Transcript Van der Waals` forces

Homework
• Private study work (bring notes to show me next
lesson);
• Read pages 62 – 65 in your text book
• Complete the summary questions on van der Waals’
(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 van
der Waals’.
Van der Waal’s and hydrogen
bonding
Friday, 10 April 2015
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…
Van der Waal’s and hydrogen bonding
Learning Objective:
• Describe the intermolecular interactions forces:
permanent dipole – dipole forces, van der Waals’ 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
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 ate three types of intermolecular forces;
• Permanent dipole-dipole interactions
• Van der Waals’ forces (induced dipole forces)
• Hydrogen bonding.
Bond Type
Ionic and covalent
bonds
Relative Strength
1000
Hydrogen bonds
50
Dipole-dipole forces
10
Van der Waals’ 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.
Van der Waals’ forces
van der Waals’ forces (or induced dipole-dipole
interactions) act between all molecules, whether they
are polar or non-polar.
• They are the weakest intermolecular force.
• They act between very small, temporary dipoles in
neighbouring molecules.
Van der Waals’ 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?
Van der Waals’ forces
Temporary dipoles
What will happen if two molecules
or atoms are near each other and
one has a temporary dipole?
Van der Waals forces
Van der Waal’s forces are attractions between
temporary dipoles.
• What factors might affect the
strength of the van der Waals
forces?
• The greater the number of
electrons  the larger the
induced dipole  the greater
the van der Waals forces.
Van der Waals’ forces
– Boiling Points
• Van der Waals’ 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
• Van der Waals’ forces
increase
• Boiling point increases
If there were no van der Waals’ forces it would be impossible to
liquefy the noble gasses or non polar molecules.
Strength of van
der
Waals
forces
200
This is illustrated
by the boiling
points of group 7
elements.
150
boiling point (°C)
The strength of van
der Waals 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 van der Waals forces
The points of contact between molecules also
affects the strength of van der Waals 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 van der
Waals 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 bond: is a strong dipole-dipole
attraction between:
• An electron-deficient hydrogen atom on one
molecule (O-Hδ+ or N-H δ+) and;
• A lone-pair of electrons on a highly
electronegative atom on a different molecule. (HOδ- or H-Nδ-)
Occurs in O-H and 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 and
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 van der Waals’ 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
Van der Waals’ 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
Page 65 in text book
Questions
1) Which of the following molecules have
hydrogen bonding?
a) H2S b) CH4 c) CH3OH d) PH3 e) NO2 f) CH3NH2
2)
a)
b)
c)
d)
Draw diagrams showing H-bonds between:
2 molecules of water
2 molecules of ammonia
2 molecules of ethanol
1 molecule of water and 1 molecule of ethanol
van der Waals’ exam questions
Chlorine, bromine and iodine are halogens
commonly used in school and college
experiments.
Describe how van der Waals’ 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)
6 marks
6 marks
Van der Waal’s and hydrogen bonding
Learning Objective:
• Describe the intermolecular interactions forces:
permanent dipole – dipole forces, van der Waals’ 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