Heat of Vaporization


The enthalpy of vaporization, (symbol \Delta{}H_{\mathrm{vap}}), also known as the heat of vaporization or heat of evaporation, is the energy required to transform a given quantity of a substance from a liquid into a gas at a given pressure (often atmospheric pressure).

It is often measured at the normal boiling point of a substance; although tabulated values are usually corrected to 298 K, the correction is often smaller than the uncertainty in the measured value.

The heat of vaporization is temperature-dependent, though a constant heat of vaporization can be assumed for small temperature ranges and for Tr<<1.0. The heat of vaporization diminishes with increasing temperature and it vanishes completely at the critical temperature (Tr=1) because above the critical temperature the liquid and vapor phases no longer exist, since the substance is a supercritical fluid.

Units

Values are usually quoted in J/mol or kJ/mol (molar enthalpy of vaporization), although kJ/kg or J/g (specific heat of vaporization), and older units like kcal/mol, cal/g and Btu/lb are sometimes still used, among others.

Physical model for vaporization


A simple physical model for the liquid-gas phase transformation has been proposed recently.[1] It is suggested that the energy required to free an atom from the liquid is equivalent to the energy needed to overcome the surface resistance of the liquid. The model allows calculating the latent heat by multiplying the maximum surface area covering an atom (Fig. 1) with the surface tension and the number of atoms in the liquid. The calculated latent heat of vaporization values for the investigated 45 elements agrees well with experiments.

Enthalpy of condensation

The enthalpy of condensation (or heat of condensation) is by definition equal to the enthalpy of vaporization with the opposite sign: enthalpy changes of vaporization are always positive (heat is absorbed by the substance), whereas enthalpy changes of condensation are always negative (heat is released by the substance).

Thermodynamic background

Molar enthalpy of zinc above 298.15 K and at 1 atm pressure, showing discontinuities at the melting and boiling points. The enthalpy of melting (ΔH°m) of zinc is 7323 J/mol, and the enthalpy of vaporization (ΔH°v) is 115 330 J/mol.

The enthalpy of vaporization can be viewed as the energy required to overcome the intermolecular interactions in the liquid (or solid, in the case of sublimation). Hence helium has a particularly low enthalpy of vaporization, 0.0845 kJ/mol, as the van der Waals forces between helium atoms are particularly weak. On the other hand, the molecules in liquid water are held together by relatively strong hydrogen bonds, and its enthalpy of vaporization, 40.65 kJ/mol, is more than five times the energy required to heat the same quantity of water from 0 °C to 100 °C (cp = 75.3 J K−1 mol−1). Care must be taken, however, when using enthalpies of vaporization to measure the strength of intermolecular forces, as these forces may persist to an extent in the gas phase (as is the case with hydrogen fluoride), and so the calculated value of the bond strength will be too low. This is particularly true of metals, which often form covalently bonded molecules in the gas phase: in these cases, the enthalpy of atomization must be used to obtain a true value of the bond energy.

An alternative description is to view the enthalpy of condensation as the heat which must be released to the surroundings to compensate for the drop in entropy when a gas condenses to a liquid. As the liquid and gas are in equilibrium at the boiling point (Tb), ΔvG = 0, which leads to:

\Delta\,_v S = S_{gas} - S_{liquid} = \Delta\,_v H/T_b

As neither entropy nor enthalpy vary greatly with temperature, it is normal to use the tabulated standard values without any correction for the difference in temperature from 298 K. A correction must be made if the pressure is different from 100 kPa, as the entropy of a gas is proportional to its pressure (or, more precisely, to its fugacity): the entropies of liquids vary little with pressure, as the compressibility of a liquid is small.

These two definitions are equivalent: the boiling point is the temperature at which the increased entropy of the gas phase overcomes the intermolecular forces. As a given quantity of matter always has a higher entropy in the gas phase than in a condensed phase (\Delta\,_v S is always positive), and from

\Delta\,G = \Delta\,H - T\Delta\,S,

the Gibbs free energy change falls with increasing temperature: gases are favored at higher temperatures, as is observed in practice.

Vaporization enthalpy of electrolyte solutions

Estimation of the enthalpy of vaporization of electrolyte solutions can be simply carried out using equations based on the chemical thermodynamic models, such as Pitzer model[2] or TCPC model.[3]

Selected values

Elements

Template:Periodic table (enthalpy of vaporisation)

Other common substances

Enthalpies of vaporization of common substances, measured at their respective standard boiling points:

Compound Boiling Point at normal pressure Heat of vaporization
(kJ mol-1)
Heat of vaporization
(kJ kg−1)
Acetone 329-330 K, 56-57 °C, 133-134 °F 31.3 538.9
Aluminium 2792 K, 2519 °C, 4566 °F 294.0 10500
Ammonia 240 K, −33.34 °C, -28 °F 23.35 1371
Butane 272-274 K, -1°C, 30-34 °F 21.0 320
Diethyl ether 307.8 K, 34.6 °C, 94.3 °F 26.17 353.1
Ethanol 352 K, 78.37 °C, 173 °F 38.6 841
Hydrogen 20.271 K, -252.879 °C, -423.182 °F 0.46 451.9
Iron 3134 K, 2862 °C, 5182 °F 340 6090
Isopropyl alcohol 356 K, 82.6 °C, 181 °F 44.0 732.2
Methane 109-113 K, -164--160 °C, -263--256 °F 8.17 480.6
Methanol 338 K, 64.7 °C, 148 °F 35.3 1104
Propane 230.9-231.11 K,-42--42 °C, -44--44 °F 15.7 356
Phosphine 185 K, -87.7 °C, -126 °F 14.6 429.4
Water 373.15K 40.68 2260

See also

References

  • CODATA Key Values for Thermodynamics
  • Kugler HK & Keller C (eds) 1985, Gmelin handbook of inorganic and organometallic chemistry, 8th ed., 'At, Astatine', system no. 8a, Springer-Verlag, Berlin, ISBN 3-540-93516-9, pp. 116–117
  • NIST Chemistry WebBook
  • Sears, Zemansky et al., University Physics, Addison-Wesley Publishing Company, Sixth ed., 1982, ISBN 0-201-07199-1
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