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Teacher's Notes
Bonding
Between
Molecules
Grades: 11-12
Duration: 14 mins
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Chapter
Timing
The biscuit molecule
0:13
Dipole Dipole Bonds
2:50
Hydrogen Bonds in Water
4:49
Hydrogen bonding and life
7:33
Hydrogen bonding and DNA
9:02
Normal Dipole bond
10:13
Fluctuating Dipole Bond
10:46
In the first video of this two part series we
examined the ways in which atoms form bonds
with other atoms. The bond that forms is the
result of the sharing of the outer electrons of the
atoms. The kind of bond that forms depends on
the level of sharing. If the electron is completely
lost or completely gained then the atoms form
permanently charged ions. Positively charged
ions are electrostatically attracted to negatively
charged ions and the bond between them is an
ionic bond. In other bonds the electrons are not
completely lost or gained, but are shared
between the atoms involved. When two
electrons are shared between two atoms as
neighbours, the bond is a covalent bond. When
the electrons are shared with all neighbouring
atoms the bond is a metallic bond.
There are many different pure substances
and they all have distinct properties.
The combination of different bonds and
different atoms gives rise to different
structures. The structures that result from
these combinations include:
Ionic solids
Simple and complex molecules
Metals.
Also the chemical and physical properties of
each substance are mainly determined by
the bonds within the groups of atoms and
by the bonds between the groups. Further,
if the only bonds that exist were bonds
within molecules, then they would only exist
as gases. The size and the arrangement of
the atoms or molecules involved also
influences the properties.
Chemical reactivity depends on how
strongly the outer electrons of the atoms
involved are held and on how readily bonds
can be changed. The most obvious physical
property of a substance is its state at room
temperature, which is governed by its
melting and boiling points.
Just as there is a gradual transition of
bonding between atoms of different
elements across the periodic table, there is
a transition in the kinds of bonds between
molecules. The molecules listed in the table
above have about the same molecular mass
yet their boiling points are quite different.
This difference is the result of forces of
attraction producing different bonds
between the molecules. The forces are
electrical in nature and are grouped
together under the name Van der Waal
forces.
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Two types are generally recognized: dipoledipole forces and dispersion or London forces.
(Van der Waal 1873, dipole-dipole (Keesom)
1912, London 1930.)
At the higher end the bonds must be
substantial, as heat has to be put in to
separate the molecules to form a gas. At the
lower end considerable heat energy has to be
removed before the molecules can bond
together to form a liquid.
Dipole-dipole bonds.
Some molecules that are electrically neutral
can still have a permanent uneven distribution
of charges. One part will show a net positive
charge and the other will show a net negative
charge, and the molecule appears to have a
positive and a negative pole. The molecule is
said to be polar. We can imagine such
molecules aligning themselves in an electric
field between two charged plates. They
normally align with each other to form chains
and the strength of the chain is governed by
the size and distribution of the charges on the
molecules. The molecules are attracted to each
other by electrostatic forces called dipoledipole forces. The bond formed between the
molecules is called a dipole-dipole bond.
When hydrogen is one of the elements involved
then the dipole-dipole bond between the
molecules takes on some special
characteristics. The bond is termed a hydrogen
bond. It is important to remember that
although hydrogen bonds are stronger than
normal dipole-dipole bonds, they are not as
strong as the bonds within molecules and take
far less energy to break. A hydrogen bond is
less than one tenth the strength of a covalent
bond but their greater strength makes them
special.
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Hydrogen bonds in water.
Water is unusual in several ways and it is its
unusual combination of properties that make life
possible. One unusual property is its high boiling
point compared to other hydrogen compounds
with similar molecular mass. The bonds between
the molecules of water must be stronger and
more difficult to break. This also plays a role in
water being able to dissolve most ionic solids.
Another unusual property is its decrease in
density as it freezes. This is the opposite of the
expected result. Normally the solid is denser
than the liquid and will sink. To explain why ice
floats in water we have to look at the bonding
that builds the solid crystal lattice from the
water molecules. The bonding is due to
hydrogen bonds. Hydrogen also forms
compounds with other common non-metals. But
they do not show the same special bonding or
the same special properties. There are other
examples where the influence of hydrogen
bonding is definitely worth examining.
Hydrogen bonding and life.
Molecules of the common organic acid, acetic
acid, are able to join in pairs to form a dimer.
Two hydrogen bonds form temporarily to hold
the dimer together. And they are sufficiently
strong to hold some dimers together even as a
vapour. This confused chemists for a while until
they understood the role of hydrogen bonding
between organic acid molecules. Our most
important organic acid is DNA.
Hydrogen bonding and DNA.
The two strands of DNA are held together
across matching bases by hydrogen bonds and
the structure forms and behaves as a unit. The
unit can be pulled apart without damaging the
molecular strands.
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The amazing, complex process of replication
can then occur and this is the basis of the
continuance of life. Hydrogen bonding at its
best?
Fluctuating dipole bonds.
We have examined how dipole-dipole
bonding holds molecules together,
particularly as they form liquids and solids.
But not all molecules are polar. Carbon
dioxide is an example of a non-polar
molecule. All hydrocarbons are also non
polar. How do they form as a liquid or solid
at all, particularly when an electron cloud
surrounds the whole structure quite evenly?
Surely the molecules would continue to
repel each other even when they are
almost stationary at very low temperatures.
The very weak forces that allow non-polar
molecules such as methane to liquefy and
even freeze are due to “dispersion” forces.
The bonding is the result of temporary
fluctuating zones of positive and negative
charge on a molecule, which induce
distortions in the electrical fields of
neighbouring molecules. These tiny
variations of the electron clouds tend to
synchronize and the molecules experience
and the two molecules experience an
attractive force. Although they are
constantly changing the force is always
attractive and the molecules bond together.
As such forces seem to depend on the
number of electrons in the structure that
results from this attraction, helium would
have to have the lowest freezing point. It
does…one degree above absolute zero!
Key Words:
Atoms, molecules, repeating, crystal, crystal
structure, covalent, ionic, temperature,
weak attractive force, solids, liquids, gases,
gradual transition of bonding between
atoms, polar molecules, properties of
materials, charged plates, electronegative,
dipole, clouds of electrons, repel, condense,
electrostatic forces, dipole-dipole forces,
hydrogen bond, the properties of water,
strength, melting point and bond strength,
tetrahedron structure, electron pair, acetic
acid dimer, non polar, movement that is
temperature, synchronized variations, weak
intermolecular forces.
Credits
Produced and Written by
John Davis
Animation
Graeme Whittle
Consultants
John Willis
Dave Stanton
Chris Wiecek
Editor
Phil Sheppard
Sound Mix
Phillip McGuire
Executive Producer
John Davis
Copyright
and Orders:
CLASSROOM VIDEO (2004)
Classroom Video
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Questions
1. There are two different bonds in the biscuit structure. How can we distinguish between the two? Is there a third?
2. Describe the links between speed, velocity, temperature and heat.
3. What is the relationship between speed of movement and the breaking of bonds? Why does this relation occur in
substances?
4. Make a statement to clarify the difference between the two kinds of bonds that are present in the biscuit crystal.
5. One relationship that may exist between atoms is that there are no chemical bonds at all. Find some examples of
this relationship. There are two other extreme relationships that exist. Name them and give some details about both
kinds of bonds.
6. What is a polar molecule? Why aren’t all molecules polar?
7. Why would polar objects line up in an electric field?
8. List the events that have to occur for a dipole-dipole bond to form between two molecules. From this list make a
summary statement on dipole-dipole bonds. Why are they only one tenth the strength of a normal covalent bond?
9. Is there a difference between hydrogen bonding and dipole-dipole bonding? If so, what is the difference? Why make
the distinction?
10. List and then discuss the ways in which life, living and water are connected.
11. A water molecule contains three atoms. How are the atoms bonded together? Describe its shape.
12. How does the shape of a water molecule influence the distribution of charge on the molecule? How might this
distribution influence the way a water molecule interacts with its neighbours?
13. How do we know that water molecules hold each other more strongly than other molecules of the same numbers
of atoms? Give some examples.
14. Compare the melting points of water H2O, HF, and NH3. Why is there such a difference in properties between each
of these substances? Graphs are useful here in seeing the differences. You may find graphs the boiling points of
groups 4A and 6A hydrides vs. their molecular mass interesting.
15. How do acetic acid molecules form dimers? Why not call such a structure a “molecule” rather than a dimer?
Compare the boiling points of ethanol and methoxymethane, both C2H6O. Why are they so different?
16. How would the formation of the dimer influence the properties of acetic acid?
17. What is the difference in strength between hydrogen bonds and covalent bonds? Account for the difference.
18. Where in the DNA structure do hydrogen bonds play an important role? Re-examine your answer to Q 15 and
make a comment about the term “DNA molecule”
19 Try to explain the fact that dihydrogen oxide (3 atoms) and sulphur dioxide (3 atoms) have such different melting
points.
20. How can non-polar, electrically neutral molecules attract each other? Is it possible for the atoms of the inert gas
elements to attract each other? If so, what kinds of forces produce the bonds?
21. How do we know whether any bonding occurs at all with the inert gases? Examine a table of the boiling points of
the halogens and the inert gases. What trends do you see?
22. To compare the influence of dispersive forces and dipole-dipole forces on the boiling points of molecules we need
to examine two molecules that are the same size with the same number of electrons. Compare the shapes, charge
distribution and the boiling points of ethane and mono fluro methane. Explain their similarities and differences. How
are they different from methane?