The Virial Theorem - Basic Energetics: Intermolecular Interactions - Liquid-State Physical Chemistry: Fundamentals, Modeling, and Applications (2013)

Liquid-State Physical Chemistry: Fundamentals, Modeling, and Applications (2013)

3. Basic Energetics: Intermolecular Interactions

3.8. The Virial Theorem*

Even without considering the details of the dynamics of a system of particles, we can connect the average kinetic energy with the average potential energy. This is achieved via a useful theorem, first introduced by Clausius [18], of which we present here a short derivation. To this purpose, it must be first noted that for a bounded function G(t) with time t as variable, the time average ⟨g⟩ of g = dG/dt is zero, since

(3.33) c3-math-0033

Attention is now turned to the kinetic energy c3-math-5040 for a set of particles with mass mi, position coordinates ri, velocity vi, and momentum pi = mivi. Because Ti is a homogeneous function of the second degree in vi, we have by Euler's theorem (see Appendix B)

(3.34) c3-math-0034

Since ∂Ti/∂vi = pi, we further obtain

(3.35) c3-math-0035

If the system has a finite volume, pi·ri is bounded and thus we have ⟨dpi·ri/dt⟩ = 0. The kinetic energy T is linked to the potential energy Φ via Newton's Second Law reading fi = miai with force fi and acceleration ai. Using fi = −∂Φ/∂ri, miai = dpi/dt, and introducing the virial (of force) c3-math-5045, we obtain the virial theorem

(3.36) c3-math-0036

where the second step can be made as long as c3-math-5030. Frequently, Φ is a homogeneous function of degree n in r, and from Euler's theorem we then have

(3.37) c3-math-0037

Since the total energy U = ⟨T⟩ + ⟨Φ⟩, we have ⟨Φ⟩ = 2U/(n + 2) and ⟨T⟩ = nU/(n + 2). For harmonic potentials n = 2 resulting in ⟨T⟩ = ⟨Φ= ½U. For the Coulomb potential n = −1 and since T ≥ 0, this implies U ≤ 0. So, for an electrostatic system to be stable the energy must be negative. Finally, we note that the virial can be used in the derivation of the equation of state (see Chapter 6).


1) Note again that r is a vector and r is a (column matrix of a) set of elements ri or ri.

2) The unit of charge is coulomb [C].

3) The unit of dipole moment is coulomb × meter [C m], but frequently (for historical reasons) the Debye unit [D] is used. 1[D] = 3.336 × 10−30 [C m].

4) Because the molecule contains several charges we have to sum over them. We indicate here the charges by the subscript i and the molecule by the superscript (i). Only if we sum explicitly over charges is the molecule is indicated by the superscript; otherwise, we use a subscript.

5) Isotropy implies equal response for all directions. When the polarization becomes anisotropic, these interactions become orientation-dependent and therefore temperature-dependent.

6) The unit of polarizability is C2 m2 J−1. Incorporating (4πε0)−1 in α via α′ = α/(4πε0) renders the dimension of α′ to be m3. Since, according to electrostatics, the polarizability of a perfectly conducting sphere of radius r in vacuum is given by 4πε0r3, the molecular radius r can be estimated from r3 = α′. Hence, α′ is known as the polarizability volume.

7) For an overview of London's work on interactions, see Ref. [19]. The term “dispersion forces” was coined by London to indicate the analogy of the expressions derived to those which appear in the dispersion formula for the polarizability of a molecule when acted upon by an alternating field.

8) We note that, since electromagnetic radiation has a finite velocity c, a second molecule feels the potential of the first molecule at a distance r only after a time r/c. This retardation effect becomes important only for distances larger than about 10 nm. The effect is important for the mutual interaction between macroscopic bodies and between a single particle and a solid surface.

9) See Section 16.2 and Section 15.2, respectively.


1 Böttcher, C.J.F. (1973) Theory of Dielectric Polarization, vol. I, 2nd edn, Elsevier, Amsterdam. See also Hirschfelder et al. (1964).

2 Pitzer, K.S. (1959) Adv. Chem. Phys., 1, 59.

3 Kihara, T. (1953) Rev. Mod. Phys., 25, 831.

4 (a) Hammam, S.D. and Lambert, J.A. (1954) Aust. J. Chem., 7, 1; (b) Hildebrand, J. and Scott, R.L. (1962) Regular Solutions, Prentice-Hall, Englewood Cliffs.

5 Morse, P.M. (1929) Phys. Rev., 20, 57.

6 (a) Kohler, F. (1954) Monatsh. Chem., 88, 857; (b) Kihara, T. (1976) Intermolecular Forces, John Wiley & Sons, Ltd, Chichester.

7 Rose, J.H., Smith, J.R., Guinea, F., and Ferrante, J. (1984) Phys. Rev., B29, 2963.

8 See Pimentel and McLellan (1960).

9 See Grabowski (2006).

10 Scott, R.L. (1971) Physical Chemistry, vol. VIIIA, Academic Press, New York, Ch. 1, p. 11.

11 (a) Axilrod, B.M. and Teller, E. (1943) J. Chem. Phys., 11, 299; (b) Axilrod, B.M. (1949) J. Chem. Phys. 17, 299 and 19, 71; (c) Midzuno, Y. and Kihara, T. (1956) J. Phys. Soc. Jpn, 11, 1045.

12 (a) Rowlinson, J.S. (1965) Discuss. Faraday Soc., 40, 19; (b) Rowlinson, J.S. and Swinton, F.L. (1982) Liquids and Liquid Mixtures, 3rd edn, Butterworth, London.

13 Dymond, J.H. and Alder, B.J. (1968) Chem. Phys. Lett., 2, 54.

14 Berendsen, H.J.C. (2007) Simulating the Physical World, Cambridge University Press, p. 187.

15 Sherwood, A.E. and Prausnitz, J.M. (1964) J. Chem. Phys., 41, 429.

16 Parson, J.M., Siska, P.E., and Lee, Y.T. (1972) J. Chem. Phys., 56, 1511.

17 Barker, J.A., Fisher, R.A., and Watts, R.O. (1971) Mol. Phys., 21, 657.

18 Clausius, R.J.E. (1870) Philos. Mag. Ser. 4, 40, 122.

19 London, F. (1937) Trans. Faraday Soc., 33, 8.

20 Atkins, P.W. (2002) Physical Chemistry, 7th edn, Oxford

21 Butt, H.-J., Graf, K., and Kappl, M. (2006) Physics and Chemistry of Interfaces, 2nd edn, Wiley-VCH.

22 Israelachvili, J. (1991) Intermolecular and Surface Forces, 2nd edn, Academic Press.

Further Reading

Grabowski, S.J. (2006) Hydrogen Bonding – New Insights, Springer, Dordrecht.

Hirschfelder, J.O., Curtiss, C.F., and Bird, R.B. (1954) Molecular Theory of Gases and Liquids, John Wiley & Sons, Inc., New York (corrected edition, 1964).

Kaplan, I.G. (2006) Intermolecular Interactions – Physical Picture, Computational Methods and Model Potentials, John Wiley & Sons, Ltd, Chichester.

Maitland, G.C., Rigby, M., Smith, E.B., and Wakeham, W.A. (1981) Intermolecular Forces – Their Origin and Determination, Clarendon, Oxford.

Margenau, H. and Kestner, N.R. (1971) Theory of Intermolecular Forces, Pergamon, Oxford.

Parsegian, V.A. (2006) Van der Waals Forces: A Handbook for Biologists, Chemists, Engineers and Physicists, Cambridge University Press, Cambridge.

Pimentel, G.C. and McLellan, A.L. (1960) The Hydrogen Bond, W.H. Freeman and Company, New York.

Stone, A.J. (1996) The Theory of Intermolecular Forces, Clarendon, Oxford.