Chemistry 11 Enriched
Reference p. 354-358
Unit VII
Problem set 7.7 Intermolecular Forces
For many substances, the forces between an atom or molecule and its neighbouring atoms or molecules tell a
lot about its properties. It explains things like melting and boiling points, evaporation rates and solubility.
There are a number of different intermolecular forces. The following is generally in decreasing strength.
1.
Ion – Ion Interactions
The strongest of the intermolecular forces arise from the electrostatic attraction of a negatively
charged ion with its neighbouring positively charged ion.
The resulting solid is called a lattice. The stronger the ionic interaction
the harder it is to break the lattice apart.
The melting point of ionic solids (like sodium chloride – shown here) are
generally very high – around 1000 oC.
The interactions between the ions can be described by Coulombs law,
which we have been considering when looking at the interaction of
electrons with the nucleus of an atom.
F = k Q1Q2 / d2
http://www.docbrown.info/page04/4_72bond2.htm
2.
So, the larger the ions are, the farther apart they are and so the less the
force of attraction. Conversely, ions that have 2+ and 2- charges will
form a stronger lattice than those with 1+ and 1-.
Dipole – Ion Interactions
A dipole results whenever a molecule has polar bonds which do not cancel each other out due to
symmetry. Consider the CO2 molecule:
O=C=O
http://chemistry.stackexchange.com/questions/1107/why-is-ho-a-dipole
The C=O bonds are polar, since the electronegativity difference
between the C and O is 3.44 – 2.55 = 0.85.
Due to this difference, the electrons are pulled away from the C toward
the O’s. Because the bond angle is 180o, however, these lopsided pulls
cancel each other out. Overall, this means that the molecule is NOT
polar.
(For physics students, imagine the polar bonds as vectors – for CO2, the
resultant vector is 0).
The diagram to the left shows the electron density mapping for both
CO2 which has polar bonds but is not a polar molecule, and for water,
which also has polar bonds and which is a polar molecule.
Water is polar because the polarity of the bonds does not cancel.
These polarities actually add together to create overall polarity to the
molecule. This polarity of a molecule is called a “dipole”.
Chemistry 11 Enriched
Reference p. 354-358
Unit VII
If a molecule has a dipole, it means that it has an end that is more negatively charged and an end that is more
positively charged. It is these charges that interact with ions.
Consider the dissolving of ionic salts that we looked at in the
solutions unit. The ions that form the ionic solid can be pulled
apart from each other due to the interaction they have with
the polar water molecules. In the diagram to the left, you can
see that the positive H end of the water molecules surround
the Cl- ions and the negative O end of the water molecules
surround the Na+ ions.
It is when the water dipole and ion interaction is stronger than
the ion-ion interaction that the solid will dissolve.
http://chemwiki.ucdavis.edu/Under_Construction/chem1/Solution_Chemistry/Solution_Types_and_Energetics
3.
Dipole – Dipole interactions.
In the same way that the negative side of water can attract a positive ion, it can also attract the positive side
of another water molecule, or the positive side of another polar molecule. It is for this reason that polar
molecules are usually soluble in water – the molecule interacts strongly with the water molecules, so they get
pulled apart from each other and spread evenly throughout the solution.
It is for this reason that ammonia (NH3) is soluble in water, for example.
In this diagram, the dipole points upward, toward the lone pair. The negative (top) side of
the molecule is attracted to the positive H side of water allowing it to dissolve.
Non polar substances, however, are not very soluble at all in water (see Dipole – Induced
Dipole interactions below for more information about this).
There is a special class of very strong dipole-dipole interactions called Hydrogen bonding. H bonding only
occurs with the most electronegative elements, where these elements are attached to a H atom.
N, O and F
The presence of these strong interactions helps to explain why the melting and boiling points of water are
much higher than for other molecules of similar size that are not polar. CH4 has a melting point of -182oC,
compared to water’s 0 oC. A lot more energy has to be added to the water to allow its particles to pull apart
far enough to slip past each other than is needed for CH4, since the CH4 molecules are not polar and thus are
not as strongly attracted to each other.
H bonding also explains the following trend in boiling points:
Substance Boiling Pnt Substance Boiling Pnt Substance
H2O
100
H2S
-60
H2Se
HF
20
HCl
-85
HBr
Boiling Pnt
-40
-66
Substance
H2Te
HI
Boiling Pnt
-2
-35
Chemistry 11 Enriched
Reference p. 354-358
H3N
H4C
-33
-164
Unit VII
H3P
H4Si
-88
-112
H3As
H4Ge
-63
-88
H3Sb
H4Sn
-17
-52
Notice how for the compounds of O, F and N the boiling point is very high, then it drops dramatically to the
element below, then slowly increases as you progress down the family.
For CH4, however, and its family, there is only a steady increase (as will be explained in the last section of this
problem set – which will also explain why for the 2nd to 3rd to 4th element in a family the boiling point
increases).
The higher than expected boiling point for these compounds is due to Hydrogen Bonding – a very strong
version of dipole-dipole interactions.
4.
Dipole – Induced Dipole Interactions
When a polar molecule comes into contact
with a non-polar one, the dipole from the polar
molecule will affect the electron cloud on the
non-polar molecule and shift its electron
density.
Consider the interaction of O2 (non-polar) with
water. The negative side of a water molecule
will repel the electrons of the O2 on the side
closest to it (as shown) or the positive side of
water could attract the electrons on the O2 on
the side closest to it (not shown). Either way, an attraction is created and the molecules can come together.
This helps to explain why O2 can dissolve in water (although it is not very soluble compared to polar molecules
or ionic salts). The short term polarity that exists on the O2 due to the presence of water is called an Induced
Dipole.
5.
Induced Dipole – Induced Dipole
It is also the case that for very short periods of time that the electron
density may be on one side of a non-polar molecule, creating a very
short-lived dipole. This short-lived induced dipole can then induce a
dipole in its neighbouring molecule – as we saw with dipole-induced
dipoles.
It is this induced dipole-induced dipole interaction that explains why if
the temperature is lowered enough, that even particles like He or H2
can condense and even freeze. (The freezing point of He is very close
to absolute zero, so this interaction is very, very weak for He).
Generally, the larger the molecule, the more easily its electron cloud
can be distorted and the longer an induced dipole will exist. The
ability of a substance to distort is called its polarizability. The longer it
exist, the stronger it will be – so this explains why H2Te requires a
higher temperature to melt than H2Se which also requires a higher
temperature to melt than H2S – the trend we saw above when looking
at hydrogen bonding. The H2Te molecule has a larger, more easily
Chemistry 11 Enriched
Reference p. 354-358
Unit VII
disturbed electron cloud which allows for longer and stronger induced dipoles and a stronger interaction
between one molecule and its neighbour.
Questions for further understanding:
p. 360-362 #5b,7-12
7.1
Rank the following in order of increasing strength:
i) a covalent bond
ii) London dispersion forces
iii) dipole ion interactions
iv) dipole – dipole interactions
v) ion – ion interactions
7.2
What is the kind of interaction that exists between non-polar molecules?
7.3
Which of each pair would have higher polarizability?
a) NH3 or PH3
b) HF or HI
c) CO2 or SiS2
7.4
Which of the following molecules can have hydrogen bonding?
a) CH3F
b) H2O
c) NH2CH3
d) CH3OH
Answers:
7.1
ii (especially small molecules) < iv < iii < v < i
7.2
London Dispersion Forces only
7.3
PH3, HI, SiS2
7.4
b), c) and d) - a) does not have an H bonded to the F.