IB Chemistry home > Syllabus 2016 > Organic chemistry > Volatility

Syllabus ref: 10.1

The word 'volatile' is usually employed in everyday life to mean 'unstable'. However, in chemistry terms this is not the case.

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:

Dipole-dipole interactions are themselves sub-divided into:

Of the three types of force, Van der Waal's 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 Waal's 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:

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 Waal's 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

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

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