Friday, 12 June 2015

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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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