POLARITY (CHEMISTRY): POLAR MOLECULES AND EXAMPLES - CHEMISTRY - 2025

The chemical polarity is a property characterized by the presence of marked heterogeneous distribution of electron densities in a molecule. In its structure, therefore, there are negatively charged regions (δ-), and others positively charged (δ +), generating a dipole moment.

The dipole moment (µ) of bond is a form of expression of the polarity of a molecule. It is usually represented as a vector whose origin is in the charge (+) and its end is located in the charge (-), although some chemists represent it inversely.

Electrostatic potential map for the water molecule. Source: Benjah-bmm27 via Wikipedia.

The upper image shows the electrostatic potential map for water, H ₂ O. The reddish region (oxygen atom) corresponds to the one with the highest electron density, and it can also be seen that it stands out on the blue regions (hydrogen atoms).

As the distribution of said electron density is heterogeneous, it is said that there is a positive and a negative pole. That is why we talk about chemical 'polarity', and dipole moment.

Dipole moment

The dipole moment µ is defined by the following equation:

µ = δ · d

Where δ is the electric charge of each pole, positive (+ δ) or negative (–δ), and d is the distance between them.

The dipole moment is usually expressed in debye, represented by the symbol D. One coulomb · meter equals 2.998 · 10 ²⁹ D.

The value of the dipole moment of the bond between two different atoms is in relation to the difference in electronegativities of the atoms that form the bond.

For a molecule to be polar, it is not enough to have polar bonds in its structure, but it must also have an asymmetric geometry; in such a way that it prevents the dipole moments from vectorly canceling each other.

Asymmetry in the water molecule

The water molecule has two OH bonds. The geometry of the molecule is angular, that is, shaped like a "V"; therefore, the dipole moments of the bonds do not cancel each other out, but rather the sum of them is produced pointing towards the oxygen atom.

The electrostatic potential map for H ₂ O reflects this.

If the angular molecule HOH is observed, the following question may arise: is it really asymmetric? If an imaginary axis is drawn through the oxygen atom, the molecule will divide into two equal halves: HOOH.

But, it is not so if the imaginary axis is horizontal. When this axis now divides the molecule back into two halves, you will have the oxygen atom on one side, and the two hydrogen atoms on the other.

For this reason, the apparent symmetry of H ₂ O ceases to exist, and it is therefore considered an asymmetric molecule.

Polar molecules

Polar molecules must meet a series of characteristics, such as:

-Distribution of electric charges in the molecular structure is asymmetric.

-They are usually soluble in water. This is because polar molecules can interact by dipole-dipole forces, where water is characterized by having a large dipole moment.

In addition, its dielectric constant is very high (78.5), which allows it to keep the electrical charges separate, increasing its solubility.

-In general, polar molecules have high boiling and melting points.

These forces are constituted by the dipole-dipole interaction, the London dispersive forces and the formation of hydrogen bonds.

-Due to their electrical charge, polar molecules can conduct electricity.

Examples

SW

Sulfur dioxide (SO ₂). Oxygen has an electronegativity of 3.44, while the electronegativity of sulfur is 2.58. Therefore, oxygen is more electronegative than sulfur. There are two S = O bonds, the O having a δ- charge and the S having a δ + charge.

Because it is an angular molecule with S at the vertex, the two dipole moments are oriented in the same direction; and therefore, they add up, making the SO ₂ molecule polar.

CHCl

Chloroform (HCCl ₃). There is one CH bond and three C-Cl bonds.

The electronegativity of C is 2.55, and the electronegativity of H is 2.2. Thus, carbon is more electronegative than hydrogen; and therefore, the dipole moment will be oriented from H (δ +) towards C (δ-): C ^δ- -H ^{δ +}.

In the case of C-Cl bonds, C has an electronegativity of 2.55, while Cl has an electronegativity of 3.16. The dipole vector or dipole moment is oriented from C to Cl in the three C ^{δ +} -Cl ^δ- bonds.

As there is an electron-poor region around the hydrogen atom and an electron-rich region made up of the three chlorine atoms, CHCl ₃ is considered a polar molecule.

HF

Hydrogen fluoride has only one HF bond. The electronegativity of H is 2.22 and the electronegativity of F is 3.98. Therefore, fluorine ends up with the highest electron density, and the bond between both atoms is best described as: H ^{δ +} -F ^δ-.

NH

Ammonia (NH ₃) has three NH bonds. The electronegativity of N is 3.06 and the electronegativity of H is 2.22. In all three bonds, the electron density is oriented towards nitrogen, being even higher by the presence of a pair of free electrons.

The NH ₃ molecule is tetrahedral, with the N atom occupying the vertex. The three dipole moments, corresponding to NH bonds, are oriented in the same direction. In them, δ- is located in N, and δ + in H. Thus, the bonds are: N ^δ- -H ^{δ +}.

These dipole moments, the asymmetry of the molecule, and the free pair of electrons on nitrogen, make ammonia a highly polar molecule.

Macromolecules with heteroatoms

When the molecules are very large, it is no longer certain to classify them as apolar or polar themselves. This is because there may be parts of its structure with both apolar (hydrophobic) and polar (hydrophilic) characteristics.

These types of compounds are known as amphiphiles or amphipathics. Because the apolar part can be considered electron-poor relative to the polar part, there is a polarity present in the structure, and amphiphilic compounds are considered to be polar compounds.

A macromolecule with heteroatoms can be generally expected to have dipole moments, and thus chemical polarity.

Heteroatoms are understood to be those that are different from those that make up the skeleton of the structure. For example, the carbon skeleton is biologically the most important of all, and the atom with which carbon forms a bond (in addition to hydrogen) is called a heteroatom.

References

1. Whitten, Davis, Peck & Stanley. (2008). Chemistry. (8th ed.). CENGAGE Learning.
2. Prof. Krishnan. (2007). Polar and nonpolar compounds. St. Louis Community College. Recovered from: users.stlcc.edu
3. Murmson, Serm. (March 14, 2018). How to Explain Polarity. Sciencing. Recovered from: sciencing.com
4. Helmenstine, Anne Marie, Ph.D. (December 05, 2018). Polar Bond Definition and Examples (Polar Covalent Bond). Recovered from: thoughtco.com
5. Wikipedia. (2019). Chemical polarity. Recovered from: en.wikipedia.org
6. Quimitube. (2012). Covalent bond: bond polarity and molecular polarity. Recovered from: quimitube.com