Lesson Explainer: Polar Bonding Chemistry

In this explainer, we will learn how to identify polar bonds within molecules and assess their effect on the polarity of the molecule as a whole.

Polar molecules have an asymmetrical distribution of electron density because one of the covalently bonded atoms has a higher electronegativity value than the other. Electronegativity measures the tendency of an atom to attract a bonding pair of electrons from a chemical bond. In 1932, Linus Pauling famously introduced an accurate electronegativity scale that could be used to quantify the ability of any element to withdraw electron density from bonding pairs of electrons.

The following diagram shows a single molecule of hydrogen chloride. The image shows that the bonding electron density (colored cloud) is concentrated around the more electronegative chlorine atom.

Pauling electronegativity values can be used to understand why some atoms form simple molecular compounds, while other atoms bond together and form giant ionic lattices. Elements will usually bond together and form simple covalently bonded compounds when they have the same or very similar electronegativity values. Elements will usually bond ionically and form giant ionic structures when one of the elements has a very low electronegativity value and the other has a much higher electronegativity value.

The absolute difference of electronegativity values can be defined as where and are the electronegativity values of two chemically bonded elements. Elements usually bond together and form covalent bonds when the difference of electronegativity values is less than 1.7. Elements usually bond together and form more ionic bonds when the difference of electronegativity values is greater than 1.7.

The absolute difference of electronegativity values also determines if simple molecular compounds will form nonpolar or polar covalently bonded molecules. Elements will usually bond together and make nonpolar covalent molecules when the difference of their electronegativity values is less than 0.4. Elements will usually bond together and make polar covalent molecules when the difference of their electronegativity values is greater than 0.4 but less than 1.7.

Bonding electrons are distributed unequally in ionic compounds but are shared slightly more equally in polar covalently bonded compounds. Bonding pairs of electrons are shared even more equally in nonpolar covalent compounds, and they are shared completely evenly in pure covalent compounds. Pure covalent compounds always contain atoms that have the same electronegativity values.

For example, hydrogen gas molecules () contain two identical hydrogen atoms that both have Pauling electronegativity numbers of 2.20. Neither one of the hydrogen atoms is capable of withdrawing more electron density than the other hydrogen atom. The electron density, as shown below, is shared completely evenly between the two bonded hydrogen atoms, and the bond is described as being pure covalent or completely nonpolar.

Bonding pairs of electrons are shared unequally in polar covalent compounds, such as hydrogen chloride () molecules. Hydrogen has a relatively low electronegativity number of 2.20, and chlorine has a much higher electronegativity number of 3.16. In a molecule of , the chlorine atom is much more effective at withdrawing the bonding pair of electrons, and the covalent bond is almost always highly polar. The negatively charged electron density shifts over toward the chlorine atom and it ends up having a partial negative electrostatic charge (), while the hydrogen atom ends up having a partial positive electrostatic charge ().

Example 4: Understanding How Electron Density Is Distributed in Hydrogen Chloride Molecules

Which of the following diagrams shows a correct representation of the electron cloud that surrounds an molecule?

Answer

Hydrogen has an electronegativity number of 2.20, and chlorine atoms have a much higher electronegativity number of 3.16. The negatively charged electron density shifts over toward the highly electronegative chlorine atom and it ends up having a partial negative electrostatic charge (), while the covalently bonded hydrogen atom ends up having a partial positive electrostatic charge (). A larger cloud of electron density would therefore be expected around the chlorine atom with a smaller cloud of electron density around the hydrogen atom. Such a distribution is shown best in option C.

The electron density distribution is asymmetric about the center of the covalent bond. One side of the hydrogen chloride molecule has a partial positive electrostatic charge, while the other side has a partial negative electrostatic charge. There is a large difference in electron density at the two ends of the hydrogen chloride molecules, and there is an electric dipole moment that runs from the hydrogen atom through to the chlorine atom. The quantitative measure of the dipole is called the dipole moment, and it is represented by the letter .

It is important to realize that a molecule cannot automatically be assumed to be polar just because it contains at least one polar covalent bond. Bonded atoms with significantly different electronegativity values will generate unidirectional electric dipole moments, but molecules can contain other unidirectional electric dipole moments, and two or more of these vectors can cancel each other out.

The following figure shows how the three polar covalent bonds of boron trifluoride molecules can cancel each other out. The molecule has three polar bonds but no overall dipole moment because the molecule has a highly symmetric shape. The polar bonds effectively oppose each other, and the molecule ends up having no overall dipole moment. Molecules will usually not have an overall dipole moment if they have highly symmetric shapes

Chemists usually use dipole arrows to show which direction the bonding pair of electrons is being pulled toward in polar molecules. The “+” sign shows which atom has the partial positive electrostatic charge, and the arrowhead shows which side has the partial negative electrostatic charge. The following figure shows how a single dipole arrow can be used to represent the electric dipole moment that runs along the bond in hydrogen chloride.

We must consider molecular geometries when we are trying to determine the overall polarity of a chemical compound that contains at least one highly electronegative atom and at least one weakly withdrawing atom. Water () and carbon dioxide () are both simple molecular compounds that contain three atoms each, but only one of these molecules has an overall dipole moment because the molecules have very different shapes. Both of these compounds and their dipoles are shown in the image below.

Carbon dioxide is a symmetric molecule, and the two opposing dipole moment arrows cancel each other out. Water has an angular (bent) structure, and the two bond dipole moment arrows reinforce each other, and this produces one large dipole moment that runs from the hydrogen atoms through to the oxygen atom’s lone pair electrons. Water molecules end up having an overall dipole moment, and carbon dioxide molecules end up having an overall dipole moment of zero.

Ammonia is another polar molecule. It has one nitrogen atom bonded to three hydrogen atoms and the molecular formula. Its nitrogen–hydrogen bonds are all polar because they have an absolute difference of electronegativity that equals 0.84. The ammonia molecule has an overall dipole moment because it is asymmetrical. The three arrows of the bond’s dipole moment reinforce each other, and there is a dipole moment from the plane of hydrogen atoms to the plane of lone pair electrons on the nitrogen atom.

The bonds in water are more polar than the bonds in ammonia due to differences in values. The bonds have a more pronounced polarity because they have , and the bonds have . Water molecules have unusual physical properties because they have highly polar bonds. They have high melting points and boiling points primarily because they have highly polar bonds.

The same line of reasoning can be applied to medium-sized simple molecular compounds to determine if they have an overall dipole moment and to calculate the size of any dipole moment that does exist. Carbon tetrachloride () contains one carbon atom that is covalently bonded to four other highly electronegative chlorine atoms. The molecule, illustrated below, has four identical covalent bonds, each of which is a polar bond. Even though carbon tetrachloride has four polar covalent bonds, it does not have an overall dipole moment because the four opposing covalent-bond dipole moment arrows cancel each other out. The molecule does not have an overall dipole moment because of its highly symmetric shape.