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 electro
negativity 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 electro negativity. 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 electro negativity
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.
REFERENCE
SparkNotes Editors. (n.d.). SparkNote on
Introduction to Chemical Bonding. Retrieved February 25, 2017, from
http://www.sparknotes.com/chemistry/bonding/intro/
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