Enthalpy 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 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 below T_r<<1.0. The heat of vaporization diminishes with increasing temperature and it vanishes completely at the critical temperature (T_r=1) because above the critical temperature the liquid and vapor phases no longer co-exist.

1 Units
2 Physical model for vaporization
3 Enthalpy of condensation
4 Thermodynamic background
5 Vaporization enthalpy of electrolyte solutions
6 Selected values
- 6.1 Elements
- 6.2 Other common substances
7 See also
8 References

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

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 (c_p = 75.3 J K⁻¹mol⁻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 (T_b), Δ_vG = 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

Enthalpies of vaporization of the elements in kJ/mol, measured at their respective normal boiling points

Group →	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
↓ Period
1	H 0.44936																	He 0.0845
2	Li 145.92	Be 292.40											B 489.7	C 355.8	N 2.7928	O 3.4099	F 3.2698	Ne 1.7326
3	Na 96.96	Mg 127.4											Al 293.4	Si 300	P 12.129	S 1.7175	Cl 10.2	Ar 6.447
4	K 79.87	Ca 153.6	Sc 314.2	Ti 421	V 452	Cr 344.3	Mn 226	Fe 349.6	Co 376.5	Ni 370.4	Cu 300.3	Zn 115.3	Ga 258.7	Ge 330.9	As 34.76	Se 26.3	Br 15.438	Kr 9.029
5	Rb 72.216	Sr 144	Y 363	Zr 581.6	Nb 696.6	Mo 598	Tc 660	Ru 595	Rh 493	Pd 357	Ag 250.58	Cd 100	In 231.5	Sn 295.8	Sb 77.14	Te 52.55	I 20.752	Xe 12.636
6	Cs 67.74	Ba 142	*	Hf 575	Ta 743	W 824	Re 715	Os 627.6	Ir 604	Pt 510	Au 334.4	Hg 59.229	Tl 164.1	Pb 177.7	Bi 104.8	Po 60.1	At 114	Rn 16.4
7	Fr n/a	Ra 37	**	Rf n/a	Db n/a	Sg n/a	Bh n/a	Hs n/a	Mt n/a	Ds n/a	Rg n/a	Cn n/a	Uut n/a	Uuq n/a	Uup n/a	Uuh n/a	Uus n/a	Uuo n/a

* Lanthanides				La 414	Ce 414	Pr n/a	Nd n/a	Pm n/a	Sm n/a	Eu n/a	Gd n/a	Tb n/a	Dy n/a	Ho n/a	Er n/a	Tm n/a	Yb n/a	Lu n/a
Actinides**				Ac n/a	Th 514.4	Pa n/a	U n/a	Np n/a	Pu n/a	Am n/a	Cm n/a	Bk n/a	Cf n/a	Es n/a	Fm n/a	Md n/a	No n/a	Lr n/a


0–10 kJ/mol	10–100 kJ/mol	100–300 kJ/mol	>300 kJ/mol

Other common substances

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

Compound	Heat of vaporization (kJ mol^-1)	Heat of vaporization (kJ kg⁻¹)
Ammonia	23.35	1371
Butane	21.0	320
Ethanol	38.6	841
Hydrogen	0.46	451.9
Methane	8.19	760
Methanol	35.3	1104
Propane	15.7	356
Phosphine	14.6	429.4
Water	40.65	2257

References

^ Garai, J. (2009). "Physical model for vaporization". Fluid Phase Equilibria 283: 89–77. doi:10.1016/j.fluid.2009.06.005. edit
^ X. Ge, X. Wang. Estimation of Freezing Point Depression, Boiling Point Elevation and Vaporization enthalpies of electrolyte solutions. Ind. Eng. Chem. Res. 48(2009)2229-2235. http://pubs.acs.org/doi/abs/10.1021/ie801348c (Correction: 2009, 48, 5123)http://pubs.acs.org/doi/abs/10.1021/ie900434h
^ X. Ge, X. Wang. Calculations of Freezing Point Depression, Boiling Point Elevation, Vapor Pressure and Enthalpies of Vaporization of Electrolyte Solutions by a Modified Three-Characteristic Parameter Correlation Model. J. Sol. Chem. 38(2009)1097-1117.http://www.springerlink.com/content/21670685448p5145/