bonding in atoms The variation of force between two atoms with inter atomic distance is shown schematically in fig 1. At large distance of separation, the significant force is attractive. At closer distance a repulsive force also becomes significant and increases rapidly with decreasing distance of separation. At one point the attractive and repulsive forces exactly balance. At this point the potential energy of the system of two atoms is minimum, and is called the `bond energy'.

According to strength, the bonds can be classified as primary and secondary bonds. Primary bonds have bond energies between 1 to 10 eV/bond. Ionic, covalent and metallic bonds are primary bonds. Amongst these the metallic bonds are the weakest. Secondary bonds have energies in the range of 0.01 to 0.5 eV/bond. Van der Waal bond and hydrogen bonds are examples of secondary bonds.

Ionic bonds : Ionic bonding forms between two oppositely charged ions. The ions are produced by transfer of electrons from an atom of one element to an atom of another element. We consider the example of sodium chloride. Sodium (Z = 11) with electronic configuration 1s²2s²2p⁶3s¹ has its first ionization potential 5.1 eV/atom. The outermost electron is removed by supplying this much of energy.

Na  -à Na⁺ + e^-

The released electron is accepted by Chlorine atom to fill its outermost vacant state.

Cl + e^- -à Cl^-

The electron affinity of Chlorine is 3.8 eV/atom. Therefore the electron transfer results with a net increase of potential energy E = 1.3 eV/pair. This increase is more than compensated by the energy decrease due to electrostatic attraction between oppositely charged ions.

Covalent bonding: It occurs by sharing two electrons between neighboring atoms. A good overlap of the orbitals is necessary, so that the shared electrons are close to both the nuclei. It results with a net decrease of potential energy. This occurs when there exist vacant electron states in the outermost orbital of the bonding atoms. Depending upon how much the number of outer electrons differ from close shell configuration an atom is limited in the number of covalent bonds it can make. For example in carbon, which has four outermost electrons, can be involved in four bonds at tetrahedral angles 109.5^o, which is seen in crystalline diamond and innumerable other organic compounds.

Overlapping orbitals are directionally oriented, and not spherically symmetric. This gives directionality of covalent bonds. Sharing of electrons and formation of covalent bonds readily occur between atoms which have unfilled p orbitals. The p orbitals are directional in nature and therefore permit efficient overlapping of orbitals in the direction of maximum probability density. The overlapping can either be end-to- end (s bond) or lateral (p bond), which are shown in the figure 2.

Metallic bonding : Metallic characteristics are shown by elements which are left to the fourth column in the periodic table. In these elements the sharing of electrons between neighbouring atoms is not possible, as there are not enough electrons to produce inert gas configuration around each atom. The electrons form a common pool, called the free electron gas, in which all atoms have contributed their outer electrons. These electrons have freedom to move anywhere within the crystal and act like all pervasive, mobile glue holding the positive core together. This freedom makes the metallic bond non directional.

Secondary bonding : Many molecules posses permanent dipole moment. This arises because in a molecule centres of positive charge and negative charge do not coincide. The best example is water molecule. Two water molecules form weak bond by orienting positive and negative charged ends. This bond is directional.

The bonding between atoms of inert gases, when they are in solid crystalline form, at very low temperature, is called Van der Waal bonds. The momentary fluctuations in charge distribution around atom result in weak fluctuating dipole moment. The electric field of this imbalance can induce dipole moment in the neighbouring atom, in such a way as to attract it. This dipole induced dipole attraction is non directional in nature. The Van der Waal bond energies are very small.