The enthalpy of vaporization, also known as the heat of vaporization or heat of evaporation,
is the enthalpy change required to transform a given quantity of a substance from a liquid
into a gas at a given 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 reduced temperature Tr<<1.0. The heat
of vaporization diminishes with increasing temperature and it vanishes completely at
the critical temperature 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,
although kJ/kg or J/g, 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 was proposed in 2009
by Jozsef Garai. 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 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.
Another model which utilizes the data set from Jozsef Garai's model shows that the liquid–gas
phase change can be explained in terms of kinetic theory by considering that the energy
required for vaporization is extracted from all six of the vaporizing molecule's neighbours.
This includes a required rethink of the probability of vaporization, and has consequences to the
Clausius-Clapeyron equation. Moreover, it does resolve the issue of the latent heat
of vaporization being significantly greater than the thermal energy exchanged between
molecules, i.e. at boiling point the latent heat for water is approximately 13.2 times
kT Enthalpy of condensation
The enthalpy of condensation is by definition equal to the enthalpy of vaporization with
the opposite sign: enthalpy changes of vaporization are always positive, whereas enthalpy changes
of condensation are always negative. Thermodynamic background
The enthalpy of vaporization can be written as
It is equal to the increased internal energy of the vapor phase compared with the liquid
phase, plus the work done against ambient pressure. The increase in the internal energy
can be viewed as the energy required to overcome the intermolecular interactions in the liquid.
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. 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, 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, ΔvG = 0,
which leads to:
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: 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, and from
, 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 or TCPC model.
Selected values Elements
Other common substances Enthalpies of vaporization of common substances,
measured at their respective standard boiling points:
See also Enthalpy of fusion
Enthalpy of sublimation Joback method
References
CODATA Key Values for Thermodynamics Kugler HK & Keller C 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