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Brown's Chemistry Advanced A Level Notes - Theoretical–Physical
Advanced Level
Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 8
Part 8.3 Comparative boiling point plot
graphs for selected
homologous series of organic molecules
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INDEX for Part 8.
Phase equilibria–vapour
pressure, boiling point and intermolecular forces
Index of ALL my chemical equilibrium
context revision notes Index
ALL my advanced A
level theoretical
chemistry revision study notes
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 forces?
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.
Three 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.
-
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 (δ–N–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 than alkanes due to
the extra intermolecular force of 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)
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).
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 carbon lattice (sublimation
at 1–4
carbon atoms at a time?).
WHAT NEXT?
INDEX for Part 8.
Phase equilibria–vapour
pressure, boiling point and intermolecular forces
Index of ALL my chemical equilibrium
context revision notes Index
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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