Dipole moment is the result of charge separation in any system. Dipole moment may be found in both ionic and covalently bonded molecules. The electronegativity difference between chemically bonded atoms or elements is the major source of the creation of dipole moment.

The separation of positive and negative charges in a compound is referred to as its polar character. Bond dipole moment is a metric for determining the polarity of a chemical bond between two atoms in a molecule. Since it has both magnitude and direction, the bond dipole moment is considered as a vector quantity. Read further to find more.

Dipole Moment

Although covalent bonding involves electron sharing, the two bonded atoms do not share the electrons equally. Therefore, there will always be one atom that attracts the electrons in the bond more strongly than the other atom unless the bond connects two atoms of the same element.

The ability of an atom to attract electrons in the presence of another atom is a measurable property called electronegativity and will produce a dipole moment.

Dipole moments are generally found in Polar Covalent Bonds. A covalent bond with an unequal sharing of electrons and the electronegativity difference within the range of \(0.1-2\) is called a polar covalent bond.
A covalent bond with an equal share of electrons and an electronegativity difference of zero is called a nonpolar covalent bond.

What is a Dipole Moment?

In Polar Covalent bonds, the electrons spend more time around the more non-metallic atom, which results in the unequal sharing of electron pairs between two participating atoms. There is a charge separation in such a bond, with one atom being slightly more positive and the other being more negative. The charge separation present in a molecule of a polar covalent compound is called a dipole moment.

An arrow with a cross represents the dipole at one end. The cross is near the end of the partially positive molecule, and the arrowhead is near the partially negative end of the molecule.

The bond dipole moment has both magnitude and direction. Hence, is a vector quantity.
Let us consider the example of hydrogen chloride \(\rm{(HCl)}\) molecule. In \(\rm{HCl}\), both hydrogen and chlorine atoms require one more electron to form an inert gas electronic configuration. As chlorine has a higher electronegativity than hydrogen atoms, it pulls the shared pair of electrons more towards itself.

Consequently, the bonding electrons in hydrogen chloride are shared unequally, resulting in a polar covalent bond between hydrogen and chlorine atom. This unequal sharing of the bonding pair of electrons results in a partial negative charge \((\rm{δ}^-)\) on the chlorine atom and a partial positive charge \((\rm{δ}^+)\) on the hydrogen atom. The symbol \(\rm{δ}\) (Greek lowercase delta) denotes these fractional charges.

The larger the differences between the electronegativities of the bonded atoms, the higher is the polarity of the bonds present between them. Thus, the dipoles are separated from each other by a distance, which is commonly denoted by ‘\(\rm{d}\)’.

Characteristics of Dipole Moment

1. In a polyatomic molecule, the dipole moment of a single bond is known as the bond dipole moment. This bond dipole moment is different from the resultant dipole moment of the entire molecule. The resultant dipole moment of the entire molecule is the vector sum of all the individual bond dipoles present in that particular molecule.
2. It is a vector quantity, i.e. it has magnitude as well as definite directions.
3. When two oppositely acting bond dipoles cancel each other, the resultant dipole moment is zero.
4. A small arrow denotes it with a crossed tail towards the negative end and its head towards the positive end by convention. This arrow symbolizes the shift of electron density in the molecule.

This arrow is placed parallel to the line joining the charge centres and pointed towards the negative end of the dipole.

Dipole Moment Formula

The polarity between the atoms results from the electronegativity difference of the involved atoms. Hence the degree of polarity of each bond in different molecules is different. For example, the \(-\rm{OH}\) bond in Water and \(-\rm{NH}\) bond in ammonia are different. This degree of polarity in a polar covalent bond is measured in terms of a physical quantity called the dipole moment.

A dipole moment of a molecule can be expressed as the product of the magnitude of the charge and the distance between the centers of the positive and negative charges. It is denoted by the Greek letter ‘\(\rm{µ}\)’.

Dipole Moment \((\rm{µ})\) = Charge \((\rm{?})\) * distance of separation \((\rm{d})\)

In chemical bonding, a bond dipole moment can be expressed as-
\( \rm{µ} = \rm{?} . \rm{d}\)
Where: \( \rm{µ}\) is the bond dipole moment,
\((\rm{?})\) is the magnitude of the partial charges \((\rm{δ}^+)\) and \((\rm{δ}^-)\),
And \(\rm{d}\) is the distance between \((\rm{δ}^+)\) and \((\rm{δ}^-)\) and is known as the bond length.
The bond dipole moment \((\rm{µ})\) is a vector quantity whose direction is parallel to the bond axis.
Dipole moment \(\rm{µ}\) is measured in Debye units denoted by the letter ‘\(\rm{D}\)’.
\(1\, \rm{Debye}\) is that dipole moment when the charges of the order of \(10^{-10}\,\rm{esu}\) are separated by the distance of the order \(10^{-18}\,\rm{cm}\).
\(1\, \rm{Debye}\) \(= 10^{-18}\,\rm{esu.cm}\)
\(1\, \rm{Debye}\) \( = 3.33564 \times {10^ { – 28}} {\text{C}}. {\text{cm}}\)
\(1\, {\text{D}} = 3.33564 \times {10^ { – 30}} {\text{C}}. {\text{m}}\); where \(\rm{C}\) is Coulomb and m denotes a meter.

Role of Dipole Moment in determining the Shape of Molecules

In chemistry, the Dipole Moments’ knowledge helps us decide the geometry and shape of the molecules.

For example, in the triatomic molecule of carbon dioxide \((\rm{CO}_2)\), the \(\rm{C-O}\) bonds are polar because oxygen is more electronegative than Carbon. However, the dipole moment of carbon dioxide is found to be zero. This is possible only if we assume the molecule to be linear, in which the two \(\rm{C-O}\) bonds are oriented in opposite directions at \(180\) degrees. Thus, the dipole moment of the \(\rm{C-O}\) bond on one side cancels that of the same bond on the other side.

Similarly, the geometry of Mercuric halides \((\rm{HgX}_2)\) and carbon disulfide \((\rm{CS}_2)\) is linear as they exhibit zero dipole moment. A triatomic molecule such as Water \((\rm{H}_2 \rm{O})\) exhibits a net dipole moment of \(1.84\,\rm{D}\). This indicates that the molecule’s shape is non-linear. There are two dipolar bonds in this molecule, while the net dipole moment is the resultant of the individual bond dipole moment, which is known as the molecular dipole moment. There are \(3\) polar \(\rm{N-H}\) bonds in ammonia molecule, each carrying a dipole moment of \(0.9\,\rm{Debyes}\). The resultant dipole moment is experimentally found to be \(1.46\,\rm{Debyes}\). This shows that the hydrogen atoms in the ammonia molecule are not aligned symmetrically with the nitrogen atoms. At the same time, the high value of dipole moment suggests the molecule to be triangular pyramidal.
The dipole moment values also help us to differentiate between ortho, meta, and para derivatives of a compound.
For example, in \(\rm{o}\)-dichlorobenzene, \(\rm{p}\)-dichlorobenzene and \(\rm{m}\)-dichlorobenzene. The resultant dipole moments in ortho and meta isomers are the combined dipole moments of two polar \(\rm{C-Cl}\) bonds.
The para isomer has zero dipole moment because the \(\rm{C-Cl}\) dipoles are opposite in direction, and they cancel each other’s effect. Hence, the dipole moment values can help us to identify the isomers.

Bond Polarity

If a molecule has more dipole moments than the other, it is more polar than other molecules. Thus, dipole moment tells us the degree of polarity in a polar covalent bond.

Consider 2 polar covalent molecules, hydrogen fluoride \((\rm{HF})\) and hydrogen iodide \((\rm{HI})\). The dipole moment of hydrogen fluoride is approximately \(1.92\,\rm{Debye}\), while the hydrogen iodide bond is \(0.58\,\rm{Debye}\). These values suggest that the polarity of bonds in \(\rm{HF}\) is greater than \(\rm{HI}\). This is because the electronegativity of Fluorine is greater than iodine. Hence Fluorine pulls the shared pair of electrons more towards itself. Thus, higher values of dipole moment indicate a higher degree of polarity.

As we know, the degree of polarity decides the type of bond. The dipole moment data helps to determine the ionic and covalent character of a bond in a molecule.

\({\text{Ionic character}} = \frac{{{\text{Observed Dipole Moment}}}}{{{\text{Theoretical Dipole Moment}}}} \times 100\%\)

For diatomic molecules, such as \(\rm{Cl}_2\), \(\rm{O}_2\), \(\rm{F}_2\), the dipole moment was zero as there is mutual sharing of electrons between the atoms. Whereas in polar molecules such as \(\rm{HCl}\), the percentage of ionic character was calculated to be around \(17\%\). This indicates that the \(\rm{HCl}\) molecule is \(83\%\) covalent and \(17\%\) ionic.

Significance of Dipole Moment

1. Compounds having dipoles possess properties intermediate between those of purely covalent and purely ionic compounds.
2. For simple molecules, the values of the dipole moment are approximately equal to the difference between the electronegativities of the atoms.
   For example, the dipole moment of \(\rm{HF} = 1.92\,\rm{D}\), which is approximately equal to the electronegativity difference between Fluorine and a hydrogen atom.
3. For complex molecules, the dipole moment of the different bonds can either reinforce one another or nullify each other.
   For example, let us consider the molecule of ammonia \((\rm{NH}_3)\). The dipole moment of each \(\rm{N-H}\) bond is \(0.9\,\rm{D}\), while the net dipole moment, which is the result of individual dipole moments, is \(1.46\,\rm{D}\). On the other hand, in a molecule such as \(\rm{CF}_4\), each \(\rm{C-F}\) bond is polar as the electronegativity of Carbon is \(2.5\) while that of Fluorine is \(4.0\). However, the \(\rm{CF}_4\) molecule is nonpolar. This difference is because all the four bonds are arranged symmetrically, and their dipole moments cancel each other.
4. The dipole moment values also help us differentiate between ortho, meta, and para derivatives of a compound.

Examples of Dipole Moment

The examples of dipole moments are:

1. Homonuclear Diatomic Molecules

Homonuclear diatomic molecules like Nitrogen, oxygen, and Chlorine have zero dipole moment due to the symmetrical charge distributions and similar electronegativity and ionization energy.

2. Heteronuclear Diatomic Molecules

Hydrogen bromide, hydrogen iodide, Hydrogen fluoride, Hydrogen Chloride have non-zero dipole moments that indicate the unsymmetrical charge distribution between two bonding atoms in the molecules.

Due to the difference in electronegativity of the constituent atoms in heteronuclear diatomic molecules, the electron pair is not equally shared in hydridized orbital. It is shifted towards the more electronegative atom.

3. Triatomic Molecules

The dipole moment of Carbon dioxide \((\rm{CO}_2)\): In carbon dioxide, oxygen being more electronegative than Carbon pulls the shared pair of electrons more towards itself. This results in a slight positive charge over the carbon atom and a slight negative charge over the oxygen atom. As the dipole moment due to two \(\rm{CO}\) bonds are opposite in direction, they nullify each other’s effect. Hence, the dipole moment is zero for carbon dioxide.

The dipole moment of Water \((\rm{H}_2 \rm{O})\): The water molecule exhibits a dipole moment of \(1.84\,\rm{D}\). The oxygen atom in the water molecule contains two lone pairs of electrons. These lone pairs of electrons have a resultant dipole moment in the same direction as the net dipole moment due to individual \(\rm{OH}\) bonds.

Dipole moment of \(\rm{BeCl}_2\): In Berrylium chloride, the individual \(\rm{Be-Cl}\) bonds are polar due to the high electronegativity of the chlorine atom. In contrast, the resultant dipole moment of the beryllium chloride molecule is zero. This is because the dipole moments due to the individual bonds are opposite in direction. Hence, they nullify each other’s effect and results in zero dipole moment.

4. Polyatomic Molecules

The dipole moment of \(\rm{BF}_3\): Let us consider the example of \(\rm{BF}_3\). The \(\rm{B-F}\) bonds in \(\rm{BF}_3\) are polar because Fluorine is more electronegative than Boron and hence pulls the shared pair of electrons more towards itself, resulting in a dipole.

As a result, Boron acquires a slight positive charge, whereas the Fluorine atom acquires a slight negative charge. The three fluorine atoms occupy the vertices of an equilateral triangle with the Boron atom at its center. As \(\rm{B-F}\) bonds are polar, the dipole moments of any two \(\rm{B-F}\) bond is equal and opposite to the third \(\rm{B-F}\) bond. The three dipole moments thus cancel each other, and the net dipole moment is zero.

Dipole moment of \(\rm{NH}_3\) and \(\rm{NF}_3\): In both molecules, the central nitrogen atom is covalently bonded to three other atoms and possesses a lone pair of electrons. The lone pair exerts dipole moments directed away from the nitrogen atom.

In ammonia, the dipole moment due to the \(\rm{N-H}\) bond is directed towards the nitrogen atom. The dipole moments due to lone pair of electrons and that due to polar \(\rm{N-H}\) bonds are added on and directed in an upward direction away from the nitrogen atom.

However, in \(\rm{NF}_3\), the dipole moment due to the \(\rm{N-F}\) bond is directed away from the nitrogen atom towards the Fluorine atom. The resultant dipole moment due to the \(\rm{N-F}\) bonds is directed downwards away from the Nitrogen. The lone pair exerts dipole moments in an upward direction away from the nitrogen atom. Hence, the resultant dipole moment due to the lone pair and the \(\rm{N-F}\) bonds are nullified to a certain extent.

Thus, Dipole Moment of \(\rm{NF}_3 < \rm{NH}_3\)

The dipole moment of \(\rm{CCl}_4\) and \(\rm{CHCl}_3\): In a molecule such as \(\rm{CCl}_4\), each \(\rm{C-Cl}\) bond is polar as the electronegativity of Carbon is \(2.5\) while that of Chlorine is \(3.9\). However, the \(\rm{CCl}_4\) molecule is nonpolar. This difference is because all the four bonds are arranged symmetrically, and their dipole moments cancel each other.

Summary

In this article, we have studied the dipole moment and its various consequences in detail. We also explored how the dipole moment plays a vital role in determining the different molecule’s shape, geometry, and symmetry. Some of the important points about dipole moment are

1. The bond dipole moment is the dipole moment of a single bond in a polyatomic molecule, which is distinct from the molecule’s overall dipole moment.
2. It’s a vector quantity, which means it has both magnitude and defined directions.
3. It can also be 0 because the two oppositely acting bond dipoles can cancel each other out as a vector quantity.
4. A little arrow with its tail on the negative centre and its head on the positive centre is used to represent it.
5. The dipole moment is symbolised in chemistry by a modest variant of the arrow sign. On the positive centre, there is a cross, and on the negative centre, there is an arrowhead. This arrow represents the molecule’s change in electron density.
6. In the case of a polyatomic molecule, the dipole moment is the vector sum of all bond dipoles present in the molecule.

Frequently Asked Questions on Dipole Moment

Q.1. What are the applications of dipole moment?
Ans: Dipole moment is used to calculate the percentage ionic character, bond angle, electric polarization, and residual charge on the atoms in the molecules. It also helps to determine the size or shape of molecules and the arrangements of chemical bonds in the molecules.

Q.2. How to identify whether a molecule consists of dipole moment or not?
Ans: When there is a difference in the electronegativity of two atoms involved in a bond, a dipole moment happens. The larger the electronegativity difference between the two atoms, the larger the bond’s dipole moment and polarity.

Q.3. Why is cis-\(1,2\)-dichloroethene have more dipole moment than trans-\(1,2\)-dichloroethene?
Ans: In cis-isomer, two similar groups, i.e. chlorine atoms, are present on the same side of the double bond, whereas, in the trans-isomer, two similar are present on the opposite side of the double bond.

Q.4. The electronegativity difference between Carbon and oxygen in \(\rm{CO}\) is very large, but the dipole moment of carbon monoxide very low. Why?
Ans: Carbon monoxide has a coordinate covalent bond. In this bond, the charge density available on the oxygen atom is back donated to the carbon atom. Hence, \(\rm{CO}\) forms a coordinate bonding directed towards the carbon atom. As a result, the dipole moment in \(\rm{CO}\) is directed towards the oxygen atom. As both directions are somewhat opposite, the net dipole moment is lowered to a greater extent.

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