INTRODUCTION TO CHEMICAL BONDING
Chemical reactions involve
the making and breaking of bonds. It is essential that we know what bonds are
before we can understand any chemical reaction. To understand bonds, we will
first describe several of their properties. The bond strength tells us how hard it is to break a
bond. Bond lengths give
us valuable structural information about the positions of the atomic
nuclei. Bond dipoles inform
us about the electron distribution around the two bonded atoms. From bond
dipoles we may derive electronegativity data
useful for predicting the bond dipoles of bonds that may have never been made
before.
From these properties of
bonds we will see that there are two fundamental types of bonds--covalent and
ionic. Covalent bonding represents a situation of about equal sharing of the
electrons between nuclei in the bond. Covalent bonds are formed between atoms
of approximately equal electronegativity. Because each atom has near equal pull
for the electrons in the bond, the electrons are not completely transferred
from one atom to another. When the difference in electronegativity between the
two atoms in a bond is large, the more electronegative atom can strip an
electron off of the less electronegative one to form a negatively charged anion
and a positively charged cation. The two ions are held together in an ionic
bond because the oppositely charged ions attract each other as described by
Coulomb's Law.
Ionic compounds, when in
the solid state, can be described as ionic lattices whose shapes are dictated
by the need to place oppositely charged ions close to each other and similarly
charged ions as far apart as possible. Though there is some structural
diversity in ionic compounds, covalent compounds present us with a world of
structural possibilities. From simple linear molecules like H2 to
complex chains of atoms like butane (CH3CH2CH2CH3),
covalent molecules can take on many shapes. To help decide which shape a
polyatomic molecule might prefer we will use Valence Shell Electron Pair
Repulsion theory (VSEPR). VSEPR states that electrons like to stay as far away
from one another as possible to provide the lowest energy (i.e. most stable)
structure for any bonding arrangement. In this way, VSEPR is a powerful tool
for predicting the geometries of covalent molecules.
The development of quantum
mechanics in the 1920's and 1930's has revolutionized our understanding of the
chemical bond. It has allowed chemists to advance from the simple picture that
covalent and ionic bonding affords to a more complex model based on molecular
orbital theory. Molecular orbital theory postulates the existence of a set of
molecular orbitals, analogous to atomic orbitals, which are produced by the
combination of atomic orbitals on the bonded atoms. From these molecular
orbitals we can predict the electron distribution in a bond about the atoms.
Molecular orbital theory provides a valuable theoretical complement to the
traditional conceptions of ionic and covalent bonding with which we will start
our analysis of the chemical bond.
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