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Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 8.3

 8.3 Comparative boiling point plots–graphs for selected homologous series of organic molecules

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

(c) doc b KS4 Science GCSE/IGCSE Notes on reversible reactions and chemical equilibrium

 

Part 8 sub–index: 8.1 Vapour pressure origin and examples * 8.2.1 Introduction to Intermolecular Forces * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecules * 8.3 Boiling point plots for six organic homologous series * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * 8.6 Evidence and theory for hydrogen bonding in simple covalent hydrides * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp 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


8.3 Comparative Boiling Point Plots for six Organic Homologous Series

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

and comments on the intermolecular forces (intermolecular bonding) involved

Read first 8.2.1 A summary of Van der Waals forces and an introduction to intermolecular forces

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

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

This is the best comparison that you can, 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.

TOP and LINKSThree general trends are immediately discernable.

  • With increased molecular mass or electrons, the boiling point increases i.e. what you would expect from increasing instantaneous dipole – induced dipole attractive forces.

  • 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.

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

    • i.e. octane C8H18 melting point –57oC and boiling point 126oC, melts and boils at considerably higher temperatures than octadecane C18H38with a melting point of 27oC and a boiling point of 316oC.

  • 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.

  • 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.

    • Primary halogenoalkanes (n–alkyl halides) have slightly higher boiling points than alkanes because of the carbon–chlorine polar bond (δ+C–Clδ) 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 (δN–Hδ+) 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.

    • 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=Oδ–) have about the same effect as the hydrogen bonding in the amines.

    • Alcohols (alkanols) show significantly higher boiling points due to hydrogen bonding BUT only for the lower members.

    • 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.

      • 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 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 Mr = 451, 258 electrons, bpt 467oC (740K), ΔHsub(s=>g)  = 271 kJ mol–1

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

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8.3 Plot (b) Boiling Point versus the molecular mass

8.3 Plot (b) Boiling Point 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 plot (a).

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

8.3 Plot (c) Boiling Point 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.

  • The ultimate molecular structure held together solely by the weakest of the Van der Waals forces, i.e. the instantaneous dipole – induced dipole  intermolecular forces, is carbon in the form of graphite (trillions of atoms per molecule!) which only sublimes at over 3500oC! The total intermolecular attractive force is so great that graphite won't melt at normal pressure and above 3500oC the thermal vibration in the molecular lattice is then great enough to disrupt it and the solid vaporises directly in molecular chunks of the lattice (sublimation at 1–4 carbon atoms at a time?).

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WHAT NEXT?

Part 8 sub–index: 8.1 Vapour pressure origin and examples * 8.2.1 Introduction to Intermolecular Forces * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecules * 8.3 Boiling point plots for six organic homologous series * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * 8.6 Evidence and theory for hydrogen bonding in simple covalent hydrides * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp 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

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