Introduction:

Whenever elements combine, compounds are formed. The forces which hold atoms altogether in compounds are termed as chemical bonds. The combination of chemical elements to form a compound is a chemical reaction. A few elements are very reactive and exist in nature only in combined states, example: sodium in sodium chloride (that is, common salt) and calcium in calcium trioxocarbonate (that is, marble). Some of the elements are relatively imreactive and exit hardly ever as free elements. They are termed as noble or rare elements, example: no compound of helium, neon and argon exists. Most of the elements have intermediate reactivity and exist as free elements and also in chemical compounds example: carbon takes place as graphite and diamond and also in petroleum and lots of organic compounds. There are a few non metallic elements which exist only as diatomic molecules in the Free State. Such elements as well take place in combined states.

Li (2, 1), Na (2, 8, 1), K (2, 8, 8, 1) all encompass similar configuration having one electron each in their outermost shell. They are metals. F (2, 7) and C1 (2, 8, 7) all require one electron to complete their outermost electron shell. They are non-metals. The noble or inert elements He (2), Ne (2, 8), Ar (2, 8, 8) all contain complete shell arrangement of electrons. The electron arrangement in the stable ions of metals and non-metals as well illustrate that the complete shell of electrons is a stable configuration example: Na (2, 8), F^- (2, 8) and O^2- (2, 8). 

In the chemical bonding thus elements tend to get the inert or noble gas configuration. There are numerous kinds of bonds however; we will mainly focus on the three kinds:     

• Electrovalent
• Co-valent
• Co-ordinate covalent.

Electrovalent (Ionic) Bonding:

If atoms act together only for bond formation, then the outermost portion of the atoms are in contact and as a result only the outer electrons in the outermost shell (that is, valence electrons) are involved. The outermost shell electron arrangement is thus very significant in finding out the kind of bond. Electrovalent bonding comprises electron transfer from the valence shell of one atom to the valence shell of the other.

One atom loses electrons to become the positively charged and the other gains electrons to become negatively charged. The negatively and positively charged ions are termed as anions and cations correspondingly. The ionic bond outcomes from the attraction between these oppositely charged ions. This kind of bonding is generally between the metals and non-metals.

Example: Li and F

Li → 2, 1

F → 2, 7

Li → Li⁺ (2) + e

F, +e → F^- (2, 8)

Li + F^- → Li⁺ 

The formula of LiF illustrated above is the electron dot formula (that is, Lewis structure). The brackets around the fluorine are intended to describe that all the eight electrons are the exclusive property of the fluoride ion (F).

The gain and loss of electrons result in inert gas configurations for the anion and cation. Apart from for helium, He (2) the inert gas configuration corresponds to eight electrons in the outer shell. The electronic theory of valency as postulated via Kossel and Lewis in the year 1716 was prompted via the remarkable stability of the rare gas elements. This stability is related by the presence in the atoms of a group of eight electrons in the outer shell. This totality appears to be the source of stability in rare gases.

The tendency for atoms to encompass eight electrons in their outermost shell (is illustrated by the octet rule). The octet rule defines that:

- Atoms tend to lose or gain electrons until there are eight (8) electrons in their valence shell.

It will be noted that the rule doesn't always hold in cases like these, other stable configurations elucidate ion stability. The number of bonds to a specific atom based on the number of electrons gained or lost to achieve stable configurations for instance.

Ca (2, 8, 8, 2) → Ca²⁺ (2, 8, 8) + 2e

2Cl (2, 8, 7) + 2e → 2 Cl^- (2, 8, 8)

Ca²⁺ + 2Cl^- → Ca²⁺

Ca²⁺ + 2Cl^- → CaCl₂

Structure of electrovalent compounds:

This is wrong to speak of a molecule of an ionic compound such as NaCl or LiF. Ionic compounds are generally solid comprising of regular arrangement of equivalent number of negative and positive charges. For Lif and NaCl there will be equivalent numbers of cations and anions.

This regular arrangement of cations and anions in the solid crystal is termed as the lattice. The structure of sodium chloride is described in the figure given below:

Fig: Lattice structure in sodium chloride

Properties of electrovalent compounds:

The general properties of electrovalent compounds are as follows:

1) Generally crystalline solids.

2) Usually soluble in water however in general insoluble in organic solvents such as ether or kerosene.

3) Generally high melting point compounds.

4) Good conductors of electricity whenever molten or in aqueous solution however not if solid.

This is simpler to describe the binding forces in the union between sodium ion and chloride ion, in the formation of sodium chloride as their opposite ionic charges attract one other. It's though hard to understand the way of bondage between the non-ionic or non polar atoms.

Covalent Bonding:

This is simpler to describe the binding forces in the union between sodium ion and chloride ion, in the formation of sodium chloride as their opposite ionic charges attract one other. It's though hard to understand the mode of bonding between the non-ionic or non polar atoms.

Lewis in the year 1916 came up with a tenable description, proposing that non ionic molecular compounds occur from the sharing of electrons among atoms, resultant in a form of bonding which was termed as the covalent bond. This kind of bonding comprises sharing of electron pairs instead of complete transfer. The shared electron pair is/are contributed via the two atoms involved in the bonding. The binding force yields from the attraction of the shared electron pairs via the nuclei of the atoms comprised in the bonding. This kind of bonding is between non metals. Observe the given illustrations.

Fig: Formation of covalent molecules

An electron pair comprises of a bond and two pairs comprise a double bond as in the oxygen molecule. The number of electron pairs shared based on the number of electrons each atom should share to accomplish an inert gas configuration. The Covalent compounds form molecules and based on the intermolecular forces between the molecules they might be gas (O₂, H₂, HCl) or liquids (Br₂, H₂O) or low boiling solids (like candle wax).

Properties of covalent compounds:

The general properties of the covalent compounds are as follows:

1) Generally liquids or gases.

2) Usually low melting whenever they are solids.

3) Usually not very soluble in water.

4) Usually soluble in the organic solvents.

5) Non conductors of heat and electricity except they dissolve to form ions example: HCl

Co-ordinate (Dative) Covalent Bonding:

In a covalent bond, the shared electrons are donated and controlled via both atoms which are comprised in the bonding. This is not the case having co-ordinate covalent bond. One atom contributes the electron pairs however both atoms control the donated pair/s.

Once a coordinate covalent bond is made it is not different from the ordinary covalent bond. The electron Pair is attracted through both nuclei of the bonded atoms. For this kind of bond to be formed, one atom should encompass a lone pair/s of electrons. (That is, Lone pairs are electron pairs which are not employed in bonding to other atoms) The other atom should encompass a vacancy in its valence shell to accept the lone pair. The bond formation as well yields in inert gas configuration for both the atoms.

Let us examine the formation of the hydronium ion and the ammonium ion, H₃O⁺ and NH₄⁺

Fig: formation of hydronium ion and ammonium ion

Co-ordinate covalent bonding is very common by means of metal complexes. The molecules donating the electron pairs are termed as ligands and the metal ion the central atom.

Example: The formation of ammonia complex with copper ions in solution, 'The octet rule'.

Fig: formation of ammonia complex with copper ions

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