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Basic Thermodynamics
P. Duthil^1 Institut de Physique Nucléaire d’Orsay, IN2P3-CNRS/Université de Paris Sud, Orsay, France
Abstract The goal of this paper is to present a general thermodynamic basis that is useable in the context of superconductivity and particle accelerators. The first part recalls the purpose of thermodynamics and summarizes its important concepts. Some applications, from cryogenics to magnetic systems, are covered. In the context of basic thermodynamics, only thermodynamic equilibrium is considered.
Keywords : thermodynamics, expansion of gas, heat machines, energy, phase transition.
1 Introduction
Thermodynamics provides macroscopic descriptions of the states of complex systems and of their behaviours when they interact or are constrained under various circumstances. A thermodynamic system is a precisely specified macroscopic region of the Universe. It is limited by boundaries of a particular nature (real or not) with particular properties. All space in the Universe outside the thermodynamic system is known as the surroundings , the environment, or the reservoir. By (European) convention, the scalar quantities exchanged between the system and the surroundings are counted negatively if they leave the system and positively in the opposite case. A thermodynamically isolated system is completely isolated from its environment: it does not exchange energy (heat, work) or matter with the environment. If matter is able to flow in or out of its boundaries, the system is said to be open. Otherwise it is said to be closed (by real boundaries). Processes that affect the interior of the region of the system are studied using general principles: the four laws of thermodynamics. Although historically established before modern physics, most of these laws are based on the fact that a system is governed by general statistical principles which apply to the random motion of its particles. Hence, two approaches are possible when studying a system by consideration of its thermodynamics: a very general and practical one is to establish balances of the transfers between the system and its surroundings; the other is to apply statistical laws to the system. The latter approach is suitable for simple systems, and may also be used to clear up some of the more abstract concepts within thermodynamics.
A unique macroscopic description of a system requires thermodynamics to use a number of macroscopic parameters: the pressure p , the volume V , the magnetization M ˆ , or the applied magnetic field H ˆ^ , for example. Some other fundamental macroscopic parameters are defined by the laws of thermodynamics: the temperature T , the total internal energy U and the entropy S. Knowing (experimentally) the minimum number of parameters required for this unique description fixes the number of independent state variables chosen for the macroscopic description. Thermodynamics then expresses constraints on the system by use of some relationships between these parameters. Parameters that are proportional to the size of the system are extensive variables and are written in the following in upper-case characters. Parameters that do not depend on the mass of the system are intensive quantities and are written in lower-case characters. An exception is made for the temperature, an intensive quantity denoted by T. For homogeneous systems, extensive parameters can
(^1) duthil@ipno.in2p3.fr
be related to intensive variables via the relation X = m x ⋅ , where m^ is the mass. They then refer to
bulk properties of the system. An example is the volume V (m^3 ) and the specific volume v^ (m^3 ·kg−1).
2 Equilibrium and equation of state
A thermodynamic system is in thermodynamic equilibrium when all its state variables remain constant with time: there is no net flow of matter or energy, no phase changes, and no unbalanced potentials (or driving forces) within the system. A system that is in thermodynamic equilibrium experiences no change when it is isolated from its surroundings. It should be stressed that thermodynamic equilibrium implies steady state, but that steady state does not always induce thermodynamic equilibrium (e.g. steady heat diffusion in a support). For different combined subsystems forming one thermodynamic system at thermodynamic equilibrium, the extensive parameters of each subsystem are summed (volume, for example) and the intensive parameters are equal (pressure, temperature). At equilibrium, a functional relationship between the parameters (state variables) of a system can be expressed as an equation of state.
As an example, a gas is completely described at equilibrium by the parameters p , V , and T. The equation of state, of the form f^ (^ p V T ,^ ,^^ )^ =^0 , describes an equilibrium surface on which a point represents an equilibrium state. The isotherms of a (dilute) gas may be described according to Boyle’s law (or the Boyle–Mariotte law) by pV^ / n^ =^ cst , where n denotes the number of moles. An experimental temperature scale can be defined by any function of the temperature associated with these isotherms. The simplest function is pV^ =^ nRT , known as the ideal gas law , which defines the isotherms as hyperbolic curves in the p^ −^ V plane—see Fig. 1(a). Using an air thermometer, Amontons suggested in 1702 that the zero on this temperature scale would be for the zero pressure limit. Taking into account the interaction between the molecules, the van der Waals equation
( p + a / v ²)( v − b )= rT is also an equation of state for a gas: it modifies the pressure term by
considering the attraction effect – a between molecules; it also corrects the volume by considering what is really accessible by one mole of molecules: a volume b is indeed not available due to the repulsion effect of the molecules. It is accurately applicable for a larger range of pressures. In the van
der Waals equation, there is a critical isotherm for which (^ ∂ p^^ /^ ∂ V^^ )^ T = ∂(^2 p^ /^ ∂ V^2 )^ T =^0 at a critical
point C having coordinates p C^ =^ a^ / (27^ b^2 ), V C^ =^3 nb , and T C^ =^8 a^ / (27^ Rb ). Tending to this critical
point, the van der Waals hyperbolic isotherms experience an inflection point—see Fig. 1(b). Below this critical point, two inflection points are revealed (furthermore, for a given volume and temperature, the calculated pressure can be negative). This domain coincides with the liquid phase of the fluid (see Section 15).
Fig. 1 : Equilibrium surfaces given by the ideal gas (a) and the van der Waals (b) equations of state
where v^ is the specific volume (unit: m^3 ·kg−1) of the fluid. The work which may be performed by the
system on some mechanical device is often called the ‘shaft’ work (useful work), and is the difference
between the total work δ W and the flow work δ^ W flow:
δ W shaft = δ W − δ W flow = V d p. (2)
4.2 Internal energy
The internal energy U (unit: J) is the sum of the kinetic and potential energies of all the physical
particles of the system (and of the rest-mass energies) at some microscale: U^ =^ ∑ E c,micro^ +∑ E p,micro.
It is a state function and can thus be defined by macroscopic parameters. For example, for a non- magnetic compressible fluid, if p and V are fixed, U^ =^ U^ (^ p V ,^ )is also fixed.
4.3 First law of thermodynamics and energy balances
Between two thermodynamic equilibria, δ^ U^ =^ δ^ W^ +^ δ Q ; for a reversible process, d U^ =^ δ W^^ +^ δ Q , where W is the exchange work (mechanical, electrical, magnetic interaction, etc.) and Q^ is the exchange heat. For a cyclic process during which the system evolves from an initial state I to an identical final state F : Δ^ U^ =^ U^ F −^ UI =^0.
Between two thermodynamic equilibria, the total energy change is given by
Δ E = Δ E c,macro + Δ E p,macro + Δ U = W + Q. E c,macro and E p,macro are the kinetic and potential energies of
the system at a macroscopic scale, and they are null if the system does not move or is not in a force field. This leads to: Δ U = W + Q. (3)
If work is due to external pressure forces only, then W = W epf = ∫ − p d V and Δ U = W epf + Q. For an
isochoric process, the volume remains constant and Δ U^ =^ Q. This relation is the basis of calorimetry.
For an open system, the work δ W flowachieved by the pressure forces to make the fluid circulate
in and out is distinguished from the useful work performed by other forces (mechanical, magnetic,
electrical, etc.) and denoted δ W u. We note that for the system of Fig. 2(b) δ W u = δ W shaft. The first law
then yields:
[ ]
outlet
Δ E = Δ E c,macro + Δ E p,macro + Δ U + δ m u inlet = W u + W flow + Q.
This can be rewritten as follows:
c,macro p,m [ ]
outlet Δ E + Δ E (^) acro + Δ U + δ m u + pV inlet = W u + Q.
Another function of state, called enthalpy , is then defined as H = U + pV (unit: J). (4)
Thus, h being the enthalpy per unit mass (J·kg −1), we obtain [ ]
outlet Δ E = Δ E c,macro (^) + Δ E p,macro (^) + Δ U + δ m h (^) inlet = W u + Q.
For a cyclic process ( Δ U = 0 ), if Δ^ E c,macro^ =^ Δ^ E p,macro=^0 ,
Δ H = W u + Q. (5)
In the case of an isobaric process (for which the pressure remains constant) and if no other type of work is performed other than the pressure work, Δ H^ =^ Q.
The energy balance for an open system at steady state or undergoing a cyclic process is written as follows:
u outlet boundaries inlet boundaries u outlet boundaries inlet boundaries
(unit: J),
(unit: W).
W Q mh mh
W Q mh mh
5 Entropy and second law of thermodynamics
For a system considered to be between two successive states, there is a function of state S called entropy , unit J·K−1, such that:
Δ S syst = Δ S = Δ S e^ + S i, (7)
where Δ^ S^ e= ∫ δ Q^ / T relates to the heat exchange and S i is an entropy production term:
S^ i^ = Δ S (^) syst + Δ S surroundings. For a reversible process, S i^ = 0 ; for an irreversible process, S i^ > 0. For an
adiabatic ( δ^ Q =^0 ) and reversible process ( S i^ = 0 ) the thermodynamic transformation is isentropic :
Δ S = 0.
5.1 Entropy of an isolated system
An isolated system has no exchange with the surroundings; thus Δ S e= 0 and Δ^ S syst^ =^ S i≥^0. At
thermodynamic equilibrium, the isolated system experiences no change, and its state variables remain
constant with time. Hence, Δ S syst = 0 and S i^ = 0.
The statistical weight Ω denotes the number of possible microstates of a system: the position of the atoms or molecules, or the distribution of the internal energy. The different possible microstates correspond to (and are consistent with) the same macrostate described by the macroscopic parameters ( p , V , etc_._ ). The probability of finding the system in one microstate is the same as that of finding it in another microstate. Thus, the probability that the system is in a given macrostate must be proportional to Ω. The entropy can be expressed as follows:
S = k B ln( Ω), (8)
where k B^ =^ 1.38 10⋅^ −^23 J·K−1^ is the Boltzmann constant. Considering two systems A and B , the number
of microstates of the combined system A ∪ B is Ω^ A ⋅^ Ω B and its entropy is S^ =^ S^ A +^ SB : the entropy is
additive. Similarly, it is proportional to the mass of the system, and is thus an extensive parameter, i.e.
S = m s ⋅ , where s is the specific entropy (unit: J·K−1·kg−1).
5.2 Principle of increase of entropy
One statement of the second law of thermodynamics is that “entropy of an isolated system tends to increase”. The system changes in response to this tendency. Such a change is said to be ‘spontaneous’. If one imagines two possible macrostates for an isolated system, corresponding to two values of the number of possible microstates Ω , the system will experimentally be found in the macrostate for which the number of microstates is the greater, thus having larger entropy. We consider, for example, a gas in a box, and assume that the gas is on one side of the box at some initial time corresponding to
Ω = Ω 1. If the gas is free to fill up the whole box, corresponding to Ω = Ω 2 , it will ‘spontaneously’
go to that state. Indeed, Ω 2 >>^ Ω 1 , and there is a higher probability of finding the gas throughout the
whole box than of staying on one side of the box. This principle applies to isolated systems only. Otherwise, one would have to consider exchanges with the surroundings, and thus the change of entropy of the surroundings. Even though not very convenient, it is always possible to transform a system into an isolated one by extending its boundaries, keeping in mind that the whole Universe is an isolated system.
8 Relating entropy to functions of state
8.1 Entropy with internal energy
U and S are state functions. Any macrostate at equilibrium can be described by a relationship between these functions and the independent state variables. For example, a simple compressible system is completely characterized by its energy and its volume, and U^ =^ f^ (^ S V ,^ ). Therefore, for a reversible process ( S i^ = 0 ), the change of the state functions can be written as
d d d d d V S S
U U U
U S V T S V
S V V
∂^ ∂^ ∂
∂^ ∂^ ∂
Identifying this equation with the expression of the first law of thermodynamics for a closed system
yields p^ = − ∂(^ U^^ /^ ∂ V^ ) S and thus
d U = T d S − p d V. (10)
The kinetic temperature of an ideal gas can be derived from the expression of its internal energy and from the motion of its molecules. This kinetic temperature can be related to the thermodynamic temperature using the Boltzmann distribution. It follows that the equation of state of an ideal gas can be written as
pV = Nk T B = nN k T A B , (11)
where N is the number of molecules of the gas and N (^) A = 6.022 10⋅ 23 mol−^1 is Avogadro’s number. Stating that R = N k A B, we find again pV = nRT , and thus we relate the experimental temperature defined in Section 2 to both the thermodynamic and kinetic temperature.
8.2 Thermodynamic potentials
We consider a non-isolated system in thermal contact with its surroundings ( EXT ) which is at
temperature TEXT , with a large heat capacity such that heat flows from the surroundings into the system without changing TEXT. The second principle states that for the system combined with its surroundings the total change of entropy is given by δ STOT = δ S + δ SEXT ≥ 0 with δ SEXT = −δ Q / TEXT. For the considered system having a constant volume, δ U = δ Q. Thus, we obtain δ ( U − TEXT S ) ≤ 0 : this spontaneous change is related to a decrease of U − TEXT S.
A new thermodynamic equilibrium ( T = TEXT ) will thus be reached for the minimum of the (Helmholtz) free energy of the system F = U − TS. Similarly, for a system in thermal contact with its surroundings and held at constant pressure, the Gibbs free energy G = U − TS + pV tends to a minimum at thermodynamic equilibrium. The free energy and the Gibbs free energy are thermodynamic potentials.
Thermodynamic potentials are specific functions of some state variables: the natural variables of the potential. When these natural variables are kept constant, a closed thermodynamic system evolves towards equilibrium if the potentials diminish. The thermodynamic equilibrium corresponds to a minimum of one of the thermodynamic potentials. The internal energy U = f ( S V , ) is such a thermodynamic potential of natural variables S and V. For a mechanical system (no heat exchange), we saw that an increase of energy is the result of a gain of mechanical work which is the product of a force by a small displacement. We can generalize this concept by stating that the increase in energy of a thermodynamic system is the result of the addition of several products of generalized thermodynamic forces which induce thermodynamic displacements from equilibrium. The thermodynamic force is driven by an intensive variable which is conjugated to an extensive variable
representing the thermodynamic displacement: p and V are such conjugated variables, as are T and
S. If S and V are kept constant, U will tend to a minimum as the considered closed system tends to
thermodynamic equilibrium. The other thermodynamic potentials tending to a minimum for a closed system at equilibrium are:
the enthalpy, if p and the extensive parameters (such as S ) are held constant,
H = U + pV ⇒ d H = T d S + V d p ;^ (12)
the Helmholtz free energy, if T and the extensive parameters (such as V ) are held constant,
F = U − TS ⇒ d F = − S d T − p d V ;^ (13)
the Gibbs free energy if p , T , and the extensive parameters are held constant,
G = U − TS + pV ⇒ d G = − S d T + V d p.^ (14)
9 Maxwell relations
For a compressible system U = f ( S V , ), the differential form d U is exact if the two mixed second
partial derivatives are equal: 2 2
S V V S
U U U U
V S S V V S S V
As (^) d U = T d S − p d V , we deduce an appropriate relation between the first derivatives of T and p :
S V
T p V S
∂^ ∂
∂^ ∂
By the same argument, we obtain from the different potentials the useful Maxwell relations:
( , ) ,
S V
S p
T p U U S V V S
T V H H S p p S
∂^ ∂
∂^ ^ ∂
V T
T p
p S F F T V T V
S V G G T p p T
∂^ ∂
∂^ ^ ∂
10 Heat capacities
The amount of heat that must be added to a system reversibly to change its temperature is the heat
capacity: C^ =^ δ Q^ / dT (unit: J·K −1). The conditions under which heat is supplied, at constant volume
V or constant pressure p , must be specified:
V V V
S U
C T
T T
∂^ ∂
∂^ ∂
, (^) p p p
C T S^ H
T T
= ^ ∂^ ^ =^ ∂
It turns out that
for V = cst , ( , ) 0 d
T CV
S T V T
T
= ∫ and Δ U = ∫ CV d T ;
for (^) p = cst , ( , ) 0 d
T (^) Cp S T p T T
= ∫ and Δ H = ∫ C p d T.
Q = ∫ C dT i is known as the ‘sensible heat’ (see the lecture ‘Materials properties at low temperature’ in
these proceedings).
usually subcooled below the inversion temperature by combining adiabatic expansion and heat removal within turbines and heat exchangers. It is finally liquefied by use of a Joule–Thomson expansion. Note that the maximum inversion temperatures of hydrogen, nitrogen and oxygen are 202 K, 623 K and 761 K, respectively.
Fig. 3: Inversion temperature of helium-4 as a function of (^) p and (^) T. Constant enthalpy lines are dashed; the
solid line is the inversion curve.
11.3 Thermodynamic transforms of an ideal gas
Table A.1 of Appendix A lists different reversible thermodynamic processes involving an ideal gas. It provides a plot of the considered evolution in a p − V diagram, expresses the heat and work exchanged, and gives the changes in internal energy, enthalpy and entropy.
12 Heat machines
The basic function of a heat machine is to convert heat into work or, conversely, to produce a heat transfer by means of supplied work. This work can be electrical, mechanical, etc., but the second principle states that its production by a heat engine requires at least two heat reservoirs at two different temperatures: the machine is driven with some heat supplied by the hotter reservoir, produces work, and dumps waste heat to a colder reservoir (see Fig. 4). A heat pump is driven by an amount of work and aims to transfer heat from a colder reservoir to a hotter one. It can therefore be used to warm or to cool. In the latter case, the heat pump is a refrigerator. A dithermal machine is a heat machine
working by way of two heat reservoirs: one hot at temperature T H and one cold at T C. Over one
thermodynamic cycle, the system in the final state recovers its initial states. Thus,
energy balance (first law): ∆ U = U (^) f − U (^) i = 0 ⇒ W + Q C (^) + Q H= 0 ;
entropy balance (second law): ∆ S = ∆ S e^ + S i^ = 0 ⇒ Q C (^) / T C (^) + Q H (^) / T H + S i= 0.
It follows that:
i Q H (^) = − Q (^) C (^) − W = − T /H T C (^) ⋅ Q C (^) − T S H. (21)
Fig. 4: Principle of a heat machine: engine (top) and heat pump or refrigerator (bottom)
From the two expressions given for the energy and entropy balances, we can plot the domain of
the existence of heat machines in a Q H − Q C diagram, called a Raveau diagram and shown in
Figs. 5(a) and (b).
Fig. 5: Raveau diagrams of a heat engine (a) and a heat pump (b). Domains of existence of both types of heat machines are indicated by pale grey zones. The domains of impossible cycles (as in not respecting the second law) are indicated in dark grey.
12.1 Heat engine
As the work is produced (given) by the engine, we have (^) W < 0 and, if T H (^) > T C, then Q H (^) > 0 and
Q C < 0 (see Fig. 5(a)). The engine is supplied by heat coming from the hot reservoir and rejects excess
heat at the cold reservoir. The efficiency η of an engine can be defined by the (positive) ratio of the
amount of produced work to the energy that is supplied. The supply being the heat coming from the hot source, we have
H C C H H H H
W W Q Q Q
η Q Q Q Q
i.e., making use of equation (21),
η= 1 − T C (^) / T H (^) − T S C^ i/ Q H (^) ≤ ηC = 1 − T C (^) / T H< 1 since S^ i^ ≥^0. (22)
The maximum possible efficiency ηC is obtained in the case of no irreversibility ( S i^ = 0 ) and is called the Carnot efficiency , referring to the ideal Carnot cycle presented in Section 12.3.
This fully reversible cycle performs the maximum efficiency obtainable with a heat machine and thus serves as a reference for comparing efficiencies of heat machines, taking into account the temperatures of their heat reservoirs. The relative efficiencies of heat machines are thus defined as follows:
engine ( )
C^ i r C H C H
T S
Q T T
heat pump
Heating C H r Heating (^) i C Heating C H C C
COP T T
COP
COP T T T S Q
refrigerator
Cooling H C r Cooling (^) i C Cooling H C C H
COP T T
COP
COP T T T S Q
Carnot efficiencies of a heat pump and a refrigerator are plotted in Fig. 7 versus the ratio of the
temperatures of their reservoirs. As COP tends to very small values as the cold temperature T C tends
to zero, the cryogenic refrigeration community often employs the inverse of the COP and speaks in terms of ‘watt to supply per watt of produced cooling power’.
Fig. 7: COP values of a heat pump and a refrigerator based on the ideal Carnot cycle
The vapour compression cycle (see Fig. 8) achieved by our domestic refrigerators has a relatively good COP compared with the Carnot cycle. This intrinsically good performance comes from the fact that (i) the vaporization of a saturated liquid and the liquefaction of a saturated vapour are two isothermal process (heat is, however, transferred irreversibly); and (ii) the isenthalpic expansion of a saturated liquid is sensitively closed to an isentropic expansion. However, this cycle cannot be used in cryogenic applications; it relies on the phase change of a working fluid, and thus the temperatures of the heat sources cannot be very different. Even when implemented in cascade, the continuity of the process towards a low temperature cannot be insured as the existence of critical fluids (for which a liquid–gas phase transition is possible) does not cover continuously the domain of cryogenic temperature.
Fig. 8: Vapour compression heat machine
13 Maximum available work
Considering an isothermal ( T = cst ) and reversible thermodynamic process, dU = δ W + T d S and thus δ W = d U − T d S = d F. The work produced by a system ( δ W < 0 ) is equal to the reduction of its free energy. For the reversible process, this work is the maximum that can be provided. Similarly, for
an open system, d^ H^ =^ δ W shaft^ +^ T^ d S and thus δ W^ shaft^ =^ d^ H^ −^ T^ d^ S^ =^ d G. The maximum produced
work (other than due to the external pressure forces) is equal to the reduction of the Gibbs free energy. In these cases, all the energy is free to perform useful work because there is no entropy production due to irreversibilities. Sources of entropy production come from heat transfers (due to temperature difference), friction (due to the movements of solid or fluid components), dissipative structure induced by fluid motions, matter diffusion, electric resistance, etc.
Heat and work are not equivalent: they do not have the same ‘thermodynamic grade’. Whereas mechanical or electric work can be integrally converted into heat, integrally converting heat into work is impossible according to the second law. Moreover, transfers are generally irreversible, which implies that thermodynamic processes have a direction (heat flows from a hot to a cold body, mass transfers from high to low pressures, etc.). The exergy allows us to ‘rank’ energies by involving the concept of ‘usable’ or ‘available’ energy: it is the maximum useful work that brings a system into equilibrium with its surroundings at TEXT^ (generally the room temperature). Exergy is defined as
Ex = U − U EXT + pEXT ( V − V EXT ) − TEXT ( S − SEXT ), (29)
such that it tends to zero when the energy U , the volume V , and the entropy S of the system tend to those of the surroundings: U^ EXT , VEXT^ , and SEXT^. Exergy is a state function of the system relative to its surroundings, not of the system only.
Considering a closed system in thermal and pressure contact with its surroundings, we can deduce from the first and second law that
Δ( U + pEXT V − TEXT S )= W u − TEXT S i,
where W u is the useful work produced by the system (other than the work due to pressure forces). This work is maximum for S i^ = 0 , and thus
W u ≤ Δ( U + pEXT V − TEXT S ) = Δ Ex. (30)
The maximum produced work is equal to the reduction of the exergy (^) Δ Ex. If we define the Gibbs free energy of our system by imposing T = TEXT and p = pEXT , we see that Δ Ex = Δ G. We can thus determine the exergy by use of the thermodynamic potentials:
d Ex = d H − TEXT d S ; d Ex = d F − pEXT d V ; d Ex = d G.
The Gibbs phase rule gives the number of degrees of freedom D (number of independent intensive variables) of a system of mixed substances. In the absence of a chemical reaction or external constraint, the variance of the system depends on the number of (chemically independent) constituents c (^) and the number of phases φ (^) : D = c + 2 − φ. (33)
The variance of a monophasic (in a single phase) pure substance in thermodynamic equilibrium is D = 2 ; a full description of the system requires both pressure and temperature to be set. At the triple point J, where the substance is at equilibrium in three phases, D = 0. For a pure substance, this
unique point is defined by a unique couple ( p J , T J ). Triple points of different pure substances are used
as metrological references for the temperature scale and for the calibration of reference temperature
sensors. When two phases coexist at equilibrium, the variance is equal to 1. Thus, setting p or T
variable will fix the other one. This is the case at point P (placed on the saturated curve [JC]^ ) on Fig. 9: the gas (or vapour) and liquid phases coexist in equilibrium. The condition for this equilibrium is that the free Gibbs energy G tends to a minimum. Consider a small quantity of matter δ m that transfers from the liquid to the gas phase, then the change in the total free Gibbs energy is δ m g ( (^) gas (^) − g liq), which is minimum for g (^) gas = g liq. If we now consider a neighbouring point Q (^) on the saturated liquid–vapour curve, we can write, per mass unit:
for the gaseous phase (^) gas ( ) (^) gas ( ) gas^ d gas d (^) gas ( ) (^) gas d (^) gasd T^ p
g g g Q g P p T g P v p s T p T
∂^ ^ ∂
for the liquid phase (^) liq ( ) (^) liq ( ) liq^ d liq d (^) liq ( ) (^) liq d (^) liqd T^ p
g g g Q g P p T g P v p s T p T
∂^ ^ ∂
which yields, by subtraction,
0 = ( v gas − v liq )d p − ( s gas − s liq )d T ,
where v gas and v liq (unit: m^3 ·kg −1) are the specific volumes (the reciprocals of the mass densities) of the gas and liquid phases.
The slope of the saturated vapour tension curve is thus given by the Clausius–Clapeyron equation :
gas liq liq vap gas liq gas liq
d ( ) d ( ) ( )
p s s l T v v T v v
= − =^ →
where l liq →vap = T s ( gas − s liq) ( > 0 ; unit: J·kg−1) is the specific latent heat for the liquid → gas
transition (vaporization). It is the quantity of heat required to transform, at constant temperature and pressure, a unit mass of liquid into vapour. We can express this specific latent heat by use of the enthalpy difference per mass unit between the vapour and the liquid. Indeed, during the reversible phase transition process,
[ ] [ ]
vap vap vap l liq (^) → vap = q = (^) ∫liq d u + p d v = (^) ∫liq d u + d( pv ) = (^) ∫liq d h = h vap (^) − h liq.^ (35)
Quantification of the vaporization latent heat from a T − s diagram is shown in Fig. 10. It can be
deduced that l liq^ → vap tends to zero as we approach the critical point. Also, l liq^ → vap can be expressed as
a function of T by neglecting v liq^ (as v gas^ >>^ v liq) and using an equation of state to express d /d p T.
Fig. 10: Evaluation of the latent heat of the liquid–gas transition in a T − s diagram
The same argument can be applied to solid–liquid and solid–gas phase transitions defining fusion (melting) and sublimation (deposition) latent heats, respectively. These phase transitions are
characterized by a discontinuity in the first derivatives of the Gibbs free energy S = −( ∂ G /∂ T ) p : at
equilibrium, the entropy of the liquid phase is not the same as that of the gaseous phase. Thus, these phase transitions are first-order transitions.
At the critical point ( T = T C ; p = p C), however, v gas = v liq and d p / d T is finite so the Clausius–
Clapeyron equation states that l liq →vap = 0 = ( s gas − s liq). No latent heat is involved during the transition
and S = −( ∂ G /∂ T ) p is continuous. But second derivatives of Gibbs free energy are not continuous,
and the transition is of second order.
15 Thermodynamics of magnetic systems: type I superconductor phase transition
In a vacuum, the magnetic field B ˆ (unit: T ≡ V⋅s⋅m−² ≡ N⋅A−^1 ⋅m−^1 ) is proportional to the excitation (or
applied) magnetic field H ˆ (unit: V·s·A−^1 ·m−^1 ≡ A⋅m−^1 ): Β ˆ^ = μ 0 Η ˆ. The proportional factor is the
permeability of free space, μ 0 = 4 π ⋅ 10 −^7 (N⋅A−^2 ). In a material, B ˆ^ = μ 0 ( Η ˆ^ + M ˆ) , where M ˆ^ = χ Η ˆ
(unit: A⋅m−^1 ) is the magnetization, which represents how strongly a region of material is magnetized
by the application of H ˆ ; χ is the magnetic susceptibility. Thus, Β ˆ^ = μ 0 (1 + χ ) Η ˆ^ =μ μ 0 r Η ˆ , where μr
is the relative magnetic permeability.
Fig. 11: Type I superconductor. (a) Evolution of the magnetization with the excitation magnetic field; (b) Phase diagram; (c) Evolution of the Gibbs free energy with the excitation magnetic field below the critical temperature.
We consider a long cylinder sample of type I superconductor placed in a solenoid producing the
excitation magnetic field H ˆ = n i ⋅. Note that in the Meissner state, i.e. below a critical temperature
However, the heat capacity C^ =^ δ Q^ / d^ T^ =^ T^ d^ S^ / d^ T^ = − T^^ (^ ∂^2 G^ /∂ T^2 )^ p , a second derivative of the
Gibbs energy, is not continuous:
( C norm − C sup )= − 4 μ 0 V H ˆ CM^2 / T CM. (40)
Therefore at T = T CMthe superconducting → normal transition is a second-order transition.
16 Second-order phase transitions
The liquid–gas transition of a fluid and the superconducting–normal state of a type I superconductor are first-order transitions, except when the temperature reaches the critical temperature: no latent heat is involved and they become of second-order transitions as the second derivatives of the Gibbs free energy are not discontinuous.
Continuity of the first derivatives of the Gibbs free energy (per unit of mass) during a second-
order transition between phases I and II implies that d s I = d s II and d v I = d v II. Using Eqs. (15), (16)
and (20) we find that
d (^) i i^ d i^ d pi^ d i^ d pi d (^) i i d p (^) T p
s s c v c s T p T T T v p T p T T T
α
∂^ ^ ∂^ ∂
∂^ ∂^ ∂
Similarly, using Eq. (20) and introducing the coefficient of isothermal compressibility
χ T = − 1 / V ( δ V / δ p ) T , we obtain
d (^) i i^ d i d (^) i i d (^) i Ti d p (^) T
v v v T p v T v p T p
∂ ^ ∂
∂^ ∂
Hence, from the equality of the derivatives on the coexistence curve, and as there is no change in volume during the transition, we obtain, for the Ehrenfest formulae,
I II I II
d
d ( )
p c p cp
T Tv α α
and
I II I II
d
d T T
p
T
α α χ χ
An example of a second-order transition is the normal–superfluid transition of liquid helium-4.
The equilibrium curve of these two phases in a p − T diagram is called the lambda line. This name
originates from the shape of the heat capacity curve plotted versus temperature, which has the form of the Greek letter lambda near the transition, as can be seen in Fig. 12(a). The lambda point is the triple
point where the normal phase, the superfluid liquid phase, and the vapour phase coexist. As T
diminishes to T λ , the heat capacity of He-I (normal state) increases. At T = T λ, the heat capacity is
discontinuous, as shown in Fig. 12(b). Below T λ , the heat capacity reaches a maximum and decreases
as T diminishes.
Fig. 12: Specific heat of helium-4 as a function of temperature near the lambda point: (a) large temperature
scale; (b) very close to T λ (from [1]).
It should be noted that higher-order phase transitions can be found in nature, even though they are less abundant than first- or second-order ones. Their order is defined as the minimum integer i
such that (^ ∂ i^^ G^ /^ ∂ x^ i^ )I^ ≠^ (∂ i^ G^ /^ ∂ xi )II.
References
[1] J.A. Lipa et al ., Phys. Rev. Lett. 76 (1995) 944.
Appendix A: Thermodynamic processes for an ideal gas
Table A.1: Summary of the properties of some reversible processes involving an ideal gas
Process Isochoric Isobaric Isothermal Adiabatic Polytropic
Fixed parameter V^ =^ cst^ p^ =^ cst^ T^ =^ cst^ δq^ = 0
p–v diagram
Relationship between state variables
B B A A
p T p T
=
Charles law
B B A A
v T v T
=
Gay–Lussac law
p v A A (^) = p v B B (^) = nRT A
Mariotte law
p v A A (^) p v B B γ = γ Laplace law γ = c (^) p / cV
A A B B p v z^ = p vz 1 < z < γ
Δ U = U B − U A (J) B^ A
ncV ( T − T ) ncV ( T B (^) − T A) ncV ( T B (^) − T A) = 0 ncV ( T B (^) − T A) ncV ( T B (^) − T A)
Δ H = H B − H A (J) B^ A
nc (^) p ( T − T ) ncp ( T B (^) − T A) ncp ( T B (^) − T A) = 0 nc (^) p ( T B (^) − T A) nc (^) p ( T B (^) − T A)
Heat Q AB (J)
∆ U (^) ∆ H B A
ln
V nRT V
0 ( (^) B A) v 1 nc z T T z
− γ − −
Work W AB (J)
0
( ) ( )
B A B A
p V V nR T T
− − = −
B A A
nRT ln V V
− (^) ∆ U 1 ( (^) B A) v 1 nc T T z
γ − (^) − −
Δ S = S B − S A (J·K −1)
B A
ncV ln T T
B B A A
c (^) p ln T^ cp ln V T V
= B A
nR ln V V^0
B A
ln 1
nR n^ V V
γ γ
− −
