tothe surroundings. Likewise, the gas could pushthe piston up and do work on the surroundings.In short, heat and work are two different modesof altering the state of a thermodynamic systemand changing its internal energy.

The notion of heat should be carefullydistinguished from the notion of internal energy.Heat is certainly energy, but it is the energy intransit. This is not just a play of words. Thedistinction is of basic significance. The state ofa thermodynamic system is characterised by itsinternal energy, not heat. A statement like ‘agas in a given state has a certain amount ofheat’ is as meaningless as the statement that‘a gas in a given state has a certain amountof work’. In contrast, ‘a gas in a given statehas a certain amount of internal energy’ is aperfectly meaningful statement. Similarly, thestatements ‘a certain amount of heat issupplied to the system’ or ‘a certain amountof work was done by the system’ are perfectlymeaningful.

To summarise, heat and work inthermodynamics are not state variables. Theyare modes of energy transfer to a systemresulting in change in its internal energy,which, as already mentioned, is a state variable.

In ordinary language, we often confuse heatwith internal energy. The distinction betweenthem is sometimes ignored in elementary

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physics books. For proper understanding ofthermodynamics, however, the distinction iscrucial.

Fig. 12.3 (a) Internal energy U of a gas is the sumof the kinetic and potential energies of itsmolecules when the box is at rest. Kineticenergy due to various types of motion(translational, rotational, vibrational) is tobe included in U. (b) If the same box ismoving as a whole with some velocity,the kinetic energy of the box is not to beincluded in U.

Fig. 12.4 Heat and work are two distinct modes ofenergy transfer to a system that results inchange in its internal energy. (a) Heat isenergy transfer due to temperaturedifference between the system and thesurroundings. (b) Work is energy transferbrought about by means (e.g. moving thepiston by raising or lowering some weightconnected to it) that do not involve such atemperature difference.

12.5 FIRST LAW OF THERMODYNAMICS

We have seen that the internal energy U of asystem can change through two modes of energy

transfer : heat and work. Let

ΔQ = Heat supplied to the system by thesurroundings

ΔW = Work done by the system on thesurroundings

ΔU = Change in internal energy of the systemThe general principle of conservation of

energy then implies that

ΔQ = ΔU + ΔW (12.1)

i.e. the energy (ΔQ) supplied to the system goesin partly to increase the internal energy of thesystem (ΔU) and the rest in work on theenvironment (ΔW). Equation (12.1) is known asthe First Law of Thermodynamics. It is simplythe general law of conservation of energy appliedto any system in which the energy transfer fromor to the surroundings is taken into account.

Let us put Eq. (12.1) in the alternative form

ΔQ – ΔW = ΔU (12.2)

Now, the system may go from an initial stateto the final state in a number of ways. Forexample, to change the state of a gas from(P1, V1) to (P2, V2), we can first change thevolume of the gas from V1 to V2, keeping itspressure constant i.e. we can first go the state(P1, V2) and then change the pressure of thegas from P1 to P2, keeping volume constant, totake the gas to (P2, V2). Alternatively, we canfirst keep the volume constant and then keepthe pressure constant. Since U is a statevariable, ΔU depends only on the initial andfinal states and not on the path taken by thegas to go from one to the other. However, ΔQand ΔW will, in general, depend on the pathtaken to go from the initial to final states. Fromthe First Law of Thermodynamics, Eq. (12.2),it is clear that the combination ΔQ – ΔW, ishowever, path independent. This shows thatif a system is taken through a process in whichΔU = 0 (for example, isothermal expansion ofan ideal gas, see section 12.8),

ΔQ = ΔW

i.e., heat supplied to the system is used upentirely by the system in doing work on theenvironment.

If the system is a gas in a cylinder with amovable piston, the gas in moving the piston

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does work. Since force is pressure times area,and area times displacement is volume, workdone by the system against a constant pressureP is

ΔW = P ΔV

where ΔV is the change in volume of the gas.Thus, for this case, Eq. (12.1) gives

ΔQ = ΔU + P ΔV (12.3)

As an application of Eq. (12.3), consider thechange in internal energy for 1 g of water whenwe go from its liquid to vapour phase. Themeasured latent heat of water is 2256 J/g. i.e.,for 1 g of water ΔQ = 2256 J. At atmosphericpressure, 1 g of water has a volume 1 cm3 inliquid phase and 1671 cm3 in vapour phase.

Therefore,

ΔW =P (Vg –Vl ) = 1.013 ×105 ×(1670)×10–6 =169.2 J

Equation (12.3) then gives

ΔU = 2256 – 169.2 = 2086.8 J

We see that most of the heat goes to increasethe internal energy of water in transition fromthe liquid to the vapour phase.

12.6 SPECIFIC HEAT CAPACITY

Suppose an amount of heat ΔQ supplied to asubstance changes its temperature from T toT + ΔT. We define heat capacity of a substance(see Chapter 11) to be

T

QS

ΔΔ= (12.4)

We expect ΔQ and, therefore, heat capacity Sto be proportional to the mass of the substance.Further, it could also depend on thetemperature, i.e., a different amount of heat maybe needed for a unit rise in temperature atdifferent temperatures. To define a constantcharacteristic of the substance andindependent of its amount, we divide S by themass of the substance m in kg :

1

S Qs

m m T

(12.5)

s is known as the specific heat capacity of thesubstance. It depends on the nature of thesubstance and its temperature. The unit ofspecific heat capacity is J kg–1 K–1.

If the amount of substance is specified interms of moles μ (instead of mass m in kg ), wecan define heat capacity per mole of thesubstance by

1S QC

T

(12.6)

C is known as molar specific heat capacity ofthe substance. Like s, C is independent of theamount of substance. C depends on the natureof the substance, its temperature and theconditions under which heat is supplied. Theunit of C is J mo1–1 K–1. As we shall see later (inconnection with specific heat capacity of gases),additional conditions may be needed to defineC or s. The idea in defining C is that simplepredictions can be made in regard to molarspecific heat capacities.

Table 12.1 lists measured specific and molarheat capacities of solids at atmospheric pressureand ordinary room temperature.

We will see in Chapter 13 that predictions ofspecific heats of gases generally agree withexperiment. We can use the same law ofequipartition of energy that we use there topredict molar specific heat capacities of solids.Consider a solid of N atoms, each vibratingabout its mean position. An oscillator in onedimension has average energy of 2 × ½ k

BT

= kBT. In three dimensions, the average energy

is 3 kBT. For a mole of a solid, the total energy is

U = 3 kBT × N

A = 3 RT

Now, at constant pressure, ΔQ = ΔU + P ΔV ≅ΔU, since for a solid ΔV is negligible. Therefore,

CQ

T

U

TR= = =

Δ

Δ

Δ

Δ3 (12.7)

Table 12.1 Specific and molar heat capacitiesof some solids at roomtemperature and atmosphericpressure

As Table 12.1 shows, the experimentallymeasured values which generally agrees with

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predicted value 3R at ordinary temperatures.(Carbon is an exception.) The agreement isknown to break down at low temperatures.

Specific heat capacity of water

The old unit of heat was calorie. One caloriewas earlier defined to be the amount of heatrequired to raise the temperature of 1g of waterby 1°C. With more precise measurements, it wasfound that the specific heat of water variesslightly with temperature. Figure 12.5 showsthis variation in the temperature range 0 to100 °C.

Fig. 12.5 Variation of specific heat capacity of waterwith temperature.

For a precise definition of calorie, it was,therefore, necessary to specify the unittemperature interval. One calorie is definedto be the amount of heat required to raise thetemperature of 1g of water from 14.5 °C to15.5 °C. Since heat is just a form of energy,it is preferable to use the unit joule, J.In SI units, the specific heat capacity of wateris 4186 J kg–1 K–1 i.e. 4.186 J g–1 K–1. The socalled mechanical equivalent of heat definedas the amount of work needed to produce1 cal of heat is in fact just a conversion factorbetween two different units of energy : calorieto joule. Since in SI units, we use the unit joulefor heat, work or any other form of energy, theterm mechanical equivalent is nowsuperfluous and need not be used.

As already remarked, the specific heatcapacity depends on the process or theconditions under which heat capacity transfertakes place. For gases, for example, we candefine two specific heats : specific heatcapacity at constant volume and specificheat capacity at constant pressure. For an

ideal gas, we have a simple relation.

Cp – C

v = R (12.8)

where Cp and C

v are molar specific heat

capacities of an ideal gas at constant pressureand volume respectively and R is the universalgas constant. To prove the relation, we beginwith Eq. (12.3) for 1 mole of the gas :

ΔQ = ΔU + P ΔV

If ΔQ is absorbed at constant volume, ΔV = 0

vv v

Q U UC

T T T (12.9)

where the subscript v is dropped in the laststep, since U of an ideal gas depends only ontemperature. (The subscript denotes thequantity kept fixed.) If, on the other hand, ΔQis absorbed at constant pressure,

pp p p

Q U V

C PT T T

(12.10)

The subscript p can be dropped from thefirst term since U of an ideal gas depends onlyon T. Now, for a mole of an ideal gas

PV = RT

which gives

p

V

P RT

(12.11)

Equations (12.9) to (12.11) give the desiredrelation, Eq. (12.8).

12.7 THERMODYNAMIC STATE VARIABLESAND EQUATION OF STATE

Every equilibrium state of a thermodynamicsystem is completely described by specificvalues of some macroscopic variables, alsocalled state variables. For example, anequilibrium state of a gas is completelyspecified by the values of pressure, volume,temperature, and mass (and composition ifthere is a mixture of gases). A thermodynamicsystem is not always in equilibrium. For example,a gas allowed to expand freely against vacuumis not an equilibrium state [Fig. 12.6(a)]. Duringthe rapid expansion, pressure of the gas may

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not be uniform throughout. Similarly, a mixtureof gases undergoing an explosive chemicalreaction (e.g. a mixture of petrol vapour andair when ignited by a spark) is not anequilibrium state; again its temperature andpressure are not uniform [Fig. 12.6(b)].Eventually, the gas attains a uniformtemperature and pressure and comes tothermal and mechanical equilibrium with itssurroundings.

Fig. 12.6 (a) The partition in the box is suddenlyremoved leading to free expansion of thegas. (b) A mixture of gases undergoing anexplosive chemical reaction. In bothsituations, the gas is not in equilibrium andcannot be described by state variables.

In short, thermodynamic state variablesdescribe equilibrium states of systems. Thevarious state variables are not necessarilyindependent. The connection between the statevariables is called the equation of state. Forexample, for an ideal gas, the equation of stateis the ideal gas relation

P V = μ R T

For a fixed amount of the gas i.e. given μ, thereare thus, only two independent variables, say Pand V or T and V. The pressure-volume curvefor a fixed temperature is called an isotherm.Real gases may have more complicatedequations of state.

The thermodynamic state variables are of twokinds: extensive and intensive. Extensivevariables indicate the ‘size’ of the system.Intensive variables such as pressure and

temperature do not. To decide which variable isextensive and which intensive, think of arelevant system in equilibrium, and imagine thatit is divided into two equal parts. The variablesthat remain unchanged for each part areintensive. The variables whose values get halvedin each part are extensive. It is easily seen, forexample, that internal energy U, volume V, totalmass M are extensive variables. Pressure P,temperature T, and density ρ are intensivevariables. It is a good practice to check theconsistency of thermodynamic equations usingthis classification of variables. For example, inthe equation

ΔQ = ΔU + P ΔV

quantities on both sides are extensive*. (Theproduct of an intensive variable like P and anextensive quantity ΔV is extensive.)

12.8 THERMODYNAMIC PROCESSES

12.8.1 Quasi-static processConsider a gas in thermal and mechanicalequilibrium with its surroundings. The pressureof the gas in that case equals the externalpressure and its temperature is the same asthat of its surroundings. Suppose that theexternal pressure is suddenly reduced (say bylifting the weight on the movable piston in thecontainer). The piston will accelerate outward.During the process, the gas passes throughstates that are not equilibrium states. The non-equilibrium states do not have well-definedpressure and temperature. In the same way, ifa finite temperature difference exists betweenthe gas and its surroundings, there will be arapid exchange of heat during which the gaswill pass through non-equilibrium states. Indue course, the gas will settle to an equilibriumstate with well-defined temperature andpressure equal to those of the surroundings. Thefree expansion of a gas in vacuum and a mixtureof gases undergoing an explosive chemicalreaction, mentioned in section 12.7 are alsoexamples where the system goes through non-equilibrium states.

Non-equilibrium states of a system are difficultto deal with. It is, therefore, convenient toimagine an idealised process in which at everystage the system is an equilibrium state. Such a

* As emphasised earlier, Q is not a state variable. However, ΔQ is clearly proportional to the total mass ofsystem and hence is extensive.

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process is, in principle, infinitely slow-hence thename quasi-static (meaning nearly static). Thesystem changes its variables (P, T, V ) so slowlythat it remains in thermal and mechanicalequilibrium with its surroundings throughout.In a quasi-static process, at every stage, thedifference in the pressure of the system and theexternal pressure is infinitesimally small. Thesame is true of the temperature differencebetween the system and its surroundings. Totake a gas from the state (P, T ) to another state(P ′ , T ′ ) via a quasi-static process, we changethe external pressure by a very small amount,allow the system to equalise its pressure withthat of the surroundings and continue theprocess infinitely slowly until the systemachieves the pressure P ′ . Similarly, to changethe temperature, we introduce an infinitesimaltemperature difference between the system andthe surrounding reservoirs and by choosingreservoirs of progressively different temperaturesT to T ′ , the system achieves the temperature T ′ .

Fig. 12.7 In a quasi-static process, the temperatureof the surrounding reservoir and theexternal pressure differ only infinitesimallyfrom the temperature and pressure of thesystem.

A quasi-static process is obviously ahypothetical construct. In practice, processesthat are sufficiently slow and do not involveaccelerated motion of the piston, largetemperature gradient, etc. are reasonablyapproximation to an ideal quasi-static process.We shall from now on deal with quasi-staticprocesses only, except when stated otherwise.

A process in which the temperature of thesystem is kept fixed throughout is called anisothermal process. The expansion of a gas ina metallic cylinder placed in a large reservoir offixed temperature is an example of an isothermalprocess. (Heat transferred from the reservoir tothe system does not materially affect thetemperature of the reservoir, because of its verylarge heat capacity.) In isobaric processes thepressure is constant while in isochoricprocesses the volume is constant. Finally, ifthe system is insulated from the surroundingsand no heat flows between the system and thesurroundings, the process is adiabatic. Thedefinitions of these special processes aresummarised in Table. 12.2

Table 12.2 Some special thermodynamicprocesses

We now consider these processes in some detail :

Isothermal process

For an isothermal process (T fixed), the ideal gasequation gives PV = constanti.e., pressure of a given mass of gas varies inverselyas its volume. This is nothing but Boyle’s Law.

Suppose an ideal gas goes isothermally (attemperature T ) from its initial state (P1, V1) tothe final state (P2, V 2). At any intermediate stagewith pressure P and volume change from V toV + ΔV (ΔV small)

ΔW = P Δ V

Taking (ΔV → 0) and summing the quantityΔW over the entire process,

d 2

1

VW = P V

V

Ind2

2

11

V V= RT RT

VV

VV

(12.12)

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where in the second step we have made use ofthe ideal gas equation PV = μ RT and taken theconstants out of the integral. For an ideal gas,internal energy depends only on temperature.Thus, there is no change in the internal energyof an ideal gas in an isothermal process. TheFirst Law of Thermodynamics then implies thatheat supplied to the gas equals the work doneby the gas : Q = W. Note from Eq. (12.12) thatfor V2 > V1, W > 0; and for V2 < V1, W < 0. Thatis, in an isothermal expansion, the gas absorbsheat and does work while in an isothermalcompression, work is done on the gas by theenvironment and heat is released.

Adiabatic process

In an adiabatic process, the system is insulatedfrom the surroundings and heat absorbed orreleased is zero. From Eq. (12.1), we see thatwork done by the gas results in decrease in itsinternal energy (and hence its temperature foran ideal gas). We quote without proof (the resultthat you will learn in higher courses) that foran adiabatic process of an ideal gas.

P V γ = const (12.13)

where γ is the ratio of specific heats (ordinaryor molar) at constant pressure and at constantvolume.

γ =CpCv

Thus if an ideal gas undergoes a change inits state adiabatically from (P1, V1) to (P2, V2) :

P1 V1γ = P2 V2

γ (12.14)

Figure12.8 shows the P-V curves of an idealgas for two adiabatic processes connecting twoisotherms.

Fig. 12.8 P-V curves for isothermal and adiabaticprocesses of an ideal gas.

We can calculate, as before, the work done inan adiabatic change of an ideal gas from thestate (P1, V1, T1) to the state (P2, V2, T2).

2

1

d V

W = P VV

2 2

11

d

1

– +1V VV V= constant constant

VV

V

=⎥⎥⎦

⎤

⎢⎢⎣

⎡×

1–1

1–2

1 –

1

)– (1tan

γγγ VV

tcons(12.15)

From Eq. (12.34), the constant is P1V1γ or P2V2

γ

2 2 1 11 1

2 1

1

1 –

P V P VW =

V V

1 22 2 1 1

)1

1 – –

R(T – T= P V P V

(12.16)

As expected, if work is done by the gas in anadiabatic process (W > 0), from Eq. (12.16),T2 < T1. On the other hand, if work is done onthe gas (W < 0), we get T2 > T1 i.e., thetemperature of the gas rises.

Isochoric process

In an isochoric process, V is constant. No workis done on or by the gas. From Eq. (12.1), theheat absorbed by the gas goes entirely to changeits internal energy and its temperature. Thechange in temperature for a given amount ofheat is determined by the specific heat of thegas at constant volume.

Isobaric process

In an isobaric process, P is fixed. Work done bythe gas is

W = P (V2 – V1) = μ R (T2 – T1) (12.17)

Since temperature changes, so does internalenergy. The heat absorbed goes partly toincrease internal energy and partly to do work.The change in temperature for a given amountof heat is determined by the specific heat of thegas at constant pressure.

Cyclic process

In a cyclic process, the system returns to itsinitial state. Since internal energy is a statevariable, ΔU = 0 for a cyclic process. From

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Eq. (12.1), the total heat absorbed equals thework done by the system.

12.9 HEAT ENGINES

Heat engine is a device by which a system ismade to undergo a cyclic process that resultsin conversion of heat to work.(1) It consists of a working substance–the

system. For example, a mixture of fuelvapour and air in a gasoline or diesel engineor steam in a steam engine are the workingsubstances.

(2) The working substance goes through a cycleconsisting of several processes. In some ofthese processes, it absorbs a total amountof heat Q1 from an external reservoir at somehigh temperature T1.

(3) In some other processes of the cycle, theworking substance releases a total amountof heat Q2 to an external reservoir at somelower temperature T2.

(4) The work done (W ) by the system in a cycleis transferred to the environment via somearrangement (e.g. the working substancemay be in a cylinder with a moving pistonthat transfers mechanical energy to thewheels of a vehicle via a shaft).

The basic features of a heat engine areschematically represented in Fig. 12.9.

Fig. 12.9 Schematic representation of a heat engine.The engine takes heat Q1 from a hotreservoir at temperature T1, releases heatQ2 to a cold reservoir at temperature T2and delivers work W to the surroundings.

The cycle is repeated again and again to getuseful work for some purpose. The discipline ofthermodynamics has its roots in the study of heatengines. A basic question relates to the efficiencyof a heat engine. The efficiency (η) of a heatengine is defined by

1QW = η (12.18)

where Q1 is the heat input i.e., the heatabsorbed by the system in one complete cycle

and W is the work done on the environment ina cycle. In a cycle, a certain amount of heat (Q2)may also be rejected to the environment. Then,according to the First Law of Thermodynamics,over one complete cycle,

W = Q1 – Q2 (12.19)

i.e.,

1Q2Q

− = 1η (12.20)

For Q2 = 0, η = 1, i.e., the engine will have100% efficiency in converting heat into work.Note that the First Law of Thermodynamics i.e.,the energy conservation law does not rule outsuch an engine. But experience shows thatsuch an ideal engine with η = 1 is never possible,even if we can eliminate various kinds of lossesassociated with actual heat engines. It turnsout that there is a fundamental limit on theefficiency of a heat engine set by an independentprinciple of nature, called the Second Law ofThermodynamics (section 12.11).

The mechanism of conversion of heat intowork varies for different heat engines. Basically,there are two ways : the system (say a gas or amixture of gases) is heated by an externalfurnace, as in a steam engine; or it is heatedinternally by an exothermic chemical reactionas in an internal combustion engine. Thevarious steps involved in a cycle also differ fromone engine to another. For the purpose of generalanalysis, it is useful to conceptualise a heatengine as having the following essentialingredients.

12.10 REFRIGERATORS AND HEAT PUMPS

A refrigerator is the reverse of a heat engine.Here the working substance extracts heat Q2

from the cold reservoir at temperature T2, someexternal work W is done on it and heat Q1 isreleased to the hot reservoir at temperature T1

(Fig. 12.10).

Fig. 12.10 Schematic representation of a refrigeratoror a heat pump, the reverse of a heatengine.

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A heat pump is the same as a refrigerator.What term we use depends on the purpose ofthe device. If the purpose is to cool a portion ofspace, like the inside of a chamber, and highertemperature reservoir is surrounding, we callthe device a refrigerator; if the idea is to pumpheat into a portion of space (the room in abuilding when the outside environment is cold),the device is called a heat pump.

In a refrigerator the working substance(usually, in gaseous form) goes through thefollowing steps : (a) sudden expansion of the gasfrom high to low pressure which cools it andconverts it into a vapour-liquid mixture, (b)absorption by the cold fluid of heat from theregion to be cooled converting it into vapour, (c)heating up of the vapour due to external workdone on the system, and (d) release of heat bythe vapour to the surroundings, bringing it tothe initial state and completing the cycle.

The coefficient of performance (α) of arefrigerator is given by

2Q

W (12.21)

where Q2 is the heat extracted from the coldreservoir and W is the work done on thesystem–the refrigerant. (α for heat pump isdefined as Q1/W) Note that while η by definitioncan never exceed 1, α can be greater than 1.By energy conservation, the heat released to thehot reservoir is

Q1 = W + Q2

i.e.,2

1 2–Q

Q Q

(12.22)

In a heat engine, heat cannot be fullyconverted to work; likewise a refrigerator cannotwork without some external work done on thesystem, i.e., the coefficient of performance in Eq.(12.21) cannot be infinite.

12.11 SECOND LAW OF THERMODYNAMICS

The First Law of Thermodynamics is the principleof conservation of energy. Common experienceshows that there are many conceivableprocesses that are perfectly allowed by the FirstLaw and yet are never observed. For example,nobody has ever seen a book lying on a tablejumping to a height by itself. But such a thing

Pioneers of Thermodynamics

Lord Kelvin (William Thomson) (1824-1907), born in Belfast, Ireland, isamong the foremost British scientists of the nineteenth century. Thomsonplayed a key role in the development of the law of conservation of energysuggested by the work of James Joule (1818-1889), Julius Mayer (1814-1878) and Hermann Helmholtz (1821-1894). He collaborated with Jouleon the so-called Joule-Thomson effect : cooling of a gas when it expandsinto vacuum. He introduced the notion of the absolute zero of temperatureand proposed the absolute temperature scale, now called the Kelvin scalein his honour. From the work of Sadi Carnot (1796-1832), Thomson arrivedat a form of the Second Law of Thermodynamics. Thomson was a versatilephysicist, with notable contributions to electromagnetic theory andhydrodynamics.

Rudolf Clausius (1822-1888), born in Poland, is generally regarded asthe discoverer of the Second Law of Thermodynamics. Based on the workof Carnot and Thomson, Clausius arrived at the important notion of entropythat led him to a fundamental version of the Second Law ofThermodynamics that states that the entropy of an isolated system cannever decrease. Clausius also worked on the kinetic theory of gases andobtained the first reliable estimates of molecular size, speed, mean freepath, etc.

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would be possible if the principle of conservationof energy were the only restriction. The tablecould cool spontaneously, converting some of itsinternal energy into an equal amount ofmechanical energy of the book, which wouldthen hop to a height with potential energy equalto the mechanical energy it acquired. But thisnever happens. Clearly, some additional basicprinciple of nature forbids the above, eventhough it satisfies the energy conservationprinciple. This principle, which disallows manyphenomena consistent with the First Law ofThermodynamics is known as the Second Lawof Thermodynamics.

The Second Law of Thermodynamics gives afundamental limitation to the efficiency of a heatengine and the co-efficient of performance of arefrigerator. In simple terms, it says thatefficiency of a heat engine can never be unity.According to Eq. (12.20), this implies that heatreleased to the cold reservoir can never be madezero. For a refrigerator, the Second Law says thatthe co-efficient of performance can never beinfinite. According to Eq. (12.21), this impliesthat external work (W ) can never be zero. Thefollowing two statements, one due to Kelvin andPlanck denying the possibility of a perfect heatengine, and another due to Clausius denyingthe possibility of a perfect refrigerator or heatpump, are a concise summary of theseobservations.

Second Law of Thermodynamics

Kelvin-Planck statementNo process is possible whose sole result is theabsorption of heat from a reservoir and thecomplete conversion of the heat into work.

Clausius statement

No process is possible whose sole result is thetransfer of heat from a colder object to a hotterobject.

It can be proved that the two statementsabove are completely equivalent.

12.12 REVERSIBLE AND IRREVERSIBLEPROCESSES

Imagine some process in which a thermodynamicsystem goes from an initial state i to a finalstate f. During the process the system absorbsheat Q from the surroundings and performswork W on it. Can we reverse this process and

bring both the system and surroundings to theirinitial states with no other effect anywhere ?Experience suggests that for most processes innature this is not possible. The spontaneousprocesses of nature are irreversible. Severalexamples can be cited. The base of a vessel onan oven is hotter than its other parts. Whenthe vessel is removed, heat is transferred fromthe base to the other parts, bringing the vesselto a uniform temperature (which in due coursecools to the temperature of the surroundings).The process cannot be reversed; a part of thevessel will not get cooler spontaneously andwarm up the base. It will violate the Second Lawof Thermodynamics, if it did. The free expansionof a gas is irreversible. The combustion reactionof a mixture of petrol and air ignited by a sparkcannot be reversed. Cooking gas leaking from agas cylinder in the kitchen diffuses to theentire room. The diffusion process will notspontaneously reverse and bring the gas backto the cylinder. The stirring of a liquid in thermalcontact with a reservoir will convert the workdone into heat, increasing the internal energyof the reservoir. The process cannot be reversedexactly; otherwise it would amount to conversionof heat entirely into work, violating the SecondLaw of Thermodynamics. Irreversibility is a rulerather an exception in nature.

Irreversibility arises mainly from two causes:one, many processes (like a free expansion, oran explosive chemical reaction) take the systemto non-equilibrium states; two, most processesinvolve friction, viscosity and other dissipativeeffects (e.g., a moving body coming to a stop andlosing its mechanical energy as heat to the floorand the body; a rotating blade in a liquid comingto a stop due to viscosity and losing itsmechanical energy with corresponding gain inthe internal energy of the liquid). Sincedissipative effects are present everywhere andcan be minimised but not fully eliminated, mostprocesses that we deal with are irreversible.

A thermodynamic process (state i → state f )is reversible if the process can be turned backsuch that both the system and the surroundingsreturn to their original states, with no otherchange anywhere else in the universe. From thepreceding discussion, a reversible process is anidealised notion. A process is reversible only ifit is quasi-static (system in equilibrium with the

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surroundings at every stage) and there are nodissipative effects. For example, a quasi-staticisothermal expansion of an ideal gas in acylinder fitted with a frictionless movable pistonis a reversible process.

Why is reversibility such a basic concept inthermodynamics ? As we have seen, one of theconcerns of thermodynamics is the efficiencywith which heat can be converted into work.The Second Law of Thermodynamics rules outthe possibility of a perfect heat engine with 100%efficiency. But what is the highest efficiencypossible for a heat engine working between tworeservoirs at temperatures T1 and T2 ? It turnsout that a heat engine based on idealisedreversible processes achieves the highestefficiency possible. All other engines involvingirreversibility in any way (as would be the casefor practical engines) have lower than thislimiting efficiency.

12.13 CARNOT ENGINE

Suppose we have a hot reservoir at temperatureT1 and a cold reservoir at temperature T2. Whatis the maximum efficiency possible for a heatengine operating between the two reservoirs andwhat cycle of processes should be adopted toachieve the maximum efficiency ? Sadi Carnot,a French engineer, first considered this questionin 1824. Interestingly, Carnot arrived at thecorrect answer, even though the basic conceptsof heat and thermodynamics had yet to be firmlyestablished.

We expect the ideal engine operating betweentwo temperatures to be a reversible engine.Irreversibility is associated with dissipativeeffects, as remarked in the preceding section,and lowers efficiency. A process is reversible ifit is quasi-static and non-dissipative. We haveseen that a process is not quasi-static if itinvolves finite temperature difference betweenthe system and the reservoir. This implies thatin a reversible heat engine operating betweentwo temperatures, heat should be absorbed(from the hot reservoir) isothermally andreleased (to the cold reservoir) isothermally. Wethus have identified two steps of the reversibleheat engine : isothermal process at temperatureT1 absorbing heat Q1 from the hot reservoir, andanother isothermal process at temperature T2

releasing heat Q2 to the cold reservoir. To

complete a cycle, we need to take the systemfrom temperature T1 to T2 and then back fromtemperature T2 to T1. Which processes shouldwe employ for this purpose that are reversible?A little reflection shows that we can only adoptreversible adiabatic processes for thesepurposes, which involve no heat flow from anyreservoir. If we employ any other process that isnot adiabatic, say an isochoric process, to takethe system from one temperature to another, weshall need a series of reservoirs in thetemperature range T2 to T1 to ensure that at eachstage the process is quasi-static. (Rememberagain that for a process to be quasi-static andreversible, there should be no finite temperaturedifference between the system and the reservoir.)But we are considering a reversible engine thatoperates between only two temperatures. Thusadiabatic processes must bring about thetemperature change in the system from T1 to T2

and T2 to T1 in this engine.

Fig. 12.11 Carnot cycle for a heat engine with anideal gas as the working substance.

A reversible heat engine operating betweentwo temperatures is called a Carnot engine. Wehave just argued that such an engine must havethe following sequence of steps constituting onecycle, called the Carnot cycle, shown in Fig.12.11. We have taken the working substance ofthe Carnot engine to be an ideal gas.

(a) Step 1 → 2 Isothermal expansion of the gastaking its state from (P1, V1, T1) to(P2, V2, T1).

The heat absorbed by the gas (Q1) from thereservoir at temperature T1 is given by

PHYSICS312

Eq. (12.12). This is also the work done (W1 → 2)by the gas on the environment.

W1 → 2 = Q1 = μ R T1 ln ⎟⎠⎞⎜

⎝⎛

1

2VV

(12.23)

(b) Step 2 → 3 Adiabatic expansion of the gasfrom (P2, V2, T1) to (P3, V3, T2)Work done by the gas, usingEq. (12.16), is

1 2

12 3

R T TW

(12.24)

(c) Step 3 → 4 Isothermal compression of thegas from (P3, V3, T2) to (P4, V4, T2).

Heat released (Q2) by the gas to the reservoirat temperature T2 is given by Eq. (12.12). Thisis also the work done (W3 → 4) on the gas by theenvironment.

32 2

4

ln3 4

VW Q RT

V

(12.25)

(d) Step 4 → 1 Adiabatic compression of thegas from (P4, V4, T2) to (P1,V1, T1).

Work done on the gas, [using Eq.(12.16)], is

1 2

-14 1

T TW R

(12.26)

From Eqs. (12.23) to (12.26) total work doneby the gas in one complete cycle is

W = W1 → 2 + W2 → 3 – W3 → 4 – W4 → 1

= μ RT1 ln ⎟⎠⎞⎜

⎝⎛

1

2VV

– μ RT2 ln ⎟⎠⎞⎜

⎝⎛

4

3VV

(12.27)

The efficiency η of the Carnot engine is

1 2

1 1

W Q

Q Q

In

1

In

3

42

1 2

1

V

VTT V

V

(12.28)

Now since step 2 → 3 is an adiabatic process,

11

1 2 2 3T V T V

i.e.

1/( 1)

2 2

3 1

V T=

V T

(12.29)

Similarly, since step 4 → 1 is an adiabaticprocess

1 1

2 4 1 1T V T V

i.e.

1/ 1

1 2

4 1

V T=

V T

(12.30)

From Eqs. (12.29) and (12.30),

1

2

4

3VV

= VV

(12.31)

Using Eq. (12.31) in Eq. (12.28), we get

1 2

1

T

T - (Carnot engine) (12.32)

We have already seen that a Carnot engineis a reversible engine. Indeed it is the onlyreversible engine possible that works betweentwo reservoirs at different temperatures. Eachstep of the Carnot cycle given in Fig. 12.11 canbe reversed. This will amount to taking heat Q2

from the cold reservoir at T2, doing work W onthe system, and transferring heat Q1 to the hotreservoir. This will be a reversible refrigerator.

We next establish the important result(sometimes called Carnot’s theorem) that(a) working between two given temperatures T1

and T2 of the hot and cold reservoirs respectively,no engine can have efficiency more than that ofthe Carnot engine and (b) the efficiency of theCarnot engine is independent of the nature ofthe working substance.

To prove the result (a), imagine a reversible(Carnot) engine R and an irreversible engine Iworking between the same source (hot reservoir)and sink (cold reservoir). Let us couple theengines, I and R, in such a way so that I actslike a heat engine and R acts as a refrigerator.Let I absorb heat Q1 from the source, deliverwork W ′ and release the heat Q1- W′ to the sink.We arrange so that R returns the same heat Q1

to the source, taking heat Q2 from the sink andrequiring work W = Q1 – Q2 to be done on it.

THERMODYNAMICS 313

SUMMARY

1. The zeroth law of thermodynamics states that ‘two systems in thermal equilibrium with athird system are in thermal equilibrium with each other’. The Zeroth Law leads to theconcept of temperature.

2. Internal energy of a system is the sum of kinetic energies and potential energies of themolecular constituents of the system. It does not include the over-all kinetic energy ofthe system. Heat and work are two modes of energy transfer to the system. Heat is theenergy transfer arising due to temperature difference between the system and thesurroundings. Work is energy transfer brought about by other means, such as movingthe piston of a cylinder containing the gas, by raising or lowering some weight connectedto it.

3. The first law of thermodynamics is the general law of conservation of energy applied toany system in which energy transfer from or to the surroundings (through heat andwork) is taken into account. It states that

ΔQ = ΔU + ΔW

where ΔQ is the heat supplied to the system, ΔW is the work done by the system and ΔUis the change in internal energy of the system.

Now suppose ηR < ηI i.e. if R were to actas an engine it would give less work outputthan that of I i.e. W < W ′ for a given Q1. With Racting like a refrigerator, this would meanQ2 = Q1 – W > Q1 – W ′ . Thus on the whole,the coupled I-R system extracts heat(Q1 – W) – (Q1 – W ′ ) = (W ′ – W ) from the coldreservoir and delivers the same amount of workin one cycle, without any change in the sourceor anywhere else. This is clearly against theKelvin-Planck statement of the Second Law ofThermodynamics. Hence the assertion ηI > ηR

is wrong. No engine can have efficiency greater

than that of the Carnot engine. A similarargument can be constructed to show that areversible engine with one particular substancecannot be more efficient than the one usinganother substance. The maximum efficiency ofa Carnot engine given by Eq. (12.32) isindependent of the nature of the systemperforming the Carnot cycle of operations. Thuswe are justified in using an ideal gas as a systemin the calculation of efficiency η of a Carnotengine. The ideal gas has a simple equation ofstate, which allows us to readily calculate η, butthe final result for η, [Eq. (12.32)], is true forany Carnot engine.

This final remark shows that in a Carnotcycle,

2

1

2

1TT

= QQ

(12.33)

is a universal relation independent of the natureof the system. Here Q1 and Q2 are respectively,the heat absorbed and released isothermally(from the hot and to the cold reservoirs) in aCarnot engine. Equation (12.33), can, therefore,be used as a relation to define a truly universalthermodynamic temperature scale that isindependent of any particular properties of thesystem used in the Carnot cycle. Of course, foran ideal gas as a working substance, thisuniversal temperature is the same as the idealgas temperature introduced in section 12.11.

I

R

W

Fig. 12.12 An irreversible engine (I) coupled to areversible refrigerator (R). If W ′ > W, thiswould amount to extraction of heatW ′ – W from the sink and its fullconversion to work, in contradiction withthe Second Law of Thermodynamics.

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4. The specific heat capacity of a substance is defined by

sm

Q

T=

1 ΔΔ

where m is the mass of the substance and ΔQ is the heat required to change itstemperature by ΔT. The molar specific heat capacity of a substance is defined by

1 QC

T

where μ is the number of moles of the substance. For a solid, the law of equipartitionof energy gives

C = 3 R

which generally agrees with experiment at ordinary temperatures.

Calorie is the old unit of heat. 1 calorie is the amount of heat required to raise thetemperature of 1 g of water from 14.5 °C to 15.5 °C. 1 cal = 4.186 J.

5. For an ideal gas, the molar specific heat capacities at constant pressure and volumesatisfy the relation

Cp – Cv = R

where R is the universal gas constant.

6. Equilibrium states of a thermodynamic system are described by state variables. Thevalue of a state variable depends only on the particular state, not on the path used toarrive at that state. Examples of state variables are pressure (P ), volume (V ), temperature(T ), and mass (m ). Heat and work are not state variables. An Equation of State (likethe ideal gas equation PV = μ RT ) is a relation connecting different state variables.

7. A quasi-static process is an infinitely slow process such that the system remains inthermal and mechanical equilibrium with the surroundings throughout. In aquasi-static process, the pressure and temperature of the environment can differ fromthose of the system only infinitesimally.

8. In an isothermal expansion of an ideal gas from volume V1 to V

2 at temperature T the

heat absorbed (Q) equals the work done (W ) by the gas, each given by

Q = W = μ R T ln ⎟⎟⎠⎞

⎜⎜⎝⎛

1

2VV

9. In an adiabatic process of an ideal gas

PVγ = constant

wherep

v

C

C

Work done by an ideal gas in an adiabatic change of state from (P1, V1, T1) to (P2, V2, T2)is

– 11 2 R T T

W

10. Heat engine is a device in which a system undergoes a cyclic process resulting inconversion of heat into work. If Q1 is the heat absorbed from the source, Q2 is the heatreleased to the sink, and the work output in one cycle is W, the efficiency η of the engineis:

1 2

1 1

QW

Q Q

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11. In a refrigerator or a heat pump, the system extracts heat Q2 from the cold reservoir and

releases Q1 amount of heat to the hot reservoir, with work W done on the system. The

co-efficient of performance of a refrigerator is given by

21

22

QQ

Q

W

Q

−== α

12. The second law of thermodynamics disallows some processes consistent with the FirstLaw of Thermodynamics. It states

Kelvin-Planck statement

No process is possible whose sole result is the absorption of heat from a reservoir andcomplete conversion of the heat into work.

Clausius statement

No process is possible whose sole result is the transfer of heat from a colder object to ahotter object.

Put simply, the Second Law implies that no heat engine can have efficiency η equal to1 or no refrigerator can have co-efficient of performance α equal to infinity.

13. A process is reversible if it can be reversed such that both the system and the surroundingsreturn to their original states, with no other change anywhere else in the universe.Spontaneous processes of nature are irreversible. The idealised reversible process is aquasi-static process with no dissipative factors such as friction, viscosity, etc.

14. Carnot engine is a reversible engine operating between two temperatures T1 (source) and

T2 (sink). The Carnot cycle consists of two isothermal processes connected by two

adiabatic processes. The efficiency of a Carnot engine is given by

1

2

T

T 1 −= η (Carnot engine)

No engine operating between two temperatures can have efficiency greater than that ofthe Carnot engine.

15. If Q > 0, heat is added to the system

If Q < 0, heat is removed to the system

If W > 0, Work is done by the system

If W < 0, Work is done on the system

Quantity Symbol Dimensions Unit Remark

Co-efficienty of volume αv [K–1] K–1 αv = 3 α1

expansion

Heat supplied to a system ΔQ [ML2 T–2] J Q is not a statevariable

Specific heat s [L2 T–2 K–1] J kg–1 K–1

Thermal Conductivity K [MLT–3 K–1] J s–1 K–1 H = – KA dd

tx

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POINTS TO PONDER

1. Temperature of a body is related to its average internal energy, not to the kinetic energyof motion of its centre of mass. A bullet fired from a gun is not at a higher temperaturebecause of its high speed.

2. Equilibrium in thermodynamics refers to the situation when macroscopic variablesdescribing the thermodynamic state of a system do not depend on time. Equilibrium ofa system in mechanics means the net external force and torque on the system are zero.

3. In a state of thermodynamic equilibrium, the microscopic constituents of a system arenot in equilibrium (in the sense of mechanics).

4. Heat capacity, in general, depends on the process the system goes through when heat issupplied.

5. In isothermal quasi-static processes, heat is absorbed or given out by the system eventhough at every stage the gas has the same temperature as that of the surroundingreservoir. This is possible because of the infinitesimal difference in temperature betweenthe system and the reservoir.

EXERCISES

12.1 A geyser heats water flowing at the rate of 3.0 litres per minute from 27 °C to 77 °C.If the geyser operates on a gas burner, what is the rate of consumption of the fuel ifits heat of combustion is 4.0 × 104 J/g ?

12.2 What amount of heat must be supplied to 2.0 × 10–2 kg of nitrogen (at roomtemperature) to raise its temperature by 45 °C at constant pressure ? (Molecularmass of N2 = 28; R = 8.3 J mol–1 K–1.)

12.3 Explain why(a) Two bodies at different temperatures T1 and T2 if brought in thermal contact do

not necessarily settle to the mean temperature (T1 + T2 )/2.(b) The coolant in a chemical or a nuclear plant (i.e., the liquid used to prevent

the different parts of a plant from getting too hot) should have high specificheat.

(c) Air pressure in a car tyre increases during driving.(d) The climate of a harbour town is more temperate than that of a town in a desert

at the same latitude.

12.4 A cylinder with a movable piston contains 3 moles of hydrogen at standard temperatureand pressure. The walls of the cylinder are made of a heat insulator, and the pistonis insulated by having a pile of sand on it. By what factor does the pressure of thegas increase if the gas is compressed to half its original volume ?

12.5 In changing the state of a gas adiabatically from an equilibrium state A to anotherequilibrium state B, an amount of work equal to 22.3 J is done on the system. If thegas is taken from state A to B via a process in which the net heat absorbed by thesystem is 9.35 cal, how much is the net work done by the system in the latter case ?(Take 1 cal = 4.19 J)

12.6 Two cylinders A and B of equal capacity are connected to each other via a stopcock.A contains a gas at standard temperature and pressure. B is completely evacuated.The entire system is thermally insulated. The stopcock is suddenly opened. Answerthe following :(a) What is the final pressure of the gas in A and B ?(b) What is the change in internal energy of the gas ?(c) What is the change in the temperature of the gas ?(d) Do the intermediate states of the system (before settling to the final equilibrium

state) lie on its P-V-T surface ?

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12.7 A steam engine delivers 5.4×108J of work per minute and services 3.6 × 109J of heatper minute from its boiler. What is the efficiency of the engine? How much heat iswasted per minute?

12.8 An electric heater supplies heat to a system at a rate of 100W. If system performswork at a rate of 75 joules per second. At what rate is the internal energy increasing?

12.9 A thermodynamic system is taken from an original state to an intermediate state bythe linear process shown in Fig. (12.13)

Fig. 12.13Its volume is then reduced to the original value from E to F by an isobaric process.Calculate the total work done by the gas from D to E to F

12.10 A refrigerator is to maintain eatables kept inside at 90C. If room temperature is 360C,calculate the coefficient of performance.