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We know, then, what makes atoms stick together, but what limits the
proportions in which they can combine and what factors control molecular
shape? The simplest and most basic factor arises from the relative sizes
of the different atoms. The concept of atomic size is, however, somewhat
fuzzy: the size of an atom is controlled by the extent of its electron
charge cloud, and the density of electrons around an atom does not suddenly
reduce to zero; it gradually decays. Nevertheless it has proved possible
to assign approximate radii to atoms (as shown on the left). And purely
geometrical
factors will limit the ways in which we can pack atoms together. These factors
become most obvious when we consider packing in crystals; thus in crystal
structures containing, for example, the large caesium ion, there are often
eight other atoms surrounding each caesium; whereas the smaller lithium
ion has room for only four. We return to these considerations when we discuss
the structure of crystals.
The periodic table of the elements showing typical chemical radii for each element.
The next point is that atoms have well defined combining powers (the
chemical concept of VALENCE) which arises from specific features
of their electronic structure; in particular the number of electrons in
the outermost shell relative to the total number of electrons that can
be present in that shell. Indeed, the concept of valence springs from one
of the oldest and most powerful ideas of theoretical chemistry - the 'electron
pair' bond. Chemical bonds (like those between the two hydrogen atoms in
H2) often have a pair of associated electrons.
And when there are different numbers (four or six) we often refer to them
as 'double' or 'triple' bonds (for example,oxygen has a double and nitrogen
a triple bond). Chemical bonding can be thought of in terms of the drive
by atoms to achieve stable configurations in which all their outermost
orbitals are full, which can be effected by teaming up with other atoms
with whichthey 'share' electrons in pair bonds. Thus the atom carbon has
4 electrons in its outermost shell which has 'room' for a total of 8; carbon
has a valence of 4 (as in the compound methane, CH4). Oxygen
has 6 electrons and again room for a total of 8; hence it has a valence
of 2 as in water, H2O. The value of this approach is shown by
the fact that atoms in which the outermost shells of electrons are full
are chemically inactive. These are the famous "rare gases" (many
of which were isolated by Sir William Ramsay working at University College
London at the beginning of the twentieth century).
Thus helium and neon have no chemistry; the atoms bond neither with themselves
nor with any other species. Argon and krypton show very limited chemistry
because it turns out that theoutermost shell of electrons can, with difficulty,
be expanded. In one of the remarkable developments of chemistry in the 1960s,
it was found by Neil Bartlett and coworkers that Xenon has a quite extensive
chemistry as it is relatively easy to expand its outermost shell. Some of
the compounds it forms are shown on the left.
Compounds of Xenon: XeF2, XeF4 and XeF6
The more traditional chemical
concepts of valence and pair bonds are fully compatible with the approach
to chemical bonding in terms of electron density redistribution which we
presented earlier. The former allows us to understand why atoms form bonds;
the latter guides us as to the constraints on the molecular structures that
can form. A key feature of chemical bonding to which we now turn concerns
the fact that it can have specific directional requirements. Thus the
carbon atom in the majority of its compounds bonds to four atoms which are close
to the corners of a tetrahedron; although in a significant proportion of
cases it adopts a rather different bonding pattern with three surrounding
atoms close to the corners of an equilateral triangle and insome gases
a linear arrangement is adopted. These specific geometricalrequirements
(discussed in greater detail in the Appendix) follow naturally from the
criteria that the resulting assembly of atoms should have the lowest possible
energy. More specifically, they can be understood in terms of the different
shapes of the atomic electron density charge clouds and by the ways in which
these can interact. Chemical bonding occurs where the atomic charge clouds
interact and overlap with each other. And different shapes of the charge
clouds.
Atomic orbitals: 1s and 2s (top row), 2p (middle row) and 3d orbials (lower row)
Butane, benzene, ethylene and acetlylene
Just as the differently shaped
atomic charge clouds are labelled s, p and d, so the different shaped molecular
charge clouds (often referred to as molecular orbitals) are labelled sigma,
pi and d. Examples of molecules containing these different types of bonds
are shown above. These different types of bonds have different properties,
sigma bonds concentrate electron density directly between the nuclei where
the interactions between electron and nuclei are strongest, unlike pi bonds
where the maximum in the density is away from the internuclear axis.
Sigma
bonds have therefore to be more stable and pi bonded molecules morereactive.
There are subtle geometrical differences as well. We can rotate about sigma
bonds but, as shown below, rotation about a pi bond destroys the
orbital overlap. Pi bonded molecules like C2H4 are
therefore more rigid than sigma bonded frameworks.
Double bonds (lower) cannot rotate without disrupting the
interaction of pi orbitals
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