This theory started from the idea that, since the atom of any one of the noble gases has a complete outer shell of eight electrons this is associated in some way that with stability and inactivity. Hence it was suggested that reaction between the atoms of other elements involved a rearrangement of their electrons to give these stable octets. Such rearrangement might involve (a) either the giving or receiving of electrons, or (b) the sharing of electrons.
   
The giving or receiving of electrons : Electrovalency

The electronic configuration of the sodium atom is 2.8.1 that of chlorine 2.8.7. Hence transfer of one electron (the outermost) from an atom of sodium to an atom of chlorine gives sodium the configuration 2.8, and chlorine 2.8.8. These are the configurations the inert gases neon and argon respectively:

By this transfer, a positive sodium ion (a cation) and a negative chloride ion (an anion) are formed; and it is easy to see that electrostatic attraction between these two ions will hold them together. In fact, a crystal of solid sodium chloride is made up of sodium and chloride ions held together in this way. This kind of binding, involving the giving or receiving of electrons is called electrovalency (sodium and chloride have an electrovalency of one) and the bond between the sodium ion, Na⁺, and the chloride ion, Cl^-, is often said to be ionic. When sodium chloride, or any ionic crystal, dissolves in water, the crystal breaks up and dissolves, the ions are separated, i.e. there is dissociation, and each ion becomes surrounded by a shell of water molecules (a hydrated ion). If an ionic compound is melted, then the ions become mobile, and hence the melt can conduct electricity like a solution. The forces between ions are strong ones and therefore ionic crystals are not easily vaporized.

The formation of an ionic bond is essentially similar for the other alkali metals; but as we go from lithium to caesium, the increasing number of electron shells between the nucleus and the outermost electron makes removal of the latter much easier; and since the chemistry of the alkali metals is essentially based on their tendency to form ions, it is understandable that the reactivity of these metals increases from lithium to caesium, i.e. as the ease with which ions are formed increases.)

Like the alkali metals, the alkaline earth metals form ionic compounds, e.g.

Here a calcium atom loses two electrons to give the divalent calcium cation Ca₂₊; and although we never observe the formation of monovalent calcium, Ca⁺ under ordinary conditions, yet it is easy to see that the removal of a second electron from the monovalent calcium ion, Ca⁺, to give the divalent ion, Ca²⁺, will be more difficult than the removal of the first electron. Hence the reactivity of the alkaline earth metals is generally less than that of the alkali metals. Similar arguments apply to the formation of divalent anions, for example, the oxide ion, O₂^-, and sulphide ion, S₂^- are generally less readily formed than the halide ions, F^-, Cl^-, Br^-.
   
The sharing of electrons: Covalency

This is well illustrated by the formation of methane, CH4, from one carbon atom and four hydrogen atoms

Here the carbon atom ‘achieves’ the octet of neon, i.e. by the sharing of its four electrons with those from the hydrogen atoms, similarly each hydrogen atom achieves the helium structure by the same sharing. Each electron pair constitutes a covalent bond, and these are directed (in the case of carbon methane, towards the corners of a tetrahedron). The shorthand representation of a covalent bond between the two atoms A and B, i.e. A; B is A – B. in covalent bond between the atoms A and B, i.e.  , A – B. In covalent compounds, no ions are produced; a molecule is formed as an entity, and in a covalent compound as we see it, be it solid, liquid or gas, these molecules, made up internally of covalent-bound atoms, are in turn held together by intermolecular forces which are usually far weaker than covalent or ionic bonds.

Since these intermolecular forces are weak, covalent compounds are usually volatile, and are often gases or liquids because the molecules are easily driven apart on heating. Only in a few cases do covalent bonds extend throughout a crystal, e.g. in diamond or certain form of boron nitride, in these cases, we obtain a very hard non-volatile solid of very high melting point. Some other examples of covalent bonding are given by the chlorine molecule, ammonia and ethylene.

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