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(and any others!) if you find them too alien. The general principles behind the
chapter—why molecules have the structures they do—are obviously so important that we cannot
omit this essential material but you should try to grasp the principles without worrying too much
about the equations. The ideas of atomic orbitals overlapping to form bonds, the molecular orbitals
that result, and the shapes that these orbitals impose on organic molecules are at least as central for
biochemistry as they are for organic chemistry. Please do not be discouraged but enjoy the challenge.
Introduction
You may recognize the model above as DNA, the molecule that carries the genetic information for all
life on earth. It is the exact structure of this compound that determines precisely what a living thing
4Structure of molecules
Connections
Building on:
• How organic structures are drawn ch2
• Evidence used to determine organic
structure ch3
Arriving at:
• How we know that electrons have
different energies
• How electrons fit into atomic orbitals
• How atomic orbitals combine to make
molecular orbitals
• Why organic molecules have linear,
planar, or tetrahedral structures
• Connection between shape and
electronic structure
• A true system of molecular orbital
energies for simple molecules
• Why such rigour is not possible for
typical organic molecules
• Predicting the locations of lone pairs
and empty orbitals
• Interaction between theory and
experiment
Looking forward to:
• Mechanisms depend on molecular
orbitals ch5
• Conjugation ch7
• 1H NMR involves molecular orbitals
ch11
• Reactivity derives from energies of
molecular orbitals ch3
is—be it man or woman, frog, or tree—and even more subtle characteristics such as what colour eyes
or hair people have.
What about this model?
You may also have recognized this molecule as
buckminsterfullerene, a form of carbon that received
enormous interest in the 1980s and 1990s. The
question is, how did you recognize these two com-
pounds? You recognized their shapes. All molecules
are simply groups of atoms held together by electrons
to give a definite three-dimensional shape. What
exactly a compound might be is determined not only
by the atoms it contains, but also by the arrange-
ment of these atoms in space—the shape of the
molecule. Both graphite and buckminsterfullerene
are composed of carbon atoms only and yet their
properties, both chemical and physical, are completely
different.
There are many methods available to chemists
and physicists to find out the shapes of molecules.
One of the most recent techniques is called
Scanning Tunnelling Microscopy (STM), which is
the closest we can get to actually ‘seeing’ the atoms
themselves.
Most techniques, for example, X-ray or elect-
ron diffraction, reveal the shapes of molecules indi-
rectly.
In Chapter 3 you met some of the spec-
troscopic methods frequently used by
organic chemists to determine the shape of
molecules. Spectroscopy would reveal the
structure of methane, for example, as
tetradral—the carbonatom in the centre of
a regular tetra-hedron with the hydrogen
atoms at the corners. In this chapter we are
going to discuss why compounds adopt the
shapes that they do.
This tetrahedral structure seems to be
very important—other molecules, both
organic and inorganic, are made up of many
tetrahedral units. What is the origin of this
tetrahedral structure? It could simply arise
from four pairs of electrons repelling each
other to get as far as possible from each
other. That would give a tetrahedron.
82 4 . Structure of molecules
graphite
H
HH
H H
C
H
H
H
the H atoms form
a tetrahedron
C
methane is
tetrahedralmethane is tetrahedral
H
H
H
H
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The dark brown blobs in this STM
picture recorded at a temperature
of 4 K are individual oxygen atoms
adsorbed on a silver surface. The
light blobs are individual ethylene
(ethene) molecules. Ethylene will
only adsorb on silver if adjacent to
an oxygen atom. This is an atomic
scale view of a very important
industrial process—the
production of ethylene oxide from
ethylene and oxygene using a
silver catalyst.
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The picture on the right is an X-ray
structure of a catenane—a
molecule consisting of two
interlocking rings joined like two
links in a chain. The key to the
synthesis depends on the self-
stacking of the planar structures
prior to ring closure.
This simple method of deducing the structure of molecules is called Valence Shell Electron Pair
Repulsion Theory (VSEPRT). It says that all electron pairs, both bonding and nonbonding, in the
outer or valence shell of an atom repel each other. This simple approach predicts (more or less) the
correct structures for methane, ammonia, and water with four electron pairs arranged tetrahedrally
in each case.
VSEPRT seems to work for simple structures but surely there must be more to it than this? Indeed
there is. If we really want to understand why molecules adopt the shapes they do, we must look at the
atoms that make up the molecules and how they combine. By the end of this chapter, you should be
able to predict or at least understand the shapes of simple molecules. For example, why are the bond
angles in ammonia 107°, while in hydrides of the other elements in the same group as nitrogen, PH3,
AsH3, and SbH3, they are all around 90°? Simple VSEPRT would suggest tetrahedral arrangements
for each.
Atomic structure
You know already what makes up an atom—protons, neutrons, and electrons. The protons and
neutrons make up the central core of an atom—the nucleus—while the electrons form some
sort of cloud around it. As chemists, we are concerned with the electrons in atoms and more impor-
tantly with the electrons in molecules: chemists need to know how many electrons there are in a sys-
tem, where they are, and what energy they have. Before we can understand the behaviour of electrons
in molecules, we need to look closely at the electronic structure of an atom. Evidence first, theory
later.
Atomic emission spectra
Many towns and streets are lit at night by sodium vapour lamps. You will be familiar with their warm
yellow-orange glow but have you ever wondered what makes this light orange and not white? The
normal light bulbs you use at home have a tungsten filament that is heated white hot. You know that
this white light could be split by a prism to reveal the whole spectrum of visible light and that each
of the different colours has a different frequency that corresponds to a distinct energy. But where
does the orange street light come from? If we put a coloured filter in front of our white light, it
would absorb some colours of the spectrum and let other colours through. We could make orange
light this way but that is not how the street lights work—they actually generate orange light and
orange light only. Inside these lights is sodium metal. When the light is switched on, the sodium
metal is slowly vaporized and, as an electric current is passed through the sodium vapour, an orange
light is emitted. This is the same colour as the light you get when you do a flame test using a sodium
compound.
The point is that only one colour light comes from a sodium lamp and this must have one specif-
ic frequency and therefore one energy. It doesn’t matter what energy source is used to generate the
light, whether it be electricity or a Bunsen burner flame; in each case light of one specific energy is
given out. Looking at the orange sodium light through a prism, we see a series of very sharp lines
with two particularly bright orange lines at around 600 nm. Other elements produce similar spec-
tra—indeed two elements, rubidium and cesium, were discovered by Robert Bunsen after studying
such spectra. They are actually named after the presence of a pair of bright coloured lines in their
spectra—cesium from the Latin caesius meaning bluish grey and rubidium from the Latin rubidus
meaning red. Even hydrogen can be made to produce an atomic spectrum and, since a hydrogen
atom is the simplest atom of all, we shall look at