Misplaced Pages

This article has multiple issues. Please help improve it or discuss these issues on the talk page. (Learn how and when to remove these messages) This article may require copy editing for grammar, style, cohesion, tone, or spelling. You can assist by editing it. (December 2024) (Learn how and when to remove this message) This article may require cleanup to meet Misplaced Pages's quality standards. The specific problem is: Numerous WP:MOS violations that need to be addressed, needs wikification. Please help improve this article if you can. (December 2024) (Learn how and when to remove this message) This article may need to be rewritten to comply with Misplaced Pages's quality standards. You can help. The talk page may contain suggestions. (December 2024) This article needs additional citations for verification. Please help improve this article by adding citations to reliable sources. Unsourced material may be challenged and removed. Find sources: "Van der Waals equation" – news · newspapers · books · scholar · JSTOR (December 2024) (Learn how and when to remove this message) (Learn how and when to remove this message)

The van der Waals equation is a mathematical formula that describes the behavior of real gases. It is named after Dutch physicist Johannes Diderik van der Waals. It is an equation of state that relates the pressure, temperature, and molar volume in a fluid. However, it can be written in terms of other, equivalent, properties in place of the molar volume, for example specific volume, or number density. The equation modifies the ideal gas law in two ways. First its particles have a finite diameter, whereas the ideal gas consists of point particles with no extension. Second, its particles interact with one other, whereas the particles of an ideal gas move as though they were alone in the volume.

The surface calculated from the ideal gas equation of state is drawn in Fig. A. This universal (all ideal gases are represented by it) surface is normalized so that the black dot, with coordinates ${\displaystyle p_{0},v_{0},T_{0}}$, appears at the location (1,1,1) of the 3 dimensional plot space. This device makes it easy to compare this surface with the one generated by the van der Waals equation in Fig. C. Figures A and C are drawn with the same scales and limits; they also present the two surfaces from the same viewpoint to make the comparison easier. Whereas the ideal gas surface is rather plain, the van der Waals surface has an interesting fold.

Figures B and C show different views of the surface calculated from the van der Waals equation. The fold seen on this surface is what enables the equation to predict the phenomenon of liquid--vapor phase change. This fold develops from a critical point defined by specific, critical, values of pressure, temperature, and molar volume. The surface is plotted using dimensionless variables that are formed by the ratio of each property to its respective critical value. This locates the critical point at the coordinates (1,1,1) of the space. When drawn using these dimensionless properties, this surface is, like that of the ideal gas, also universal. Moreover, it represents all real substances thttps://en.wikipedia.org/Van_der_Waals_equationo a remarkably high degree of accuracy. This principle of corresponding states, developed by van der Waals from his equation, has become one of the fundamental ideas in the thermodynamics of fluids.

The boundary of the fold on the surface is marked, on each side of the critical point, by the spinodal curve, identified in Fig. B, and seen in Figs. B and C. However, this curve does not define the location of the phase change. That place is given by the saturation curve, a curve that is not specified by the properties of the surface alone. The saturation curve is the locus of saturated liquid and vapor states which, being in equilibrium with each other, can coexist. The saturated liquid and vapor curves are identified in Fig. B. Together they comprise the saturation (or coexistence) curve seen in Figs. B and C. The inset in Fig. B shows the mixture states which are a combination of the saturated liquid and vapor states that exist at each end of the line, intersections of the mixture line and its isotherm. However, these mixture states are not part of the surface generated by the van der Waals equation; they are not solutions of the equation.

Although it is hard to imagine today when the quantum nature of physics is learned about early in life, it was only in 1909 that the scientific debate about the nature of matter, discrete or continuous, was finally settled. Indeed, at the time van der Waals created his equation, which he based on the idea that fluids are composed of discrete particles, few scientists believed that such particles really existed. They were regarded as purely metaphysical constructs that added nothing useful to the knowledge obtained from the results of experimental observations. However, the theoretical explanation of the critical point, which had been discovered a few years earlier, and later its qualitative and quantitative agreement with experiments cemented its acceptance in the scientific community. Ultimately these accomplishments won him the 1910 Nobel prize in physics. Today the equation is recognized as an important model of phase change processes. Van der Waals also adapted his equation so that it applied to a binary mixture of fluids. He, and others, then used the modified equation to discover a host of important facts about the phase equilibria of such fluids. This application, expanded to treat multi-component mixtures, has extended the predictive ability of the equation to fluids of industrial and commercial importance. In this arena it has spawned many similar equations in a continuing attempt by engineers to improve their ability to understand and manage these fluids; it remains relevant to the present.

Behavior of the equation

One way to write the van der Waals equation is:

$${\displaystyle p={\frac {RT}{v-b}}-{\frac {a}{v^{2}}}}$$

where ${\displaystyle p}$ is pressure, ${\displaystyle T}$ is temperature, and ${\displaystyle v=VN_{\text{A}}/N}$ is molar volume. In addition ${\displaystyle N_{\text{A}}}$ is the Avogadro constant, ${\displaystyle V}$ is the volume, and ${\displaystyle N}$ is the number of molecules (the ratio ${\displaystyle N/N_{\text{A}}}$ is a physical quantity with base unit mole (symbol mol) in the SI). Finally ${\displaystyle R=N_{\text{A}}k}$ is the universal gas constant, ${\displaystyle k}$ is the Boltzmann constant, and ${\displaystyle a}$ and ${\displaystyle b}$ are experimentally determinable, substance-specific constants.

As noted previously, when van der Waals created his equation the idea that fluids were composed of rapidly moving particles was believed by very few scientists. Moreover, those who thought so had no knowledge of the atomic/molecular structure. The simplest conception, and the easiest to model mathematically was a hard sphere and that is what van der Waals used. In that case two particles of diameter, ${\displaystyle \sigma }$, would come into contact when their centers were a distance ${\displaystyle \sigma }$ apart, hence the center of the one was excluded from a spherical volume equal to ${\displaystyle 4\pi \sigma ^{3}/3}$ about the other. That is 8 times ${\displaystyle V_{0}}$, the volume of each particle of radius ${\displaystyle \sigma /2}$, but there are 2 particles which gives 4 times the volume per particle. The total excluded volume is then ${\displaystyle B=4NV_{0}}$, 4 times the volume of all the particles. Van der Waals and his contemporaries used an alternative, but equivalent, analysis based on the mean free path between molecular collisions that gave this result. From the fact that the volume fraction of particles, ${\displaystyle f=(V-B)/V}$ must be positive, van der Waals noted that as ${\displaystyle N}$ becomes larger the factor 4 must decrease (for spheres there is a known minimum ${\displaystyle 0<f_{\mbox{min}}<f\leq 1}$), but he was never able to determine the nature of the decrease. The constant ${\displaystyle b=N_{\text{A}}B/N}$ in the equation above has dimension molar volume, . The constant ${\displaystyle a}$ expresses the strength of the hypothesized interparticle attraction. Van der Waals only had as a model Newton's law of gravitation, in which two particles attracted in proportion to the product of their masses. Thus he argued that in his case the attractive pressure was proportional to the square of the density. The proportionality constant, ${\displaystyle a}$, when written in the form used above, has dimension pressure times molar volume squared, which is also molar energy times molar volume.

The intermolecular force was later conveniently described by the negative derivative of a pair potential function. For spherically symmetric particles this is most simply a function of separation distance with a single characteristic length, ${\displaystyle \sigma }$, and a minimum energy, ${\displaystyle -\varepsilon }$ (with ${\displaystyle \varepsilon \geq 0}$). Two of the many such functions that have been suggested are shown in the accompanying plot.

A modern theory based on statistical mechanics produces the same result for ${\displaystyle b=4N_{\text{A}}}$ obtained by van der Waals and his contemporaries. This result is valid for any pair potential for which the increase in ${\displaystyle \varphi >0}$ is sufficiently rapid. This includes the hard sphere model for which the increase is infinitely rapid and the result is exact. Indeed the Sutherland potential most accurately models van der Waals' conception of a molecule. It also includes potentials that do not represent hard sphere force interactions provided that the increase in ${\displaystyle \varphi {\mbox{ for }}r<\sigma }$ is fast enough, but then it is approximate; increasingly better the faster the increase. In that case ${\displaystyle \sigma }$ is only an "effective diameter" of the molecule. This theory also produces ${\displaystyle a=IN_{\text{A}}\varepsilon b}$ where ${\displaystyle I}$ is a number that depends on the shape of the potential function, ${\displaystyle \varphi (r)/\varepsilon }$. However, this result is only valid when the potential is weak, namely, when the minimum potential energy is very much smaller than the thermal energy, ${\displaystyle \varepsilon \ll kT}$.

In his book (see references and ) Boltzmann wrote equations using ${\displaystyle V/M}$ (specific volume) in place of ${\displaystyle VN_{\text{A}}/N}$ (molar volume) used here, Gibbs did as well, so do most engineers. Also the property, ${\displaystyle V/N=1/\rho _{N}}$ the reciprocal of number density, is used by physicists, but there is no essential difference between equations written with any of these properties. Equations of state written using molar volume contain ${\displaystyle R}$, those using specific volume contain ${\displaystyle R/{\bar {m}}}$ (the substance specific ${\displaystyle {\bar {m}}=mN_{\text{A}}}$ is the molar mass with ${\displaystyle m}$ the mass of a single particle), and those written with number density contain ${\displaystyle k}$.

Once ${\displaystyle a}$ and ${\displaystyle b}$ are experimentally determined for a given substance, the van der Waals equation can be used to predict the boiling point at any given pressure, the critical point (defined by pressure and temperature values, ${\displaystyle p_{\text{c}}}$, ${\displaystyle T_{\text{c}}}$ such that the substance cannot be liquefied either when ${\displaystyle p>p_{\text{c}}}$ no matter how low the temperature, or when ${\displaystyle T>T_{\text{c}}}$ no matter how high the pressure; ${\displaystyle p_{c},T_{c}}$ uniquely define ${\displaystyle v_{c}}$), and other attributes. These predictions are accurate for only a few substances. For most simple fluids they are only a valuable approximation. The equation also explains why superheated liquids can exist above their boiling point and subcooled vapors can exist below their condensation point.

The graph on the right follows from the intersection of the surface shown in Figs. B and C and 4 constant pressure planes. Each intersection produces a curve in the ${\displaystyle T,v}$ plane corresponding to the value of the pressure chosen.

On the red isobar (another name for a constant pressure curve), ${\displaystyle p=2p_{\text{c}}}$, the slope is positive over the entire range, ${\displaystyle b\leq v<\infty }$ (although the plot only shows a finite quadrant). This describes a fluid as homogeneous for all ${\displaystyle T}$, and is characteristic of all supercritical isobars ${\displaystyle p>p_{\text{c}}.}$

The green isobar, ${\displaystyle p=0.2\,p_{\text{c}}}$, has a region of negative slope. This region consists of states that are unstable and therefore never observed (for this reason this region is shown dotted gray). The green curve thus consists of two disconnected branches; a vapor on the right, and a denser liquid on the left. For this pressure, at a temperature, ${\displaystyle T_{s}}$, specified by mechanical, thermal, and material equilibrium, and shown as green circles on the curve, the boiling (saturated) liquid, ${\displaystyle v_{f}}$, (the left circle) and condensing (saturated) vapor, ${\displaystyle v_{g}}$, (the right circle) coexist. Due to gravity the denser liquid appears below the vapor, and a meniscus is seen between them. This heterogeneous combination of coexisting liquid and vapor is the phase change. Heating the fluid in this state increases the fraction of vapor in the mixture; its ${\displaystyle v}$, an average of ${\displaystyle v_{f}}$ and ${\displaystyle v_{g}}$ weighted by this fraction, increases while ${\displaystyle T_{s}}$ remains the same. This is shown as the dotted gray line, because it does not represent a solution of the equation; however, it does describe the observed behavior. The points above ${\displaystyle T_{s}}$, superheated liquid, and those below it, subcooled vapor, are metastable; a sufficiently strong disturbance causes them to transform to the stable alternative. Consequently they are shown dashed.

All this describes a fluid as a stable vapor for ${\displaystyle T>T_{\text{s}}}$, a stable liquid for ${\displaystyle T<T_{\text{s}}}$, and a mixture of liquid and vapor at ${\displaystyle T=T_{\text{s}}}$, that also supports metastable states of subcooled vapor and superheated liquid. It is characteristic of all subcritical isobars ${\displaystyle 0<p<p_{\text{c}}}$, where ${\displaystyle T_{\text{s}}}$ is a function of ${\displaystyle p}$.

The orange isobar is the critical one on which the minimum and maximum are equal. The critical point lies on this isobar.

The black isobar is the limit of positive pressures, although drawn solid none of its points represent stable solutions, they are either metastable (positive or zero slope) or unstable (negative slope). Interestingly, states of negative pressure (tension) exist. They lie below the black isobar, and although they are not not drawn in this figure, they form those parts of the surfaces seen in Figs. B and C that lie below the zero pressure plane. In this, ${\displaystyle T,v}$, plane they have a parabola like shape and like the zero pressure isobar their states are all either metastable (positive or zero slope) or unstable (negative slope).

Relationship to the ideal gas law

The ideal gas law follows from the van der Waals equation whenever ${\displaystyle v}$ is sufficiently large (or correspondingly whenever the molar density, ${\displaystyle \rho =1/v}$, is sufficiently small), Specifically

• when ${\displaystyle v\gg b}$, then ${\displaystyle v-b}$ is numerically indistinguishable from ${\displaystyle v}$,
• and when ${\displaystyle v\gg (a/p)^{1/2}}$, then ${\displaystyle p+a/v^{2}}$ is numerically indistinguishable from ${\displaystyle p}$.

Putting these two approximations into the van der Waals equation when ${\displaystyle v}$ is large enough that both inequalities are satisfied reduces it to

$${\displaystyle p=RT/v\quad {\mbox{or in terms of }}V{\mbox{ and }}N\quad pV=NkT}$$

which is the ideal gas law. This is not surprising since the van der Waals equation was constructed from the ideal gas equation in order to obtain an equation valid beyond the limit of ideal gas behavior.

What is truly remarkable is the extent to which van der Waals succeeded. Indeed, Epstein in his classic thermodynamics textbook began his discussion of the van der Waals equation by writing, "In spite of its simplicity, it comprehends both the gaseous and the liquid state and brings out, in a most remarkable way, all the phenomena pertaining to the continuity of these two states". Also in Volume 5 of his Lectures on Theoretical Physics Sommerfeld, in addition to noting that "Boltzmann described van der Waals as the Newton of real gases", also wrote "It is very remarkable that the theory due to van der Waals is in a position to predict, at least qualitatively, the unstable states" that are associated with the phase change process.

Utility of the equation

The equation has been, and remains very useful because:

• its specific heat at constant volume, ${\displaystyle c_{v}}$, can be shown to be a function of ${\displaystyle T}$ only, and its thermodynamic properties, internal energy ${\displaystyle u}$, entropy ${\displaystyle s}$, as well as the specific heat at constant pressure ${\displaystyle c_{p}}$ have simple analytic expressions [this is also true of enthalpy ${\displaystyle h=u+pv}$, Helmholtz free energy ${\displaystyle f=u-Ts}$, and Gibbs free energy ${\displaystyle g=u+pv-Ts=f+pv=h-Ts}$]

• Its coefficient of thermal expansion, ${\displaystyle \alpha =(\partial _{T}v|_{p})/v}$ has a simple analytic expression [this is also true of its isothermal compressibility, ${\displaystyle \kappa _{T}=-(\partial _{p}v|_{T})/v}$]
• it explains the existence of the critical point and the liquid–vapor phase transition including the observed metastable states
• it establishes the law of corresponding states
• its Joule–Thomson coefficient and associated inversion curve, which were instrumental in the development of the commercial liquefaction of gases, have simple analytic expressions.

In addition its vapor pressure curve (also called the coexistence, or saturation, curve) has a simple analytic solution. It depicts the liquid metals, Mercury and Cesium, quantitatively, and describes most real fluids qualitatively. Consequently it can be regarded as one member of a family of equations of state, that depend on a molecular parameter such as the critical compressibility factor, ${\displaystyle Z_{\text{c}}=p_{\text{c}}v_{\text{c}}/(RT_{\text{c}})}$, or the Pitzer (acentric) factor, ${\displaystyle \omega =-\log-1}$, where ${\displaystyle p_{s}/p_{\text{c}}=p_{s}(T/T_{\text{c}},\omega )}$ is a dimensionless saturation pressure, and log is the logarithm base 10. Consequently, the equation plays an important role in the modern theory of phase transitions.

All this makes it a worthwhile pedagogical tool for physics, chemistry, and engineering lecturers, in addition to being a useful mathematical model which can aid student understanding.

History

In 1857 Rudolf Clausius published The Nature of the Motion which We Call Heat. In it he derived the relation ${\displaystyle p=(N/V)m{\overline {c^{2}}}/3}$ for the pressure, ${\displaystyle p}$, in a gas, composed of particles in motion, with number density ${\displaystyle N/V}$, mass ${\displaystyle m}$, and mean square speed ${\displaystyle {\overline {c^{2}}}}$. He then noted that using the classical laws of Boyle and Charles one could write ${\displaystyle m{\overline {c^{2}}}/3=kT}$ with ${\displaystyle k}$ a constant of proportionality. Hence temperature was proportional to the average kinetic energy of the particles. This article inspired further work based on the twin ideas that substances are composed of indivisible particles, and that heat is a consequence of the particle motion; movement that evolves in accordance with Newton's laws. The work, known as the kinetic theory of gases, was done principally by Clausius, James Clerk Maxwell, and Ludwig Boltzmann. At about the same time J. Willard Gibbs also contributed, and advanced it by converting it into statistical mechanics.

This environment influenced Johannes Diderik van der Waals. After initially pursuing a teaching credential, he was accepted for doctoral studies at the University of Leiden under Pieter Rijke. This led, in 1873, to a dissertation that provided a simple, particle based, equation that described the gas–liquid change of state, the origin of a critical temperature, and the concept of corresponding states. The equation is based on two premises, first that fluids are composed of particles with non-zero volumes, and second that at a large enough distance each particle exerts an attractive force on all other particles in its vicinity. These forces were called by Boltzmann van der Waals cohesive forces.

In 1869 Irish professor of chemistry Thomas Andrews at Queen's University Belfast in a paper entitled On the Continuity of the Gaseous and Liquid States of Matter, displayed an experimentally obtained set of isotherms of carbonic acid, H${\displaystyle _{2}}$CO${\displaystyle _{3}}$, that showed at low temperatures a jump in density at a certain pressure, while at higher temperatures there was no abrupt change; the figure can be seen here. Andrews called the isotherm at which the jump just disappeared the critical point. Given the similarity of the titles of this paper and van der Waals subsequent thesis one might think that van der Waals set out to develop a theoretical explanation of Andrews' experiments; however, this is not what happened. Van der Waals began work by trying to determine a mollecular attraction that appeared in Laplace's theory of capillarity, and only after establishing his equation he tested it using Andrews results.

By 1877 sprays of both liquid oxygen and liquid nitrogen had been produced, and a new field of research, low temperature physics, had been opened. The van der Waals equation played a part in all this especially with respect to the liquefaction of hydrogen and helium which was finally achieved in 1908. From measurements of ${\displaystyle p_{1},T_{1}}$ and ${\displaystyle p_{2},T_{2}}$ in two states with the same density, the van der Waals equation produces the values,

$${\displaystyle b=v-{\frac {R(T_{2}-T_{1})}{p_{2}-p_{1}}}\qquad {\mbox{and}}\qquad a=v^{2}{\frac {p_{2}T_{1}-p_{1}T_{2}}{T_{2}-T_{1}}}.}$$

Thus from two such measurements of pressure and temperature one could determine ${\displaystyle a}$ and ${\displaystyle b}$, and from these values calculate the expected critical pressure, temperature, and molar volume. Goodstein summarized this contribution of the van der Waals equation as follows:

All this labor required considerable faith in the belief that gas–liquid systems were all basically the same, even if no one had ever seen the liquid phase. This faith arose out of the repeated success of the van der Waals theory, which is essentially a universal equation of state, independent of the details of any particular substance once it has been properly scaled. ... As a result, not only was it possible to believe that hydrogen could be liquefied. but it was even possible to predict the necessary temperature and pressure.

Van der Waals was awarded the Nobel Prize in 1910, in recognition of the contribution of his formulation of this "equation of state for gases and liquids".

As noted previously, modern day studies of first order phase changes make use of the van der Waals equation together with the Gibbs criterion, equal chemical potential of each phase, as a model of the phenomenon. This model has an analytic coexistence (saturation) curve expressed parametrically, ${\displaystyle p_{s}=f_{p}(y),T_{s}=f_{T}(y)}$ (the parameter ${\displaystyle y}$ is related to the entropy difference between the two phases), that was first obtained by Plank, was known to Gibbs and others, and was later derived in a beautifully simple and elegant manner by Lekner. A summary of Lekner's solution is presented in a subsequent section, and a more complete discussion in the Maxwell construction.

Critical point and corresponding states

Figure 1 shows four isotherms of the van der Waals equation (abbreviated as vdW) on a pressure, molar volume plane. The essential character of these curves is that:

1. at some critical temperature, ${\displaystyle T=T_{\text{c}}}$ the slope is negative, ${\displaystyle \partial p/\partial v|_{T}<0}$, everywhere except at a single point, the critical point, ${\displaystyle p=p_{\text{c}},v=v_{\text{c}}}$, where both the slope and curvature are zero, ${\displaystyle \partial p/\partial v|_{T}=\partial ^{2}p/\partial v^{2}|_{T}=0;}$
2. at higher temperatures the slope of the isotherms is everywhere negative (values of ${\displaystyle p,T}$ for which the equation has 1 real root for ${\displaystyle v}$);
3. at lower temperatures there are two points on each isotherm where the slope is zero (values of ${\displaystyle p}$, ${\displaystyle T}$ for which the equation has 3 real roots for ${\displaystyle v}$)

Evaluating the two partial derivatives in 1) using the vdW equation and equating them to zero produces, ${\displaystyle v_{\text{c}}=3b,T_{\text{c}}=8a/(27Rb)}$, and using these in the equation gives ${\displaystyle p_{\text{c}}=a/27b^{2}}$.

This calculation can also be done algebraically by noting that the vdW equation can be written as a cubic in ${\displaystyle v}$, which at the critical point is,

$${\displaystyle p_{\text{c}}v^{3}-(p_{\text{c}}b+RT_{\text{c}})v^{2}+av-ab=0.}$$

Moreover, at the critical point all three roots coalesce so it can also be written as

$${\displaystyle (v-v_{\text{c}})^{3}=v^{3}-3v_{\text{c}}v^{2}+3v_{\text{c}}^{2}v-v_{\text{c}}^{3}=0}$$

Then dividing the first by ${\displaystyle p_{\text{c}}}$, and noting that these two cubic equations are the same when all their coefficients are equal gives three equations, ${\displaystyle b+RT_{\text{c}}/p_{\text{c}}=3v_{\text{c}}\quad a/p_{\text{c}}=3v_{\text{c}}^{2}\quad ab/p_{\text{c}}=v_{\text{c}}^{3}}$, whose solution produces the previous results for ${\displaystyle p_{\text{c}},v_{\text{c}},T_{\text{c}}}$.

Using these critical values to define reduced properties ${\displaystyle p_{r}=p/p_{\text{c}},T_{r}=T/T_{\text{c}},v_{r}=v/v_{\text{c}}}$ renders the equation in the dimensionless form used to construct Fig. 1

$${\displaystyle p_{r}={\frac {8T_{r}}{3v_{r}-1}}-{\frac {3}{v_{r}^{2}}}}$$

This dimensionless form is a similarity relation; it indicates that all vdW fluids at the same ${\displaystyle T_{r}}$ will plot on the same curve. It expresses the law of corresponding states which Boltzmann described as follows:

All the constants characterizing the gas have dropped out of this equation. If one bases measurements on the van der Waals units , then he obtains the same equation of state for all gases. ... Only the values of the critical volume, pressure, and temperature depend on the nature of the particular substance; the numbers that express the actual volume, pressure, and temperature as multiples of the critical values satisfy the same equation for all substances. In other words, the same equation relates the reduced volume, reduced pressure, and reduced temperature for all substances.

Obviously such a broad general relation is unlikely to be correct; nevertheless, the fact that one can obtain from it an essentially correct description of actual phenomena is very remarkable.

This "law" is just a special case of dimensional analysis in which an equation containing 6 dimensional quantities, ${\displaystyle p,v,T,a,b,R}$, and 3 independent dimensions, , , (independent means that "none of the dimensions of these quantities can be represented as a product of powers of the dimensions of the remaining quantities", and =), must be expressible in terms of 6 − 3 = 3 dimensionless groups. Here ${\displaystyle v^{*}=b}$ is a characteristic molar volume, ${\displaystyle p^{*}=a/b^{2}}$ a characteristic pressure, and ${\displaystyle T^{*}=a/(Rb)}$ a characteristic temperature, and the 3 dimensionless groups are ${\displaystyle p/p^{*},v/v^{*},T/T^{*}}$. According to dimensional analysis the equation must then have the form ${\displaystyle p/p^{*}=\Phi (v/v^{*},T/T^{*})}$, a general similarity relation. In his discussion of the vdW equation Sommerfeld also mentioned this point. The reduced properties defined previously are ${\displaystyle p_{r}=27(p/p^{*})}$, ${\displaystyle v_{r}=(1/3)(v/v^{*})}$, and ${\displaystyle T_{r}=(27/8)(T/T^{*})}$. Recent research has suggested that there is a family of equations of state that depend on an additional dimensionless group, and this provides a more exact correlation of properties. Nevertheless, as Boltzmann observed, the van der Waals equation provides an essentially correct description.

The vdW equation produces ${\displaystyle Z_{\text{c}}=p_{\text{c}}v_{\text{c}}/(RT_{\text{c}})=3/8}$, while for most real fluids ${\displaystyle 0.23<Z_{\text{c}}<0.31}$. Thus most real fluids do not satisfy this condition, and consequently their behavior is only described qualitatively by the vdW equation. However, the vdW equation of state is a member of a family of state equations based on the Pitzer (acentric) factor, ${\displaystyle \omega }$, and the liquid metals, Mercury and Cesium, are well approximated by it.

Thermodynamic properties

The properties molar internal energy, ${\displaystyle u}$, and entropy, ${\displaystyle s}$, defined by the first and second laws of thermodynamics, hence all thermodynamic properties of a simple compressible substance, can be specified, up to a constant of integration, by two measurable functions, a mechanical equation of state, ${\displaystyle p=p(v,T)}$, and a constant volume specific heat, ${\displaystyle c_{v}(v,T)}$.

Internal energy and specific heat at constant volume

The internal energy is given by the energetic equation of state,

$${\displaystyle u-C_{u}=\int \,c_{v}(v,T)\,dT+\int \,\left\,dv=\int \,c_{v}(v,T)\,dT+\int \,\left\,dv}$$

where ${\displaystyle C_{u}}$ is an arbitrary constant of integration.

Now in order for ${\displaystyle du(v,T)}$ to be an exact differential, namely that ${\displaystyle u(v,T)}$ be continuous with continuous partial derivatives, its second mixed partial derivatives must also be equal, ${\displaystyle \partial _{v}\partial _{T}u=\partial _{T}\partial _{v}u}$. Then with ${\displaystyle c_{v}=\partial _{T}u}$ this condition can be written simply as ${\displaystyle \partial _{v}c(v,T)=\partial _{T}}$. Differentiating ${\displaystyle p/T}$ for the vdW equation gives ${\displaystyle T^{2}\partial _{T}(p/T)]=a/v^{2}}$, so ${\displaystyle \partial _{v}c_{v}=0}$. Consequently ${\displaystyle c_{v}=c_{v}(T)}$ for a vdW fluid exactly as it is for an ideal gas. To keep things simple it is regarded as a constant here, ${\displaystyle c_{v}=cR}$ with ${\displaystyle c}$ a number. Then both integrals can be easily evaluated and the result is

$${\displaystyle u-C_{u}=cRT-a/v}$$

This is the energetic equation of state for a perfect vdW fluid. By making a dimensional analysis (what might be called extending the principle of corresponding states to other thermodynamic properties) it can be written simply in reduced form as,

$${\displaystyle u_{r}-{\mbox{C}}_{u}=cT_{r}-9/(8v_{r})}$$

where ${\displaystyle u_{r}=u/(RT_{\text{c}})}$ and ${\displaystyle {\mbox{C}}_{u}}$ is a dimensionless constant.

Enthalpy

The enthalpy is ${\displaystyle h=u+pv}$, and the product ${\displaystyle pv}$ is just ${\displaystyle pv=RTv/(v-b)-a/v}$. Then

${\displaystyle h}$ is simply$${\displaystyle h-C_{u}=RT-2a/v}$$

This is the enthalpic equation of state for a perfect vdW fluid, or in reduced form,

$${\displaystyle h_{r}-{\mbox{C}}_{u}=T_{r}-9/(4v_{r})\quad {\mbox{where}}\quad h_{r}=h/(RT_{\text{c}}c)}$$

Entropy

The entropy is given by the entropic equation of state:

$${\displaystyle s-C_{s}=\int \,c_{v}(T)\,{\frac {dT}{T}}+\int \,{\frac {\partial p}{\partial T}}\,dv}$$

Using ${\displaystyle c_{v}=cR}$ as before, and integrating the second term using ${\displaystyle \partial _{T}p=R/(v-b)}$ we obtain simply

$${\displaystyle s-C_{s}=R\ln}$$

This is the entropic equation of state for a perfect vdW fluid, or in reduced form,

$${\displaystyle s_{r}-{\mbox{C}}_{s}=\ln}$$

Helmholtz free energy

The Helmholtz free energy is ${\displaystyle f=u-Ts}$ so combining the previous results

$${\displaystyle f=C_{u}+cT-a/v-T\{C_{s}+R\ln\}}$$

This is the Helmholtz free energy for a perfect vdw fluid, or in reduced form

$${\displaystyle f_{r}={\mbox{C}}_{u}+cT_{r}-9/(8v_{r})-T_{r}\{{\mbox{C}}_{s}+\ln\}}$$

Gibbs free energy

The Gibbs free energy is ${\displaystyle g=h-Ts}$ so combining the previous results gives

$${\displaystyle g-C_{u}=\{c+v/(v-b)-C_{s}-\ln\}RT-2a/v}$$

This is the Gibbs free energy for a perfect vdW fluid, or in reduced form

$${\displaystyle g_{r}-{\mbox{C}}_{u}=\{c+3v_{r}/(3v_{r}-1)-{\mbox{C}}_{s}-\ln\}T_{r}-9/(4v_{r})}$$

Thermodynamic derivatives: α, κ_T and c_p

The two first partial derivatives of the vdW equation are

$${\displaystyle \left.{\frac {\partial p}{\partial T}}\right)_{v}={\frac {R}{v-b}}={\frac {\alpha }{\kappa _{T}}}\quad {\mbox{and}}\quad \left.{\frac {\partial p}{\partial v}}\right)_{T}=-{\frac {RT}{(v-b)^{2}}}+{\frac {2a}{v^{3}}}=-{\frac {1}{v\kappa _{T}}}}$$

Here ${\displaystyle \kappa _{T}=-v^{-1}\partial _{p}v}$, the isothermal compressibility, is a measure of the relative increase of volume from an increase of pressure, at constant temperature, while ${\displaystyle \alpha =v^{-1}\partial _{T}v_{p}}$, the coefficient of thermal expansion, is a measure of the relative increase of volume from an increase of temperature, at constant pressure. Therefore,

$${\displaystyle \kappa _{T}={\frac {v^{2}(v-b)^{2}}{RTv^{3}-2a(v-b)^{2}}}\quad {\mbox{and}}\quad \alpha ={\frac {Rv^{2}(v-b)}{RTv^{3}-2a(v-b)^{2}}}}$$

In the limit ${\displaystyle v\rightarrow \infty }$ ${\displaystyle \alpha =1/T}$ while ${\displaystyle \kappa _{T}=v/(RT)}$. Since the vdW equation in this limit becomes ${\displaystyle p=RT/v}$, finally ${\displaystyle \kappa _{T}=1/p}$. Both of these are the ideal gas values, which is consistent because, as noted earlier, the vdW fluid behaves like an ideal gas in this limit.

The specific heat at constant pressure, ${\displaystyle c_{p}}$ is defined as the partial derivative ${\displaystyle c_{p}=\partial _{T}h|_{p}}$. However, it is not independent of ${\displaystyle c_{v}}$, they are related by the Mayer equation, ${\displaystyle c_{p}-c_{v}=-T(\partial _{T}p)^{2}/\partial _{v}p=Tv\alpha ^{2}/\kappa _{T}}$. Then the two partials of the vdW equation can be used to express ${\displaystyle c_{p}}$ as,

$${\displaystyle c_{p}(v,T)-c_{v}(T)={\frac {R^{2}Tv^{3}}{RTv^{3}-2a(v-b)^{2}}}\geq R}$$

Here in the limit ${\displaystyle v\rightarrow \infty }$, ${\displaystyle c_{p}-c_{v}=R}$, which is also the ideal gas result as expected; however the limit ${\displaystyle v\rightarrow b}$ gives the same result, which does not agree with experiments on liquids.

In this liquid limit we also find ${\displaystyle \alpha =\kappa _{T}=0}$, namely that the vdW liquid is incompressible. Moreover, since ${\displaystyle \partial _{T}p=-\partial _{T}v/\partial _{p}v=\alpha /\kappa _{T}=\infty }$, it is also mechanically incompressible, that is ${\displaystyle \kappa _{T}\rightarrow 0}$ faster than ${\displaystyle \alpha }$.

Finally ${\displaystyle c_{p},\alpha }$, and ${\displaystyle \kappa _{T}}$ are all infinite on the curve ${\displaystyle T=2a(v-b)^{2}/(Rv^{3})=T_{\text{c}}(3v_{r}-1)^{2}/(4v_{r}^{3})}$. This curve, called the spinodal curve, is defined by ${\displaystyle \kappa _{T}^{-1}=0}$, and is discussed at length in the next section.

Stability

According to the extremum principle of thermodynamics ${\displaystyle dS=0}$ and ${\displaystyle d^{2}S<0}$, namely that at equilibrium the entropy is a maximum. This leads to a requirement that ${\displaystyle \partial p/\partial v|_{T}<0}$. This mathematical criterion expresses a physical condition which Epstein described as follows:

"It is obvious that this middle part, dotted in our curves , can have no physical reality. In fact, let us imagine the fluid in a state corresponding to this part of the curve contained in a heat conducting vertical cylinder whose top is formed by a piston. The piston can slide up and down in the cylinder, and we put on it a load exactly balancing the pressure of the gas. If we take a little weight off the piston, there will no longer be equilibrium and it will begin to move upward. However, as it moves the volume of the gas increases and with it its pressure. The resultant force on the piston gets larger, retaining its upward direction. The piston will, therefore, continue to move and the gas to expand until it reaches the state represented by the maximum of the isotherm. Vice versa, if we add ever so little to the load of the balanced piston, the gas will collapse to the state corresponding to the minimum of the isotherm"

While on an isotherm ${\displaystyle T>T_{\text{c}}}$ this requirement is satisfied everywhere so all states are gas, those states on an isotherm, ${\displaystyle T<T_{\text{c}}}$ which lie between the local minimum, ${\displaystyle v_{\rm {min}}}$, and local maximum, ${\displaystyle v_{\rm {max}}}$, for which ${\displaystyle \partial p/\partial v|_{T}>0}$ (shown dashed gray in Fig. 1), are unstable and thus not observed. This is the genesis of the phase change; there is a range ${\displaystyle v_{\rm {min}}\leq v\leq v_{\rm {max}}}$, for which no observable states exist. The states for ${\displaystyle v<v_{\rm {min}}}$ are liquid, and for ${\displaystyle v>v_{\rm {max}}>v_{\rm {min}}}$ are vapor; the denser liquid lies below the vapor due to gravity. The transition points, states with zero slope, are called spinodal points. Their locus is the spinodal curve that separates the regions of the plane for which liquid, vapor, and gas exist from a region where no observable homogeneous states exist. This spinodal curve is obtained here from the vdW equation by differentiation (or equivalently from ${\displaystyle \kappa _{T}=\infty }$) as

$${\displaystyle T_{\rm {sp}}=2a{\frac {(v-b)^{2}}{Rv^{3}}}=T_{\text{c}}{\frac {(3v_{r}-1)^{2}}{4v_{r}^{3}}}\qquad p_{\rm {sp}}={\frac {a(v-2b)}{v^{3}}}=p_{\text{c}}{\frac {(3v_{r}-2)}{v_{r}^{3}}}}$$

A projection of this space curve is plotted in Fig. 1 as the black dash dot curve. It passes through the critical point which is also a spinodal point.

Saturation

Although the gap in ${\displaystyle v}$ delimited by the two spinodal points on an isotherm (e.g. ${\displaystyle T_{r}=7/8}$ shown in Fig. 1) is the origin of the phase change, the spinodal points do not represent its full extent, because both states, saturated liquid and saturated vapor coexist in equlilbrium; they both must have the same pressure as well as the same temperature. Thus the phase change is characterized, at temperature ${\displaystyle T_{s}}$, by a pressure ${\displaystyle p_{\rm {min}}<p_{s}<p_{\rm {max}}}$ that lies between that of the minimum and maximum spinodal points, and with molar volumes of liquid, ${\displaystyle v_{f}<v_{\rm {min}}}$ and vapor ${\displaystyle v_{g}>v_{\rm {max}}}$. Then from the vdW equation applied to these saturated liquid and vapor states

$${\displaystyle p_{s}={\frac {RT_{s}}{v_{f}-b}}-{\frac {a}{v_{f}^{2}}}\quad {\mbox{and}}\quad p_{s}={\frac {RT_{s}}{v_{g}-b}}-{\frac {a}{v_{g}^{2}}}}$$

These two vdW equations contain 4 variables, ${\displaystyle p_{\text{s}},T_{\text{s}},v_{f},v_{g}}$, so another equation is required in order to specify the values of 3 of these variables uniquely in terms of a fourth. Such an equation is provided here by the equality of the Gibbs free energy in the saturated liquid and vapor states, ${\displaystyle g_{f}=g_{g}}$. This condition of material equilibrium can be obtained from a simple physical argument as follows: the energy required to vaporize a mole is from the second law at constant temperature ${\displaystyle q_{\rm {vap}}=T(s_{g}-s_{f})}$, and from the first law at constant pressure ${\displaystyle q_{\rm {vap}}=h_{g}-h_{f}}$. Equating these two, rearranging, and recalling that ${\displaystyle g=h-Ts}$ produces the result.

The Gibbs free energy is one of the 4 thermodynamic potentials whose partial derivatives produce all other thermodynamics state properties; its differential is ${\displaystyle dg=\partial _{p}g\,dp+\partial _{T}g\,dT=v\,dp-s\,dT}$. Integrating this over an isotherm from ${\displaystyle p_{s},v_{f}}$ to ${\displaystyle p_{s},v_{g}}$, noting that the pressure is the same at each endpoint, and setting the result to zero yields

$${\displaystyle g_{g}-g_{f}=\sum _{j=1}^{3}\int _{p_{j}}^{p_{j}+1}\,v_{j}\,dp=-\int _{v_{f}}^{v_{g}}\,p\,dv+p_{s}(v_{g}-v_{f})=0}$$

Here because ${\displaystyle v}$ is a multivalued function, the ${\displaystyle p}$ integral must be divided into 3 parts corresponding to the 3 real roots of the vdW equation in the form, ${\displaystyle v(p,T)}$ (this can be visualized most easily by imagining Fig. 1 rotated ${\displaystyle 90^{\circ }}$); the result is a special case of material equilibrium. The last equality, which follows from integrating ${\displaystyle v\,dp=d(pv)-p\,dv}$, is the Maxwell equal area rule which requires that the upper area between the vdW curve and the horizontal through ${\displaystyle p_{\text{s}}}$ be equal to the lower one. This form means that the thermodynamic restriction that fixes ${\displaystyle p_{s}}$ is specified by the equation of state itself, ${\displaystyle p=p(v,T)}$. Using the equation for the Gibbs free energy obtained previously for the vdW equation applied to the saturated vapor state and subtracting the result applied to the saturated liquid state produces,

$${\displaystyle RT_{s}\left-2a\left({\frac {1}{v_{g}}}-{\frac {1}{v_{f}}}\right)=0}$$

This is a third equation that along with the two vdW equations above can be solved numerically. This has been done given a value for either ${\displaystyle T_{s}}$ or ${\displaystyle p_{s}}$, and tabular results presented; however, the equations also admit an analytic parametric solution obtained most simply and elegantly, by Lekner. Details of this solution may be found in the Maxwell Construction; the results are

$${\displaystyle T_{rs}(y)=\left({\frac {27}{8}}\right){\frac {2f(y)}{g(y)^{2}}}\quad p_{rs}=27{\frac {f(y)^{2}}{g(y)^{2}}}}$$

$${\displaystyle v_{rf}=\left({\frac {1}{3}}\right){\frac {1+f(y)e^{y}}{f(y)e^{y}}}\qquad \qquad \qquad \quad v_{rg}=\left({\frac {1}{3}}\right){\frac {1+f(y)e^{-y}}{f(y)e^{-y}}}}$$ where

$${\displaystyle f(y)={\frac {y\cosh y-\sinh y}{\sinh y\cosh y-y}}\qquad \qquad g(y)=1+2f(y)\cosh y+f(y)^{2}}$$

and the parameter ${\displaystyle 0\leq y<\infty }$ is given physically by ${\displaystyle y=(s_{g}-s_{f})/(2R)}$. The values of all other property discontinuities across the saturation curve also follow from this solution. These functions define the coexistence curve which is the locus of the saturated liquid and saturated vapor states of the vdW fluid. The curve is plotted in Fig. 1 and Fig. 2, two projections of the state surface. These curves and the numerical results referenced earlier agree exactly, as they must.

Referring back to Fig. 1 the isotherms for ${\displaystyle T_{r}<1}$ are discontinuous. Considering ${\displaystyle T_{r}=7/8}$ as an example, it consists of the two separate green segments. The solid segment above the green circle on the left, and below the one on the right correspond to stable states, the dots represent the saturated liquid and vapor states that comprise the phase change, and the two green dotted segments below and above the dots are metastable states, superheated liquid and subcooled vapor, that are created in the process of phase transition, have a short lifetime, then devolve into their lower energy stable alternative.

In his treatise of 1898 in which he described the van der Waals equation in great detail Boltzmann discussed these states in a section titled "Undercooling, Delayed evaporation"; they are now denoted subcooled vapor, and superheated liquid. Moreover, it has now become clear that these metastable states occur regularly in the phase transition process. In particular processes that involve very high heat fluxes create large numbers of these states, and transition to their stable alternative with a corresponding release of energy can be dangerous. Consequently there is a pressing need to study their thermal properties.

In the same section Boltzmann also addressed and explained the negative pressures which some liquid metastable states exhibit (for example ${\displaystyle T_{r}=0.8}$ of Fig. 1). He concluded that such liquid states of tensile stresses were real, as did Tien and Lienhard many years later who wrote "The van der Waals equation predicts that at low temperatures liquids sustain enormous tension...In recent years measurements have been made that reveal this to be entirely correct."

Even though the phase change produces a mathematical discontinuity in the homogeneous fluid properties, for example ${\displaystyle v}$, there is no physical discontinuity. As the liquid begins to vaporize the fluid becomes a heterogeneous mixture of liquid and vapor whose molar volume varies continuously from ${\displaystyle v_{f}}$ to ${\displaystyle v_{g}}$ according to the equation of state

$${\displaystyle v=v_{f}+x(v_{g}-v_{f})\qquad x=N_{g}/(N_{f}+N_{g})}$$

where ${\displaystyle 0\leq x\leq 1}$ is the mole fraction of the vapor. This equation is called the lever rule and applies to other properties as well. The states it represents form a horizontal line connecting the same colored dots on an isotherm, but not shown in Fig. 1 as noted already since it is a distinct equation of state for the heterogeneous combination of liquid and vapor components.

Extended corresponding states

The idea of corresponding states originated when van der Waals cast his equation in the dimensionless form, ${\displaystyle p_{r}=p(v_{r},T_{r})}$. However, as Boltzmann noted, such a simple representation could not correctly describe all substances. Indeed, the saturation analysis of this form produces ${\displaystyle p_{rs}=p_{s}(T_{r})}$, namely all substances have the same dimensionless coexistence curve. In order to avoid this paradox an extended principle of corresponding states has been suggested in which ${\displaystyle p_{r}=p(v_{r},T_{r},\phi )}$ where ${\displaystyle \phi }$ is a substance dependent dimensionless parameter related to the only physical feature associated with an individual substance, its critical point.

The most obvious candidate for ${\displaystyle \phi }$ is the critical compressibility factor ${\displaystyle Z_{\text{c}}=p_{\text{c}}v_{\text{c}}/(RT_{\text{c}})}$, but because ${\displaystyle v_{c}}$ is difficult to measure accurately, the acentric factor developed by Kenneth Pitzer, ${\displaystyle \omega =-{\mbox{log}}_{10}-1}$, is more useful. The saturation pressure in this situation is represented by a one parameter family of curves, ${\displaystyle p_{rs}=p_{s}(T_{r},\omega )}$. Several investigators have produced correlations of saturation data for a number of substances, the best is that of Dong and Lienhard,

${\displaystyle \ln p_{rs}=5.37270(1-1/T_{r})+\omega (7.49408-}$ $${\displaystyle \qquad 11.181777T_{r}^{3}+3.68769T_{r}^{6}+17.92998\,\ln T_{r})}$$

which has an rms error of ${\displaystyle \pm 0.42}$ over the range ${\displaystyle 1\leq T_{r}\leq 0.3}$

Figure 3 is a plot of ${\displaystyle p_{rs}}$ vs ${\displaystyle T_{r}}$. for various values of ${\displaystyle \omega }$ as given by this equation. The ordinate is logarithmic in order to show the behavior at pressures far below the critical where differences among the various substances (indicated by varying values of ${\displaystyle \omega }$) are more pronounced.

Figure 4 is another plot of the same equation showing ${\displaystyle T_{r}}$ as a function of ${\displaystyle \omega }$ for various values of ${\displaystyle p_{rs}}$. It includes data from 51 substances, including the vdW fluid, over the range ${\displaystyle -0.4<\omega <0.9}$. This plot shows clearly that the vdW fluid (${\displaystyle \omega =-0.302}$) is a member of the class of real fluids; indeed it quantitatively describes the behavior of the liquid metals cesium (${\displaystyle \omega =-0.267}$) and mercury (${\displaystyle \omega =-0.21}$) whose values of ${\displaystyle \omega }$ are close to the vdW value. However, it describes the behavior of other fluids only qualitatively, because specific numerical values are modified by differing values of their Pitzer factor, ${\displaystyle \omega }$.

Joule–Thomson coefficient

The Joule–Thomson coefficient, ${\displaystyle \mu _{J}=\partial _{p}T|_{h}}$, is of practical importance because the two end states of a throttling process (${\displaystyle h_{2}=h_{1}}$) lie on a constant enthalpy curve. Although ideal gases, for which ${\displaystyle h=h(T)}$, do not change temperature in such a process, real gases do, and it is important in applications to know whether they heat up or cool down.

This coefficient can be found in terms of the previously described derivatives as,

$${\displaystyle \mu _{J}={\frac {v(\alpha T-1)}{c_{p}}}}$$

so when ${\displaystyle \mu _{J}}$ is positive the gas temperature decreases when it passes through a throttle, and if it is negative the temperature increases. Therefore the condition ${\displaystyle \mu _{J}=0}$ defines a curve that separates the region of the ${\displaystyle T,p}$ plane where ${\displaystyle \mu _{J}>0}$ from the region where it is less than zero. This curve is called the inversion curve, and its equation is ${\displaystyle \alpha T-1=0}$. Using the expression for ${\displaystyle \alpha }$ derived previously for the van der Waals equation this is

$${\displaystyle {\frac {2a(v-b)^{2}-RTv^{2}b}{RTv^{3}-2a(v-b)^{2}}}=0\quad {\mbox{or}}\quad 2a(v-b)^{2}-RTv^{2}b=0}$$

Note that for ${\displaystyle v\gg b}$ there will be cooling for ${\displaystyle 2a>RTb}$ or in terms of the critical temperature ${\displaystyle T<27T_{\text{c}}/4}$. As Sommerfeld noted, "This is the case with air and with most other gases. Air can be cooled at will by repeated expansion and can finally be liquified."

In terms of ${\displaystyle b/v}$ the equation has a simple positive solution ${\displaystyle b/v=1-{\sqrt {RTb/(2a)}}}$ which, for ${\displaystyle b/v=0}$ produces, ${\displaystyle T=2a/(Rb)=27T_{c}/4}$. Using this to eliminate ${\displaystyle v}$ from the vdW equation then gives the inversion curve as

$${\displaystyle {\frac {p}{p^{*}}}=-1+4\,\left({\frac {T}{2T^{*}}}\right)^{1/2}-3\,\left({\frac {T}{2T^{*}}}\right),}$$

where, for simplicity, ${\displaystyle a,b,R}$ have been replaced by ${\displaystyle p^{*},T^{*}}$.

The maximum of this, quadratic, curve occurs, with ${\displaystyle z^{2}=T/(2T^{*})}$, for

$${\displaystyle D_{z}p/p^{*}=4-6(T/2T^{*})^{1/2}=0}$$

which gives ${\displaystyle (T/2T^{*})_{\rm {max}}^{1/2}=2/3}$, or ${\displaystyle T_{\rm {max}}=8T^{*}/9}$, and the corresponding ${\displaystyle p_{\rm {max}}=p^{*}/3}$. The zeros of the curve ${\displaystyle 3z^{2}-4z+1=0}$, are, making use of the quadratic formula, ${\displaystyle z=(4\pm {\sqrt {16-12}}\,)/6}$, or ${\displaystyle z=1/3}$ and ${\displaystyle 1}$ (${\displaystyle T/T^{*}=2/9=0.{\overline {2}}}$ and ${\displaystyle 2}$). In terms of the dimensionless variables, ${\displaystyle T_{r},p_{r}}$ the zeros are at ${\displaystyle T_{r}=3/4}$ and ${\displaystyle 27/4}$, while the maximum is ${\displaystyle p_{r{\rm {max}}}=9}$, and occurs at ${\displaystyle T_{r{\rm {max}}}=3}$. A plot of the curve is shown in green in Fig. 5. Sommerfeld also displays this plot, together with a curve drawn using experimental data from H₂. The two curves agree qualitatively, but not quantitatively. For example the maximum on these two curves differ by about 40% in both magnitude and location.

Figure 5 shows an overlap between the saturation curve and the inversion curve plotted there. This region is shown enlarged in the right hand graph of the figure. Thus a van der Waals gas can be liquified by passing it through a throttle under the proper conditions; real gases are liquified in this way.

Compressibility factor

Real gases are characterized by their difference from ideal by writing ${\displaystyle pv=ZRT}$. Here ${\displaystyle Z}$, called the compressibility factor, is expressed either as ${\displaystyle Z(p,T)}$ or ${\displaystyle Z(\rho ,T)}$. In either case

$${\displaystyle \lim _{p\rightarrow 0}Z=1,\quad \lim _{\rho \rightarrow 0}Z=1;}$$

${\displaystyle Z}$ takes the ideal gas value. In the second case ${\displaystyle Z(\rho ,T)=p(\rho ,T)/\rho RT}$, so for a van der Waals fluid the compressibility factor is simply ${\displaystyle Z=1/(1-b\rho )-a\rho /(RT)}$, or in terms of reduced variables

$${\displaystyle Z={\frac {3}{3-\rho _{r}}}-{\frac {9\rho _{r}}{8T_{r}}}}$$

where ${\displaystyle 0\leq \rho _{r}=1/v_{r}\leq 3}$. At the critical point, ${\displaystyle T_{r}=\rho _{r}=1}$, ${\displaystyle Z=Z_{\text{c}}=3/2-9/8=3/8}$.

In the limit ${\displaystyle \rho \rightarrow 0}$, ${\displaystyle Z=1}$; the fluid behaves like an ideal gas, a point noted several times earlier. The derivative ${\displaystyle \partial _{\rho }Z|_{T}=b}$ is never negative when ${\displaystyle a/bRT=T^{*}/T\leq 1}$, namely when ${\displaystyle T/T^{*}\geq 1}$ (${\displaystyle T_{r}\geq 27/8}$). Alternatively when ${\displaystyle T/T^{*}<1}$ the initial slope is negative, it becomes zero at ${\displaystyle b\rho =1-(T/T^{*})^{1/2}}$, and is positive for larger ${\displaystyle b\rho \leq 1}$ (see Fig. 6). In this case the value of ${\displaystyle Z}$ passes through ${\displaystyle 1}$ when ${\displaystyle b\rho _{B}=1-T_{B}/T^{*}}$. Here ${\displaystyle T_{B}=(27T_{\text{c}}/8)(1-b\rho _{B})}$ is called the Boyle temperature. It varies between ${\displaystyle 27T_{\text{c}}/8\geq T_{B}\geq 0}$, and denotes a point in ${\displaystyle T,\rho }$ space where the equation of state reduces to the ideal gas law. However the fluid does not behave like an ideal gas there, because neither its derivatives ${\displaystyle (\alpha ,\kappa _{T}){\mbox{ nor }}c_{p}}$ reduce to their ideal gas values, other than where ${\displaystyle b\rho _{B}\ll 1,\,T_{B}\sim 27T_{\text{c}}/8}$ the actual ideal gas region.

Figure 6 shows a plot of various isotherms of ${\displaystyle Z(\rho ,T_{r})}$ vs ${\displaystyle \rho _{r}}$. Also shown are the spinodal and coexistence curves described previously. The subcritical isotherm consists of stable, metastable, and unstable segments, and are identified the same as they were in Fig. 1. Also included are the zero initial slope isotherm and the one corresponding to infinite temperature.

By plotting ${\displaystyle Z(\rho _{r},T_{r})}$ vs ${\displaystyle p_{r}(\rho _{r},T_{r})}$ using ${\displaystyle \rho _{r}}$ as a parameter, one obtains the generalized compressibility chart for a vdW gas, which is shown in Fig. 7. Like all other vdW properties, this is not quantitatively correct for most gases but it has the correct qualitative features as can be seen by comparison with this figure which was produced from data using real gases. The two graphs are similar, including the caustic generated by the crossing isotherms; they are qualitatively very much alike.

Virial expansion

Statistical mechanics suggests that ${\displaystyle Z}$ can be expressed by a power series called a virial expansion,

$${\displaystyle Z(\rho ,T)=1+\sum _{k=2}^{\infty }\,B_{k}(T)(\rho )^{k-1}}$$

The functions ${\displaystyle B_{k}(T)}$ are the virial coefficients; the ${\displaystyle k}$th term represents a ${\displaystyle k}$ particle interaction.

Expanding the term ${\displaystyle (1-b\rho )^{-1}}$ in the compressibility factor of the vdW equation in its infinite series, convergent for ${\displaystyle b\rho <1}$, produces

$${\displaystyle Z(\rho ,T)=1+b\rho +\sum _{k=3}^{\infty }\,B_{k}(T)(b\rho )^{k-1}.}$$

The corresponding expression for ${\displaystyle Z(\rho _{r},T_{r})}$ when ${\displaystyle \rho _{r}<3}$ is

$${\displaystyle Z(\rho _{r},T_{r})=1+^{-1}](\rho _{r}/3)+\sum _{k=2}^{\infty }\,(\rho _{r}/3)^{k-1}.}$$

These are the virial expansions, one dimensional and one dimensionless, for the van der Waals fluid. The second virial coefficient is the slope of ${\displaystyle Z(\rho _{r},T_{r})}$ at ${\displaystyle \rho _{r}=0}$. Notice that it can be positive or negative depending on whether or not ${\displaystyle T_{r}>{\mbox{ or}}<27/8}$, which agrees with the result found previously by differentiation.

For molecules that are non attracting hard spheres, ${\displaystyle a=0}$, the vdW virial expansion becomes simply

$${\displaystyle Z(\rho )=(1-b\rho )^{-1}=1+\sum _{k=2}^{\infty }(b\rho )^{k-1},}$$

which illustrates the effect of the excluded volume alone. It was recognized early on that this was in error beginning with the term ${\displaystyle (b\rho )^{2}}$. Boltzmann calculated its correct value as ${\displaystyle (5/8)(b\rho )^{2}}$, and used the result to propose an enhanced version of the vdW equation

$${\displaystyle (p+a/v^{2})(v-b/3)=RT.}$$

On expanding ${\displaystyle (v-b/3)^{-1}}$, this produced the correct coefficients thru ${\displaystyle (b/v)^{2}}$ and also gave infinite pressure at ${\displaystyle b/3}$, which is approximately the close packing distance for hard spheres. This was one of the first of many equations of state proposed over the years that attempted to make quantitative improvements to the remarkably accurate explanations of real gas behavior produced by the vdW equation.

Mixtures

In 1890 van der Waals published an article that initiated the study of fluid mixtures. It was subsequently included as Part III of a later published version of his thesis. His essential idea was that in a binary mixture of vdw fluids described by the equations

$${\displaystyle p_{1}={\frac {RT}{v-b_{11}}}-{\frac {a_{11}}{v^{2}}}\quad {\mbox{and}}\quad p_{2}={\frac {RT}{v-b_{22}}}-{\frac {a_{22}}{v^{2}}}}$$

the mixture is also a vdW fluid given by

$${\displaystyle p={\frac {RT}{v-b_{x}}}-{\frac {a_{x}}{v^{2}}}}$$ where $${\displaystyle a_{x}=a_{11}x_{1}^{2}+2a_{12}x_{1}x_{2}+a_{22}x_{2}^{2}\quad {\mbox{and}}\quad b_{x}=b_{11}x_{1}^{2}+2b_{12}x_{1}x_{2}+b_{22}x_{2}^{2}}$$

Here ${\displaystyle x_{1}=N_{1}/N}$, and ${\displaystyle x_{2}=N_{2}/N}$, with ${\displaystyle N=N_{1}+N_{2}}$ (so that ${\displaystyle x_{1}+x_{2}=1}$) are the mole fractions of the two fluid substances. Adding the equations for the two fluids shows that ${\displaystyle p\neq p_{1}+p_{2}}$, although for ${\displaystyle v}$ sufficiently large ${\displaystyle p\approx p_{1}+p_{2}}$ with equality holding in the ideal gas limit. The quadratic forms for ${\displaystyle a_{x}}$ and ${\displaystyle b_{x}}$ are a consequence of the forces between molecules. This was first shown by Lorentz, and was credited to him by van der Waals. The quantities ${\displaystyle a_{11},\,a_{22}}$ and ${\displaystyle b_{11},\,b_{22}}$ in these expressions characterize collisions between two molecules of the same fluid component while ${\displaystyle a_{12}=a_{21}}$ and ${\displaystyle b_{12}=b_{21}}$ represent collisions between one molecule of each of the two different component fluids. This idea of van der Waals was later called a one fluid model of mixture behavior.