Vaporisation

Volatile means to be turned into a vapour easily. Liquids turn into vapour at all temperatures, because, according to the Maxwell-Boltzmann energy distribution curve, there are always some particles with enough energy to escape the body of liquid. The more volatile a substance is the easier is the vaporisation process:

liquid vapour

Volatility is a function of the ambient temperature, as liquids have a higher vapour pressure as the temperature increases. Thus, even a relativelly non-volatile liquid produces more vapour at higher temperatures.

The volatility of a covalent liquid depends on the strength of the intermolecular forces, as these must be overcome in order for the particles to change from liquid state to gaseous state.


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Intermolecular forces

With organic compounds we are primariliy dealing with covalent molecules. These are attracted to one another by one or both of two forces:

  • Van der Waals (induced dipole) forces
  • Dipole - dipole interactions

Dipole-dipole interactions are themselves sub-divided into:

  • Permanent dipole - dipole attractions
  • Hydrogen bonding

Of the three types of force, Van der Waals forces are present between all molecules, but are the weakest forces when the molecules are small. Permanent dipole-dipole forces are stronger than Van der Waals attractions but weaker than hydrogen bonding.


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Dispersion forces (induced dipole - dipole attractions)

All covalent molecules have these forces. They are caused by vibrations within the molecule that produce temporary dipoles. These temporary dioles vibrate sympathetically in neighbouring molecules leaving the partial positive part of the dipole next to the partial negative side of the dipole and vice-versa.

dispersion forces act between ALL molecules regardless of their structure. The strength of the force depends on two structural features:

  • the relative molecular mass of the compound
  • the shape of the compound

Relative molecular mass effect

For compounds with low relative molecular mass, the dispersion forces are very weak and can be easily overcome. Hence low RMM non-polar molecules are gases at room temperature. As the relative molecular mass increases, so do the dispersion forces and in molecules with large relative molecular mass, such as polymers, the Van der Waals forces are sufficiently large to keep the compound in the solid state even at temperatures well above room temperature.

Compound

formula

Relative mass

boiling point /ºC

Methane

CH4

16

-161

Ethane

C2H6

28

-89

Eicosane

C20H42

282

343

Shape of the molecule (degree of branching)

When comparing two compounds that have similar relative molecular mass the shape of the molecule becomes important. Branched molecules are more spherical in shape and have a lower surface area. This means that there is less surface for the dispersion force to act between and the overall force between the molecules is smaller. Hence, greater branching reduces the intermolecular force and consequently the boiling point.

Compound

formula

Relative mass

boiling point /ºC

pentane

C5H12

72

36

methylbutane

C5H12

72

28

dimethylpropane

C5H12

72

10


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Permanent dipole - dipole interactions

These arise when a molecule has a permanent dipole caused by the presence of atoms of electronegativity different from carbon and hydrogen. Carbon and hydrogen are neither electronegative nor electropositive, having electronegativity values of 2.5 and 2.1 respectively.

A chlorine atom, for example, has an electronegativity of 3.0. Consequently chlorine attracts electrons when bonded to carbon and develops a partial negative charge, leaving a partial positive charge on the carbon. This is called a permanent dipole. These permanent dipoles attract a molecule to neighbouring molecules.

Permanent dipole-dipole forces are stronger than dispersion forces in small molecules.


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Hydrogen bonding

This is a special case of dipole-dipole interaction. It happens when a molecule contains hydrogen attached to oxygen or nitrogen. These two elements are highly electronegative and draw the electrons away from the hydrogen atom. As hydrogen only has one electron in the first place, the effect exposes the hydrogen nucleus, causing a high partial charge density. This makes the N-H, or O-H bond, very polar and, in consequence, dipole-dipole forces are very strong (about one tenth of a normal covalent bond).

Compounds with hydrogen bonding

  • Alcohols
  • Carboxylic acids
  • Amines
  • Amides

The hydrogen bonding in carboxylic acids is particularly strong, as two hydrogen bonds per molecule can form. It is strong enough to cause ethanoic acid, for example, to dimerise to (CH3COOH)2 even in the vapour state.


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Enthalpy of vaporisation

The enthalpy of vaporisation is the amount of energy needed to turn 1 mole of a liquid into vapour.

H2O(l) H2O(g)            ∆Hvap = 41.1 kJ mol-1

It depends on the intermolecular forces present in the liquid. If the intermolecular forces are strong then the enthalpy of vaporisation is high.

The enthalpy of vaporisation allows up to assess the strength of the intermolecular forces. It can be measured experimentally and the values compared to give us an idea of the forces involved in the liquid state

Table of Vaporisation enthalpies

Compound

  ∆Hvap /kJ mol-1

Intermolecular forces

methane

8.2

dispersion

ethane

14.7

dispersion

propane

19.0

dispersion

butane

22.4

dispersion

ethanal

25.8

dipole-dipole

propanal

28.3

dipole-dipole

propanone

29.1

dipole-dipole

methanol

37.4

H-bonding

ethanol

38.6

H-bonding

ethanoic acid

51.6

H-bonding


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Boiling point

The boiling point of a liquid is the temperature at which its vapour pressure is equal to the atmospheric pressure. It is therefore directly related to the vapour pressure of a liquid.

As the temperature increases so the vapour pressure increases. Volatile liquids have higher vapour pressures and therefore boil at a lower temperature.

Boiling points are quoted for standard pressure 100.0 Nm-2

Table of Vaporisation enthalpies and b.p

Compound

  ∆Hvap /kJ mol-1

boiling point /ºC

methane

8.2

-161

ethane

14.7

-89

propane

19.0

-42

butane

22.4

0

ethanal

25.8

21

propanal

28.3

49

propanone

29.1

56

methanol

37.4

65

ethanol

38.6

79

ethanoic acid

51.6

118

Colourful


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