Chemical Bonding

CHEMICAL BONDAGE
4.1 Introduction
The union of two or more atoms involving redistribution of electrons either by transfer or sharing between themselves, so that all of them acquire the stable noble gas configuration of minimum energy is known as chemical bonding.
The attraction between atoms within a molecule is called a chemical bond.
Electronic theory of valency :
The electronic theory leads to the following types of combination between atoms.

1) Ionic or Electrovalent Bond : Involving transfer of one or more electrons from one atom to another. 2) Covalent Bond : Mutual sharing of electrons between two atoms.

Octet rule : Formation of many ionic and covalent compounds can be explained on the basis of the octet rule which states that "Atoms of elements tend to acquire an octet of electrons (i.e. eight electrons) in their outermost shell through combination with other atoms." In doing so they tend to achieve electronic configuration of the nearest inert gas. However there are some limitations with the octet rule as it fails to explain incomplete and expanded octets.

The Ionic Bond

An ionic bond is formed by transfer of electrons from one atom to another atom in the molecule.
Consider the formation of potassium chloride. An electron is transferred from potassium atom to chlorine atom. The resulting K⁺ and Cl- ions possessing configuration of argon (2,8,8) combine to form an electrovalent compound as shown below :



Figures in the parenthesis indicate electron distribution in different orbitals (Electronic configuration) in ground state.

Another example :

The total number of electrons lost by Lithium must equal the total number of electrons gained by Bromine. Thus the number of Lithium ions produced is the same as the number of bromide ions produced. These ions attract each other to form crystals as shown below.


When positive ions (cations) and negative ions (anions) come closer to each other, they are held by electrostatic forces of attraction.
The electrostatic attraction between the oppositely charged ions is called the ionic bond.

Ion types are summarized as shown in Table 5

Table 5 - Types of Ions
Ion type	Electronic configuration of outer shell	Examples
Noble gas configuration	ns2 np6	Na+, Ca2+ Ai3+ Cl-, O2 -
Pseudo noble gas configuration	ns2 np6 nd10	Ag+, Zn2+
Inert pair	(n-1)s2(n-1)p6(n-1)d10 ns2	Sn2+, Bi3+

Ionic Radius :

The force of attraction or repulsion, F, between two charges q_C (charge on cation) and q_A (charge on anion) separated by a distance d (sum of ionic radii of two ions) is given by Coulomb’s law as
Where D = dielectric constant.
d = radius of cation + radius of anion as shown in figure.

Figure 7 Ionic radii

Table 6 Ionic Radii (in 0 A)
CATIONS						ANIONS
Group I A		Group II A		Group III A		Group VI A		Group VII A
E	R	E	R	E	R	E	R	E	R
Li+	0.60	Be2-	0.31			O2-	1.4	F-	1.36
Na-	0.95	Mg2+	0.65	A13+	0.50	S2-	1.84	Cl-	1.81
K+	1.33	Ca2+	0.99	Ga3-	0.62	Se2-	1.98	Br-	1.95
Rb+	1.48	Sr2+	1.13	In3+	0.81	Te2-	2.21	r-	2.16
Cs+	1.69	Ba2-	1.35	Ti3+	0.95

A positive ion is always smaller than the neutral atom ( Na is 1.57⁰A and Na⁺ is 0.95⁰A ) and a negative ion is always bigger than the neutral atom ( F is 0.72⁰A and F- is 1.36⁰A )
Within a period ( as shown in the Table ) isoelectronic positive ions ( having same number of electrons ) show a decrease in ionic radium from left to right due to increase in nuclear charge ( Na⁺ > Mg²+ > Al³+ ). Similarly for isoelectronic negative ions, ionic radium decreases from left to right ( O^{2 -} > F^- ).
Within a group of periodic table, similarly charged ions increase in size from top to bottom
( Li⁺ < Na⁺ < K⁺ or F ^- < Cl^- < Br^- < I ^- ) due to more electron shells.
The above table can be used to compare the strengths of ionic bonds in two compound in same medium.
Example : According to Coulomb’s law

The strength of the CaO bond is more than that of NaBr. Higher the strength, greater will be the melting point and higher will be the dissociation constant.

Covalent Bond

A Covalent Bond is formed due to sharing of electrons between two atoms.

When both the atoms ( similar or dissimilar ) taking part in a chemical combination are short of electrons to complete the nearest inert gas configuration, the combination between them takes place by sharing electrons. Each atom contributes one electron to form a common pair which then is shared by both.
The Bond established between atoms by this process of sharing is known as a covalent bond.
Chlorine molecule is represented as in Figure 8.

A covalent linkage is expressed by a dash (-) in writing formula. This system of representation of formula of a substance is known as electron dot (lewis) and dash formula.
Other examples of covalent compounds where dissimilar atoms combine are, NH₃ and CH₄. Their formation is illustrated as follows :

Sometimes atoms combine by sharing two or more electron pairs as shown below :

The sharing of two electron pairs between two atoms is referred to as double bond.

The sharing of three electron pairs between two atoms is referred to as triple bond.


Limitations of the Octet Rule :

The Octet rule is not adequate as it fails to explain the following points of bond formation.

1. Incomplete octet : For atoms of compounds possessing less than eight electrons.

2) Expanded octet : For atoms of compounds possessing more than eight electrons. Covalency :
The number of electron pairs which an atom can share with other atoms is known as the covalency of that element.

Atoms	H	Cl	O	N	C	P
Covalency	1	1	2	3	4	3

Polar Bonds

Polar Bonds are formed when one of the elements attracts the shared electrons more strongly than the second element. These bonds are somewhere between covalent and pure ionic bonds in character. There is asymmetry in the distribution of electron density. Usually in common expression we speak of electrons being shifted towards the more electronegative partner.

The unequal sharing of electrons leads to fractional electrical charges on atoms represented by the Greek letter delta

Transition between Ionic and Covalent Bonding :

Chemical reactions and physical measurement support the belief that most chemical compounds are intermediate in character between the purely ionic and the purely covalent.
Ionic bonds are found in compounds between the metals of low ionization potential and non-metals of high electron affinity.
A pure covalent bond is found only in molecules formed from two identical atoms such as Cl₂.
Most compounds fit somewhere between these two extremes.
If two spherical ions of opposite charges are brought together, the positively charged ion attracts and reforms the electron cloud of the anion. The electron cloud of the anion is drawn toward the cation, and, in extreme cases such ion deformation can lead to formation molecules with predominant covalent bonds.


Thus Aluminium chloride is a predominantly covalent material and not pure ionic material.


Electronegativity
Relative abilities of an atom to draw the bonding electrons towards its nuclei is expressed in terms of electronegativity.
It is a measure of how powerfully a bonded atom attracts the electrons in the bond.
An electronegativity scale has been developed by Linus Pauling to describe the attraction of the elements in a chemical bond for shared electrons.

H 2.2	Table 7 The Electronegativity Scale ( H = 2 .2 )
Li 1	Be 1.5	Slight difference in these figures have been reported in different citations	B 2.0	C 2.5	N 3.0	O 3.5	F 4.0
Na 0.9	Mg 1.2		Al 1.5	Si 1.8	P 2.1	S 2.5	Cl 3.1
K 0.8	Ca 1.0		Ga 1.7	Ge 1.8	As 2.1	Se 2.4	Br 3.0
Rb 0.8	Sr 1.0 Ba 0.9			Sn 1.9	Sb 2	Te 2.1	I 2.6
Cs 0.7				Pb 2.3	Bi 2	Po 2	At 2.2

The importance of these quantities lies in the fact that they at least qualitatively reflect the chemical behavior of the elements. The greater the separation between two atoms in the electronegativity scale the greater is the ionic character of a bond between them. The bond between two highly electronegative atoms is covalent as in NO or CCl₄. The bond linking two elements of low electronegativity is a metallic one as in Brass (Cu + Zn). The bond type between a highly electronegative and a low electronegative element is ionic as in KCl.

 Also the greater the separation of two elements in the electronegativity scale, the stronger will be the bond between them.

Table 8 Electronegativity and type of Bond

Electronegativity difference	Covalent character	Ionic character	Bond type
0.0 0.4 0.8 1.0	100 97 88 82	0 3 12 18	Covalent
1.2 1.6 2.0	75 60 40	25 40 60	Polar
2.4 2.6 3.0	32 26 10	68 74 90	Ionic

Based on the above two tables of electro-negativity the bond character of a molecule can be predicted.

Example :
HF 4.0 - 2 . 2 = 1.8
      F        H
The difference in electronegativity is 1.8. Based on the above table it indicates that it will have 50% covalent and 50% ionic character. Hence the bond type is polar.
In case of CO₂ it will be
3.5 - 2.5 = 1
O       C
Having electronegativity of 1 means it is a covalent bond.





Other Bonds
Metallic Bond :

In metals, the atoms are linked by a special type of bond called ’metallic bond’. It is suggested that the atoms in a metal free some of their electrons and become positive ions. The matrix of positive ions remains embedded in a sea of mobile electrons. The free electrons really cement the positive metal ions and explain why metals are highly conductive of heat or electricity.
Hydrogen Bond :
Hydrogen atoms occuring in certain groups such as -OH carry some positive charge and are thus attracted to electronegative atoms like oxygen and nitrogen thereby making a link, called hydrogen bond, between two oxygen or nitrogen atoms.

This hydrogen bond is electrostatic in nature and is not very strong. In ice crystals we have a tetrahedral structure, each oxygen atom is surrounded by four hydrogen atoms at the corner of the tetrahedron. Two of these are covalently bond and other two form a hydrogen bond. Van der Waal’s forces
These forces are the same as those existing between molecules in a real gas leading to its non-ideality. The forces are really interatomic or intermolecular, originating from dipole character. The forces are essentially weak. They exist between lattice layers of graphite or mica, as also between chains of selenium and tellurium.

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