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In 1923, Kazimierz Fajans formulated rules for predicting whether the bond is predominantly ionic or covalent. Fajans’ rule can be explained via the concept of chemical bonding and its types. This article covers the main description of Fajans’ rule and its postulates. Let’s see a brief about chemical bonding before delving into Fajans’ rule.
Chemical bonding:
Among all atoms, only noble gases will exist freely. But all the other atoms will easily combine with atoms to form different molecules. This process of combining two or more atoms will lead to chemical bonding.
In this process, two or more atoms will transfer or share electrons among themselves or with the outer shell of other atoms. The atoms will achieve a stable noble gas configuration with minimum energy.
There are three main types of chemical bonds, as follows:
Ionic bond
Covalent bond
Coordinate covalent bond
Ionic bond:
When one or more electrons transfer from the valence shell of one atom to the valence shell of another atom is called an ionic bond. Their electronic configuration will be closer to the noble gas configuration.
The atom which gives up electrons will become positively charged and it is called a positive ion.
The atom which takes up electrons from another atom will become negatively charged and is called a negative ion.
For example: Na+ + Cl–(g) → NaCl(s)
Covalent bond:
The two atoms share the electrons mutually amongst each other. This type of mutual sharing of electrons is called a covalent bond.
G.N.Lewis introduced the concept of the covalent bond in 1916.
When the two atoms share electrons amongst each other to their outermost shell, the electronic configuration of the two atoms will be closer to that of the electronic configuration of noble gases.
Example: H+ + Cl– → H–Cl
Coordinate covalent bond:
A coordinate covalent bond is also called a dative bond. Lewis also suggests this type of bond. When both the electrons are shared from a single atom to form the bond, it is called a coordinate covalent bond.
Fajans’ rule:
Fajans formulated rules to predict whether a bond is predominantly ionic or covalent. It is not always that a bond must be completely covalent or completely ionic. The covalent bond often may have an ionic character and an ionic bond will have a covalent character. This type of bonding is called the polarisation of ions.
Polarisation of ions occurs when two ions with opposite charges come in contact with each other, the anion and cation attract each other. At the same time, there will be repulsion between their nucleus, which has a positive charge and their electrons, which are of negative charge. As a result of this attraction and repulsion between the anion and cation, it will change the shape of the ion from spherical to a distorted shape (distortion), deformation or polarisation occurs in the anions.
The distortion of anions is polarised towards cation, and the distortion of cation is polarised towards anion. Among these, the distortion of anions occurs mostly than the polarisation of cations. There will be a high electron charge concentration between the two nuclei due to the polarisation of ions. As a result, this leads to the formation of a bond between ionic and covalent. This bond is called a polar covalent bond, which is more stable than a pure covalent bond.
The distortion of ions is shown below.
What is Fajans’ rule?
To predict whether the bond is predominantly covalent or predominantly ionic, Fajans suggested some rules regarding the polarisation of ions which are listed below:
When the size of the cation is smaller, the polarising power will be higher. This means that there will be polarisation of anion to a great extent. The cations with a positive charge will only have a small surface area. This results in high charge density, leading to the distortion of anions to a greater extent, i.e. formation of covalent bonds. There will be low polarising power for the cation, which has a larger size, i.e., ionic bond formation.
The larger the size of the anion, the higher will be the polarisability. This type of anion will be deformed easily and it will be polarised by a cation easily, i.e. forming a covalent bond. When the size of an anion is small, it forms an ionic bond.
A high charge on both cations and anions will result in effective polarisation. Polarisation is caused by the electrostatic force between anion and cation and the electrostatic force would be increased when the ions have a high charge on them. There will be more covalent bond formation. When the charge is less, there will be the formation of an ionic bond.
When the cations have inert gas pseudo configuration (ns2p6d10) or inert pair configuration (d10(n+1)s2), they will have high polarising power due to a high effective nuclear charge. While cations having noble gas configuration (ns2p6) will have low polarising power due to a low effective nuclear charge.
Effects of polarisation:
Polarisation of anion results in the formation of polar covalent bonds, which gives more stability. Hence, bromides and iodides have high lattice energies and high stabilities.
An increase in the formation of a covalent bond will decrease the solubility of ionic compounds in polar solvents.
An increase in covalent character will also lead to a decrease in the hardness of ionic compounds.
Explanation of Fajans’ rule via an example:
The melting point of KCl is higher than AgCl, but on comparing the radii of Ag+ & K+, they are almost the same. This is due to the molecule having a low melting point and having a low ionic character. So, the polarising power is based on the cation. The K+ has a configuration closer to the noble gas configuration and the Ag+ will have a configuration closer to the pseudo noble gas configuration. Thus, Ag+ will have more ionic character and hence its melting point would also be high.
Conclusion:
Fajans’ rule is important to find the nature of the predominant bond character of ions. This article explained Fajans’ rules with the help of an example. Fajans’ rules were also listed to identify the polarisation of ions and thus predict the nature of chemical bonds between elements.
