Homework
• Private study work (bring notes to show me next
lesson);
• Read pages 62 – 65 in your text book
• Complete the summary questions on London (page
63)
• Look at the following websites
• http://www.chemguide.co.uk/atoms/bonding/electr
oneg.html#top
• http://www.chemnotes.org.uk/f321.html
• Topic 5, concentrate on hydrogen bonding and
London.
London forces and hydrogen
bonding
Tuesday, 15 March 2016
Thinking skills – Try to use key scientific words where possible.
Stage 1: This picture definitely shows
me…
Stage 2: I think
this picture
shows me…
Stage 4: The
questions I need
to ask about this
picture are…
Stage 3: This picture does not show
me…
London forces and hydrogen bonding
Learning Objective:
• Describe the intermolecular interactions forces: permanent dipole
– dipole forces, London forces and hydrogen bonding
Learning Outcomes:
• understand the nature of intermolecular forces resulting from the
following interactions:
i London forces (instantaneous dipole – induced dipole)
ii permanent dipoles
iii hydrogen bonds
• understand the interactions in molecules, such as H2O, liquid NH3
and liquid HF, which give rise to hydrogen bonding
• understand the following anomalous properties of water resulting
from hydrogen bonding:
i its relatively high melting temperature and boiling temperature
ii the density of ice compared to that of water
• be able to predict the presence of hydrogen bonding in molecules
analogous to those mentioned above
Intermolecular Forces
Strength of Bonds and Forces:
• Ionic and covalent bonds are strong.
• Ionic bonds hold ions together in a lattice so
that at room temperature all ionic compounds
are solid.
• Covalent bonds hold atoms together by
sharing electrons. Many covalent compounds
are small molecules with strong covalent
bonds within them. These are intramolecular forces.
Intermolecular Forces
Intermolecular Forces: is an attractive force
between neighbouring molecules.
• Intermolecular forces are weak compared to covalent
bonds.
• Intermolecular forces act between different
molecules. They are caused by weak attractive forces
between very small dipoles in different molecules.
• Intra-molecular bonds act within one molecule.
Intermolecular Forces
Intermolecular Forces:
There are three types of intermolecular forces;
• Permanent dipole-dipole interactions
• London forces (induced dipole forces)
• Hydrogen bonding.
Bond Type
Ionic and covalent
bonds
Relative Strength
1000
Hydrogen bonds
50
Dipole-dipole forces
10
London forces
1
Permanent dipole-dipole
interactions
A permanent dipole-dipole force: a weak attractive
force between permanent dipoles in neighbouring polar
molecules.
Polar molecules have a permanent dipole.
The permanent dipole of one molecule attracts the
permanent dipole of another.
Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the
molecules may align so there is electrostatic attraction
between the opposite charges on neighbouring molecules.
Permanent
dipole–dipole
forces (dotted
lines) occur in
hydrogen chloride
(HCl) gas.
The permanent dipole–dipole forces are approximately
one hundredth the strength of a covalent bond.
London forces
London forces (or induced dipole-dipole interactions) act
between all molecules, whether they are polar or nonpolar.
• They are the weakest intermolecular force.
• They act between very small, temporary dipoles in
neighbouring molecules.
London forces
• Electrons are always moving in an atom.
• Would it be possible for a non-polar
molecule or atom to produce a dipole?
• Why or why not?
London forces
Temporary dipoles
What will happen if two molecules
or atoms are near each other and
one has a temporary dipole?
London forces are attractions between temporary
dipoles.
• What factors might affect the
strength of the London forces?
• The greater the number of
electrons  the larger the
induced dipole  the greater
the London forces.
London forces –
Boiling Points
• London forces are the only attractions between
non-polar molecules.
Noble Gas
No. of
electrons
He
-269
2
Ne
-246
10
Ar
-186
18
Kr
-153
36
Xe
-108
54
Rn
-62
86
• No. of e- increases
• London forces increase
• Boiling point increases
If there were no London forces it would be impossible to liquefy the
noble gasses or non polar molecules.
Strength of200London forces
This is illustrated
by the boiling
points of group 7
elements.
150
boiling point (°C)
The strength of
London forces
increases as
molecular size
increases.
100
50
0
-50
-100
-150
-200
F2
Cl2
Br2
I2
element
Atomic radius increases down the group, so the outer electrons
become further from the nucleus. They are attracted less
strongly by the nucleus and so temporary dipoles are easier to
induce.
Strength of London forces
The points of contact between molecules also
affects the strength of London forces.
butane (C4H10)
boiling point = 272 K
2-methylpropane (C4H10)
boiling point = 261 K
Straight chain alkanes can pack closer together than
branched alkanes, creating more points of contact
between molecules. This results in stronger London
forces.
How can a gecko’s feet stick to almost any surface?
Write down your ideas.
HYDROGEN GRAPH
• In groups, try to come up with an explanation
for the pattern of each graph.
• The boiling point of compounds of hydrogen and
group 4 elements. Why does it increase?
How hydrogen bonding affects boiling points
Look at the the
boiling points of
compounds of
hydrogen and
group 5,6 and 7
elements.
What is
unusual?
DIFFERENCE IN
ELECTRONEGATIVITY
H
2.1
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Al
Si
1.5
1.8
Na
0.9
K
0.8
Mg
1.2
P
2.1
S
2.5
Cl
2.9
Br
2.2
what is hydrogen bonding?
• They occur when hydrogen is bonded to either
oxygen, nitrogen or fluorine.
• There is a large difference in electronegativity
between H and O,N,F which results in a strong
permanent dipole.
• Hydrogen bonding is the intermolecular force
that occurs between molecules containing these
permanent dipoles.
• Hydrogen bonding is the strongest type of
intermolecular force.
Hydrogen Bonding
Hydrogen bonding: A strong intermolecular force
between a δ+ hydrogen atom covalently bonded to
F, O or N and a lone pair of electrons on the δ- F, O
or N atom of a nearby molecule.
Occurs in O-H, N-H bonds and H-F
Hydrogen bonding
In molecules with OH
or NH groups, a lone
pair of electrons on
nitrogen or oxygen is
attracted to the slight
positive charge on the
hydrogen on a
neighbouring molecule.
hydrogen
bond
lone pair
Hydrogen bonding makes the melting and boiling points of
water higher than might be expected. It also means that
alcohols have much higher boiling points than alkanes of a
similar size.
• Why do hydrogen bonds
only form between O-H, N-H
and F-H?
It might help to sketch the
shape and dipoles of the
molecules containing these
atoms (such as H2O)
What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a
larger dipole occurs than in other polar bonds.
This is because these atoms are highly
electronegative due to their high
nuclear charge and small size. When
these atoms bond to hydrogen,
electrons are withdrawn from the H
atom, making it slightly positive.
The H atom is very small so the positive charge is more
concentrated, making it easier to link with other molecules.
Hydrogen bonds are therefore particularly strong examples of
permanent dipole–dipole forces.
DEMONSTRATE
hydrogen bonding between HF
molecules
• Following a similar procedure, show by
way of diagram, the hydrogen bonding
between:
HELP SHEET AVAILABLE TO
a) two water molecules
b) two ammonia molecules
HELP YOU STRUCTURE
YOUR ANSWER
• Extension: Why do you think water is
considered to be a ‘perfect example’ of
hydrogen bonding?
Why do icebergs
float?
Hydrogen Bonds
• What happens to the
volume of water when it
freezes?
• How does this differ from
other liquids?
• What causes this?
•
•
Ice has open lattice, H-bonds hold water molecules apart.
When ice melts, H2O molecules move closer together.
How much does volume increase?
Comparison of:
Liquid
water
Mass = 100 g
Volume =
100 mL
Ice
Mass = 100 g
Volume = ?
mL
Density = 1.0 Density =
g/mL
0.92 g/mL
Many simple
molecular
structures are
gases at room
temperature but
H2O is a liquid –
why?
• Other molecules are held together
by London forces.
• However, water molecules are
held together by hydrogen bonds
which are stronger and harder to
overcome.
Bond type
Ionic and covalent bonds
Hydrogen bonds
Dipole-dipole forces
London forces
Relative
strength
1000
50
10
1
Other
• H-bonds also give water relatively
high surface tension and viscosity.
• H-bonds are important in organic
compounds containing O-H and N-H
bonds (alcohols, carboxylic acids etc)
• They are responsible for shape of
proteins and even DNA.
DNA
DNA
DNA Double Helix
Hydrogen bonding explains the higher boiling
point of hydrogen fluoride compared to the
other hydrogen halide compounds.
Exam questions (new sample material)
Exam questions (new sample material)
Exam questions
1 mark each (total two marks)
Exam questions
Exam questions
London exam questions (old OCR)
Chlorine, bromine and iodine are halogens
commonly used in school and college
experiments.
Describe how London forces arise
(3 marks)
Exam question
(3 marks)
Exam questions
a) State and explain the trend in the boiling
points of chlorine, bromine and iodine.
(3 marks)
b) The halogen astatine does not exist in large
enough quantities to observe any of its
reactions.
Why would astatine be expected to react
similarly to other halogens?
(1 mark)
a)
London and hydrogen bonding
Learning Objective:
• Describe the intermolecular interactions forces:
permanent dipole – dipole forces, London forces and
hydrogen bonding
Learning Outcomes:
• State the different types of intermolecular bonding
• Describe intermolecular forces in terms of permanent and
instantaneous dipoles (including hydrogen bonding)
• Draw diagrams to describe these effects
• Explain the trend in boiling point due to these forces
• Describe and explain the anomalous properties of H2O
resulting from hydrogen bonding, eg:
the density of ice compared with water,
its relatively high freezing point and
boiling point