CLASSICAL TREATMENT OF BONDING
STRUCTURE OF ATOM
The
views on atomic structure which are accepted today have developed from
the classical Rutherford-Bohr theory. According to his theory, the atom
is made of a central positively charged nucleus containing positively
charged particles called Protons, and neutral particles called Neutrons,
both having unit mass. The nucleus is surrounded by negatively charged
particles called Electrons which carry one unit negative charge and
negligible weight.
The
electrons are said to revolve around the nucleus in fixed orbits or
energy levels. While the electron moves in such a level, it possesses a
definite quantity of energy and it neither emits nor absorbs energy. The
electrons are arranged in the orbits so that the maximum number of
electrons in the various orbits staring from the one nearest the nucleus
is 2, 8, 18, 32, 18, 8.The outermost orbit of the electrons in
different atoms [except those of inert gases], is incomplete and the
electrons in it are known as the Valence Electrons.
WHY ATOMS COMBINE TO FORM MOLECULES?
The
classical concept of formation of molecules, proposed by Lewis and
Kossel, is based upon the electronic structure of atoms. The atoms of
inert gases have either 2 or 8 electrons in the outermost orbit. These
gases do not enter into chemical combination and, therefore, are assumed
to have complete or stable orbits. The atoms of all other elements have
incomplete outermost orbits and tend to complete them by chemical
combination with other atoms. G. N. Lewis proposed that it is the urge
of atoms to complete their outermost orbits of electrons as in the inert
gases, which is responsible for chemical combination. In other words,
chemical combination between two atoms results from the redistribution
of electrons between them so that both the atoms complete their
outermost orbits or acquire stable electronic configuration possessed by
the inert gases.
TYPES OF BONDS
There are three basic ways in which chemical combination occurs:
- Ionic or electrovalent bond
- Covalent bond
- Coordinate bond
Ionic or Electrovalent Bond
Ionic
or electrovalent bonds are formed by transfer of valence electrons from
one atom to another. This type of bond unites two atoms one of which
has excess electrons than the stable number [2 or 8], and the other is
short of electrons. Sodium chloride is a typical compound formed in this
way.
Here
the sodium atom [2, 8, 1] transfers its excess electron to chlorine
atom [2, 8, 7], and thus both attain a stable inert gas type electronic
configuration. Sodium acquires the electronic configuration of Neon [2,
8] and becomes positively charged. Chlorine acquires the electronic
configuration of Argon [2, 8, 8], and becomes negatively charged. These
oppositely charged ions held together by electrostatic force of
attraction. This type of bond is commonly found in inorganic compounds.
Electrovalent bond |
Ionic
or electrovalent compounds are non-volatile, soluble in water and
possess high melting points. Their aqueous solutions conduct electric
current.
Covalent Bond
Covalent
bonds are formed by mutual sharing of electrons. This type of bond
unites two atoms, both of which are short of electrons. The two atoms
contribute one electron each and then share the resulting pair of
electrons. Hydrogen is the simplest compound formed in this way.
Covalent bond |
Here
the two electrons are shared, and give to each hydrogen atom the
configuration of helium. This type of bond is termed covalent bond and
is indicated by a thin line. Covalent bonds are commonly found in
organic compounds.
Covalent
compounds are volatile, generally insoluble in water but soluble in
organic solvents. They possess low melting and boiling points. Their
solutions do not conduct electric current.
Coordinate Bond
Coordinate
bond is also formed by mutual sharing of electrons but in this case the
two electrons that are shared come from the same atom. A coordinate
bond unites two atoms, one of which has a spare pair of electrons and
the other is short of a pair of electrons. The first atom (donor atom)
contributes one pair (lone pair) of electrons and the second atom
(acceptor atom) accepts it. After the formation of the bond, the lone
pair of electrons is held in common. The coordinate bond is represented
by an arrow, pointing away from the donor atom. An excellent
illustration of the coordinate bond is found in the boron hydride-
ammonia complex.
Coordinate bond |
VALENCE OF CARBON
The atomic number of carbon is 6 and its atomic weight is 12. Its electronic configuration is shown in the figure below:
Electronic configuration of carbon |
It
has 4 electrons in the last orbit and tends to gain 4 more electrons by
forming 4 covalent bonds with other hydrogen atoms. Thus, the
structural formula of the simplest hydrocarbon methane (CH4) can be written as:
Metahne |
Similarly
in all organic molecules carbon atom is tetravalent. That is, it has a
valence of 4. According to Lee Bel and van’t Hoff the four valencies of
carbon do not lie in one plane. They are directed towards the corners of
a regular tetrahedron so that the angle between any two valencies is
109°28’.
Four valencies of carbon |
Carbon-Carbon Single Bond
Carbon
atom has the wonderful property of uniting with other carbon atoms
through covalent bonds. This serves to construct the carbon structure of
organic molecules. Thus the molecules of hydrocarbons, ethane and
propane contain two and three carbon atoms respectively linked by
covalent bonds.
Single bond |
Carbon-Carbon Double bond
In
some compounds, two of the valencies of a carbon atom may be satisfied
by union with the two valencies of another carbon atom. Thus in ethylene
the two carbons are united by two covalent bonds. Such as union
involving two covalent bonds between adjacent carbons is called double
bond.
Double bond |
Carbon-carbon triple bond
Sometimes
two adjacent carbons are linked together by three covalent bonds. Such a
union involving three covalent bonds between adjacent carbon atoms is
called triple bond. Thus acetylene molecule is represented as:
Triple bond |
Cyclic structures
We
have given some examples of substances where the molecules consist of
carbon atoms joined together in chains that are free at both ends. There
are numerous compounds known where carbon atoms join to form closed
rings. These are called cyclic compounds or ring compounds.
Cyclic structures |
BOND LENGTHS
When
two atoms are bonded by a covalent bond, the distance between the
centres of the two nuclei is called bond length. Bond lengths are
measured by X-ray crystallography and by microwave spectroscopy. The
unit of bond length is Angstrom (1A°=10-8 cm). For most bonds the values are 1 to 2 A°. Some typical bond lengths are:
BOND
|
BOND LENGTH (A°)
|
C-H
|
1.09
|
C-C
|
1.54
|
C=C
|
1.34
|
C≡C
|
1.20
|
O-H
|
0.96
|
BOND ENERGIES
Bond
energy or bond strength is defined as the amount of energy required to
break a bond in a molecule. Bond energies depend upon the type of bond
as well as the structural environment in which the bond is situated.
They are determined by quantitative measurements of heats of chemical
reaction (calorimetry) and by spectroscopic methods. The unit of bond
energy is kcal/mole. Some typical bond energies are:
BOND
|
BOND ENERGY (kcal/mole)
|
C-H
|
99
|
C-C
|
83
|
C=C
|
146
|
C≡C
|
200
|
O-H
|
111
|
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