Bond Polarity: Definition, Formula, Characteristics, Examples

When two atoms share a pair of electrons, they try to pull it towards themselves. This gives rise to the concept of Bond Polarity. Covalently bonded atoms share bonded pair of electrons among themselves. But this sharing of electrons is not ideal always. This means that the atom does not share the bonded electrons equally. Read on to learn more about it in this article.

Bond Polarity

Before learning about the polarity of bonds, we must learn about the causes behind their formation. The reasons which lead to the polarity of bonds are:

1. The electronegativity difference between the combining atoms.
2. Formation of a dipole.

What is Dipole and Dipole Moment?

A covalent bond involves the mutual sharing of valence electrons of the participating atom. This mutual sharing of electrons takes place to attain the stable electronic configuration of the neighbouring noble gas atom.

When atoms share their electrons, it tends to happen that the atom having higher electronegativity than the other pulls the shared pair of electrons more towards itself. This unequal sharing of electron pairs leads to the formation of dipoles (\(2\) poles), and the molecules exhibiting this property are called Polar molecules.

The formation of a dipole results in a charge separation in such a type of 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 the 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, it is a vector quantity.

Let us take the example of a hydrogen chloride \((\rm{HCl})\) molecule. In \((\rm{HCl})\) molecules, hydrogen and Chlorine atoms require one more electron to form an inert gas electronic configuration. Despite having a higher electronegativity than hydrogen, the chlorine atom’s attraction for electrons is insufficient to remove an electron from the hydrogen atom.

This results in the bonding electrons in hydrogen chloride being more inclined towards the chlorine atom. Consequently, a polar covalent bond is formed between hydrogen and chlorine atom. This unequal sharing of the bonding pair of electrons results in a partial negative charge \((δ^-\)) on the chlorine atom and a partial positive charge \((δ^+)\) on the hydrogen atom. The symbol \(δ\) (Greek lowercase delta) denotes these fractional charges.

The larger the electronegativity differences between 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 ‘\(d\)’.

What is Bond Polarity?

Bond Polarity is the measure of the extent to which a bond exhibits polarity. In simple words, a bond polarity is a scientific tool that gives us an idea about the nature of the bonds and the type of bonding they will undergo to form compounds.

As dipole formation is a characteristic of bond polarity, the molecules exhibiting dipoles are more precisely called Polar covalent molecules.

Due to the unequal sharing of bonding pair of electrons, the atom with a higher electronegative value develops a slightly negative charge \((-δ)\). Conversely, the atom with a less electronegative value acquires a slightly positive charge \((+δ)\). This charge separation in polar covalent bonds due to the electronegativity difference is called a dipole moment.

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.

However, in covalent molecules that incorporate equal sharing of electrons between the atoms, such a situation does not arise. This mainly occurs in molecules where the electronegativity difference between the participating atoms is zero or less than \(0.4\).

As we know, the degree of polarity decides the type of bonding in a molecule, the dipole moment data helps 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 found to be zero as there exists 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.

Bond Polarity and Electronegativity

We know, in a polar covalent bond, the two bonded atoms do not share the electrons equally, unless the bond connects two atoms of the same element. There will always be one atom that attracts the electrons in the bond more strongly than the other atom.

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.

Nonpolar Covalent Bond

A bond that is formed by the equal sharing of electrons between the combining atoms is called a nonpolar covalent bond. This phenomenon arises when there lies no difference in the electronegativities of the two combining atoms. Hence, identical pairs of atoms form a nonpolar covalent bond. For a nonpolar bond to exist, the electronegativity difference of the combining atoms should be less than \(0.5\).

Let us consider the example of the oxygen molecule. Though oxygen is highly electronegative, the covalent bond in \(\rm{O}_2\) is not Polar. This is because both atoms have the same electronegativity, and electrons are shared equally between them.

Some of the compounds that have nonpolar bonds are listed below.

Are All Molecules Possessing Bond Polarity, Polar?

We all know that water being a small molecule has a very high boiling point \([100°\rm{C}]\). This is because water is a polar molecule. On the contrary, Carbon dioxide is a nonpolar molecule despite significant bond polarity. In the \(\rm{C = O}\) bond, oxygen is more electronegative than Carbon, developing a slight negative charge, whereas the carbon atom develops a slight positive charge.

As a result, the \(\rm{CO}\) bonds become polar, and the dipole moment is directed from the carbon atom towards the oxygen atom. As there are two \(\rm{CO}\) bonds in \(\rm{CO}_2\), the dipole moment is exactly opposite to each other in the molecule and cancels each other’s effects. However, the individual \(\rm{C = O}\) bonds are polar; as a whole, carbon dioxide molecules are nonpolar. This lack of polarity influences some of carbon dioxide’s properties. (For example, carbon dioxide becomes a gas at \(−77°\rm{C}\), almost \(200°\rm{C}\) lower than the temperature at which water boils.)

Hence, we can conclude that all Polar covalent molecules possess bond Polarity, but the presence of bond polarity in covalent molecules does not impart polarity to all the molecules as a whole. A molecule can be nonpolar even if it has polar bonds. The overall molecule will have no charge or be neutral if the polarity is equal and directly opposing, as in carbon dioxide (b).

Polar Covalent Bond Vs. Nonpolar Covalent Bond

Characteristics of Dipole Moment

1. In a polyatomic molecule, the dipole moment of a single bond is known as the bond dipole moment. 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. Hence, a polyatomic molecule can be nonpolar despite having polar bonds.
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 symbolises the shift of electron density in the molecule.
   
5. This arrow is placed parallel to the line joining the charge centers and pointed towards the negative end of the dipole.

Bond Polarity Examples

The examples of bond polarity are given below:

Molecules with Bond Polarity

a. Polar Molecules with Bond Polarity

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, known as the molecular dipole moment.

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 bond is always polar. Hence, the electron pair is not equally shared in hybridized disulfide orbital and shifted to the more electronegative atom.

b. Nonpolar Molecules with Bond Polarity

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 disulphide \((\rm{CS}_2)\) is linear as they exhibit zero 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.

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.

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.

Some other examples of nonpolar molecules with polar covalent bonds are listed below:

Molecules without Bond Polarity

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

Factors on which Bond Polarity Depends

1) Relative Electronegativity of Participating Atoms

As bond polarity involves the pulling of electrons towards itself, a high electronegative element will affect the extent of the polarity of bonds. The amount of the shifting of a bond pair of electrons will depend upon the relative electronegativity of the participating atoms.

2) The Spatial Arrangement of Various Bonds in the Atom

The shared pair of electrons also experience pulling force from the other bonded and non-bonded pair of electrons. This results in different bond polarities between the same atoms but present in different molecules. E.g. Bond Polarity of the \(\rm{O-H}\) bond in a water molecule and acetic acid molecule is different. This is due to the different spatial arrangement of various bonds in the molecule.

Summary

In this article, we have studied the bond polarity 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.

FAQs on Bond Polarity

Q.1. What increases bond polarity?
Ans: The electronegativity difference between the participating atoms increases the bond polarity. The higher the difference, the greater is the extent of bond polarity. 

Q.2. How do you determine the polarity of a bond?
Ans: The polarity of a bond can be determined by calculating the electronegativity difference between the participating atoms. For example, if the difference lies within \(0.4-1.8\), then the bond is a polar covalent bond.

Q.3. Which bond has the greatest polarity?
Ans: The electronegativity difference between Hydrogen and Fluorine is the highest. Hence, the HF bond has the greatest polarity. 

Q.4. What makes bonds polar or nonpolar?
Ans: The electronegativity difference between the participating atoms makes a bond polar or nonpolar. For a bond to be polar, the electronegativity difference between the two elements needs to be between \(0.5\) to \(1.6\). If the electronegativity difference is less than \(0.5\), the bond is nonpolar. Any more than \(1.6\) and the molecules become charged ions and form ionic bonds instead.

Q.5. What are the types of bond polarity?
Ans: When the electronegativity difference is minimal or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. 

Q.6. What is the relationship between electronegativity and bond polarity?
Ans: In a polar covalent bond, bonded electrons are shifted toward the atom having a higher electronegativity value. Thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarised is the electron distribution and the larger is the partial charges of the atoms.

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