Doc Brown's Advanced A level theoretical chemistry revision notes

Part 8.3 Comparative boiling point plot graphs for six selected homologous series of linear organic molecules

[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:  boiling point comparison of organic homologous series  [page updated May 2nd 2026 *]

8.3 Comparative Boiling Point Plots for six Organic Homologous Series

In this section I am assuming you are familiar with the different types of intermolecular bonding, particularly:

Read first Summary of Van der Waals forces and introduction to intermolecular forces

Instantaneous dipole – induced dipole interactions (London forces or dispersive forces)

Permanent dipole – permanent dipole interactions (Keesom forces/orientation forces)

and Hydrogen bonding intermolecular forces

otherwise this page might prove more difficult than is necessary!

How do the boiling points of one homologous series of organic molecules compare with another?

How do we explain the differences in terms of intermolecular attractive forces and polarizability of the molecule?

but only comparing linear molecules of the six selected aliphatic homologous series

and comments on the intermolecular forces (intermolecular bonding) involved

8.3 Plot (a) Boiling Point versus the number of electrons in the molecule

The detailed discussion of boiling point trends is mainly based on this first plot of boiling point (K) versus the number of electrons in the molecule.

8.3 Plot (a) Boiling Point (K) versus the number of electrons in the molecule

Some general comments that apply to all three graph discussions (a) to (c)

This is the best comparison that you can do, because, this gives a baseline of the effect of instantaneous dipole – induced dipole interactions i.e. intermolecular forces common to all molecules with the same number of electrons in the molecule and they are all approximately linear molecules.

The larger the molecule and the greater the number of electrons, the greater the polarizability of the molecules, the greater the chance of dipole formation (δ⁺ δ^–), hence the increase in the intermolecular force between neighbouring molecules.

BUT, there may be several types of intermolecular force contributing the total effect of the Van der Waals intermolecular bonding e.g. permanent dipole - permanent dipole attractions.

Three general trends are immediately discernable.

• (1) With increased molecular mass, carbon atoms in the chain or electrons, the boiling point increases i.e. what you would expect from increasing instantaneous dipole – induced dipole attractive forces.

  • The larger the molecule, the more electrons it has and so the more polarizable is the molecule, irrespective other attractive forces arising from polar functional groups (in this case) at the end of a linear molecule.

• (2) Initially, the highly polar bonds of the lower aldehydes, alcohols and carboxylic acid, lead to relatively high boiling points compared to the less polar molecules due to contributions of permanent dipole - permanent dipole (aldehydes) and also hydrogen bonding (alcohols and carboxylic acids) to the intermolecular bond.

• (3) BUT, all the boiling point graph lines tend to converge, despite the highly polar molecules, because the carbon chain of -CH₂- units gets longer and longer, the dominant intermolecular force of the intermolecular bonds is due to the instantaneous dipole - induced dipole forces.

  • Which leads to an important 'sub–trend' influence of a polar bond giving rise to an extra contribution to the intermolecular forces considerably decreases with increase in carbon chain length.

  • Generally speaking the more polar the molecule the higher the boiling point for those molecules of similar molecular mass or number of electrons, but this is only significant for small very polar molecules.

  • BUT the predominant intermolecular force for most molecules is the instantaneous dipole – induced dipole (δ⁺ δ^–), whose partial electrical charges account for the intermolecular attractive force which increases with the size of the molecule – best considered as increase in electron charge clouds that become more polarizable.

• Alkane boiling point curve

• Why it increases with chain length is illustrated below using ball and stick AND space–filling styled molecular representations. The addition of a CH₂ unit adds another 8 electrons to the molecule making the overall molecule more polarizable because the greater total of electron clouds can leader to more/bigger dipoles created.

• i.e. octane C₈H₁₈ melting point –57^oC and boiling point 126^oC, melts and boils at considerably higher temperatures than octadecane C₁₈H₃₈ with a melting point of 27^oC and a boiling point of 316^oC.

Pictures (above and below) to illustrate the increase in intermolecular attractive forces with increase in size of molecule.  The diagrams only refer to the instantaneous dipole – induced dipole interactions between non-polar alkane hydrocarbon molecules.

• The instantaneous dipole – induced dipole forces predominate EXCEPT for a few small highly polar molecules exhibiting hydrogen bonding e.g. methanol and ethanol, methanoic and ethanoic acid, methanamide and ethanamide where the permanent dipole – permanent dipole interactions are very significant.

• Taking each molecule in order of increasing boiling point ....

  • Alkanes have the relatively lowest boiling points because only instantaneous dipole – induced dipole interactions can contribute to the intermolecular forces.

  • The steady increase in intermolecular attractive forces with increase in chain length explains why components in crude oil can be separated by fractional distillation.

  • Primary haloalkanes (halogenoalkanes, n–alkyl halides) have slightly higher boiling points than alkanes because of the carbon–chlorine polar bond (^δ+C–Cl^δ–) giving a permanent dipole – permanent dipole an extra small contribution to the intermolecular forces.

  • Primary aliphatic amines are the next highest boiling because the N–H bond (^δ–N–H^δ+) is more polar than the C–Cl bond and these permanent dipole – permanent dipole interactions will give rise to hydrogen bonding but not as strong as for alcohols.

    • The source of hydrogen bonding is RH₂N^δ–llll^δ+H-NH-R

  • Aldehydes (similar for isomeric ketones) are very similar to the linear primary aliphatic amines. It would appear here that the permanent dipole – permanent dipole intermolecular attractive forces due to the polarised carbonyl bond ^δ+C=O^δ– have about the same effect as the hydrogen bonding in the amines.

  • You would get a similar graph line for the boiling points linear series of 2-ketones (-2-ones) i.e. propanone, butan-2-one, pentan-2-on.etc. which have the same polar bond ^δ+C=O^δ–

  • Aliphatic alcohols (alkanols) show significantly higher boiling points than alkanes due to the extra intermolecular force of hydrogen bonding BUT only for the lower members.

    • The source of hydrogen bonding is R-O^δ–llll^δ+H-O-R

    • With all the series showing hydrogen bonding i.e. alcohols, amines and carboxylic acids, the effect of the hydrogen bonding contribution to the total intermolecular attractive force diminishes with increase in chain length as the instantaneous dipole - induced dipole forces become increasingly important. You can see this as many of the graph lines converge at higher carbon atom numbers.

  • Carboxylic acids (alkanoic acids) are even higher than alcohols because there are extra >C=O ... >C=0 attractions as well as hydrogen bonding or you could argue there are two sites on the molecule for hydrogen bonding, again only for the lower members, and dimers are readily formed for the lower members of the series.

    • One source of hydrogen bonding is >C=O^δ–llll^δ+H-O-

    • Note: For ethanoic acid, if you do the plot point with double the electrons i.e. as in the dimer, the point is then close to the alkane curve!

    • This is what you expect if the liquid ethanoic acid is a cyclic hydrogen bonded dimer and you assume the only attractive intermolecular force is due to the instantaneous dipole – induced dipole interactions!

• Although each interaction is minute, collectively, even the weakest of intermolecular forces can add up to give an impressive effect.

  • Non–polar alkanes are only attracted to each other via the weakest of Van der Waals forces (instantaneous dipole – induced dipole attractive interactions) but once the carbon number gets high, so does the boiling point! and enthalpy of vaporisation which eventually equals bond energies, so very high molecular mass hydrocarbons can thermally decompose before they boil!

  • Some examples of high boiling alkanes are

  • C₃₂H₆₆ M_r = 451, 258 electrons, bpt 467^oC (740K), ΔH_sub(s=>g)  = 271 kJ mol^–1

  • C₃₅H₇₂ M_r = 493, 282 electrons, bpt 490^oC (763K)

8.3 Plot (b) Boiling Point versus the molecular mass

8.3 Plot (b) Boiling Point (K) versus the molecular mass

These plots give a very similar pattern to the electron number plots, but its still best to think via the electron number of plot (a) and the polarizability of the molecule .

8.3 Plot (c) Boiling Point versus the number of carbon atoms in the molecule

8.3 Plot (c) Boiling Point (K) versus the number of carbon atoms in the molecule

• This plot disregards the number of electrons or molecular mass and is not the most useful set of graphs for comparing the boiling points of organic homologous series.

• It is however, a way of comparing the boiling points of various homologous series for the same number of carbon atoms in the molecule.

• All the plot lines show the steady increase of boiling point with carbon number as an extra –CH₂– group is successively added and the relatively diminishing effects of a permanent dipole as the instantaneous dipole – induced dipole interactions almost totally dominate with longer carbon chain lengths.

• It should be remembered that although each interaction is minute, collectively, even the weakest of intermolecular forces can add up to give an impressive effect.

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