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Polarity of Bonds

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Two atoms share electrons to form a covalent bond. Depending upon the difference in electronegativities of the bonded atoms, the bonds can be polar or non-polar as discussed below:

Non Polar Covalent Bonds

When two similar atoms share electrons to form a covalent bond, the shared pair of electrons is equally attracted by the two atoms because the electronegativity of both the atoms is equal.
Example : The bond between two hydrogen atoms in H2 molecule is non-polar.

Non polar covalent bond in H2 molecule

Some other molecules that contain non polar bonds : N2, Cl2, F2, O2 etc.

Polar Covalent Bonds

When a covalent bond is formed in a heteronuclear molecule — molecules composed of more than one type of atoms — the shared pair of electrons between the two atoms gets displaced towards the atom which is more electronegative. As a result, one end of the molecule becomes slightly negatively charged while the other end becomes slightly positively charged. In other words, positive and negative poles are developed in the molecule. This phenomenon is known as polarisation and the bond thus formed is called a polar covalent bond.
Example : The bond between hydrogen and chlorine in HCl is polar covalent bond.

Polar covalent bond in HCl

Dipole Moment

Due to polarisation, the molecule possess a dipole moment which can be defined as the product of the magnitude of negative or positive charge(q) and the distance (d) between the centres of positive and negative charges.

Mathematically,

Dipole moment μ = charge (q) × distance of separation (d)

Dipole moment is usually denoted by μ and is expressed in Debye units.

1 D = 3.33564 × 10-30 C m.

where C and m are coulomb and meter respectively.

Note : A molecule as a whole is always electrically neutral, that means the positive charge is always equal in magnitude to the negative charge.

Dipole moment is a vector quantity and is usually represented by symbol with head pointing towards the negative end.

In polyatomic molecules — molecules having more than two atoms — the dipole moment not only depends on the individual dipole moments of bonds but also on the spatial arrangement of various bonds in the molecule. The dipole moment in polyatomic molecules is the vector sum of the dipole moments of various bonds. Study the following examples for better understanding :

Dipole moment in BeF2 :

The electronegativity of fluorine is higher than that of beryllium; therefore, the shift of electron density is towards fluorine (as depicted by arrows).

Dipole moment in BeF2

The dipole moment in BeF2 is zero because the two equal bond dipoles point in opposite direction and cancel the effect of each other.

Dipole moment in BeF2 is zero

Dipole moment in BF3 :

The electronegativity of fluorine is higher than that of boron; therefore, the shift of electron density is towards F.

Dipole moment of BF3

The B-F bonds in BF3 are oriented at an angle of 120° to one another. Even so, the net dipole moment in BF3 is zero because the resultant of any two is equal and opposite to the third.

Dipole moment in BF3 is zero

Which one has higher dipole moment : NF3 or NH3

Although the electronegativity of fluorine is higher than that of hydrogen, the dipole moment of NH3 is higher than that of NF3. The reason can be understood with the help of the following figures :

Dipole moment of NH3
Dipole moment of NF3

In NH3, the shift of electron density of N-H bonds is towards nitrogen because nitrogen is more electronegative than hydrogen; hence, the orbital dipole due to lone pair and the resultant dipole moment of the N-H bonds are in the same direction.

In NF3, fluorine is more electronegative than nitrogen; hence, the orbital dipole and the resultant dipole moment of the three N-F bonds are in the opposite direction. The orbital dipole decreases the effect of the resultant dipole moment of N-F bonds resulting in the lower dipole moment of NF3.

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