CHAPTER OUTLINEil-t What Is Thermodynamics and

Why ls It Useful?

12 Basic Definitions Needed toDescribe ThermodynamicSystems

I I Thermometry

tA Equations of State and theIdeal Gas Law

l.s A Brief Introduction toReal Gases

CHAPTER T

Fundamental Conceptsof Thermodynamics

Thermodynamics provides a description of matter on a macroscopic scale. lnthis approach, matter is described in terms of bulk properties such aspressure, density, volume, and temperature. The basic concepts employed inthermodynamics, such as system, surroundings, intensive and extensivevariables, adiabatic and diathermal walls, equilibrium, temperature, andthermometry, are discussed in this chapter. The usefulness of equations ofstate, which relate the state variables of pressure, volume, and temperature, isalso discussed for real and ideal gases. r

I .l What Is Thermodynamicsand Why Is It Useful?

Thermodynamics is the branch of science that describes the behavior of matter andthe transformation between different forms of energy on a macroscopic scale, or thehuman scale and larger. Thermodynamics describes a system of interest in terms ofits bulk properties. only a few such variables are needed to describe the system, andthe variables are generally directly accessible through measurements. A thermody-namic description of matter does not make reference to its structure and behavior atthe microscopic level. For example, 1 mol of gaseous water at a sufficiently low den-sity is completely described by two of the three macroscopic variables of pressure,volume, and temperature. By contrast, the microscopic scale refers to dimensions onthe order of the size of molecules. At the microscopic level, water is a dipolar tri-atomic molecule, Hro, with a bond angle of 104.5" that forms a network of hydro-gen bonds.

In this book, you will learn first about thermodynamics and then about statisticalthermodynamics. Statistical thermodynamics (Chapters 3l through 33) uses atomicand molecular properties to calculate the macroscopic properties of matter. For exam-ple, statistical thermodynamics can show that liquid water is the stable form of aggre-gation at a pressure of I bar and a temperature of 90'C, whereas gaseous water is thestable form at 1 bar and 110'C. using statistical thermodynamics, macroscopic proper-ties are calculated from underlying molecular properties.

Given that the microscopic nature of matter is becoming increasingly well under-stood using theories such as quantum mechanics, why is a macroscopic science like

2 C HAPTER I Fundamental Concepts ofThermodynamics

thermodynamics relevant today? The usefulness of thermodynamics can be illustratedby describing four applications of thermodynamics:

r You have built a plant to synthesize NH. gas from N, and Hr. you find that theyield is insufficient to make the process profitable and decide to try to improve theNH. output by changing the temperature and/or the pressure. However, you do notknow whether to increase or decrease the values of these variables. As will beshown in Chapter 6, the ammonia yield will be higher at equilibrium if thetemperature is decreased and the pressure is increased.

' You wish to use methanol to power a car. one engineer provides a design for aninternal combustion engine that will burn methanol efficiently according to thereaction CH3OH (l) + 2O2(g) - CO, (g) + ZH'O (/). A second engineer designsan electrochemical fuel cell that caries out the same reaction. He claims that thevehicle will travel much farther if powered by the fuel cell than by the internalcombustion engine. As will be shown in Chapter 5, this assertion is correct, and anestimate of the relative efficiencies of the two propulsion systems can be made.

r You are asked to design a new battery that will be used to power a hybrid car.Because the voltage required by the driving motors is much higher than can begenerated in a single electrochemical cell, many cells must be connected in series.Because the space for the battery is limited, as few cells as possible should beused. You are given a list of possible cell reactions and told to determine thenumber of cells needed to generate the required voltage. As you will learn inchapter 11, this problem can be solved using tabulated values of thermodynamicfunctions.

' Your attempts to synthesize a new and potentially very marketable compound haveconsistently led to yields that make it unprofitable to begin production. Asupervisor suggests a major effort to make the compound by first synthesizing acatalyst that promotes the reaction. How can you decide if this effort is worth therequired investment? As will be shown in chapter 6, the maximum yield expectedunder equilibrium conditions should be calculated first. Ifthis yield is insufficient,a catalyst is useless.

1.2 Basic Definitions Needed to DescribeThermodynamic Systems

A thermodynamic system consists of all the materials involved in the process understudy. This material could be the contents of an open beaker containing reagents, theelectrolyte solution within an electrochemical cell, or the contents of a cylinder andmovable piston assembly in an engine. In thermodynamics, the rest of the universe is re-ferred to as the surroundings. If a system can exchange matter with the sur:roundings,it is called an open systeml if not, it is a closed system. Both open and closed systemscan exchange energy with the surroundings. Systems that can exchange neither matternor energy with the sur:roundings are called isolated systems.

The interface between the system and its surroundings is called the boundary. Theboundaries determine if energy and mass can be transferred between the system and thesurroundings and lead to the distinction between open, closed, and isolated systems.consider the Earth's oceans as a system, with the rest of the universe being the sur-roundings. The system-surroundings boundary consists of the solid-liquid interface be-tween the continents and the ocean floor and the water-air interface at the ocean surface.For an open beaker in which the system is the contents, the boundary surface isjust in-side the inner wall of the beaker, and passes across the open top of the beaker. In thiscase, energy can be exchanged freely between the system and surroundings through theside and bottom walls, and both matter and energy can be exchanged between the sys-

I .2 Basic Definitions Needed to Describe Thermodynamic Systems 3

tem and suffoundings through the open top boundary. The portion of the boundaryformed by the beaker in the previous example is called a wall. Walls are always bound-

aries, but a boundary need not be a wall. Walls can be rigid or moyable and permeableor nonpermeable. An example of a movable wall is the surface of a balloon.

The exchange of energy and matter across the boundary between system and sur-

roundings is central to the important concept of equilibrium. The system and surround-

ings can be in equilibrium with respect to one or more of several different systemvariables such as pressure (P), temperature (f, and concentration. Thermodynamicequilibrium refers to a condition in which equilibrium exists with respect to P, T, and

concentration. A system can be in equilibrium with its surroundings with respect to a

given variable. Equilibrium exists only if that variable does not change with time, and ifit has the same value in all parts of the system and surroundings. For example, the inte-rior of a soap bubblel (the system) and the surroundings (the room) are in equilibriumwith respect to P because the movable wall (the bubble) can reach a position where P on

both sides of the wall is the same, and because P has the same value throughout the sys-

tem and surroundings. Equilibrium with respect to concentration exists only iftransportof all species across the boundary in both directions is possible. If the boundary is a mov-able wall that is not permeable to all species, equilibriurn can exist with respect to P, butnot with respect to concentration. Because N, and O, cannot diffuse through the (ideal-

ized) bubble, the system and surroundings are in equilibrium with respect to P, but notto concentration. Equilibrium with respect to temperature is a special case that is dis-

cussed next.Temperature is an abstract quantity that is only measured indirectly, for example, by

measuring the volume of mercury confined to a narrow capillary, the electromotive force

generated at the junction of two dissimilar metals, or the electrical resistance of a platinumwire. At the microscopic level, temperature is related to the mean kinetic energy of mole-cules. Although each of us has a sense of a 'itemperature scale" based on the qualitative de-

scriptors /zofandcold,weneed amore quantitative and transferable measure of temperature

that is not grounded in individual experience. To make this discussion more concrete, weconsider a dilute gas under conditions in which the ideal gas law ofEquation ( L 1 ) describes

the relationship among P,T, andthe molar density p = nlV wrth sufficient accuracy:

p=pRT (1.1)

The ideal gas law is discussed further in Section 1.4.

In thermodynamics, temperature is the property of a system that determines if the

system is in thermal equilibrium with other systems or the surroundings. Equation (1.1)

can be rewritten as follows:

(t.2)

This equation shows that for ideal gas systems having the same molar density, a pres-

sure gauge can be used to compare the systems and determine which of I, or T, isgreater. Thermal equilibrium between systems exists if Pr: Pr.for gaseous systems

with the same molar density.We use the concepts of temperature and thermal equilibrium to characterize the

walls between a system and its surroundings. Consider the two systems with rigid wallsshown in Figure 1.1a. Each system has the same molar density and is equipped with a

pressure gauge. If we bring the two systems into direct contact, two limiting behaviorsare observed. If neither pressure gauge changes, as in Figure 1.lb, we refer to the wallsas being adiabatic. Because Pr+ P2, the systems are not in thermal equilibrium and,

therefore, have different temperatures. An example of a system surrounded by adiabatic

I For this example, the surface tension of the bubble is assumed to be so smail that it can be set equal to zero.

This is in keeping with the thermodynamic tradition of weightless pistons and frictionless pulleys.

P

pRT_

(c)

FrsuR,r r.:(a) Two separated systems with rigid wallsand the same molar density have differenttemperatures. (b) The two systems are

brought together so that the walls are inintimate contact. Even after a long timehas passed, the pressure in each system is

unchanged. (c) As in part (b), the twosystems are brought together so that thewalls are in intimate contact. After a

sufficient time has passed, the pressures

are equal.

sEs

eg

4 c HAPTER 'l Fundamental Concepts of Thermodynamics

walls is coffee in a Styrofoam cup with a Styrofoam lid.2 Experience shows that it isnot possible to bring two systems enclosed by adiabatic walls into thermal equilibriumby bringing them into contact, because adiabatic walls insulate against the transfer of"heat." If you push a Styrofoam cup containing hot coffee against one containing ice

water, they will not reach the same temperature. Rely on your experience at this point

regarding the meaning of heat; a thermodynamic definition will be given in Chapter 2.

The second limiting case is shown in Figure 1.1c. In bringing the systems into inti-mate contact, both pressures change and reach the same value after some time. We con-

clude that the systems have the same temperature, Tr: T, and say that they are in thermal

equilibrium. We refer to the walls as being diathermal. Two systems in contact separated

by diathermal walls reach thermal equilibrium because diathermal walls conduct heat.

Hot coffee stored in a copper cup is an example of a system surrounded by diathermal

walls. Because the walls are diathermal, the coffee will quickly reach room temperature.

The zeroth law of thermodynamics generalizes the experiment illustrated in

Figure 1.1 and asserts the existence of an objective temperature that can be used to de-

fine the condition of thermal equilibrium. The formal statement of this law is as follows:

There are four laws of thermodynamics, all of which are generalizations from experi-

ence rather than mathematical theorems. They have been rigorously tested in more than a

century of experimentation, and no violations of these laws have been found. The unfortu-

nate name assigned to the"zeroth" law is due to the fact that it was formulated after the firstlaw of themodynamics, but logically precedes it. The zeroth law tells us that we can deter-

mine if two systems are in thermal equilibrium without bringing them into contact. Imag-

ine the third system to be a thermometer, which is defined more precisely in the next section.

The third system can be used to compare the temperatures of the other two systems; if they

have the sarne temperature, they will be in thermal equilibrium if placed in contact.

1.3 ThermometryThe discussion of thermal equilibrium required only that a device exist, called a thermome-

ter, that can measure relative hotness or coldness. However, scientific work requires a quan-

titative temperature scale. For any useful thermometer, the empirical temperature, /, must be

a single-valued, continuous, and monotonic function of some thermometric system property

designated by x. Examples of thermometric properties are the volume of a liquid, the elec-

trical resistance of a metal or semiconductor, and the electromotive force generated at the

junction of two dissimilar metals. The simplest case that one can imagine is if the empirical

temperature, r, is linearly related to the value of the thermometric property, r:

t(x) = 616, (1.3)

Equation (1.4) defines a temperature scale in terms of a specific thermometric prop-

erty, once the constants a andb are fixed. The constant a determines the zero of the tem-

perature scale because t(0) : a, and the constant b determines the size of a unit oftemperature, called a degree.

One of the first practical thermometers was the mercury-in-glass thermometer. It uti-lizes the thermometric property that the volume of mercury increases monotonically

over the temperature range in which it is in the liquid state (between -38.8o and

356.7"C).It1145, Carolus Linnaeus gave this thermometer a standardized scale by ar-

bitrarily assigning the values 0 and 100 to the freezing and boiling points of water, re-

spectively. This chosen interval was divided into 100 equal degrees, and the same size

2 In this discussion, Styrofoam is assumed to be a perfect insulator.

L3 Thermometry 5

degree is used outside of the interval. Because there are 100 degrees between the twocalibration points, it is called the centigrade scale.

The centigrade scale has been superseded by the Celsius scale, which is in wide-spread use today. The Celsius scale (denoted in units of 'C) is similar to the centigradescale. However, rather than being determined by two fixed points, the Celsius scale is de-termined by one fixed reference point at which ice, liquid water, and gaseous water are inequilibrium. This point is called the triple point (see Section 8.2) and is assigned the value0.01'C. On the Celsius scale, the boiling point of water at a pressure of 1 atmosphere is99.975'C. The size of the degree is chosen to be the same as on the centigrade scale.

Although the Celsius scale is used widely throughout the world today, the numeri-cal values for this temperature scale are completely arbitrary. It would be preferable tohave a temperature scale derived directly from physical principles. There is such a scale,called the thermodynamic temperature scale or absolute temperature scale. For sucha scale, the temperature is independent of the substance used in the thermometer, and theconstant a in Equation (1.4) is zero. The gas thermometer is a practical thermometerwith which the absolute temperature can be measured. The thermometric property is thetemperature dependence of P for a dilute gas at constant V. The gas thermometer pro-vides the international standard for thermometry at very low temperatures. At interme-diate temperatures, the electrical resistance of platinum wire is the standard, and athigher temperatures, the radiated energy emitted from glowing silver is the standard.

How is the gas thermometer used to measure the thermodynamic temperature? Mea-0.1 L surements carried out by Robert Boyle in the 19th century demonstrated that the pressure

exerted by a fixed amount of gas at constant Vvaries linearly with temperature on the Cel-sius scale as shown in Figure 1.2. Atthetime of Boyle's experiments, temperatures below-30'C were not attainable in the laboratory. However, the P versus T datacatbe extrap-olated to the limiting Tvalue at which P + 0. It is found that these straight lines obtainedfor different values of yintersect at a common point on the Zaxis that lies near -2'l3oC.

The data show that at constant V, the thermometric property P varies with temper-ature as

0.2 L

0.3 L

0.4 L0.5 L0.6 L

(r.4)

where / is the temperature on the Celsius scale, and c and d are experimentally obtainedproportionality constants.

Figure 1.2 shows that all lines intersect at a single point, even for different gases. Thissuggests a unique reference point for temperature, rather than the two reference pointsused in constructing the centigrade scale. The value zero is given to the temperature atwhich P -+ 0. However, this choice is not sufficient to define the temperature scale, be-cause the size of the degree is undefined. By convention, the size of the degree on the ab-solute temperature scale is set equal to the size ofthe degree on the Celsius scale, becausethe Celsius scale was in widespread use at the time the absolute temperature scale wasformulated. With these two choices, the absolute and Celsius temperature scales are re-lated by Equation (1.5). The scale measured by the ideal gas thermometer is the absolutetemperature scale used in thermodynamics. The unit of temperature on this scale is calledthe kelvin, abbreviated K (without a degree sign):

Z(K) = T('C)+273.15 (l.s)

using the triple point of water as a reference, the absolute temperature, z(K), measuredby an ideal gas thermometer is given by

'F,=,e *dt:,

-200 -100 0 100 200 300 400

Temperature (Celsius)

FIG{JRE T.3The pressure exerted by 5.00 x 10-3 molof a dilute gas is shown as a function ofthe temperature measured on the Celsiusscale for different fixed volumes. Thedashed portion indicates that the data areextrapolated to lower temperatures thancould be achieved experimentally byBoy1e.

r(K)= TT3.rcL,pte,

(1.6)

whete P*is the pressure corresponding to the triple point of water. On this scale, the vol-ume of an ideal gas is direct$ proportional to its temperature; if the temperature is reducedto half of its initial value, Vis also reduced to half of its initial value. In practice, deviationsfrom the ideal gas law that occur for real gases must be taken into account when using a gas

6 C HAPTER I Fundamental Concepts ofThermodynamics

fl zre.z

thermometer. If data were obtained from a gas thermometer using He, Ar, and Nr for a tem-

perature very near T,o,they would exhibit the behavior shown in Figure 1.3' We see that the

iemperature only be6omes independent of P and of the gas used il the thermometer if the

dataare extrapolated to zero pressure. It is in this limit that the gas thermometer provides a

measure of the thermodynamic tempefature. Because 1 bar : 105 Pa, gas-independent

7- values are only obtained below P ' 0.01 bar. The absolute temperature is shown in

Figure 1.4 on a logarithmic scale together with associated physical phenomena.

1.4 Equations of State and the Ideal Gas LawMacroscopic models in which the system is described by a set of variables are based on

experience. It is particularly useful to formulate an equation ofstate, which relates the

state variables. Using the absolute temperature scale, it is possible to obtain an equation

of state for an ideal gas from experiments. If the pressure of He is measured as a func-

tion of the volume for different values of temperature, the set of nonintersecting hyper-

bolas shown in Figure 1.5 is obtained. The curves in this figure can be quantitatively fitby the functional form

PV =aT (1.7)

where 7is the absolute tempefature as defined by Equation (1.6), allowing a to be de-

termined, which is found to be directly proportional to the mass of gas used. It is useful

to separate out this dependence by writing a = nR, where n is the number of moles of the

gas, ind R is a constant that is independent of the size of the system. The result is the

ideal gas equation of state:

PV = NkI =nRT (1.8)

where the proportionality constants k and R are called the Boltzmann constant and the

ideal gas constant, respectively; N is the number of molecules and n is the number ofmoles of the gas. The equation of state given in Equation (1.8) is known as the ideal gas

law. Because the four variables are related through the equation of state, any three of

these variables is sufficient to completely describe the ideal gas. Note that the total num-

ber of moles-not the number of moles of the individual gas components of a gas

mixture-appears in the ideal gas law.

Of these four variables, P and Tare independent of the amount of gas, whereas Vand

n are proportional to the amount of gas. A variable that is independent of the size of the

system (for example, P and Z) is referred to as an intensive variable, and one that is

proportional to the size of the system (for example, V) is referued to as an extensive vari-

able. Equation (1.8) can be written in terms of intensive variables exclusively:

p=pRT ( 1.8a)

For a fixed number of moles, the ideal gas equation of state has only two independent

intensive variables: any two of P, T, atd p.

For an ideal gas mixture

273.0

272.8

fi$ffiLiffi.ffi €"3The temperature measured in a gas

themometer defined by Equation (1.6) is

independent of the gas used only in the

limitthatP+0.

1012 K

1010K

108 K

106 K

104 K

102 K

1K

1o-2 K

1o 4K

Ambient temperature10 a s after Big Bang

fl$ffiwRffi I.4The absolute temperature is shown on a

logarithmic scale together with the

temperature of a number of physical

phenomena.

Core of red giant star

Core of sun

Solar corona

Surface of sun

Mercury boils

H2O is liquid

Oxygen boils

Helium boils

Average temperatureof universe

3He superfluidrv =\n,RT

i

(1.e)

because the gas molecules do not interact with one another. Equation (1.9) can be rewrit-

ten in the form

P=T''R7=TP4 V 4',

4681o-4 P (Pa)

= Pt* Pz+ 4 +.., (1.10)

1.4 Equations of State and the ldeal Cas Law 7

In EQuation (1.10), p, is the partial pressure ofeach gas. This equation states thateach ideal gas exefis a pressure that is independent of the other ga.es in the mixture.We also have

(1.11)

0.2 0.4 0.6 0.8 1.0 1.2 1.4

100 x volume (m3)

:tr*ilffifr *"sliustration of the relationship between:r'.ssure and volume of 0.010 mol of He:rr fxed values of temperature, which:ff'er by 100 K.

n,RT n,RTD-'i_ V V n

P 5I4r nRT-n--'iLv v

which relates the partial pressure of a component in the mixture, p, with its mole frac-tion, x, = niln, and the total pressure, p.

In the SI system of units, pressure is measured in pascar (pa) units, where 1 pa: I N/m2. The volume is measured in cubic meters, and the temperature is measuredin kelvin. However, other units of pressure are frequently used, and these units arerelated to the Pascal as indicated in Table 1.1. In this tablL, numbers that are not ex-act have been given to five significant figures. The other commonly used unit of vol_umeistheliter(L), where 1 m3 : 103Land 1 L : I dm3 : 10-3m3.

In the sI system, the constant R that appears in the idear gas raw has the value8.314 J K-r mol 1, where the joule (J) is the unit of energy in tlhe SI system. To sim_plify calculations for other units of pressure and volume, values of the constant R withdifferent combinations of units are given in Table 1.2.

EXAMPLE PROBTEM I.IStarting o-ut on a trip into the mountains, you inflate the tires on yourautomobile to a recommended pressu re of 3.2r x 10s pa on a dai when thetemperarure is -5.00'c. you drive ro rn. u.u.rr, *rr.*lrr.

";;1,;;;.; i;28.0'c. Assume that rhe votume of the rire rtu, i'n.r.ur.J unlzJ. *ir", ir1t .final pressure in the tires? The manufacturer of the tir" .".o.r.";; il ;;;not exceed rhe recommended pressure by more rhan I0To. Hr, ,rri, ii"rr, d.*exceeded?

SolutionBecause the number oF moles is constant,

P,V. P,V, PV,T"t I- J J. D _lt I

T. T" I v.T-t ^J '.1'i

^ PiviTr V 073.15+2g-0)Pt= r7=3.2lxl05Pax' "ffi=3.50xl0spaII

This pressure is within l\Vo of rhe recommended pressure.

TABLE I.It-Inits af Pressure and Ccnversion F*ctcrs

[]nit ofPressure Symbol Nunrerical Value

,PiiScal',, " ,' ,

Alrnosphere ' .r,

Bar ...,.: ;: . ,. ::.

Torr or rni.llimeters of,Flg

For,lndspqq Sou*ulosh

lNm 2=lkgmi-2. , lratia'= rc#25 Pa (exactly) ,

.1 bar =' 105 P, "

I Torr : I0l .325/760: I33.32 pa

I psi : 6,894.8 pa

Pa

atm

bar

Torr

psl

8 CHAPTER I Fundamental Concepts ofThermodynamics

TABLE I.2T'3r* Xel*;*} G;as

il*c1staffit, Jt,

in Vari.*ws {Jnit*R: 8.314J K I mol I

R : 8.314 Pa m3 K-1 mol-lR:8.314 x 10 2LbarK 1mol-r

R : 8.206 x 10-z L atm K-1 mol-1

R : 62.36 LTorr K-r mol-1

EXAMPI.E PROBLEM I.2

Consider the composite system, which is held at 298K, shown in the following

figure. Assuming ideal gas behavior, calculate the total pressure, and the partial

pressure of each component if the barriers separating the compartments are

removed. Assume that the volume of the barriers is negligible'

He2.00 L

1.50 bar

Ne3.00 L

2.50 bar

SnlutionThe number of moles of He, Ne, and Xe is given by

PV 1.50 bar x 2.00 LflH"

flNn

RT 8.314 x 10-2 L bar K I mol-r x 298 K

PV 2.50bar x 3.00 LRT 8.314 x 10r L bar K-1 mol-1 x 298 KPV 1.00bar x 1.00 L

= 0.121moI

= 0.303 mol

= 0.0403 molnxn= u =

fl=frHn+nNe+nx"=0.464

The mole fractions are

n- 0.121-Y,, =3=:___::_=0.]6 In( n 0.464

n n,- 0.303-r. =___rq= "'-"-

=0.65.1I\e n O.464

n.. 0.0403v ---4L -0.0860n'' - n 0.464

The total pressure is given by

^ (fl,0 +riNe + n"n)RT

v= 1.92 bar

The partial pressures are given by

Pu"= xrrP =0.261x1'92bar = 0'501bar

P*n = x *"P = 0.653 x1' .92 bar = 1 .25 bar

P"n= x""P = 0.0860 xl.92bu = 0.165bar

0.464 molx8.3145 x 10-2 L bar K-r mol-l x 298 K6.00 L

Real gas ldeal gasl+

FI6tJRE I.€The potential energy of interaction of twomolecules or atoms is shown as a functionoftheir separation, r. The yellow curverbws the potential energy function for anirlesl gas. The dashed blue line indicates an

E4)roximate r value below which a morenearly exact equation of state than theileal gas law should be used. V(r) = g u1

r: rv:oand as r+ €.

1.5 A Brief lntroduction to Real Cases 9

1.5 A Brief Introduction to Real GasesThe ideal gas law provides a first look at the usefulness of describing a system interms of macroscopic parameters. However, we should also emphasize the downsideof not taking the microscopic nature of the system into account. For example, the idealgas law only holds for gases at low densities. Experiments show that Equation (1.g)is accurate to higher values of pressure and lower values of temperature for He thanfor NH.. Why is this the case? Real gases will be discussed in detail in Chapter 7.However, because we need to take nonideal gas behavior into account in Chapters 2through 6, we introduce an equation of state that is valid to higher densities in thissection.

An ideal gas is described by two assumptions: the atoms or molecules of an ideal gasdo not interact with one another, and the atoms or molecules can be treated as pointmasses. These assumptions have a limited range of validity, which can be discussed us-ing the potential energy function typical for a real gas, as shown in Figure 1.6. This fig-ure shows the potential energy of interaction of two gas molecules as a function of thedistance between them. The intermolecular potential can be divided into regions inwhich the potential energy is essentially zero (r> r,,on"xinn),negative (attractive interac-tiofl) (roon"i,io,, !, rr:o), and positive (repulsive ini"iiCtion) (r < rr:o). The distancertransitionis not uniquely defined and depends on the energy of the molecule. It can be es-timated from the relation lV(r,*,,i,ion)l= *f .

As the density is increased from very low values, molecules approach one anotherto within a few molecular diameters and experience a long-range attractive van derWaals force due to time-fluctuating dipole moments in each molecule. This strength ofthe attractive interaction is proportional to the polarizability of the electron charge in amolecule and is, therefore, substance dependent. In the attractive region, p is lower thanthat calculated using the ideal gas law. This is the case because the attractive interactionbrings the atoms or molecules closer than they would be if they did not interact. At suf-ficiently high densities, the atoms or molecules experience a short-range repulsive in-teraction due to the overlap of the electron charge distributions. Because of thisinteraction, P is higher than that calculated using the ideal gas law. We see that for a realgas, P can be either greater or less than the ideal gas value. Note that the potential be-comes repulsive for a value of r greater than zero. As a consequence, the volume of agas well above its boiling temperature approaches a finite limiting value as p increases.By contrast, the ideal gas law predicts thatV -+ 0 as p -+ -.

Given the potential energy function depicted in Figure 1.6, under what conditions isthe ideal gas equatio'n of state valid? Areal gas behaves ideally only at low densities forwhich r ) rrans,ion, and the value of rtansitionis substance dependent. Real gases will bediscussed in much greater detail in Chapter 7. However, at this point we introduce a realgas equation of stat$ecause it will be used in the next few chapters. The van der Waalsequation of state takes both the finite size of molecules and the attractive potential intoaccount. It has the form

^ nRT n2ari = --:+- -

. ,'. ,*nb.,',,V2,(1.t21

This equation of state has two parameters that are substance dependent and must be exper-imentally determined. The parameters a and b take the finite size of the molecules and thestrength of the attractive interaction into account, respectively. (Values of a and, Z, for se-lected gases are listed in Table 7 .4.) The van der Waals equation of state is more accurate incalculating the relationship between p, v, and zfor gases than the ideal gas law because aand b have been optimized using experimental results. However, there are other more ac-curate equations of state that are valid over a wider range than the van der Waals equation.such equations of state include up to l6 adjustable substance-dependent parameters.

lO CHAPTER 1 Fundamental Concepts ofThermodynamics

EXAMPLE PROBLEM I '3

Van der Waals parameters a1e generally tabulated with either of two sets of units:

a: Pa m6 mo1-2 or bar dm6 mol-2

b: m3 mol-r or dm3 mol-1

Determine the conversion factor to converl one System of units to the Other' Note

thatldms:10-3mr:lL.Solution

- bar lo6 dm6 = tobardm6 mol-2Pamo mol-r x

105 pa x

m6

103 dm3m, mol-. X--=-- - r03 dm3 mol-1

mr

In Example Problem 1.4, a comparison is made of the molar volume for N, calcu-

lated at low and high pressures, using the ideal gas and van der waals equations of state.

EXAMPLE PROBLEM I.4

a. calculate the pressure exerted by N, at 300 K for molar volumes oJ 250

and 0.100 L using the ideal gas and the van der waals eguations of state.

The values of parameters a and b for N, are 1'370 bar dm6 mol-z and

0.0387 dm3 mol-1, resPectivelY'

b. Compare the results of your calculations at the two pressures' If P

calculated using the van der waals equation of state is greater than those

calculated with the ideal gas law, we can conclude that the repulsive

interaction of the N" molJcules outweighs the attractive interaction for the

calculated value of ihe density. A similar statement can be made regarding

the attractive interaction. Is the attractive or repulsive interaction greater for

Nrat 300 K and V- : 0.100 L?

Solutiona. The pressures calculated from the ideal gas equation of state are

^ nRT lmolx8.314x10r Lbarmol-rK-tx300 K =9.9gx10-2 bar

V 25OL

^ nRT lmolx8.314xl0 2 L barnglrK-r x]90K =249bw-=- v 0.100 L

The pressures calculated from the van der Waals equation of state are

nRT n2aD- '""' -"

--

' - v-rh v2

1molx8.314x104 Lbarmol-rK-l x300 K (lmol)2 xl'310b*4*1*o1'- 250 L-1moIx0.0387 dm3 mol-l (250 L)2

(1mol)2 x 1.370 bar dm6 mol'= 9.98 x 10-2 bar

1 mol x 8.314 x 10-2 L bar mol-1K-1 x 300 KP.

0.1 L - lmol x0.0387 dm3 mol-1 (0.1L)2

= 27obu

b. Note that the result is identical with that for the ideal gas law fot v*: 250L,

and that the result calculated for V^:0.100 L deviates from the ideal gas

law result. Because preat s pideal , *" conclude that the repulsiveinteraction is

more important than the attractive interaction for this specific value of molar

volume and temPerature.

Problems I I

Vocabularyffiwolute temperature scale

mahatic

Dmtramann constant

ffisrondary

0hhius scale

'rcmigrade scale

fued system

ffiidrermal

'-llotion of state

mFdtibrium

mtgnsive variable

8qs thermometer

Meal gas constant

iirileal gas law

intensive variable

isolated system

kelvin scale

macroscopic scale

macroscopic variables

microscopic scale

mole fraction

movable wall

nonpermeable wall

open system

partial pressure

permeable wall

rigid wall

statistical thermodynamics

surroundings

system

system variables

temperature

temperature scale

thermal equilibrium

thermodynamic equilibrium

thermodynamic temperature scale

thermometer

van der Waals equation of state

wall

zeroth law of thermodynamics

Questions on Concepts

Qf.l The location of the boundary between the system andil!" surroundings is a choice that must be made by thefumodynamicist. Consider a beaker of boiling water in anflirtight room. Is the system open or closed if you place thehoundaryjust outside the liquid water? Is the system open ordo6ed if you place the'boundary just inside the walls of thermm?

QfJ Real walls are never totally adiabatic. Order thefollowing walls in increasing order with respect to their beingdiethermal: 1 -cm-thick concrete, 1 -cm-thick vacuum, 1 -cm-ttick copper, l-cm-thick cork.

Q1.3 Why is the possibility of exchange of matter or energyappropriate to the variable ofinterest a necessary conditionfor equilibrium between two systems?

Q1.4 At sufficiently high temperatures, the van der Waalsequation has the form P = RTf (V*- b). Note that theattractive part of the potential has no influence in thisexpression. Justify this behavior using the potential energydiagramcf Figure 1.6.

Q1.5 The parameter ainthe van der Waals equation isgreater for HrO than for He. What does this say about theform of the potential function in Figure 1.6 for the two gases?

Problemshoblem numbers in red indicate that the solution to theproblem is given inthe Student's Solutions Manual.Pl.l A sealed flask with a capacity of 1.00 dm3 contains5-00 g of ethane. The flask is so weak that it will burst iftre pressure exceeds 1.00 X 106 pa. At what temperaturenill the pressure of the gas exceed the burstingtemperature?

Pl.2 Consider a gas mixture in a 2.00-dm3 flask at 27.0"C.For each of the following mixtures, calculate the partialpressure ofeach gas, the total pressure, and the compositionof the mixture in mole percent:

a 1.00 g H, and 1.00 g O,b. 1.00 g N, and 1.00 g O,c. 1.00 g CH. and 1.00 g NH.

P1.3 Suppose that you measured the product pV of I mol ofa dilute gas and found that PV : 22.98 L atm at 0.C and31.18 L atm at 100'C. Assume that the ideal gas law is valid,with Z : ('C) * a, and that the value of R is not known.Determine R and a from the measurements provided.

P1.4 A compressed cylinder of gas contains 1.50 x 103 gof N, gas at a pressure of 2.00 X 107 Pa and a temperatureof 17.1'C. What volume of gas has been released intothe atmosphere if the final pressure in the cylinder is1.80 X 105 Pa? Assume ideal behavior and that the gastemperature is unchanged.

P1.5 A balloon filled with 10.50 L of Ar ar 18.0"C andI atm rises to a height in the atmosphere where the pressure is248Torr and the temperature is -30.5.C. What is the finalvolume of the balloon?

l2 c HAPTER I Fundamental Concepts ofThermodynamics

P1.6 Consider a20.0-L sample of moist air at 60"C and 1 atm

in which the partial pressure of water vapor is 0.120 atm.

Assume that dry air has the composition 78.0 mole percent N,21.0 mole percent O, and i.00 mole percentAr.

a. What are the mole percentages of each of the gases in the

sample?

b. The percent relative humidity is defined as ToRH: Pr,o f Pirowhere Prrois the partial pressure of water inthe sdmple and Pi,"o: 0.197 atm is the equilibrium vaporpressure of water at 60'C. The gas is compressed at 60'Cuntil the relative humidity is l0o7o. What volume does the

mixture contain now?

c. What fraction of the water will be condensed if the total

pressure of the mixture is isothermally increased to 200 atm?

Pl.7 A mixture of 2.50 x 10-3 g of Or,3.51 x 10-3 mol ofN, and 4.67 x 7020 molecules of CO are placed into a vessel

of volume 3.50 L at 5.20'C.

a. Calculate the total pressure in the vessel.

b. Calculate the mole fractions and partial pressures ofeach gas.

P1.8 Liquid N, has a density of 875.4 kg m t at its normal

boiling point. What volume does a balloon occupy at 18.5'C

and a pressure of 1.00 atm if 2.00 x 10-3 L of liquid N, isinjected into it?

P1.9 Arigid vessel of volume 0.500 m3 containingHrat20.5'C and a pressure of 6 1 1 x 1 03 Pa is connected to a

second rigid vessel of volume 0.750 m3 containing Ar at

3I.2'C at a pressure of 433 x 103 Pa. A valve separating the

two vessels is opened and both are cooled to a temperature of14.5'C. What is the final pressure in the vessels?

P1.10 A sample of propane (C3Hs) is placed in a closed

vessel together with an amount of O, that is 3.00 times the

amount needed to completely oxidize the propane to CO, and

HrO at constant temperature. Calculate the mole fraction ofeach component in the resulting mixture after oxidationassuming that the HrO is present as a gas.

P1.11 A glass bulb of volume 0.136 L contains 0.7031 g

of gas at 759.0 Torr and 99.5'C. What is the molar mass ofthe gas?

Pl.l2 The total pressure of a mixture of oxygen and

hydrogen is 1.00 atm. The mixture is ignited and the water is

removed. The remaining gas is pure hydrogen and exerts a

pressure of 0.400 atm when measured at the same values of Zand V as the original mixture. What was the composition ofthe original mixture in mole percent?

P1.13 A gas sample is known to be a mixture of ethane and

butane. A bulb having a 200.0-cm3 capacity is filled with the

gas to a pressure of 100.0 x 103 Pa at 20.0"C. If the weight ofthe gas in the bulb is 0.3846 g, what is the mole percent ofbutane in the mixture?

P1..14 When Julius Caesar expired, his last exhalation had a

volume of 500 cm3 and contained 1.00 mole percent argon'

Assume thatT :300 K and P : 1.00 atm at the location of

his demise. Assume further that T and P currently have the

same values throughout the Earth's atmosphere. If all of his

exhaled CO, molecules are now uniformly distributedthroughout the atmosphere (which for our calculation is taken

to have a thickness of 1.00 km), how many inhalations of500 cm3 must we make to inhale one of the Ar molecules

exhaled in Caesar's last breath? Assume the radius of the

Earth to be 6.31 x 106 m. [Hint: Calculate the number of Aratoms in the atmosphere in the simplified geometry of a plane

of area equal to that of the Earth's surface and a height equal

to the thickness of the atmosphere. See Problem f . i5 for the

dependence of the barometric pressure on the height above

the Earth's surface.

P1.15 The barometric pressure falls off with height above

sea level in the Earth's atmosphere as Pr = P9e-MiszlRr where

P, is the partial pressure at the height z, {0 is the partial

pressure of component I at sea level, g is the acceleration ofgravity, R is the gas constant, and Zis the absolute

temperature. Consider an atmosphere that has the

composition x", = 0.600 and xro, = 0.400 and that

T : 300 K. Near sea level, the total pressure is 1.00 bar.

Calculate the mole fractions of the two components at a

height of 50.0 km. Why is the composition different fromits value at sea level?

P1.16 Assume that afu has a mean molar mass of 28.9 g

mol-1 and that the atmosphere has a uniform temperature

of 25.0'C. Calculate the barometric pressure at Denver, forwhich z : 1600 m. Use the information contained inProblem P1.15.

Pl,l7 Calculate the pressure exerted by Ar for a molar

volume of 7.42 L at 300 K using the van der Waals equation

of state. The va4der Waals parameters 4 and b for Ar are

1.355 bar dm6 mol 2 and 0.0320 dm3 mol-l, respectively. Is

the attractive or repulsive portion of the potential dominant

under these conditions?

P1.18 Calculate the pressure exerled by benzene for a molar

volume of I.42 L at 190 K using the Redlich-Kwong equation

of state:

RTalnRTn2al' - i.- b JT v,{v^+b) v - nb "fr vtv +nul

The Redlich-Kwong parameters a and b for benzene are

452.0 bar dm6 mol 2 K1t2 and 0.082'71 dm3 mol 1, respec-

tively. Is the attractive or repulsive portion of the potential

dominant under these conditions?P1..19 Devise a temperature scale, abbreviated G, for whichthe magnitude of the ideal gas constant is 1 .00 J G- 1 mol- 1.

Pl.2O A mixture of oxygen and hydrogen is analyzed bypassing it over hot copper oxide and through a drying tube'

Hydrogen reduces the CuO according to the reaction CuO *H, * Cu + H2O, and oxygen reoxidizes the copper formed

according to Cu * 1/2 or + Cuo' At 25"C and 750 Ton' 100'0

cm3 of the mixture yields 84.5 cm3 of dry oxygen measured at

25'C and 750 Ton after passage over CuO and the drying

agent. What is the original composition of the mixture?

CHAPTER OUTLINEzl The Internal Energy and

the First Law ofThermodynamics

22 Work

2-r Heat

2A Heat Capacity

z5 State Functions and PathFunctions

CHAPTER 2

Heat, Work, InternalEnergy, Enthalpy,and the First Lawof Thermodynamics

ln this chapter, the internal energy, U, is introduced. The first law ofthermodynamics relates AU to the heat (q) and work (w) that flows across theboundary between the system and the surroundings. Other important conceptsintroduced include the heat capacity, the difference between state and pathfunctions, and reversible versus irreversible processes. The enthalpy, H, isintroduced as a form of energy that can be directly measured by the heat flowin a constant pressure process. We show how AU, LH, q, and w can becalculated for processes involving ideal gases. r

2.1 The Internal Energy and the First Lawof Thermodynamics

In this section, we focus on the change in energy of the system and surroundings dur-ing a thermodynamic process such as an expansion or compression of a gas. In thermo-dynamics, we are interested in the internal energy of the system, as opposed to theenergy associated with the system relative to a particular frame of reference. For exam-ple, a spinning container of gas has a kinetic energy relative to a stationary observer.However, the internal energy of the gas is defined relative to a coordinate system fixedon the container. Viewed at a microscopic level, the internal energy can take on a num-ber of forms such as

. the kinetic energy of the molecules;

. the potential energy of the constituents of the system; for example, a crystalconsisting of dipolar molecules will experience a change in its potential energy as

an electric field is applied to the system;

' the internal energy stored in the form of molecular vibrations and rotations; and. the internal energy stored in the form ofchemical bonds that can be released

through a chemical reaction.

The total of all these forms of energy for the system of interest is given the symbol uand is called the internal energy.

The first law of thermodynamics is based on our experience that energy can be nei-ther created nor destroyed, if both the system and the surroundings are taken into

Equilibrium, Change, andReversibility

.

Comparing Work forReversible and IrreversibleProcesses

Determining AU andIntroducing Enthalpy, a NewState Function

Calculating q, w, LIJ, and,LH ior Processes InvolvingIdeal Gases

The Reversible AdiabaticExpansion and Compressionof an Ideal Gas

l4 CHAPTER 2 Heat, Work, lnternal Energy, Enthalpy, and the First Law ofThermodynamics

account. This law can be formulated in a number of equivalent forms. Our initial for-

mulation of this law is stated as follows:

The,internal e4ergy, U, of an isolated system is cortslant.

This form ofthe first law looks uninteresting, because it suggests that nothing happens

in an isolated system. How can the first law tell us anything about thermodynamic

processes such as chemical reactions? When changes in U occur in a system in contact

with its surroundings, LUn ot is given by

LU,n,or = LU ry,,"* I LU ,urour,rino, = 0

Therefore. the first law becomes

LU rn"*= -LU *orundiog,

(2.r)

(2.2)

Mechanicalstops

$:E*{iRil-;.}A system is shown in which compression

work is being done on a gas. The walls are

adiabatic.

For any decre ase of (J *,ten U ,u,,ou,di,g,mvst increase by exactly the same amount. For ex-

ample, if a gas (the syitem) is cooled, and the surroundings is also a gas, the temperature

of the surroundings must increase.

How can the energy of a system be changed? There are many ways to alter U, sev-

eral of which are discussed in this chapter. Experience has shown that all changes in a

closed system in which no chemical reactions or phase changes occur can be classified

as heat, work, or a combination ofboth. Therefore, the internal energy ofsuch a system

can only be changed by the flow ofheat or work across the boundary between the sys-

tem and surroundings. For example , U for agas can be increased by heating it in a flame

or by doing compression work on it. This important recognition leads to a second and

more useful formulation of the first law:

LU = q+w (2.3)

where q and w designate heat and work, respectively. We use AU without a subscript to

indicate the change in internal energy of the system. What do we mean by heat and

work? In the following two sections, we define these important concepts and distinguish

between them.The symbol A is used to indicate a change that occurs as a result of an arbitrary

process. The simplest processes ale those in which only one of P,V, or 7 changes. A con-

stant temperature process is referred to as isothermal, and the corresponding terms for

constants P andV arc isobaric and isochoric, respectively.

2,2 WorkIn this and the next section, we discuss the two ways in which the internal energy of a

system can be changed. Work in thermodynamics is defined as any quantity of energy

that "flows" across the boundary between the system and surroundings that can be used

to change the height of a mass in the surroundings. An example is shown in Figure 2.1.

We define the system as the gas inside the adiabatic cylinder and piston. Everything else

shown in the figure is in the surroundings. As the gas is compressed, the height of the

mass in the surroundings is lowered and the initial and final volumes are defined by the

mechanical stops indicated in the figure.Consider the system and surroundings before and after the process shown in

Figure 2.1, andnote that the height of the mass in the surroundings has changed. It is this

change that distinguishes work from heat. Work has several important characteristics:

j:3ffi ffi *#iaLyi+H€$sset:::-

P,,V,

Electricalgenerator

Pi,vr

Electricalgenerator

Final state

FIGlJ&tr H-3Current produced by a generator is used toelectrolyze water and thereby do work onlbe system as shown by the lowered masslinked to the generator. The gas producedin this process does P-V work on theswroundings, as shown by the raised masson the piston.

2.2 work t 5

' work is transitory in that it only appears during a change in state of the system andsuroundings. only energy, and not work, is associated with the initial and finalstates of the systems.

r The net effect ofwork is to change Uofthe system and surroundings inaccordance with the first law. If the only change in the surroundings is that a masshas been raised or lowered, work has flowed between the system and thesurroundings.

' The quantity ofwork can be calculated from the change in potential energy ofthemass, Epotmtinr = m?h, where g is the gravitational acceleration and /z is the changein the height of the mass, rn.

r The sign convention for work is as follows. If the height of the mass in thesurroundings is lowered, w is positive; ifthe height is raised, w is negative. Inshort, w > 0 if AU > 0. It is common usage to say that if w is positive, work isdone on the system by the surroundings. Ifw is negative, work is done by thesystem on the surroundings.

How much work is done in the process shown in Figure 2.1 ? Using a definition fromphysics, work is done when an object subject to a force, F, is moved through a distance,dl, according to the path integral

, = Jwar (.2.+S

Using the definition of pressure as the force per unit area, the work done in moving themass is given by

' = Jr.ar * - ll r";, dnat = * J

r*,"*, av (2.s)

The minus sign appears because of our sign convention for work. Note that the pressurethat appears in this expression is the external pressure, P ut",,ot, which need not equal thesystem pressure, P.

An example of another important kind of work, namely, electrical work, is shown inFigure 2.2 in which the contents of the cylinder is the system. Electrical current flowsthrough a conductive aqueous solution and water undergoes electrolysis to produce H, andO, gas. The curent is produced by a generator, like that used to power a light on a bicyclethrough the mechanical work of pedaling. As current flows, the mass that drives the gen-erator is lowered. In this case, the surroundings do the electrical work on the system. As aresult, some of the liquid water is transformed to H, and or. From electrostatics, the workdone in transporting a charge,4, through an electrical potential difference, @, is

Wekctical = Q6

For a constant current, I that flows for a time, t, e = It. Therefore,

wr*"*r, = I6t

The system also does work on the surroundings through the increase in the volume ofthe gas phase at constant pressure. The total work done is

w= wpv *wet"cr,icat = tOt - Jpu,"^ot dv = I6t- p"*,"^o,!dv = IOt - piUI-vi) (2.g)

other forms of work include the work of expanding a surface, such as a soap bub-ble, against the surface tension. Table 2.I shows the expressions for work for four dif-ferent cases. Each of these different types of work poses a requirement on the wallsseparating the system and surroundings. To be able to carry out the first three types ofwork, the walls must be movable, whereas for electrical work, they must be conductive.Several examples of work calculations are given in Example problem 2. l.

(2.6)

(21)

l6 CHAPTER 2 Heat, Work, lnternal Energy, Enthalpy, and the First Law ofThermodynamics

TABLE 2.I

EXAMPLE PROBLEM 2,I

a. Calculate the work involved in expanding 20.0 L of an ideal gas to a final :

volume of 85,0 L against a constant external pressure of'2.50 bar.

b. A water bubble is expanded from a radius of 1 .00 cm to a radius of 3,25 cm.,:, The surtace tension of water:is 71.99 N m-1. How muCh work is done in the

process?

c- A current of 3.20:4is passed through a heating coil for 30,0 s. The electricalpotential across the resistoris 14.5 v. calculate the work done on{he coil.

SolutionI

a. w=-J P"*^adV =-\o"_ot!I-Vi)

I =J.50 u*r{}*(8s.0 L- 20.0L)" 19-,1 Pl: ='16.3 kJ: bar L ----=b. Afactor of 2 is included in the calculation below because'a bribble'has an

inner and an outer surface: .

, t -$ r do ='zry4r{r? - ri|)

* =Bn x7r,99 Nm*r(3,2 Sz cmz _ 1.002 "lnr1"

1!i,li,cm"

= -1.73 JIc. w = 1 ddO= 6q = I$t =14.5Vx3.20Ax30.0 s = 1.39kJ

2.3 HeatHeatl is defined in thermodynamics as the quantity of energy that flows across theboundary between the system and surroundings because of a temperature difference be-

1 Heat is perhaps the most misused term in thermodynamics as discussed by Robert Romer lAmerican Jour-nal of Physics,69 (2001), 107-1091. In common usage, it is incorrectly referred to as a substance as in thephrase "Close the door; you're letting the heat outl" An equally inappropriate term is heat capacity (discussedin Section 2.4), because it implies that materials have the capacity to hold heat, rather than the capacity to storeenergy. We use the terms heat flow or heat transfer to emphasize the trarisitory nature of heat. However, youshould not think of heat as a fluid or a substance-

Rest of universe

FIGUR,E 2,SAn isolated composite system is created inrtich the surroundings to the system ofifierest are limited in extent. The wallswrrounding the inner water bath are rigid.

2.3 Heat l7

tween the system and the surroundings. Just as for work, several important characteris-tics of heat are of importance:

' Heat is transitory in that it only appears during a change in state ofthe system andsurroundings. only energy, aad not heat, is associated with the initial and finalstates of the system and the surroundings.

o The net effect ofheat is to change the internal energy ofthe system andsurroundings in accordance with the first law. If the only change in thesurroundings is a change in temperature of a reservoir, heat has flowed betweensystem and surroundings. The quantity ofheat that has flowed is directlyproportional to the change in temperature of the reservoir.

o The sign convention for heat is as follows. If the temperature of the surroundingsis lowered, q is positive; if it is raised, q is negative. It is common usage to say thatif 4 is positive, heat is withdrawn from the surroundings and deposited in thesystem. If 4 is negative, heat is withdrawn from the system and deposited in thesurroundings.

Defining the surroundings as the rest of the universe is impractical, because it is notrealistic to search through the whole universe to see if a mass has been raised or low-ered and if the temperature of a reservoir has changed. Experience shows that in gen-eral only those parts of the universe close to the system interact with the system.Experiments can be constructed to ensure that this is the case, as shown in Figure 2.3.Imagine that you are interested in an exothermic chemical reaction that is carried outin a rigid sealed container with diathermal walls. You define the system as consistingsolely of the reactant and product mixture. The vessel containing the system is im-mersed in an inner water bath separated from an outer water bath by a container withrigid diathermal walls. During the reaction, heat flows out of the system (q < 0), andthe temperature of the inner water bath increases to l. Using an electrical heater, thetemperature of the outer water bath is increased so thdt at all times, T ou,",

: T ,nn",. B"-cause of this condition, no heat flows across the boundary between the two water baths,and because the container enclosing the inner water bath is rigid, no work flows acrossthisboundary. Therefore, LU : S + w:0 + 0: \forthe composite systemmadeupof the inner water bath and everything within it. Therefore, this composite system is anisolated system that does not interact with the rest of the universe. To determine q andw for the reactant and product mixture, we need to examine only the composite systemand can disregard the rest of the universe.

To emphasize the distinction between q and w, and the relationship betsveen q, w,and Au, we discuss the two processes shown in Figure 2.4. They are carried out in anisolated system, divided into two subsystems, I and II. In both cases, system I consistssolely of the liquid in the beaker, and everything else including the rigid adiabatic wallsis in system II. We assume that the temperature of the liquid is well below its boilingpoint and that its vapor pressure is negligibly small. This ensures that no liquid is va-porized in the process, so that both systems are closed. System II can be viewed as thesurroundings for system I and vice versa.

In Figure 2.4a, a Bunsen burner fueled by a propane canister is used to heat theliquid (system I). The boundary between systems I and II is the surface that enclosesthe liquid, and heat can flow between systems I and II all across this wall. It is ob-served that the temperature of the liquid increases in the process. From general chem-istry, we know that the internal energy of a monatomic gas increases linearly with Z.This result can be generalized to state that U is a monotonically increasing functionof Zfor a uniform single-phase system ofconstant composition. Therefore, becausea?i > 0, au, > 0.

We next consider the changes in system II. From the first law, AUo: - AU, < 0.observe that the masses of propane and oxygen have decreased, and that the massesof Hro and co, have increased as a result of the combustion of propane. None of theweights in system II have been lowered as a result of the process, and the volume ofboth systems has remained constant. We conclude that no work was done by system II

TJrermometers

l8 CHApTER 2 Heat, Work, tnternal Energy, Enthalpy, and the First Law ofThermodynamics

on systeml,thatis, r,t:0. Therefore, if A{4 : q * w>0,q> 0. Notethatbecause

the wall separating systems I and II is diathermal, T, : T' and therefore

AZo > 0. If A{' > 0, why is AUo < 0? This decrease can be attributed to the difference

of the internal energy stored in the bonds of the combustion products CO, and HrO

compared with the reactants C.H, and Or. The internal energy U is not a monotonic

function of Ifor this system because the chemical composition changes in the process.

We now consider Figure 2.4b.In this case, the boundary between systems I and IIliesjust inside the inner wall ofthe beaker, passes across the open top ofthe beaker, and

just outside of the surface of the heating coil. Note that the heating coil is entirely in sys-

tem II. Heat can flow between systems I and II all across the boundary surface. Upon

letting the mass in system II fall, electricity flows through the heating coil. It is our ex-

perience that the temperature of the liquid (system I) will increase. We again conclude

that At/, > 0. However, it is not clear what values 4 and lt have.

To resolve this issue, consider the changes in system II of this composite system.

From the first law, LUn: - LU, < 0. We see that a mass has been lowered in sys-

tem II. Can it be concluded that work has been done by system II on system I? The

lowering of the mass in system II shows that work is being done, but it is being done

only on the heating coil, which is also in systemll. This is the case because the cur-

rent flow never crosses the boundary between the systems. Therefore, w : 0, where

w is the work done on system I. However, AUr > 0, so that if w : 0, q > 0'The con-

clusion is that the increase in U, is due to heat flow from system II to system I caused

by the difference between the temperature of the heating filament and the liquid. Note

that because AU, + AUn : 0, the heat flow can be calculated from the electrical

work,4:w:l$t.Why is LUo< 0 if A& > 0? The value of U' is the sum of the internal energy of all

parts of system II. The value of Z' increases, which increases the thermal part of Urr.

However, the potential energy of system II decreases because the mass is lowered in the

process. The net effect of these two opposing changes is that AUo < 0.

These examples show that the distinction between heat and work must be made care-

fully with a clear knowledge of the position and nature of the boundary between the sys-

tem and the surroundings.

(b)

s{ffi1$Rfi 3.4Two systems, I and II, are enclosed in a

rigid adiabatic enclosure. System I consists

solely of the liquid in the beaker for each

case. System II consists of everything else

in the enclosure. (a) Process in which the

liquid is heated using a flame. (b) Heating

using a resistive coil through which an

electrical current flows.

EXAMPLE PROBLEM 2.2

A heating coil is immersed in a 100-9 sample of HrO liquid at 100'C in an open

insulated beaker on a laboratory bench at I bar pressure. In th is process. I 07o ofthe liquid is converted to the gaseous form at a pressure of I bar. A current of2.00 Aflows through the heater from a 12.0-V battery for 1.00 x 103 s to effect

the transformation. The densities of liquid and gaseous water under these

conditions are997 and 0.590 kg t-t.respectively.

a. It is often useful to replace areal process by a model that exhibits the

important features of the process. Design a model system and surroundings,

like those shown in Figures 2. I and 2.i. thatwould allow you to measure

the heat and work associated with this translormation. For the model

system. define the system and surroundings as well as the boundary

between them.

b. Calculate q and w for the process.

c. How can you define the system lor the open insulated beaker on the

laboratory bench such that the work is properly described?

Solutiona. The model system is shown 1n the following figure. The cylinder walls and

the piston lorm adiabatic walls. The external pressure is held constant by a

suitable weight.

2.4 Heat Capacity I 9

Pexternat- 1.00 atm

b. In the system shown. the heat inpur to the liquid water can be equated withthe work done on the heating coil. Therefore.

q = IOt= 2.00 A x 12.0 Vx 1.00 x103 s = 24.0kJ

As the liquid is vaporized, the volume of the system increases at a constantexternal pressure. Therefore. the work done by the system on thesurroundings is

w = p",,"n,ot(vi -v,l =to5 pa xl++]pf$* 90.0x l0-3 kg -

100.0x l0r kgl| 0.590kgm 3 997kgm 3 997kgm-: )

= l.70 kJ

Note that the electrical work done on the heating coil is much larger thanthe P*Vwork done in the expansion.

c. Deline the system as the liquid in the beaker and the volume containing onlymolecules of Hro in the gas phase. This volume will consist of disconnectedvolume elemenis dispersed in the air above the laboratory bench.

2.4 Heat CapacftyThe process shown in Figure 2.4 provides a way to quantify heat flow in terms of theeasily measured electrical work done on the heating coil, w =I6t.The response of asingle-phase system ofconstant composition to heat input is an increase in 7. This is notthe case if the system undergoes a phase change, such as the vaporization of a liquid.The temperature of an equilibrium mixture of liquid and gas at the boiling point remainsconstant as heat flows into the system. However, the mass of the gas phase increases atthe expense of the liquid phase.

The response of the system to heat flow is described by a very imporlant thermody-namic property called the heat capacity. The heat capacity is a material-dependent prop-erty defined by the relation

c= lim q =dq

Ar-+07, *T dT(2.9)

ZO C H A pT E R 2 Heat, Work, tnternal Energy, Enthalpy, and the First Law of Thermodvnamics

100 2oo 300 400

TemPerature (K)

fif&#tr{qil tr.sThe r ariation ol C"- with temperature is

shown for Clr.

where C is in the SI unit of J K r. It is an extensive quantity that, for example, dou-

blesasthemassofthesystemisdoubled.often,themolarheatcapacity,C-,isusedin calculations and it is an intensive quantity with the units of J K-r mol-r' Experi-

mentally, the heat capacity of fluids ii measured by immersing a heating coil in the

fluid anJ equating the electrical work done on the coil with the heat flow into the sam-

ple. For sotids, the heating coil is wrapped around the solid. In both cases, the exper-

imental results must be conected for heat losses to the surroundings' The significance

of the notatio n dq fot an incremental amount of heat is explained in the next section'

Thevalueoftheheatcapacitydependsontheexperimentalconditionsunderwhichit is determined. The most iommon conditions are constant volume or constant pres-

sure, for which the heat capacity is denoted Cu and Ct' re:p^ecliyely'^Values of -C'*at

zgg.lsK for pure substances are tabulated in Tables 22 and 2.3 (see Appendix B, Data

Tables), and formulas for calculatin E C r,,atother temperatures for gases and solids are

listed in Tables 2.4 and 2.5, respectively lsee Appendix B' Data Tables)'

AnexampleofhowC.-dependson7isillustratedinFigure2.5forCl'.Tomakethefunctional form of Cr-(fiirnairstandable, we briefly discuss the relative magnitudes of

C-inthesolid,tiquid,andgaseousphasesusingamicroscopicmodel.Asolidcanbe*?nrti .t

"r-" r"i

"ii"terco-nnected harmonic oscillators, and heat uptake leads to the

excitations of the collective vibrations of the solid. At very low temperatures' these vi-

brations cannot be activated, because the spacing of the vibrational energy levels is large

compared to k7. As a consequence' energy cannot be taken up by the solid' Hence' Ct-

approaches zero as Z approa"he, zero' Foi the- solid' C.,, :ises TpidY .r-vlth

Zbecause the

thermal energy availabli as Tincreases is sufficient to iCtivate the vibrations of the solid'

The heat capacity increases discontinuously as the solid melts to form a liquid' This is

the case because the liquid retains all the local vibrational modes of the solid' and more

modes with low frequencies become available upon melting. Therefore, the heat capac-

ity of the liquid is greater than that of the solid. As the liquid vaporizes, the local vibra-

tional modes present in the liquid are converted to translations that cannot take up as

much energy a-s vibrations. Therefore, C.,, decreases discontinuously at the vaporization

temperature. The heat capacity in the gaieous state increases slowly with temperature as

the vibrational modes ofihe molecule-are activated. These changes tn C,n*can,be calcu-

lated for a specific substance using a microscopic model and statistical thermodynamics'

as will be discussed in detail in Chapter 33'

once the heat capacity of a varilty of different substances has been determined, we

have a convenient way to quantify heat flow. For example, at constant pressure, the heat

flow between the system and surroundings can be written as

I

6Lc

Y330

Ecio-

'u.,1

C,i,,",, \T)dT = - I ,';*unain<' 1T tdT

T,,,,,,

By measuring the tempefature change of a thermal reservoir in the surroundings at con-

stant pressurJ, 4"cante determined. In Equation (2.10), the heat flow at constant pres-

sure has been expressed both from the perspective of the system and from the perspective

of the surroundings. A similar equation can be written for a constant volume process' Wa-

ter is a convenient choice of maierial for a heat bath in experiments because Co is nearly

constantatthevalue4.l8Jg-1K 1or75.3Jmol-1K lovertherangefrom0"tol00'c'

EXAMPLE PROBLEM 2.3

The volume of a system consisting of an ideal gas decreases at constant

pressure. As a result, the temperature of a 1.50-kg water bath in the surroundings

increases by 14.2"C. Calculate 4"fot the system'

4r"..i

J{",,

12.10)llp =

100 200 300

2.4 Heat Capacity 2t

SolutionT*,:.r

n, =- l cy+dtut' 1r)n = -cy"1"4l'e' o,trso.j

'

= -1.50 kg x 4.18 Jg:l K-1 x14.2 K = -89.1kJ

FIGUR.E 2.SNot all the heat flow into the system canbe used to increase L,U in a constant

Iressure process, because the system doeswork on the surroundings as it expands.However, no work is done for constantvolume heating.

How are C "

and C urelated for a gas? Consider the processes shown in Figure 2.6in which a fixed amount of heat flows from the surroundings into a gas. In the con-stant pressure process, the gas expands as its temperature increases. Therefore, thesystem does work on the surroundings. As a consequence, not all the heat flow intothe system can be used to increase Au. No such work occurs for the correspondingconstant volume process, and all the heat flow into the system can be used to increaseAU. Therefore, dTo < dT, for the same heat flow dq. For this reason, C, > C, forgases.

The same argument applies to liquids aad solids as long as Vincreases with Z. Nearlyall substances follow this behavior, although notable exceptions occur such as liquid wa-ter between 4" and 0'c for which the volume increases as zdecreases. However, be-cause Av upon heating is much smaller than for a gas, the difference between c" andCrfor a liquid or solid is much smaller than for a gas.

The preceding remarks about the difference between c" and crhave been qualita-tive in nature. However, the following quantitative relationship, which will be proved inChapter 3, holds for an ideal gas:

C;=.C;:', ;,',.,.,e .,e1;,-,o;;,+R (2.tr)

Constantpressureheating

lnitial state Final state

Constanlvolumeheating

lnitial state Final state

22 CHAPTER 2 Heat, Work, tnternal Energy, Enthalpy, and the First Law of Thermodynamics

2.5 State Functions and Path FunctionsAn alternate statement of the first law is that AU is independent of the path between theinitial and final states and depends only on the initial and final states. We make this state-ment plausible for the kinetic energy, and the argument can be extended to the otherforms of energy listed in Section 2.1 . Consider a single molecule in the system. Imaginethat the molecule of mass m initially has the speed v,. We now change its speed incre-mentally following the sequence v, - vz * v. - v,,. The change in the kinetic energyalong this sequence is given by

(2.12)- 1*o t). () **,r - I *r" s')

P - )*o tt')7 .")--m(v)')

Even though v, and v, can take on any arbitrary values, they still do not influence theresult. We conclude that the change in the kinetic energy depends only on the initial andfinal speed and that it is independent of the path between these values. Our conclusionremains the same if we increase the number of speed increments in the interval betweenv, and v, to an arbitrarily large number. Because this conclusion holds for all moleculesin the system, it also holds for AU.

This example supports the conclusion that AU depends only on the final and initialstates and not on the path connecting these states. Any function that satisfies this condi-tion is called a state function, because it depends only on the state of the system and notthe path taken to reach the state. This property can be expressed in a mathematical form.Any state function, for example, U, must satisfy the equation

(2.13)

This equation states that in order for AU to depend only on the initial and final statescharacterized here by i andf, the value of the integral must be independent of the path.If this is the caseo u canbe expressed as an infinitesimal quantity, du, that when inte-grated, depends only on the initial and final states. The quantity dU is called an exactdifferential. We defer a discussion of exact differentials to Chapter 3.

It is useful to define a cyclic integral, denoted by rhe symbol p. as applying to a cyclicpath such that the initial and final states are identical. For Uor'any other state function,

+rU;.*{{ =''0 (2.r4)

because the initial and final states are the same in a cyclic process.The state of a single-phase system at fixed composition is characterized by any two

of the three variables P, T, and V. The same is true of U. Therefore, for a system of fixedmass, Ucanbewritteninanyof thethreeforms U(V,T),U(P,D,orU(P,I).Imaginethata gas of fixed mass characterizedby V,and Tris confined in a piston and cylinder sys-tem that is isolated from the surroundings. There is a thermal reservoir in the sur:round-ings at a temperature Tr) Tr. We first compress the system from an initial volume V, toa final volume V, using a constant external pressure Pu,",not. The work is given by

vJI

-P*,^ot J dV = -P,n"^a(Vf -V)= *P",,"*otLV (2.I5)

vi

'. .:,i . J..,r .. r: . . :. .

AU=JdU=Uf -Ui.. ., }i':l' , rl

vJIw=-l P ,dV =J enemqt

vi

lnitial state

lntermediate state

Final state

FICUR.H A.PA lvstem consisting of an ideal gas iscnntained in a piston and cylinder;nusembly. The gas in the initial state yr,f1mr compressed to an intermediate state,

mfiercby the temperature increases to theratue f,. It is then brought into contact

'nm.itl a thermal reservoir at 7r, leading to affiurther rise in temperature. The final state

* l.'L'

2.5 State Functions and Path Functions 23

Because the height of the mass in the surroundings is lower after the compression (see

Figure 2.7), w is positive and U increases. Because the system consists of a uniform sin-gle phase, [./ is a monotonic function of Z, and 7 also increases. The change in volumeAV has been chosen such that the temperature of the system T, after the compressionsatisfies the inequality Tt <Tz <73.

We next let an additional amount of heat q flow between the system and surround-ings at constant V by bringing the system into contact with the reservoir at temperatureZ.. The final values of 7 and V after these two steps are T, and Vr.

This two-step process is repeated for different values of the mass. In each case thesystem is in the same final state characterizedby the variables Vrand 2.. The sequence

of steps that takes the system from the initial stateVrTtto the final state V'2, is refer:red

to as a path. By changing the mass, a set of different paths is generated, all of whichoriginate from the state VpTp and end in the state V,Tr. According to the first law, AUfor this two-step process is

L,U =U(Tr,Vr)-U(71,v)- q+w (2.16)

Because AU is a state function, its value for the two-step process just described is thesarne for each of the different values of the mass.

Are q and w also state functions? For this process,

w = -P"t"*o,LV \2.r7)

&fld P",,",,otis different for each value ofthe mass, or for each path. Therefore, w is alsodifferent for each path; one can take one path from V,,{ toVr,Trand a different pathfromV,T, back to V,, {. Because the work is different along these paths, the cyclic in-tegral of work is not equal to zero. Therefore, w is not a state function.

Using the first law to calculate q for each of the paths. we obtain the result

A= LU -w = LU + P",r"*otLV (2. l8)

Because AU is the same for each path, and w is different for each path, we conclude that

4 is also different for each path. Just as for work, the cyclic integral ofheat is not equalto zero. Therefore, neither qnor w are state functions, and they are called path functions.

Because both q and w are path functions, there are no exact differentials for workand heat. Incremental amounts of these quantities are denoted by dq and dw, rcther thandq and dw,lo emphasize the fact that incremental amounts of work and heat are not ex-act differentials. This is an important result that can be expressed mathematically.

*q, and *,wi (2.re)

Infact,therearenosuchquantitiesasL,q,qr,q wtdLw,wr,w,.Thisisanimportantre-sult. One cannot refer to the work or heat possessed by a system, because these conceptshave no meaning. After a process involving the transfer of heat and work between thesystem and surroundings is completed, the system and surroundings possess internal en-ergy, but neither possesses heat or work.

Although the preceding discussion may appear pedantic on first reading, it is im-pofiant to use the terms work and heat in a way that reflects the fact that they are notstate functions. Examples of systems of interest to us are batteries, fuel cells, refrigera-tors, and internal combustion engines. In each case, the utility of these systems is thatwork and/or heat flows between the system and surroundings. For example, in a refrig-erator, electrical energy is used to extract heat from the inside of the device and to re-lease it in the surroundings. One can speak of the refrigerator as having the capacity toextract heat, but it would be wrong to speak of it as having heat. In the internal com-bustion engine, chemical energy contained in the bonds of the fuel molecules and in O,is released in forming CO, and HrO. This change in internal energy can be used to rotate

ILw+ldw+w,

IJ

i

ftL,q*ldq+q,

t

24 C H A PT E R 2 Heat, Work, lnternal Energy, Enthalpy, and the First Law of Thermodynamics

^6zo-o-

@

1o

the wheels of the vehicle, thereby inducing a flow of work between the vehicle and thesurroundings. One can refer to the capability of the engine to do work, but it would beincorrect to refer to the vehicle or the engine as containing or having work.

2.6 Equilibrium, Change, and ReversibilityThermodynamics can only be applied to systems in internal equilibrium, and a require-ment for equilibrium is that the overall rate of change of all processes such as diffusionor chemical reaction be zero. How do we reconcile these statements with our calcula-tions of 4, w, and AU associated with processes in which there is a macroscopic changein the system? To answer this question, it is important to distinguish between the systemand surroundings each being in internal equilibrium, and the system and surroundingsbeing in equilibrium with one another.

We first discuss the issue of internal equilibrium. Consider a system made up of anideal gas, which satisfies the equation of state, P = nRTM . All combinations of p, v, andz consistent with this equation of state form a surface in p-v-T space as shown inFigure 2.8. All points on the surface correspond to equilibrium states of the system.Points that are not on the surface do not correspond to any physic ally realizable state ofthe system. Nonequilibrium situations cannot be represented on such a plot, because p,V , and Z do not have unique values for a system that is not in internal equilibrium.

Next, consider a process in which the system changes from an initial state character-r1elby P,, Vr

TndT to

_afinal state characterizedby prVr, and las shown in Figure 2.g.

If the rate of change of the macroscopic variables is negiigibly imatl, the system passesthrough a succession of states of internal equilibrium as it goes from the initial to the fi-nal state. Such a process is called a quasi-static process, and internal equilibrium ismaintained in this process. If the rate of change is sufficiently large, the rates of diffusionand intermolecular collisions may not be high enough to maintain the system exactly ina state of internal equilibrium. However, thermodynamic calculations for such a processare still valid as long as it is meaningful to assign a single value of the macroscopic vari-ables P, v,T, and concentration to the system undergoing change. For rapid changes, lo-cal fluctuations may occur in the values of the macroscopic variables throughout thesystem. If these fluctuations are small, quantities such as work and heat can still be cal-culated using mean values of the macroscopic variables. The same considerations holdfor the surroundings.

The concept of a quasi-static process allows one to visualize a process in which thesystem undergoes a major change in terms of a directed path along a sequence of statesin which the system and surroundings are in internal equilibrium. We next distinguish be-tween two very important classes of quasi-static processes, namely, reversible and irre-versible processes. It is useful to consider the mechanical system shown in Figxe 2.9when discussing reversible and irreversible processes. Because the two masses have thesame value, the net force acting on one end of the wire is zero, and the masses will notmove. If an additional mass is placed on either of the two masses, the system is no longerin mechanical equilibrium and the masses will move. In the limit that the incrementalmass approaches zero, the velocity at which the initial masses move approaches zero. Inthis case, one refers to the process as being reversible, meaning that the direction of theprocess can be reversed by placing the infinitesimal mass on the other side of the pulley.

Reversibility in a chemical system caa be illustrated by liquid water in equilibriumwith gaseous water surrounded by a thermal reservoir. The system and surroundings areboth at temperature Z. An infinitesimally small increase in Z results in a small increaseof the amount of water in the gaseous phase, and a small decrease in the liquid phase.An equally small decrease in the temperature has the opposite effect. Therefore, fluctu-ations in Z give rise to corresponding fluctuations in the composition of the system. Ifan infinitesimal opposing change in the variable that drives the process (temperature inthis case) causes a reversal in the direction ofthe process, the process is reversible.

12.5 10 7.5 5.0 2.5v(L)

pi6t"rRg ?.sAll combinations of pressure, volume, andtemperature consistent with 1 mol of anideal gas lie on the colore{ surface. Allcombinations of pressure and volumeconsistent with Z: 800 K and allcombinations of pressure and temperatureconsistent with a volume of 4.0 L are shownas black curves that lie in the P-V-Tsurface. The third curve coffesponds to apath between an initial state I and a finalstate/that is neither a constant temperaturenor a constant volume path.

1kg

F G:J=C p"S"!trlftnm nrasses of exactly 1 kg each are

:mmmei-red by a wire of zero mass runningmmma fictionless pulley. The system is inm-hnnical equilibrium and the masses

mc medonary.

2.7 Comparing Work for Reversible and lrreversible Processes 25

If an infinitesimal change in the driving variable does not change the direction of theprocess, one says that the process is irreversible. For example, if a large stepwise tem-perature increase is induced in the system using a heat pulse, the amount of water in the

gas phase increases abruptly. In this case, the composition of the system cannot be re-turned to its initial value by an infinitesimal temperature decrease. This relationship ischaracteristic of an irreversible process. Although any process that takes place at a rapidrate in the real world is irreversible, real processes can approach reversibility in the ap-

propriate limit. For example, by small variations in the electrical potential in an electro-chemical cell, the conversion ofreactants to products can be carried out nearly reversibly.

2.7 Comparing Work {or Reversibleand Irreversible Processes

We concluded in Section 2.5 that w is not a state function and that the work associated

with a process is path dependent. This statement can be put on a quantitative footing bycomparing the work associated with the reversible and irreversible expansion and com-pression of an ideal gas. This process is discussed next and illustrated in Figure 2.10.

Consider the following irreversible process, meaning that the internal and externalpressures are not equal. A quantity of an ideal gas is confined in a cylinder with a

weightless movable piston. The walls of the system are diathermal, allowing heat toflow between the system and surroundings. Therefore, the process is isothermal at thetemperature of the surroundings, Z. The system is initially defined by the variables 7,P', and Vr. The position of the piston is determined by P.*,.,,ot: P' which can be

changed by adding or removing weights from the piston. Because they are moved hor-izontally, no work is done in adding or removing the weights. The gas is first expandedat constant temperature by decreasing P

",,"*ot abruptly to the value P, (weights are re-

moved), where P, < Pr. A sufficient amount of heat flows into the system through thediathermal walls to keep the temperature at the constant value Z. The system is now inthe state defined by T, P, andV, where V, > V,. The system is then returned to its orig-inal state in an isothermal process by increasing P

",,",,ot abruptly to its original value

P, (weights are added). Heat flows out of the system into the surroundings in this step.

The system has been restored to its original state and, because this is a cyclic process,AU : 0. Ar" e,o,o,andw,o,otalso zero for the cyclic process? The total work associatedwith this cyclic process is given by the sum of the work for each individual step:

wtotol =\-P**-r,,LVi = w"rporrion * wco-pr"rrio, = -PdVz -Vr) - P\(v\ -V) (2'20)

= -(Pz- Pr)x(Vr-Vr) > 0 because P, < P, and V, > V,

The relationship between P andV for the process under consideration is shown graphi-cally in Figure 2.10, in what is called an indicator diagram. An indicator diagram isuseful because the work done in the expansion and contraction steps can be evaluatedfrom the appropriate area in the figure, which is equivalent to evaluating the integral* = -l Pu*-o, dV. Note that the work done in the expansion is negative because Ay > 0,

and that done in the compression is positive because Ay < 0. Because Pr.< P,., the mag-nitude of the work done in the compression process is more than that done in the ex-pansion process lnd w ,o,ot > 0. What can one say about e,o,o,? The first law states thatbecause LU : qro,o, ! w rorot

: O, erorot 1 O.

The same cyclical process is carried out in a reversible cycle. A necessary conditionfor reversibility is that P = P*n*ot at every step of the cycle. This means that P changes

during the expansion and compression steps. The work associated with the expansion is

wetqansion= -l rn**o,av = -J r av =-rRrl+=*nnrnft e.2r)

26 CHAPTER 2 Heat,Work, lnternal Energy,Enthalpy,andtheFirstLawofThermodynamics

This work is shown schematically as the red area in the indicator diagram of Figure 2. 1 1 .

If this process is reversed and the compression work is calculated, the following re-sult is obtained:

wcoarpwion = -nRrh+ (2.22)

It is seen that the magnitude of the work in the forward and reverse processes is equal.The total work done in this cyclical process is given by

v^ \'= -nRTln2- nRTInav. v''1 '2

v-v=-nRT ln2+ nRT62=0

vl

Therefore, the work done in a reversible isothermal cycle is zero. Because 6ry : q -f wis a state function, e=-w =0 for this reversible isothermal process. Looking at theheights of the weights in the surroundings at the end of the process, we find that they arethe same as at the beginning of the process. As will be shown in Chapter 5, q andw arenot equal to zero for a reversible cyclical process if Iis not constant.

w=w +we\pqnsron compt?.t\ton(2.23)

v1

v2v1

V

s3€#KK 3.i#The work for each step and the total workcan be obtained from an indicator diagram.For the compression step, w is given by thetotal area in red and yellow; for theexpansion step, w is given by the red area.

The arrows indicate the direction ofchange in Vin the two steps. The sign of wis opposite for these two processes. Thetotal work in the cycle is the yellow area.

v(L)

F#*Ajsg e"€xIndicator diagram for a reversible process.

Unlike Figure 2.10, the areas under theP-V curves are the same in the forwardand reverse directions.

EXAMPLE PROBLEM 2.4

In this example, 2.00 mol of an ideal gas undergoes isothermal expansion alongthree different paths: (1) reversible expansion from P : 25.0 bar and V : 4.50Lto Pr:4.50 bar; (2) a single-step expansion against a constant external pressureo14.50 bar. and (3) a two-step expansion consisting initially ofan expansionagainst a constant external pressure of 1 1.0 bar until P : P *ro,no, followed by an

expansion against a constant external pressure ol 4.50 Uar untii-p : P.,,",nut.

Calculate the work for each of these processes. For which of the irreversibleprocesses is the magnitude of the work greater?

SolutionThe processes are depicled in the following indicator diagram:

^15rdo\

10

o51015zo2sv{L}

We first calculate the constant temperature at which the process is carried oul.the final volume, and the intermediate volume in the two-step expansion:

P.V.T- I I

-

2.7 Comparing Work for Reversible and lrreversible Processes 27

25.0 bar x 4.50 L=6ll KnR 8.314x104Lbarmol-l K-l x2.00mol

nRTv. ---rP"

T

8.314 x 10-2 L bar mol-1 K-1 x 2.00 mol x 677 K

V.' lnl

4.50bar

nRT 8.314 x 10-2 L bar mol-r K-r x 2.00 mol x 677 KPrn 11.0bar

=25.0L

=10.2L

The work of the reversible process is given by

V,w - -nRT, ln ''V,

=-2.00 mol x8.314 Jmol-r K I x611r xm 15-9 | =-19.3x l0r J4.50 L

We next calculate the work of the single-step and two-step ireversibleprocesses:

w-.,,-,,=-P,LV =4.50b.r l0tPu"(25.00L-4.50r,r. l0 3m3stnqte extemqt DAT L

= -9.23 x 103 J

' =-r r.0 bar x lOs Pa xrl.TL - 4.50r,t l0 3 m3w*o-'"p = -Prr,"^olLU bar L

- 4.50 u* , 19t Pu x (25.00 L - t0.2L) x !-3 nn3

= -rrr r, ror, o* '\ \'J'vv ! " L

The magnitude of the work is greater for the two-step process than for thesingle-step process, but less than that for the reversible process.

Example Problem 2.4 suggests that the magnitude of the work for the multistep ex-pansion process increases with the number of steps. This is indeed the case, as shown inFigure 2.I2.Imagine that the number of steps n increases indefinitely. As n increases, thepressure difference P,",

",noL- P for each individual step decreases. In the limit that n + -,

the pressure difference P",,.,not- P + 0, and the total area ofthe rectangles in the indica-tor diagram approaches the area under the reversible curve. Therefore, the irreversibleprocess becomes reversible and the value ofthe work equals that ofthe reversible process.

By contrast, the magnitude of the irreversible compression work exceeds that of thereversible process for finite values of n and becomes equal to that of the reversibleprocess as n + -. The difference between the expansion and compression processes re-sults from the requirement that P",,",not< P at the beginning of each expansion step,whereas P.*t",,ot > P at the beginning of each compression step.

On the basis of these calculations for the reversible and irreversible cycles, we in-troduce another criterion to distinguish between reversible and irreversible processes.

Suppose that a system undergoes a change through one or a number ofindividual steps,and that the system is restored to its initial state by following the same steps in reverseorder. The system is restored to its initial state because the process is cyclical. Ifthe sur-roundings are also returned to their original state (all masses at the same height and allreservoirs at their original temperatures), the process is reversible. If the sur:roundingsare not restored to their original state, the process is irreversible.

One is often interested in extracting work from a system. For example, it is the ex-pansion of the fuel-air mixture in an automobile engine upon ignition that providesthe torque that eventually drives the wheels. Is the capacity to do work similar forreversible and irreversible processes? This question is answered using the indicator

28 cHApTER 2 Heat,work, lnternal Energy, Enthalpy, and the First Law ofThermodynamics

safr4.$iRK H.E;The work done in an expansion (red area)

is compared with the work done in a

multistep series of irreversible expansion

processes at constant pressure (yellow area)

in the top panel. The bottom panel shows

analogous results for the compression.

where the area under the black curve is

the reversible compression work. Note

that the total work done in the irreversible

expansion and compression processes

approaches that of the reversible process as

the number of steps becomes large.

lw,n, " o ir, tnl 2

lw i o* *, it ul

By contrast, for the compression step,

lw,* " ^iw "l

llw io* u, ir, t "l

diagrams of Figures 2.10 and 2.11 fot the specific case of isothermal expansion work,

noting that the work can be calculated from the alea under the P-V curve' We com-

pu." th" work for expansion from V, to Vrin a single stage at constant pressure to that

ior the reversible case. For the singie-stage expansion, the constant external pressure

is given by P"on*ot = nRT lVz. However, if the expansion is carried out reversibly, the

,yr=t"- pr"rr"i" ir always greater than this value. By comparing the areas in the indi-

cator diagrams of Figure 2.I2,it is seen that

(2.24)

(2.2s)

The reversible work is the lower bound for the compression work and the upper bound

for the expansion work. This result for the expansion work can be generalized to an im-

2.8 Determining AU and lntroducing Enthalpy, a New State Function 29

portant statement that holds for all forms of work: The maximum work that can be ex-

tracted from a process between the same initial and final stetes is that obtained underreve rs i bl e con d i t i on s.

Although the preceding statement is true, it suggests that it would be optimal to op-erate an automobile engine under conditions in which the pressure inside the cylihdersdiffers only infinitesimally from the external atmospheric pressure. This is clearly notpossible. A practical engine must generate torque on the drive shaft, and this can onlyoccur if the cylinder pressure is appreciably greater than the external pressure. Similarly,a battery is only useful if one can extract a sizable rather than an infinitesimal current.To operate such devices under useful irreversible conditions, the work output is less thanthe theoretically possible limit set by the reversible process.

2.8 Determining AU and IntroducingEnthalpy, a New State Function

How can the AU for a thermodynamic process be measured? This will be the topic ofChapter 4, in which calorimetry is discussed. However, this topic is briefly discussedhere in order to enable you to carry out calculations on ideal gas systems in the end-of-chapter problems. The first law states that A,U - q + w .Imagine that the process is car-ried out under constant volume conditions and that nonexpansion work is not possible.Because under these conditions, , = -l P"rn*odV = 0,

A,U = ev (2.26)

Equation (2.26) statesthat L'U can be experimentally determined by measuring the heatflow between the system and surroundings in a constant volume process.

What does the first law look like under reversible and constant pressure conditions?We write

dU : da p - P",,"*otdV = dqp - P dV

and integrate this expression between the initial and final states:

(2.21)

-u, = J

aq, - J r av = ep - ervr - nvi) (2.28)

Note that in order to evaluate the integral involving P, we must know the functionalrelationship P(L/), which in this case is \ = P, = P. Rearranging the last equation, weobtain

(U r + PrVr)-(U,+ P,V,)= ee (2.2e)

Because P, V, and U arc all state functions, U + PV is a state function. This new statefunction is called enthalpy and is given the symbol 11. [A more rigorous demonstrationthat 11 is a state function can be given by invoking the first law for U and applying thecriterion of Equation (3.5) to the product PV.l

II_U+PV (2.30)

As is the case for U, H has the units of energy, and it is an extensive property. Asshown in Equation (2.29), AiI for a process involving only P-V work can be deter-mined by measuring the heat flow between the system and surroundings at constantpressure:

Jau =u,

A,H = ep (2.3t)

30 C HAPTER 2 Heat, Work, lnternal Energy, Enthalpy, and the First Law of Thermodynamics

oaa,q,,+frflry,

tSto2.1 ldeal Gas under Constant

Pressure or Volume

This equation is the constant pressure analogue of Equation (2.26). Because chemicalreactions are much more frequently carried out at constant P than constant V, the energy

change measured experimentally by monitoring the heat flow is AF1 rather than AU.

2.9 Calculating e, w, LU, and AH forProcesses Involving Ideal Gases

In this section, we discuss how AU and A11, as well as q and w, can be calculated fromthe initial and final state variables if the path between the initial and final state is known.The problems at the end of this chapter ask you to calculate q, w, LU, and A11 for sim-ple and multistep processes. Because an equation of state is often needed to carry outsuch calculations, the system will generally be an ideal gas. Using an ideal gas as a sur-

rogate for more complex systems has the significant advantage that the mathematics is

simplified, allowing one to concentrate on the process rather than the manipulation ofequations and the evaluation ofintegrals.

What does one need to know to calculate AU? The following discussion is restricted

to processes that do not involve chemical reactions or changes in phase. Because Uis a state

function, AU is independent of the path between the initial and final states. To describe a

fixed amount of an ideal gas (i.e., n is constant), the values of two of the variables P,V,arrdZmust be known. Is this also ffue for AU for processes involving ideal gases? To answer

this question, consider the expansion of an ideal gas from an initial state V'7, to a finalstateV,Tr.We first assume that Uis a function of both V andT.Is this assumption correct?

Because ideal gas atoms or molecules do not interact with one another, U will not depend

on the distance between the atoms or molecules. Therefore, U is not a function of V, and

we conclude that L,U must be a function of Zonly for an ideal gas, AU = LU(T\We also know that for a temperature range over which Cu. is constant,

{tl:t4.r.=efi{r -4} (2.32)

Is this equation only valid for constant V? Because U is a function of 7 only for an idealgas, Equation (2.32) is also valid for processes involving ideal gases in which V is notconstant. Therefore, if one knows C v T r, and T r, A U can be calculated, regardless of the

path between the initial and final states.

How many variables are required to define AII for an ideal gas? We write

LH = L,U(T)+ A(PY) = LU(T)+ L(NRT)= AH(T) (2.33)

It is seen that AII is also a function of Zonly for an ideal gas. In analogy to Equation (2.32),

Ar.=4;:4(+ *'4) (2.34)

Because A11 is a function of Zonly for an ideal gas, Equation (2.34) holds for all processes

involving ideal gases, whether P is constant or not as long as it is reasonable to assume

that C. is constant. Therefore, if the initial and final temperatures are known or can be cal-

culated, and if C" and C. are known, AU and AI1 can be calculated regardless of the pathforprocesses involving ideal gases using Equations (2.32) and(2.34), as long as no chem-

ical reactions or phase changes occur. Because U and 11 are state functions, the previousstatement is true for both reversible and irreversible processes. Recall that for an ideal gas

Cr-C, =nR,sothatif one of Crand C.isknown,theothercanbereadilydetermined.Wenextnotethatthefirstlawlinks q,w,andAU. If anytwoof thesequantitiescan

be calculated, the first law can be used to calculate the third. In calculating work, oftenonly expansion work takes place. In this case one always proceeds from the equation

* = -t P",,"-o, dv (2.3s)

&##i. r,'{'{r

2.9 Calculating q, w, LlJ, and lH for Processes lnvolving ldeal Cases 3l

This integral can only be evaluated if the functional relationship between P .*,.,,ot afld Vis known. A frequently encountered case is P

",,"_ot: constant,'such that

w = -P"o"-o1(V.f -Vi) (2.36)

Because P",,n ^or

* P, the work considered in Equation (2.36) is for an irreversible process.

A second frequently encountered case is that the system and external pressure dif-fer only by an infinitesimal amount. In this case, it is sufficiently accurate to writeP

"*t",,nt: P, and the process is reversible:

##s2.2 Reversible Isothermal

Com'pression of an Ideal Gas

a I Reversible Isobaric Compressionand Expansion oI an ldeal Gas

a4 Isochoric Heating and Coolingof an Ideal Gas

2.5 Reversible Cyclic Processes

(2.31)

This integral can only be evaluated if Iis known as a function of V. The most commonlyencountered case is an isothermal process, in which Z is constant. As was seen in Sec-

tiorr2.2. for this case

'=-l"Y dv

w=-nRTldv =-nnrnLJV Vi(2.38)

In solving thermodynamic problems, it is very helpful to understand the process

thoroughly before starling the calculation, because it is often possible to obtain the valueof one or more of 4, w, and AU and AI1 without a calculation. For example, A,U : A,H: 0 for an isothermal process because AU and AII depend only on 7. For an adiabaticprocess, q : 0by definition. If only expansion work is possible, u, : 0 for a constantvolume process. These guidelines are illustrated in the following two example problems.

EXAMPLE PROBLEM 2.5

A system containing 2.50 mol of an ideal gas for which Cr,,_ : 20.19 J mol-r K 1

is taken through the cycle in the lollowing diagram in the direction indicated bythe amows. The curved path corresponds ro PV = nRT. wherc f =Tt =7..

d.oa-

a.

b.

v(L)

Calculate q, w, atd AU and A11 for each segment and for the cycle.

Calculate q, w, ard AU and A11 for each segment and for the cycle in whichthe direction of each process is reversed.

32 cHAprER 2 Heat, work, rnternar Energy, Enthalpy, and the Fjrst Law ofrhermodynamics

SolutionWe begin by asking whether we can evaluate q, w, LU ,or A11 for any of thesegments without any calculations. Because the path between ,tut", i and 3 isisothermal. Au and A11 are zero for this segment. Therefore, from the first law.4z-:t= -lf3-l . For this reason, we onry need to calculate one of these twoquantities. Because AV : 0 along the path between states 2 and 3, wr--. : 0.Therefore, LUr_, = Q2-3. Again, we only need to calculate one of th'ese twoquantities. Because the total process is cyclic, the change in any state function iszero. Therefore, AU : L,H :0 for the cycle, no matter which direction ischosen. We now deal with each segment individually.

Segment 1- 2

The values o! n, P, and V, and p " and Vrare known. Therefore, T, and T, can

be calculated using the ideal gas liw. we-use these temperatures to calculate Auas follows:

\U,-, = nC, .,,,172 - Ttl = ffCt."'

, p,v, - p,V, IttRrr

- 2'50molx20'79Jmol-lK-l

v(1 AAt^-,,.< Ar 1..1--2.50 mol x 0.083 14 L bar K-r mol-l

= 99.6kJ

x (16.6bar x 25.0 L- 16.6 bar x 1.00 L)

= -39.8 kJ

Using the first law

q = LU - w- =99.6 kJ + 39.g kJ = 139.4 kJ

We next calculate T, andthen AH,_, :

16.6 bar x 25.0L2.50mol x 0.083 l4 L barK-l mol-l

= 2.00 x 103 K

We next calculate Q = {I 6.6 bar x 1.00 L

2.50 mol x 0.083 14 L barmol-r K-l =19.9 K

LH r-, = LU r-r+ A(PY) = LfJ r-z + nR(T, - T,)

=99.6x103.I + 2.5molxg.314 Jmol-r K{ x(2000 K_ .,g.gK) = 139.4kJSegment 2+ 3

As noted above, w : 0, and

LUr r,=Qz-z=C,(\_ T2)

The process takes place at constant pressure, so

rr : -P",*-ot(v2 -v1)= -i6.6bar r l05I_m-2

x (25.0 x 10-3 m3 - 1.00 x 10-3 m3)DAI

= 2.50 mot x20.19 J mol 1K_i(79.9 K _ 2000 K)= -99.6 kJ

The numerical result is equal in magnitude, but opposite in sign to AU,-, becauseT, : T, For the same reason! LHr-, = *LH-.z.

Segment 3 , lFor this segment, A%-l

_ 0 and AH:*r = 0 as noted earlier and w3-t = _43->1.

Because this is a reversible isothermal compression,

PV^Tll'2- nR -

PV.'r- tt_

"- nR

2.9 Calculating q, w, LIJ, and AH for Processes lnvolving ldeal Cases 33

V.w. .--nRTln '-r+r v3

= 5.35 kJ

The results for the individual segments and for the cycle in the indicated directionare given in the following table. If the cycle is traversed in the reverse fashion,the magnitudes of all quantities in the table rema.in the same. but all signs change.

-2.50 mol x 8.3 t4 J mol-l K-1 x lg.g K* hr 1'09 r-19-t"'-- 25.0 x 10-3m3

Path 4 GJ) w (kJ) AU (KJ) aH(kJ)

l-+2 139.4 -39.8 99.6 139.4

2-+3 -99.6 0 -99.6 -139.4

3--+1 -5.3s s.35 0 0

Cycle 34.5 -34.5 0 0

EXAMPLE PROBLEM 2.6

ln this example.2.50 mol of an ideal gas with C,,. :12.47 J mol-r K-r isexpanded adiabatically against u

"onriun, .*r.*ui p..tture of 1.00 Uar. it"r. initiut

temperature and pressure of the gas arc 325 K and 2.50 bar, respectively. Thefinal pressure is 1.25 bar. Calculate the final temperature, q, w, L(J, and Af1.

SolutionBecause the process is adiabatic, q : 0. and L,tl : w.Therefore.

A,U = nCu,*(Tf -Ti) = -p"r,"*o1(Vy -Vi)Using the ideal gas law,

(r. r)nCu.,(T, -7,) = -nM ,l J -:L I' e*e'o'lP'

P' )/\/( nRP ,) ( nRP \

Trl nC,.^ * Y l= r,l nC,., *T I\.rtr\t'i)

,, =r,l

t,

RP"*-ot

P,

RPexrcmql

Pf

+

+

Co,*

Cu.^

8.3 14 J mol-1 K-1 x 1.00 bar

=3r5KtI12.4'l J mol-lK*1 +

2.50bar=268 K

12.47 I mol-l K-r + 8.314 J mol-1 K-1 x 1.00bar

I.25bw

We calculate LU : w from

L,lJ = nC,, ^(7, -T,l=2.5 molx I 2.41 J mol-tK-r x(268 R-325 K) = -1.78 kJv.m' I r.

Because the temperature falls in the expansion, the internal energy decreases:

LH = L(l+ A(PV) = L(J + nR(T, - T1)

=-1.78x103 J + 2.5mol x8.314 Jmol-rK-l x(268 K -325 K) =-2.96kJ

34 C H A PT E R 2 Heat, Work, lnternal Energy, Enthalpy, and the First Law of Thermodynamics

2.lo The Reversible Adiabatic Expansionand Compression of an Ideal Gas

The adiabatic expansion and compression of gases is an important meteorological process.

For example, the cooling of a cloud as it moves upward in the atmosphere can be modeled

as an adiabatic process because the heat ffansfer between the cloud and the rest of the at-

mosphere is slow on the timescale of its upward motion.Consider the adiabatic expansion of an ideal gas. Because q : O, the first law takes

the form

A,U = w or CrdT = -P"*r"^oldV

For a reversible adiabatic process, P : P ",,"*ot,

dad

CvdT =-rRr+ or. equivalenrly. Cr+=-rR+V 'T V

Integrating both sides ofthis equation between the initial and final states,

(2.4t)T, vt

If C,, is constant over the temperature interval Tr-7,, then

*ssws*@

2.6 Reversible Adiabatic Heating andCooling of an Ideal Gas

10 20 30 40 50 60

Volume (L)

FeffiL'ffiffi X,tSTwo systems containing 1 mol of N, have

the same P and Vvalues at 1 atm. The red

curye coffesponds to reversible expansionand compression about P : I atm underadiabatic conditions. The yellow curvecorresponds to reversible expansion andcompression aboutP : 1 atm underisothermal conditions.

T" V"G, hl - -nRln I,TV

t"[?J=-tr-4*[5) o,, "ruiua"nuv, +=[+j"

EXAMPLE PROBLEM 2.7

A cloud mass moving across the ocean at an altitude of 2000 m encounters a

coastal mountain range. As it rises to a height of 3500 m to pass over themountains, it undergoes an adiabatic expansion. The pressure at 2000 and

(2.3e)

(2.40)

(2.42)

1

i'jr,T=-,o'j+

Because C, - C, = nR for an ideal gas, Equation (2.42) can be written in the form

(2.43)

where 7 : C p.*lCv.*. Substituting Tr lf,: PrV, f P,V inthe previous equation, we obtain

1v,' = Ptv/ (2.44)

for the adiabatic reversible expansion or compression of an ideal gas. Note that our deriva-tion is only applicable to a reversible process, because we have assumed that P : P.,,"^ot.

Reversible adiabatic compression of a gas leads to heating, and reversible adiabaticexpansion leads to cooling, as can be concluded from Figure 2.13. Two systems con-taining 1 mol of N, gas have the same volume at P : 1 atm. Under isothermal condi-tions, heat flows out of the system as it is compressed to P > 1 atm, and heat flows intothe system as it is expanded to P < 1 atm to keep I constant. No heat flows into or outof the system under adiabatic conditions. Note in Figure 2.13 lhat in a reversible adia-batic compression originating at 1 atm, Poo*uor. ) Pi,oth"*t for all P > 1 atm. Therefore,this value of P must correspond to a value of Z for which Z > Ti"othemat Similarly, in areversible adiabatic expansion originating at 1 atm, Padiabatic 1 Pirothr*t for all P < 1 atm.Therefore, this value of P must correspond to a value of Z for which I < Ti"oth"*ot.

Questions on Concepts 35

3500 m is 0.802 and 0.602 atm, respectively. If the initial temperature of thecloud mass is 288 K, what is the cloud temperature as it passes over themountains? Assume that C"-for air is 28.86 J K-r mol-r and that air obeys the

ideal gas law If you *e ooih. mountain, should you expect rain or snow?

Solutionlr.) (v,\ (r,p.) (r.) lp)

lnl --r- l=-(n-l)lnl --- l=-(v-1)lnl , , l=-(tt-l)lnlr l-{v-l)lnl I I

l. 1i l.4l lr,P,) l. 1,1 " \Pr)(c. )| "Pm - I/ \ l:-a-ll / \(y-r), Ie ] [c',,-R ). ( P,\

---rr.| - l-----------;-_-,1 _ lt, lP I r lP. l' ,. I) "P'm \r./Cr.r-R

( 28.86 J K-r mol-r )t_11

I Za.a6 J K-r mol-r - 8.314 J K-r mol-r ' )

28.86 J K-r mol-r

29.96 JK-1 mol-1 _ 9.314 J K-l mol-lTt =0.9207Ti= 265 K

You can expect snow.

* rI o.sozut,n') = _{.0826

[ 0.602atm J

Vocabularycyclic path

enthalpy

exact differential

first law of thermodynamics

heat

heat capacity

indicator diagram

internal energy

irreversible process

isobaric

isochoric

isothermal

isothermal process

path

path function

quasi-static process

reversible process

state function

work

Questions on Coneepts

Q2.l Electrical current is passed through a resistor immersedin a liquid in an adiabatic container. The temperature of theliquid is varied by 1'C. The system consists solely of theliquid. Does heat or work flow across the boundary betweenthe system and surroundings? Justify your answer.

Q2.2 Explain how a mass of water in the surroundings canbe used to determine q for a process. Calculate q rf rhetemperature of a 1.00-kg water bath in the surroundingsincreases by 1.25'C. Assume that the surroundings are at aconstant pressure.

Q2.3 Explain the relationship between the terms exactdffi rential and state function.

Q2.4 Why is it incorrect to speak of the heat or workassociated with a system?

Q2.5 Two ideal gas systems undergo reversible expansionstarting from the same P and V. At the end of the expansion,the two systems have the same volume. The pressure in thesystem that has undergone adiabatic expansion is lower thanin the system that has undergone isothermal expansion.Explain this result without using equations.

Q2.6 A cup of water at 278 K(the system) is placed in amicrowave oven and the oven is turned on for 1 minuteduring which it begins to boil. Which of q, w, and AU arepositive, negative, or zero?

36 C H A PT E R 2 Heat, Work, lnternal Energy, Enthalpy, and the First Law of Thermodynamics

Q2.7 What is wrong with the following statement?:

Because the well-insulated house stored a lot of heat, the

temperature didn't fall much when the furnace failed. Ptewritethe sentence to convey the same information in a correct way.

Q2.8 What is wrong with the following statement?: Burnscaused by steam at 100"C can be more severe than those

caused by water at 100"C because steam contains more heat

than water Rewrite the sentence to convey the same

information in a correct way.

Q2.9 Describe how reversible and irreversible expansions

differ by discussing the degree to which equilibrium ismaintained between the system and the surroundings.

Q2.10 A chemical reaction occurs in a constant volumeenclosure separated from the sunoundings by diathermalwalls. Can you say whether the temperature of the

surroundings increases, decreases, or remains the same in thisprocess? Explain.

Problems

Problem numbers in red indicate that the solution to theproblem is given inthe Student's Solutions ManuaL

Pz.I 3.00 moles of an ideal gas at27.0oC expands

isothermally from an initial volume of 20.0 dm3 to a finalvolume of 60.0 dm3. Calculate w for this process (a) forexpansion against a constant external pressure of 1.00 x 10s Pa

and (b) for a reversible expansion.

P2.2 3.00 moles of an ideal gas are compressedisothermally from 60.0 to 20.0 L using a constant externalpressure of 5.00 atm. Calculate q, w, LU , and LH .

P2.3 A system consisting of 57.5 g of liquid water at298 K is heated using an immersion heater at a constantpressure of 1.00 bar. If a current of 1.50 A passes throughthe 10.0-ohm resistor for 150 s, what is the finaltemperature of the water? The heat capacity for water can

be found in Appendix B.

P2.4 For 1 mol of an ideal EAS, P ",,",not:

P : 200 x 103

Pa. The temperature is changed from 100oC to 25.0oC, and

Cu*: 312R. Calculate q, w, LU, and AF1.

Y2.5 Consider the isothermal expansion of 5.25 mol of an

ideal gas at 450 K from an initial pressure of 15.0 bar to a finalpressure of 3.50 bar. Describe the process that will result in the

greatest amount of work being done by the system withP*"*ot 2 3.50bar and calculate w. Describe the process that willresult in the least amount of work being done by the system withPut"-ot > 3.50 bar and calculate r.u. What is the least amount ofwork done without restrictions on the external pressure?

P2.6 Calculate A11 and AU for the transformation of 1 mol ofan ideal gas from 27.0"C and 1.00 atmto32ToC and 17.0 atm if

TCr,* = 20.9 + 0.042-in units of J K-lmol-l

P2.7 Calculate w for the adiabatic expansion of 1 mol ofan ideal gas at 411 initial pressure of 2.00 bar from an initialtemperature of fSO K to a final temperature of 300'K. Writean expression ftr the work done in the isothermal reversibleexpansion ofthe gas at 300 K from an initial pressure of2.00bar. What value of the final pressure would give thesame value of w as the first part of this problem? Assumethat C r-: 5/2R.

P2.& In the adiabatic expansion of 1 mol of an ideal gas

from an initial temperature of 25.0'C, the work done on the

surroundings is 1200 J.If Cu*: 312R, calculate q,w, LU,and L.H.

P2.9 An ideal gas undergoes an expansion from the initialstate described by P,, V ,, T to a final state described by P, V,Zin (a) a process at the constant external pressure Prand (b)

in a reversible process. Derive expressions for the largestmass that can be lifted through a height h in the surroundingsin these processes.

P2.10 An automobile tire contains air at 320 x 103 Pa at20.0'C. The stem valve is removed and the air is allowed toexpand adiabatically against the constant external pressure of100 x 103 Pa until P : P.,,",,o,.For air, Cr^: 5/2R.Calculate the final temperature. Assume ideal gas behavior.

P2.ll 3.50 moles of an ideal gas is expanded from 450 Kand an initial pressure of 5.00 bar to a final pressure of1.00 bar, and C"*: 5/2R. Calculate w for the followingtwo cases:

a. The expansion is isothermal and reversible.

b. The expansion is adiabatic and reversible.

Without resorting to equations, explain why the result to part(b) is greater than or less than the result to part (a).

P2.12 An ideal gas described by T,: 300 K, P, :1.00 bar, andV.: 10.0 L is heated at constant volume untilP : 10.0 bar. It then undergoes a reversible isothermalexpansion until P : 1.00 bar. It is then restored to itsoriginal state by the extraction ofheat at constant pressure.

Depict this closed-cycle process in a P-V diagram.Calculate w for each step and for the total process. Whatvalues for w would you calculate if the cycle were traversedin the opposite direction?

P2,13 3.00 moles of an ideal gas with Cu*: 3l2R initiallyat a temperature T,: 298 K and P, - 1.00 bar is enclosed inan adiabatic piston and cylinder assembly. The gas iscompressed by placing a 625-kg mass on the piston ofdiameter 20.0 cm. Calculate the work done in this process aad

the distance that the piston travels. Assume that the mass ofthe piston is negligible.

mt[4 Abottle at2l.0"C contains an ideal gas at a

mmffiure of 126.4 x 103 Pa. The rubber stopper closing the]Me is removed. The gas expands adiabatically against

'lk-, : 101.9 x 103 Pa, and some gas is expelled from thelftde in the process. When P : P",t",,ot, the stopper is

ryfrckly replaced. The gas remaining in the bottle slowlymns up to 21.0"C. What is the final pressure in thebottleffiina monatomic gas, for which Cr^: 3/2R, and a diatomicnn for which Cr^: 5/2R?

E[.f5 Apellet of Zn of mass 10.0 g is dropped into a flaskumrrrining dilute HrSOo at a pressure of P : 1.00 bar andlqErature of Z: 298K. What is the reaction that occurs?Odmlate w for the process.

trLf6 One mole of an ideal gas for which Cr^: 20.8 J K-1hr is heated from an initial temperature of OoC to a final

ru l ature of 275"C at constant volume. Calculate q,w, LU,d lll for this process.

nfi One mole of an ideal gas, for which Cr,._ : 312R,

ffially at20.0'C and 1.00 x 106 Pa undergoes a two-stagerrdormation. For each of the stages described in the followingllfrn 6afgufale the fixal pressure, as well as q, w, LU , and A,H.

#bo calculate q, w, LU , and L,H for the complete process.

inr" The gas is expanded isothermally and reversibly until therolume doubles.

h" Beginning at the end of the first stage, the temperature israised to 80.0'C at constant volume.

mf8 One mole of an ideal gas with Cu*: 3/2R initially atMS K and 1.00 X 10s Pa undergoes a reversible adiabaticomrpression. At the end of the process, the pressure is 1.00 xMPa. Calculate the final temperature of the gas. CaTcuIate q,

m, AU, and LH for this process.

;LI) One mole of an idealgas, for which Cr*: 3/2R, isdjected to two successive changes in state: (1) From 25.0"Crld 100 x 103 Pa, the gas is expanded isothermally against amstant pressure of 20.0 x 103 Pa to twice the initial volume.r@) "{t the end of the previous process, the gas is cooled at

nsr2nt volume from 25.0'C to -25.0'C. Calculate q, w, LU,nd iH for each of the stages. Also calculate q, w, AU, and/ld for the complete process.

VrJn The temperature of 1 mol of an ideal gas increasesfum 18.0'to 55.1"C as the gas is compressed adiabatically.Calculate q, w, L,U , and AI1 for this process assuming thatE**:312n.

rrrlergees the following two-step process: (1) From an initial*e of the gas described by T :28.0'C and P : 2.00 x 104

h- the gas undergoes an isothermal expansion against amstant external pressure of 1.00 X 104 Pa until the volumehs doubled. (2) Subsequently, the gas is cooled at constantdume. The temperature falls to -40.5'C. Calculate q,w, L,U,md Af1 for each step and for the overall process.

]222 A cylindrical vessel with rigid adiabatic walls is

ryarated into two parts by a frictionless adiabatic piston.

Problems 37

Each part contains 50.0 L of an ideal monatomic gas withC u- : 312R. Initially, T : 298 K and P, : I .00 bar in eachpart. Heat is slowly introduced into the left part using anelectrical heater until the piston has moved sufficiently to theright to result in a final pressure Pr: l.50bar in the rightpart. Consider the compression of the gas in the right part tobe a reversible process.

a. Calculate the work done on the right part in this processand the final temperature in the right part.

l

b. Calculate the final temperature in the left part and theamount of heat that flowed into this part.

P2,23 A vessel containing 1 mol of an ideal gas with P,: 1.00 bar and C

"*: 5/2R is in thermal contact with a water

bath. Treat the vessel, gas, and water bath as being in thermalequilibrium, initially at 298 K, and as separated by adiabaticwalls from the rest of the universe. The vessel, gas, and waterbath have an average heat capacity of Cr:7500 J K-1. Thegas is compressed reversibly to Pr: 10.5 bar. What is thetemperature of the system after thermal equilibrium has beenestablished?

P2,24 The heat capacity of solid lead oxide is given by

c r,*= 44.35+l.41xr01lin onlt, of J K-r mol-rK

Calculate the change in enthalpy of I mol of PbO(s) if it iscooled from 500 to 300 K at constant pressure.

P2.25 Consider the adiabatic expansion of 0.500 mol of anideal monatomic gas with Cr*: 3/2R. The initial state isdescribed by P : 3.25bn andT :300 K.

a. Calculate the final temperature if the gas undergoes a

reversible adiabatic expansion to a final pressure ofP : 1.00bar.

b. Calculate the final temperature if the same gas undergoesan adiabatic expansion against an external pressure ofP : 1.00 bar to a final pressure P : 1.00 6ar.

Explain the difference in your results for parls (a) and (b).

P2.26 An ideal gas undergoes a single-stage expansionagainst a constant external pressure P",r",,otat constanttemperature from Z, P , V., to T, Pf, Vf.

a. What is the largest mass m that can be lifted through theheight /z in this expansion?

b. The system is restored to its initial state in a single-statecompression. What is the smallest mass rn'that must-fallthrough the height ft to restore the system to its initialstate?

c. If h: 10.0 ctr', P. : 1.00 X 106 Pa, Pr: 0.500 X 106 Pa,

Z :300 K, and n : I .00 mol, calculate the values of themasses in parts (a) and (b).

P2,27 Calculate q, w, L.U, and AI1 if 1.00 mol of an idealgas with C u^

: 3/2R undergoes a reversible adiabaticexpansion from an initial volum e V, : 5 .25 m3 to a finalvolume Vr: 25.5 m3. The initial temperature is 300 K.

3g cHAprER z Heat, work, tnternal Energy, Enthalpy, and the First Law ofrhermodynamics

P2.28 A nearly flat bicycle tire becomes noticeably warmer

after it has been pumped up' Approximate this process as a

reversible adiabatic compression' Assume the initial pressure

and temperature of the air before it is put in the tire to be P,

:1.00baranaT,:2gSK.Thefinalvolumeoftheairinthetire is V": 1.00 L and the final pressure fu Pr: 5'00 bar'

Calculate the final temperature of the air in fhe tire' Assume

thatCr^:512RP2.2g One mole of an ideal gas with Cn^: 3l2R\s

""funa"A adiabatically against a constaniexternal pressure of

t.riO Uar. The initial templrature and pressure are Z : 300 K

and P, : 25.0 bar. The final pressure is Pr: 1'00 bar'

Calcuiate q, w. LU . and AH for the process'

P2.30 One mole of N, in a state definedby T,: 300 K and

V. : 2.50L undergoes in isothermal reversible expansion

until V. : 23.OL.talculate w assuming (a) that the gas is

descrided by the ideal gas 1aw and (b) that the gas is

described by the van der Waals equation of state' What is the

p"t "n, "rr-

in using the ideal gas law instead of the van der

Waals equation? The van der Waals parameters for N, are

listed in Table7.4.

web-Based simulations, Animations, and Problems

W2.1 A simulation is carried out in which an ideal gas

is heated under constant pressure or constant volume

conditions. The quantities AV (or L^P),w, AU' and AT are

determinedasafunctionoftheheatinput.Theheattakenupby the gas under constant P or Vis calculated and compared

with AU and LH.

w2.2Thereversibleisothermalcompressionandexpansionof an ideal gas is simulated for different values of Z' The

work, ru, is*calculated from the 7 and V values obtained in the

simulation. The heat, q, andlhe number of moles of gas in the

system are calculated from the results'

W2.3 The reversible isobaric compression and expansion of

anidealissimulatedfor'differentvaluesofpressuregasasheat flows to/from the surroundings' The quantities q' w' and

AU are calculated from the AI and AV values obtained in the

simulation.

W2.4 The isochoric heating and cooling of an ideal gas is

simulated for different values of volume' The number of

-oi"t of gas and LU ate calculated from the constant V value

and fromihe T and P values obtained in the simulation'

W2.5 Reversible cyclic processes are simulated in which

the cycle is either rectangular or triangular on a P-Vplot' For

"u"fri"g-"nt and for the cycle, LU, q, and w are determined'

For a gien cycle type, the ratio of work done on the

,orrottding, to the heat absorbed from the surroundings is

determined for different P and V values'

W2.6 The reversible adiabatic heating and cooling of an

ideal gas is simulated for different values of the initial

temperature. The quantity y = Cp,^lCv,ru ls-well as

-C"-andC-. aredetermined from the P,Vvalues of the simulation'

LU and LH are calculated from the V,T, and P values

obtained in the simulation.