Doc Brown's Advanced A level theoretical chemistry revision notes

Part 8.1 Phase Equilibrium – Vapour pressure – Boiling Point

[Author ©  Dr Phil Brown PhD: Doc Brown's Chemistry exam revision notes suitable for advanced A level chemistry students studying pre–university/college advanced level theoretical–physical chemistry courses:  vapour pressure & boiling point  [page updated May 2nd 2026 *]

• The liquid–gas phase equilibria situation is discussed and the vapour pressures of selected liquids are compared and discussed and the relationship between boiling point and ambient pressure.

• e.g. liquid water water vapour or H₂O_(l) H₂O_(g)
• The co–existence of water in two different phases in a dynamic equilibrium is a clear example of a phase equilibrium.
  • If water (or an aqueous solution) is poured into a dry bottle of 'dry' air and then stoppered, some of the water will then evaporate to create a vapour pressure, so water vapour and liquid water co–exist.
  • However, some of the water vapour may condense if the gas molecules hit the liquid surface but others can escape the intermolecular forces at the surface (which are imbalanced and incidentally cause the net inward pull of surface tension).
  • When the rate of evaporation = rate of condensation, a dynamic physical equilibrium exists between liquid water and water vapour.
• Note, that at a certain temperature and pressure (called the triple point), liquid water, water vapour and ice can all co–exist, a three phase system!
• The saturated vapour pressure is the maximum pressure the vapour from a liquid (or solid/solution) can exert at a given temperature if the substance alone occupied a container in which the liquid/solution/solid is in equilibrium with the vapour.
• The maximum or saturated vapour pressure that liquid water exerts at a particular temperature can be measured and it increases exponentially with increase in temperature, i.e. the equilibrium is shifted more to the right, see graph below.
• The same shaped graph line (examples below for pure liquids) is obtained with any liquid, though the relative position of the curve depends on the volatility of the liquid, which in turn depends on the strength of the inter–molecular attractive forces (see later paragraph).
• Liquids containing dissolved substances i.e. mixtures of miscible liquids or solutions of solids, will show a reduced vapour pressure compared to the pure liquid (not covered here).
• 
• Typical saturated vapour pressure curves (max. p_vap versus temperature).
  • The vapour pressure curves are shown in the graph above for: tetrachloromethane CCl₄,
  • ethanol C₂H₅OH, benzene C₆H₆, water H₂O and ethanoic acid CH₃COOH
  • The more polar the molecule and the capacity for hydrogen bonding will tend to lower the comparative vapour pressure at a given temperature and therefore raise the boiling point of the liquid.
  • Although ethanol is polar and exhibits hydrogen bonding, non-polar tetrachloromethane and benzene have far more electrons, so more polarizable, increasing the intermolecular bonding forces, giving them similar vapour pressure curves.
  • Water is a highly polar molecule exhibiting hydrogen bonding and has a very low comparative vapour pressure for such a small molecule.
  • Ethanoic acid's comparative vapour pressure is the lowest shown because it forms a dimer via hydrogen bonding and is thus effectively a larger more polarizable molecule with great intermolecular bonding forces.
• The curves illustrate the exponential rise of vapour pressure with increase in temperature.
• The increase in vapour pressure of any liquid with rise in temperature is due to the increasing average kinetic energy of the liquid molecules, so more molecules with sufficient KE can 'escape' the intermolecular attractive forces of the liquid molecules, increasing the rate of evaporation until it matches the rate of condensation i.e. a dynamic equilibrium point is reached and the liquid exerts its maximum or saturated vapour pressure.
• When the vapour pressure of a liquid matches the ambient pressure, the liquid boils as bubbles of vapour can now form in the bulk liquid.
  • At 100^oC, water vapour pressure = normal atmospheric pressure, and so water boils at normal atmospheric pressure (1 atm/760 mmHg, see dotted lines on diagram above).
  • This has consequences for brewing a cup of tea as you climb a high mountain. It gets more and more difficult to brew a good pot of tea because the boiling temperature gets lower and lower as the pressure falls with increasing height!
  • Note on some applications of vapour pressure data:
    • A knowledge of vapour pressures from data tables is important is understanding and using the technique known as steam distillation, described in section 8.5, which is a method of extracting thermally unstable compounds from a mixture by a distillation process operating at a far lower temperature than its boiling point.
    • You can also distil thermally unstable compounds at a more lower temperature than its normal boiling point at normal atmospheric pressure by means of reduced pressure distillation ('vacuum distillation'). If you reduce the pressure in the flask and condenser system, you then reduce the vapour pressure needed for the liquid to boil i.e. lower vapour pressure means a lower boiling–distillation temperature. This is actually a much cleaner and efficient process than steam distillation.
• The greater the intermolecular forces between molecules, the lower the saturated vapour pressure at a given temperature = maximum p_vap given in Pa/mm Hg 20^oC.
  • Collectively 'intermolecular forces' are also referred to as 'Van der Waals Forces' and you must be able to clearly distinguish these forces from ionic or covalent bonds, even for the unfortunately misnamed 'hydrogen bond'.
  • All intermolecular forces are essentially electrostatic in origin arising from 'Coulombic' attraction between positive and negative 'centres' in the atoms or molecules. A 'centre' might quite simply be a region more positive or more negative than another due to asymmetry in the distribution of the electron charge (clouds).
  • Details of these intermolecular forces are further discussed later in the context of comparative studies of particular series of molecules.
    • 8.2.1 A summary of Van der Waals forces and an introduction to intermolecular forces
  • The greater the intermolecular attractive bonding forces the lower the vapour pressure, the higher the boiling point of the liquid and to some extent enthalpies of vaporisation, the latter does not give as clearer a pattern as you might expect.
  • You also see, not surprisingly, higher enthalpies of vaporisation (ΔH_vap/kJ mol^–1) BUT not a good trend if you study the summary tables.
  • However you get a more clearer trend with increasingly higher boiling point (bpt at 101325Pa / 1atm / 760mmHg) as more kinetic energy is needed by the molecules to overcome the increase in intermolecular attractive forces.
  • These intermolecular forces are strongly influenced by the structure of the molecule and the extent of bond polarity can have a major effect on the boiling point of an organic liquid, but most of the attractive force for most molecules originates from transient (temporary) dipole – induced dipole interactions (explained further in the ALKANE section).
• Section 8.2.6 compares eight organic gases/liquids of similar molecular mass (58 to 60 and 64.5 for the halogenoalkane) and total number of electrons (32 to 34) whose physical properties illustrate these trends.
• All eight organic molecules, apart from chloroethane, are based on combinations of 4 C/N/O atoms plus H atoms.
• They were the best match and range I could come up with to do a comparative paper based study!
• The 8 molecules are discussed in order of increasing boiling point (bpt).
• See also
  • Section 8.3 Boiling Point Plots for six Organic Homologous Series
    • Related organic link: Guide to the structure of organic homologous series of molecules
  • Section 8.4 Other molecular case studies of boiling points related to intermolecular forces
  • Section 8.2.4 Evidence for hydrogen bonding in simple hydrides – graphs and theory

WHAT NEXT?

Part 8 sub–index: 8.1 Vapour pressure, nature, origin and examples explained * 8.2.1 Introduction to the types of intermolecular forces and examples explained (index) * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecule of similar molecular mass * 8.3 Boiling point plots of six organic homologous series - graphs and explanation * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * Evidence and theory for hydrogen bonding in simple covalent hydride * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. K_c and K_p equilibrium expressions and calculations * Part 3. Equilibrium and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series