
Doc Brown's
Chemistry Advanced Level Pre-University Chemistry Revision Study Notes for UK IB
KS5 A/AS GCE advanced level physical theoretical chemistry students US K12 grade 11 grade 12
theoretical chemistry
8.2.1 Intermolecular Forces
Intermolecular Bonding – Van der Waals forces
The different types of intermolecular force
(intermolecular bond) are described, explained and discussed with
examples i.e. Instantaneous dipole – induced dipole interaction (London
forces, dispersive/dispersion forces), permanent dipole – permanent
dipole interactions (Keesom forces, orientation forces),
Permanent
dipole – induced dipole interactions (Debye forces, induction forces)
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 and its importance in other
molecules *
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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8.2 Survey of
8 selected organic molecules
– their boiling points and intermolecular forces
8.2.1 Introduction to intermolecular forces
– Van der
Waals forces
also referred to as 'intermolecular bonding'
forces
(do NOT confuse with
chemical bonding between atoms or ions i.e. so-called ionic, covalent or
metallic bonds)
-
From the start understand that:
-
Intermolecular forces are all about
partially positive (δ+)
sites and partially negative
(δ–)
sites on molecules causing the attraction between neighbouring
molecules - though their origin can differ.
-
The fact that molecules
congregate together to form liquids and solids suggests that there
must be attractive forces between the molecules independently from
the intramolecular bonds which hold the atoms together in the
molecule.
-
The origin of each
source of intermolecular force is summarised below and discussed further
for particular molecules.
-
In the context of this
page, the word dipole
means an asymmetric distribution of electron electrical charge to
give partially positive (δ+) and partially negative
(δ–) regions in the
same molecule.
-
In a simple sense its a molecule with
a partially positive end and a partial negative charge at the other
end.
-
Electric dipoles (δ+
and δ–) may be
permanent or transient (temporary) and the molecules
discussed here are electrically neutral overall.
-
There
are always attractive forces operating between ANY particles whatever
their particle constitution in gases, liquids or solids composed of
atoms, ions or molecules.
-
They are referred to as intermolecular attractive forces
or intermolecular bonding.
-
Collectively they are often
referred to as Van der Waals forces.
-
DO NOT confuse
intermolecular bonds with the very much stronger intramolecular
bonds e.g. between atoms in a molecule like the O-H bond holding
atoms together in water, or the C-C and C-H bonds holding atoms
together in hydrocarbon molecules.
-
The total intermolecular
force is quoted as a summation of the various possible dipoles
interaction and the principal attractive forces are shown in bold for selected
molecules.
-
Wherever possible,
albeit just for a few cases, I've quoted % contributions from the three
types of intermolecular attractive force that I've been able to
obtain from internet searches or textbooks and if I couldn't match
the molecule then I may quote percentages for a similar molecule.
-
One source used by writers of research papers is A. L. McClellan,
Tables of Experimental Dipole Moments.
-
In pre–university
advanced chemistry exams I suggest you
use the terms in bold to
describe the intermolecular force component
-
Summary of the types of
intermolecular bonding forces (Van Waals forces)
-
Instantaneous
dipole – induced dipole interaction
-
Also called
London
forces or
dispersive or dispersion forces.
-
The electrons of an atom
behave in a random way within the spatial region they occupy for
their specific quantum level e.g. in 3s, 2p, 3d atomic orbitals or a
bonding molecular orbital.
-
At any given instant in
time the electron cloud will randomly distorted, giving rise
to a dipole of partial charges which then induces a dipole in a neighbouring
molecule.
- Note that these partial charges
are shown as a delta + (δ+)
or a delta – (δ–)
and they are tiny charges compared to a full single plus charge
e.g. on an Na+ sodium ion or a full single
minus charge on a Cl– chloride ion.
-
The random partial positive
charge of one dipole will attract the partial negative in the
neighbouring molecule or vice versa 
-
So even with a
completely non–polar hydrocarbon molecule (i.e. a molecule with no
significant polar bonds like alkanes and alkenes) there are still intermolecular attractive
forces.
- Even in the case of helium,
lowest boiling point of any substance, you still can get
transient dipoles because of the random behaviour of the
electrons ...
-
attractions
-
Instantaneous
dipole – induced dipole interaction increase the more electrons in
the molecule.
-
For most molecules,
this is the dominating contribution to the total intermolecular
force, but the presence of polar bonds can add a significant
contribution to this and the consequential affects on the properties
of the molecule.
-
Comparing the boiling
points and intermolecular forces operating between molecules with a
similar number of electrons does provide important insights
into their molecular behaviour.
-
However, you should be
aware that the way the electrons are distributed, both in terms of
their electronic energy levels, and their spatial distribution, can
have significant effects on the strength of instantaneous dipole –
induced dipole forces. You will see this particularly in the case
studies of
section 8.4.
-
Notes on
instantaneous dipole - induced dipole forces:
-
(i) You come across
other words other than instantaneous e.g. 'temporary', 'transient'
or even 'induced' – induced dipole attractions.
-
(ii) The molecule
does not have to be polar for this force to exist.
-
(iii) The same force
exists between ANY neighbouring molecules, whether they
are the same, different, polar or non-polar. It is a
universal intermolecular force.
-
See
comparing organic molecule
boiling points and
homologous series comparison
-
and
other case studies of
boiling points related to intermolecular forces
-
Permanent dipole –
permanent dipole interactions
-
Also called
Keesom forces or orientation forces.
-
If two atoms
constituting a bond have significantly different
electronegativities, the bond will be permanently polar and produce
a permanently polar molecule.
-
Such molecules posses
what is known as
dipole moment.
-
Therefore, as result of
this permanent dipole, these permanently polarised molecules will
attract neighbouring molecules because of this dipole moment as
well as the attraction due to instantaneous dipole – induced dipole.
-
e.g.
-
attractions
seem to make little difference to the bpt!
-
or in
carbonyl compounds δ+C=Oδ–....δ+C=Oδ– in organic carbonyl compounds e.g. aldehydes,
ketones and carboxylic acids.
-
Trichloromethane (CHCl3)
is another polar molecule
δ-Cl3Cδ+H
-
HYDROGEN BONDING
-
There is special
sub–category of permanent dipole – permanent dipole interactions
called hydrogen bonding.
-
It is a
spatially directed permanent dipole -
permanent dipole attractive force - the only intermolecular
force that is spatially and specifically directional.
-
Hydrogen bonding only
usually occurs when the three most electronegative elements (N, O
and F) are covalently bonded to a hydrogen atom (intramolecular) AND
bonded to a similar neighbouring molecule and an
intermolecular bonding force. (more on
hydrogen bonding in section 8.6).
-
In these molecules you
get one of the following three very polar bonds:–
-
(i)
δ–:N–Hδ+
e.g. in ammonia NH3, amines R-NH2, amides
RCONH2 (R = alkyl or aryl)
-
-
(ii)
δ–:O–Hδ+
e.g. in water H2O (above), alcohols ROH (above), carboxylic acids
RCOOH (R = alkyl or aryl)
-
(iii)
δ+H–Fδ–
in hydrogen fluoride HF
-
and via these highly polar bonds you get molecule to
molecule attraction via so called hydrogen bonding.
-
Note the spatially
important non-bonding pairs of electrons (:)
on the most electronegative atom.
-
These are the strongest permanent dipole permanent
dipole intermolecular forces
-
e.g. using
llll to indicate a hydrogen bond
-
δ–O–Hδ+llllδ–:O–Hδ+llllδ–:O–Hδ+llllδ–:O–Hδ+llll in water (liquid or solid ice)
-
llll δ–N–Hδ+llllδ–:N–Hδ+llllδ–:N–Hδ+ in amines
or liquid ammonia,
-
in the case of
carboxylic acids the dominant interaction is the hydrogen bonding
via
-
In hydrogen fluoride, in
all physical states you get a 'chain connection' llll
δ–F–Hδ+llllδ–F–Hδ+llllδ–F–Hδ+
-
You can also get hydrogen
bonding between these different molecules e.g.
-
δ–N–Hδ+llllδ–:O–Hδ+ or δ–O–Hδ+llllδ–:N–Hδ+ in aqueous ammonia solution (NH3(aq),
-
C-δ–O–Hδ+llllδ–:O–Hδ+
in aqueous carboxylic acid solutions (RCOOH(aq)),
-
δ–F–Hδ+llllδ–:O–Hδ+
or δ–O–Hδ+llllδ–:F–Hδ+
in hydrofluoric acid solution (HF(aq)).
-
A slightly different
and modestly intriguing case of hydrogen bonding!
-
Trichloromethane
(CHCl3) and propanone (CH3COCH3)
are both polar molecules, but do not hydrogen bond with
themselves.
-
BUT, if you mix the
two liquids, they readily dissolve in each other via
hydrogen bonding!
-
Cl3C-Hδ+llllδ-:O=C(CH3)2
-
Note that the three
electronegative chlorine atoms have such an effect on
the carbon atom (δ+),
that the hydrogen atom also acquires a sufficient (δ+)
to hydrogen bond with the oxygen atom (via a lone pair
of electrons).
-
An important exam note:
-
You must clearly show the
directional linearity of the
Xδ--Hδ+ǁǁǁ:Xδ-
arrangement of the hydrogen bond including the single X-H
covalent bond and the lone pair on the other X atom too! (X
is usually O, N or F)
-
You must do this accurately in exams
when drawing intermolecular bonding diagrams of water or alcohols because it is the only spatially
directed intermolecular force, all the rest of the other
types of intermolecular bonding forces are randomised - the δ+
and δ- electric fields acting in all directions.
-
The spatial directional
nature of the hydrogen bond is very important when studying
e.g. the crystalline structure of ice or the double helix of
DNA - the latter is held together by base pair hydrogen
bonds.
-
For detailed case studies
see my main
hydrogen bonding
page
-
Permanent dipole –
induced dipole interactions
-
Also called Debye forces or induction forces.
-
The permanently polar
bond in one molecule can induce a dipole in a neighbouring molecule,
whether the other molecule is polar or non–polar, it makes no
difference, induction happens!
-
e.g.

attractions
-
This applies to any pure
polar molecule or in a mixture of a polar and non-polar
molecules as illustrated above.
-
This tends to be the
minority contribution to a molecule's total intermolecular
bonding forces.
-
The discussions below
primarily focus on intermolecular forces 1. and 2. and 3. is usually
less significant and gets a brief mention.
-
Data abbreviations used: Mpt = melting point,
Bpt = boiling point (given in Kelvin and Celsius)
-
Here in section 8.2 examples are
chosen from the eight homologous series listed below and all the
molecules have 32–34 electrons and four larger C, N or O atoms plus
hydrogen (except chloroethane C + C + Cl, but still has 34
electrons).
-
See
comparing organic molecule boiling points and
homologous series comparison
-
and
other case studies of
boiling points related to intermolecular forces
The survey and a preliminary
summary table
hopefully justified by the arguments outlined after the
table and in on a separate page in
section
8.2.2
-
Ins =
instantaneous (temporary) dipole – induced dipole attraction (a sort
of baseline force since it applies to all molecules, in fact it
operates between ANY adjacent particles - atoms, ions or molecules).
-
WP = weaker
permanent dipole – permanent dipole attraction (doesn't seem to have
much effect on the boiling point)
-
SP stronger
permanent dipole – permanent dipole attraction (NOT H
bonding, but has a definite effect on the boiling point)
-
HB = hydrogen
bonding attraction - the strongest permanent dipole – permanent dipole
attractive force, i.e. the strongest SP and has the largest effect
on the boiling point)
-
MHB multiple hydrogen bonding
attraction sites on the molecule
(i.e. where there are at least two 'functional' groups capable of two permanent
dipole – permanent dipole interactions including hydrogen bonding,
hence
producing an even bigger effect on raising the boiling point)
-
Note that permanent dipole – induced
dipole attractive forces are not mentioned much and generally only
contribute a small portion of the total intermolecular
force.
-
Also, where I can obtain
data, I've indicated the percentage contribution of the three
types of intermolecular attraction which contribute to the total
intermolecular force i.e. the % contributions to Van der Waals
force.
-
D = Debye dipole moment
units
8.4 Table 1a. Comparing 32–34
electron species – linear organic
molecules (4 C/O/N atoms)
1.to 8. are discussed in detail
on a separate page |
MOLECULE |
formula |
Mr |
electrons |
bpt |
ΔHvap/kJmol–1 |
Dipole moment/D |
Intermolecular
forces |
1. butane |
CH3CH2CH2CH3 |
58 |
34 |
272.5K/–0.5oC |
22 |
0.00 |
Ins |
2. methoxyethane |
CH3OCH2CH3 |
60 |
34 |
280K/7oC |
21 |
1.23 |
Ins, WP |
3. chloroethane |
CH3CH2Cl |
64.5 |
34 |
285.5K/12.5oC |
25 |
2.06 |
Ins, WP |
4. propylamine |
CH3CH2CH2NH2 |
59 |
34 |
321K/48oC |
30 |
1.17 |
Ins, HB |
5. propanone |
CH3COCH3 |
58 |
32 |
329K/56oC |
29 |
2.88 |
Ins, SP |
6. propan–1–ol |
CH3CH2CH2OH |
60 |
34 |
370K/97oC |
45 |
1.69 |
Ins, HB |
7. ethanoic acid |
CH3COOH |
60 |
32 |
391K/118oC |
58 |
1.74 |
Ins,
SP, MHB |
8. ethanamide |
CH3CONH2 |
59 |
32 |
494K/221oC |
46 |
3.60 |
Ins,
SP, MHB |
WHAT NEXT?
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ignore adverts at top
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
TOP OF PAGE
and sub-index
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