4.1 Kossel- Lewis Approach to Chemical Bonding

• Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. In the case of sodium and chlorine, this can happen by the transfer of an electron from sodium to chlorine thereby giving the \(\mathrm{Na}^{+}\)and \(\mathrm{Cl}^{-}\)ions. In the case of other molecules like \(\mathrm{Cl}_2, \mathrm{H}_2, \mathrm{~F}_2\), etc., the bond is formed by the sharing of a pair of electrons between the atoms. In the process, each atom attains a stable outer octet of electrons.
• Lewis Symbols: In the formation of a molecule, only the outer shell electrons take part in chemical bonding and they are known as valence electrons. The inner shell electrons are well protected and are generally not involved in the combination (bonding) process. G.N. Lewis, an American chemist introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols. For example, the Lewis symbols for the elements of the second period are as under:

• Significance of Lewis Symbols: The number of dots around the symbol represents the number of valence electrons. This number of valence electrons helps to calculate the common or group valence of the element. The group valence of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.

Kössel, in relation to chemical bonding, drew attention to the following facts:

• In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases;
• The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of an electron by the respective atoms;
• The negative and positive ions thus formed attain stable noble gas electronic configurations. The noble gases (with the exception of helium which has a duplet of electrons) have a particularly stable outer shell configuration of eight (octet) electrons, \(n s^2 n p^6\).
• The negative and positive ions are stabilized by electrostatic attraction.
• For example, the formation of \(\mathrm{NaCl}\) from sodium and chlorine, according to the above scheme, can be explained: \(
  \begin{array}{lll}
  \mathrm{Na} & \rightarrow & \mathrm{Na}^{+}+\mathrm{e}^{-} \\
  {[\mathrm{Ne}] 3 \mathrm{~s}^1} & & {[\mathrm{Ne}]} \\
  \mathrm{Cl}+\mathrm{e}^{-} & \rightarrow & \mathrm{Cl}^{-} \\
  {[\mathrm{Ne}] 3 s^2 3 p^5} & & {[\mathrm{Ne}] 3 s^2 3 p^6 \text { or }[\mathrm{Ar}]} \\
  \mathrm{Na}^{+}+\mathrm{Cl}^{-} & \rightarrow & \mathrm{NaCl} \text { or } \mathrm{Na}^{+} \mathrm{Cl}^{-}
  \end{array}
  \)
  Similarly the formation of \(\mathrm{CaF}_2\) may be shown as:
  \(
  \begin{array}{llcl}
  \mathrm{Ca} & \rightarrow \mathrm{Ca}^{2+}+2 \mathrm{e}^{-} \\
  {[\mathrm{Ar}] 4 s^2} & \quad[\mathrm{Ar}] \\
  \mathrm{F}+\mathrm{e}^{-} & \mathrm{F}^{-} \\
  {[\mathrm{He}] 2 s^2 2 p^5} & {[\mathrm{He}] 2 s^2 2 p^6 \text { or }[\mathrm{Ne}]} \\
  \mathrm{Ca}^{2+}+2 \mathrm{~F}^{-} & \rightarrow \mathrm{CaF}_2 \text { or } \mathrm{Ca}^{2+}\left(\mathrm{F}^{-}\right)_2
  \end{array}
  \)
• The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the electrovalent bond. The electrovalence is thus equal to the number of unit charge(s) on the ion. Thus, calcium is assigned a positive electrovalence of two, while chlorine a negative electrovalence of one.
• Kössel’s postulations provide the basis for the modern concepts regarding ion formation by electron transfer and the formation of ionic crystalline compounds. 

Octet Rule

Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as the electronic theory of chemical bonding. According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.

Covalent Bond

• The Lewis-Langmuir (1919) theory can be understood by considering the formation of the chlorine molecule, \(\mathrm{Cl}_2\). The \(\mathrm{Cl}\) atom with electronic configuration, [Ne] \(3 s^2 3 p^5\), is one electron short of the argon configuration. The formation of the \(\mathrm{Cl}_2\) molecule can be understood in terms of the sharing of a pair of electrons between the two chlorine atoms, each chlorine atom contributing one electron to the shared pair. In the process, both chlorine atoms attain the outer shell octet of the nearest noble gas (i.e., argon).

• The bond formed as a result of sharing of electrons between elements is called covalent bond. When electrons shared by bonded atoms are contributed entirely by one of the bonded atoms bond formed is known as a coordinate bond. Lewis dot structures provide a picture of bonding in molecules and ions in terms of shared pair of electrons and the octet rule.
• The dots represent electrons. Such structures are referred to as Lewis dot structures.
• Covalency: It is defined as the number of electrons contributed by an atom of the element for sharing with other atoms so as to achieve noble gas configuration.

The Lewis dot structures can be written for other molecules also, in which the combining atoms may be identical or different. The important conditions being that:

• Each bond is formed as a result of sharing of an electron pair between the atoms.
• Each combining atom contributes at least one electron to the shared pair.
• The combining atoms attain the outer shell noble gas configurations as a result of the sharing of electrons.
• Thus in water and carbon tetrachloride molecules, the formation of covalent bonds can be represented as:

Single covalent bond

When two atoms share one electron pair they are said to be joined by a single covalent bond.  Example: Covalent bond between two Cl atoms as shown below.

Double bond

If two atoms share two pairs of electrons, the covalent bond between them is called a double bond. For example, in the carbon dioxide molecule, we have two double bonds between the carbon and oxygen atoms.

Similarly in ethene molecule the two carbon atoms are joined by a double bond.

Triple bond

When combining atoms share three electron pairs as in the case of two nitrogen atoms in the \(\mathrm{N}_2\) molecule and the two carbon atoms in the ethyne molecule, a triple bond is formed.

Lewis Representation of Simple Molecules (the Lewis Structures)

The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule. It does help in understanding the formation and properties of a molecule to a large extent. Writing of Lewis dot structures of molecules is, therefore, very useful. The Lewis dot structures can be written by adopting the following steps:

• The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms. For example, in the \(\mathrm{CH}_4\) molecule there are eight valence electrons available for bonding ( 4 from carbon and 4 from the four hydrogen atoms).
• For anions, each negative charge would mean the addition of one electron. For cations, each positive charge would result in the subtraction of one electron from the total number of valence electrons. For example, for the \(\mathrm{CO}_3^{2-}\) ion, the two negative charges indicate that there are two additional electrons than those provided by the neutral atoms. For \(\mathrm{NH}_4^{+}\)ion, one positive charge indicates the loss of one electron from the group of neutral atoms.
• Knowing the chemical symbols of the combining atoms and having knowledge of the skeletal structure of the compound (known or guessed intelligently), it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.
• In general, the least electronegative atom occupies the central position in the molecule/ion. For example in the \(\mathrm{NF}_3\) and \(\mathrm{CO}_3^{2-}\), nitrogen and carbon are the central atoms whereas fluorine and oxygen occupy the terminal positions.
• After accounting for the shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as lone pairs. The basic requirement being that each bonded atom gets an octet of electrons.
  Lewis representations of a few molecules/ ions are given in Table 4.1.

Example 4.1: Write the Lewis dot structure of \(\mathrm{CO}\) molecule.

Solution:

Step 1. Count the total number of valence electrons of carbon and oxygen atoms. The outer (valence) shell configurations of carbon and oxygen atoms are \(2 s^2 2 p^2\) and \(2 s^2 2 p^4\), respectively. The valence electrons available are \(4+6=10\).
Step 2. The skeletal structure of \(\mathrm{CO}\) is written as: \(\mathrm{C} \quad \quad \mathrm{O}\)
Step 3. Draw a single bond (one shared electron pair) between \(\mathrm{C}\) and \(\mathrm{O}\) and complete the octet on \(\mathrm{O}\), the remaining two electrons are the lone pair on \(\mathrm{C}\).

This does not complete the octet on carbon and hence we have to resort to multiple bonding (in this case a triple bond) between \(\mathrm{C}\) and \(\mathrm{O}\) atoms. This satisfies the octet rule condition for both atoms.

Example 4.2: Write the Lewis structure of the nitrite ion, \(\mathrm{NO}_2^{-}\).

Solution:

Step 1. Count the total number of valence electrons of the nitrogen atom, the oxygen atoms and the additional one negative charge (equal to one electron).
\(
\begin{aligned}
& \mathrm{N}\left(2 s^2 2 p^3\right), \mathrm{O}\left(2 s^2 2 p^4\right) \\
& 5+(2 \times 6)+1=18 \text { electrons }
\end{aligned}
\)
Step 2. The skeletal structure of \(\mathrm{NO}_2^{-}\)is written as : \(\mathrm{O} \quad \quad \mathrm{N} \quad \quad \mathrm{O}\)
Step 3. Draw a single bond (one shared electron pair) between the nitrogen and each of the oxygen atoms completing the octets on oxygen atoms. This, however, does not complete the octet on nitrogen if the remaining two electrons constitute lone pair on it.

Hence we have to resort to multiple bonding between nitrogen and one of the oxygen atoms (in this case a double bond). This leads to the following Lewis dot structures.

Formal Charge

Lewis dot structures, in general, do not represent the actual shapes of the molecules. In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom. It is, however, feasible to assign a formal charge on each atom.

The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It is expressed as:

The counting is based on the assumption that the atom in the molecule owns one electron of each shared pair and both the electrons of a lone pair.

Let us consider the ozone molecule \(\left(\mathrm{O}_3\right)\). The Lewis structure of \(\mathrm{O}_3\) may be drawn as:

The atoms have been numbered as 1,2 and 3. The formal charge on:

• The central \(\mathrm{O}\) atom marked 1
  \(
  =6-2-\frac{1}{2}(6)=+1
  \)
• \(\quad\) The end \(\mathrm{O}\) atom marked 2
  \(
  =6-4-\frac{1}{2}(4)=0
  \)
• The end \(\mathrm{O}\) atom marked 3
  \(
  =6-6-\frac{1}{2}(2)=-1
  \)
  Hence, we represent \(\mathrm{O}_3\) along with the formal charges as follows:

We must understand that formal charges do not indicate real charge separation within the molecule. Indicating the charges on the atoms in the Lewis structure only helps in keeping track of the valence electrons in the molecule. Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species.

Generally, the lowest energy structure is the one with the smallest formal charges on the atoms. The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighbouring atoms.

Limitations of the Octet Rule

The octet rule, though useful, is not universal. It is quite useful for understanding the structures of most of the organic compounds and it applies mainly to the second-period elements of the periodic table. There are three types of exceptions to the octet rule.

• The incomplete octet of the central atom: In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are \(\mathrm{LiCl}, \mathrm{BeH}_2\) and \(\mathrm{BCl}_3\).

• \(\mathrm{Li}, \mathrm{Be}\) and \(\mathrm{B}\) have 1,2 and 3 valence electrons only. Some other such compounds are \(\mathrm{AlCl}_3\) and \(\mathrm{BF}_3\).
• Odd-electron molecules: In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, \(\mathrm{NO}_2\), the octet rule is not satisfied for all the atoms

• The expanded octet: Elements in and beyond the third period of the periodic table have, apart from \(3 s\) and \(3 p\) orbitals, \(3 d\) orbitals also available for bonding. In a number of compounds of these elements, there are more than eight valence electrons around the central atom. This is termed as the expanded octet. Obviously, the octet rule does not apply in such cases.

  Some of the examples of such compounds are: \(\mathrm{PF}_5, \mathrm{SF}_6, \mathrm{H}_2 \mathrm{SO}_4\) and a number of coordination compounds.

  
• Interestingly, sulphur also forms many compounds in which the octet rule is obeyed. In sulphur dichloride, the S atom has an octet of electrons around it.

Drawbacks of the octet theory

• It is clear that the octet rule is based upon the chemical inertness of noble gases. However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like \(\mathrm{XeF}_2, \mathrm{KrF}_2, \mathrm{XeOF}_2\) etc.,
• This theory does not account for the shape of molecules.
• It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.