Certain elements exist in more than one oxidation state and thus show variable valency. For example, transition metals show wide range of valence states, e.g., Fe2+, Fe3+; Cr2+, Cr3+ Co2+, CO3+, (more common), Co4+ and Co5+ (less common) etc. Similarly, normal metals, e.g., Pb2+ and Pb4+ ; Sn2+ and Sn4+ also show variable valencies. Certain non-metals are also found to show more than one valence states e.g., P3+ and P5+.

Variable valency may be due to different reasons and would be discussed over here accordingly.

In transition elements (elements in which d orbitals are in the process of completion) the variable valency is due to the partially filled ‘ d’ orbitals, state of hybridization and type of reactants (usually called ligands). Such elements can involve different number of electrons in compound formation and would show variable valencies. The stability of the particular valence state would also depend upon the nature of the reacting species and the number of d electrons present.

Iron (Fe) shows normally Fe2+ and Fe3+ valence states. Atomic number of Fe is 26 and its electronic configuration would be: 1s2 2s2 2p6 3s2 3p6 3d6 4s2.

Loss of two electrons from 4s orbitals would not leave behind an inert gas configuration. Therefore, more chances of electrons to be pulled out exist. Since `cl’ orbitals would be more stable when half-filled (with d5 configuration), so loss of one electron from 3d orbital along with two electrons from 4s orbitals would give a more stable state of the ion. Therefore, iron in Fe3+ state would be more stable. Let us show the valence shell having 3d orbitals and arrange the electrons in accordance with valence bond theory (See Chapter 4). In order to get next inert gas configuration more electrons would be required which are supplied by reacting species (Lewis bases or ligands) and the ions are thus stabilized. Mostly water will be acting as Lewis base and due to this reason most of the transition metal salts exist in stable state as hydrated species.

Elements of lanthanides and actinides show variable valencies due to the involvement of ‘f’ orbitals.

Some of the `1,’ block elements show variable valency due to the involvement of an ‘inert pair’ of electrons. The ‘inert pair’ of electrons are present in the ‘s’ orbitals and do not take part in chemical reactions under the prevailing conditions. Under such conditions only `p’ orbitals would take part in bond formation. However, if the ‘inert pair’ of electrons present in `s’ orbitals is actuated to take part in bond formation, an increase in valence state by two units will be observed. It is due to the `inert pair’ of electrons that the difference in valence states of such elements is always by two units.

Although the inert pair effect is quite marked for bivalent tin compounds but tetravalent tin is more stable of the two. Hence, bivalent tin is readily converted to tetravalent state, and thus the former ion is a reducing agent.

Sn(II) — 2e- Sn(IV)

The `13′ block elements which have vacant ‘cl’ orbitals available in their electronic configuration also show variable valency by utilizing these orbitals in addition to the corresponding ‘s’ and `ps’ orbitals. Thus the involvement of vacant ‘cr orbitals would be responsible for the variation in valence state of such elements. Let us consider the first two members of Vth group namely, nitrogen and phosphorus. There is no chance for the presence and involvement of `cr ‘orbitals in nitrogen as indicated by its electronic configuration (1s2 2s2 2p3) i.e., no ‘ar orbital is available in the 2nd shell.


This type of bond is exhibited by atoms which can either lose electrons to form positively charged ions (cations) or gain electrons to form negatively charged ions (anions). The atom which can lose electrons is said to be electropositive or basic and the atom capable of gaining one or more electrons is referred. as electronegative. The more electropositive atom has always low value of ionization potential and is thus capable of losing electrons with greater ease. The electrons lost by electropositive atoms are completely transferred to other atoms which show greater electro negativities. The bond formed by complete transfer of electrons from electropositive atom to more electronegative atom is called ionic or electrovalent bond The electropositive elements in energy terms should have higher energy states than those of electronegative elements. This ‘energy difference will be responsible for the flow of electrons from higher energy states to lower energy states. The two atoms are held together by electrostatic forces of attraction acting between such atoms.

M° + X°                 M+ : X (3.1)

The energy required to completely separate the ions from a diatomic molecule is given by the following expression:

Potential energy = Electrostatic energy + Van der Waals’ energy.

The electrostatic energy — q1 q2 e  R

where q1 and q2 are charges on atoms M and X and R is the internuclear separation.

The general tendency of various atoms to form molecules is to attain inert gas configuration, being the most stable. The atoms of the inert gases have outermost p orbitals completely filled. Such a configuration will not easily lose or gain any electron because very high ionization potentials (or electro negativities) will be required to remove an electron or gain any additional electron. Let us consider potassium and chlorine atoms which would combine to form potassium chloride molecule. Potassium atom (Atomic number of K = 19) has the electronic configuration 1s2 2s2 2p6 3s2 3p6 4 • si. Chlorine (Atomic number of Cl = 17) has electronic configuration 1s2 2s2 2p6 3s2 3p5. None of these has an inert gas structure. But they possess an incomplete shell of electrons and orbitals. The nearest inert gas to both is argon having electronic configuration 1s2 2s2 2/ 3s23p6. Thu:s the lois of one electron from potassium and gain of this lost electron by chlorine would leave both the atoms with argon configuration. During this process, IC+ and would be produced and ‘the electrostatic attraction between these oppositely charged Ions should be responsible for a stable ionic bond.

Let us take some more examples to elaborate the ionic bonding situation.

  • Formation of Sodium Chloride Molecule
  • Formation of Magnesium Chloride Molecule
  • Formation of Aluminium Fluoride

Transition metal ions which do not resort to inert gas configuration attain their stability through the formation of complex ions. The ions with odd number of s or p electrons are not known, but an odd number of d electrons is found in transition metal ions.

It should be noted that important forces between atoms or groups of atoms are electrostatic in nature.