4.3: Polarity of Bonds and Molecules

Learning Objectives

• Explain how polar compounds differ from nonpolar compounds.
• Determine if a molecule is polar or nonpolar.
• Given a pair of compounds, predict which would have a higher melting or boiling point.

Bond Polarity

A covalent bond forms when atoms share electrons. However, elements do not all have the same relative attraction for electrons when they are a part of a compound.  This relative attraction for the electrons involved in a bond is known as electronegativity. Different scientists assigned electronegativity values based on different properties. The values in the different tables are not identical but follow the same trends across periods and down groups. Linus Pauling developed the first measurable scale for electronegativity and based the values on covalent bond energy calculations.

Differences in electronegativity values can be used to predict the primary nature of the bond ranging from pure covalent to ionic.  Polar covalent bonds lie on the bonding continuum between pure covalent and ionic bonds. The larger the difference in electronegativity values, the greater the attraction for the electrons by one of the atoms, the more polar the bond.

The polarity of a covalent bond can be judged by determining the difference of the electronegativities of the two atoms involved in the covalent bond, as summarized in the following table:

Electronegativity Difference Bond Type
0–0.4 pure covalent
0.5–2.0 polar covalent
>2.0 likely ionic

Nonpolar Covalent Bonds

A bond in which the electronegativity difference is less than 1.9 is considered to be mostly covalent in character. However, at this point we need to distinguish between two general types of covalent bonds. A nonpolar covalent bond is a covalent bond in which the bonding electrons are shared equally between the two atoms. In a nonpolar covalent bond, the distribution of electrical charge is balanced between the two atoms.

The two chlorine atoms share the pair of electrons in the single covalent bond equally, and the electron density surrounding the $$\ce{Cl_2}$$ molecule is symmetrical. Also note that molecules in which the electronegativity difference is very small (<0.5) are also considered nonpolar covalent. An example would be a bond between chlorine and bromine ($$\Delta$$EN $$=3.0 - 2.8 = 0.2$$).

Polar Covalent Bonds

A bond in which the electronegativity difference between the atoms is between 0.5 and 2.0 is called a polar covalent bond. A polar covalent bond is a covalent bond in which the atoms have an unequal attraction for electrons and so the sharing is unequal. In a polar covalent bond, sometimes simply called a polar bond, the distribution of electrons around the molecule is no longer symmetrical. The ΔEN difference of 2.0 as the upper limit between polar covalent and ionic is arbitrary rather than an absolute cut off and that the properties of the compound are the best indicator of the primary nature of the bond.  The differences in electronegativity are most valuable when used to predict the relative polarity of covalent bonds.

An easy way to illustrate the uneven electron distribution in a polar covalent bond is to use the Greek letter delta $$\left( \delta \right)$$.

The atom with the greater electronegativity acquires a partial negative charge, while the atom with the lesser electronegativity acquires a partial positive charge. The delta symbol is used to indicate that the quantity of charge is less than one. A polar covalent bond results in electrical charges separated by a distance so it has a measurable dipole moment units of Coulomb-meters, C-m, or more commonly expressed in Debeye, D (1 D = 3.336 x 10-30 C-m).   Dipole moments are vectors – quantities that have both magnitude and direction. An alternative notation to indicate a bond is polar and which end of the bond is partially positive and which is partially negative is to draw a vector (an arrow) with the arrow head pointing in the direction of the partially negative end of the bond. A vertical line is added to the end of the vector by the partially positive end of the bond so it appears to be a plus sign. The length of the arrow can be used to illustrate the relative polarity of the bond.

Electronegativity differences in bonding using Pauling scale. Using differences in electronegativity to classify bonds as covalent, polar covalent, or ionic.

Example $$\PageIndex{1}$$: Bond Polarity

What is the polarity of each bond?

1. C–H
2. O–H

Solution

Using Figure $$\PageIndex{1}$$, we can calculate the difference of the electronegativities of the atoms involved in the bond.

1. For the C–H bond, the difference in the electronegativities is 2.5 − 2.1 = 0.4. Thus we predict that this bond will be non polar covalent.
2. For the O–H bond, the difference in electronegativities is 3.5 − 2.1 = 1.4, so we predict that this bond will be definitely polar covalent.

Exercise $$\PageIndex{1}$$

What is the polarity of each bond?

1. Rb–F
2. P–Cl

likely ionic

polar covalent

Molecular Polarity

To determine if a molecule is polar or nonpolar, it is frequently useful to look at Lewis structures. Nonpolar compounds will be symmetric. For molecules with four or fewer total electron groups around the central atom, a symmetrical molecule is identical on all sides – the bonded atoms are identical and there are no unshared electrons on the central atom. Polar molecules are asymmetric, either containing lone pairs of electrons on a central atom or having atoms with different electronegativities bonded. This works pretty well - as long as you can visualize the molecular geometry. That's the hard part. To know how the bonds are oriented in space, you have to have a strong grasp of Lewis structures and VSEPR theory. Assuming you do, you can look at the structure of each one and decide if it is polar or not - whether or not you know the individual atom electronegativity. This is because you know that all bonds between dissimilar elements are polar, and in these particular examples, it doesn't matter which direction the dipole moment vectors are pointing (out or in).

A polar molecule is a molecule in which one end of the molecule is slightly positive, while the other end is slightly negative. A molecule with two poles is called a dipole and has a net dipole moment. A diatomic molecule that consists of a polar covalent bond, such as $$\ce{HF}$$, is a polar molecule. The two electrically charged regions on either end of the molecule are called poles, similar to a magnet having a north and a south pole. A molecule with two poles is called a dipole (see figure below). Hydrogen fluoride is a dipole.

For molecules with more than two atoms, the molecular geometry must also be taken into account when determining if the molecule is polar or nonpolar. The figure below shows a comparison between carbon dioxide and water. Carbon dioxide $$\left( \ce{CO_2} \right)$$ is a linear molecule. The oxygen atoms are more electronegative than the carbon atom, so there are two individual dipoles pointing outward from the $$\ce{C}$$ atom to each $$\ce{O}$$ atom. However, since the dipoles are of equal strength and are oriented this way, they cancel out and the overall molecular polarity of $$\ce{CO_2}$$ is zero.

Water is a bent molecule because of the two lone pairs on the central oxygen atom. The individual dipoles point from the $$\ce{H}$$ atoms toward the $$\ce{O}$$ atom. Because of the shape, the dipoles do not cancel each other out and the water molecule is polar. In the figure below, the net dipole is shown in blue and points upward.

Some other molecules are shown in the figure below. Notice that a tetrahedral molecule such as $$\ce{CH_4}$$ is nonpolar. However, if one of the peripheral $$\ce{H}$$ atoms is replaced with another atom that has a different electronegativity, the molecule becomes polar. A trigonal planar molecule $$\left( \ce{BF_3} \right)$$ may be nonpolar if all three peripheral atoms are the same, but a trigonal pyramidal molecule $$\left( \ce{NH_3} \right)$$ is polar.

Steps to Identify Polar Molecules

1. Draw the Lewis structure
2. Figure out the geometry (using VSEPR theory)
3. Visualize or draw the geometry
4. Find the net dipole moment (you don't have to actually do calculations if you can visualize it)
5. If the net dipole moment is zero, it is non-polar. Otherwise, it is polar.

Properties of Polar Molecules

Polar molecules tend to align when placed in an electric field with the positive end of the molecule oriented toward the negative plate and the negative end toward the positive plate (Figure $$\PageIndex{14}$$). We can use an electrically charged object to attract polar molecules, but nonpolar molecules are not attracted. Also, polar solvents are better at dissolving polar substances, and nonpolar solvents are better at dissolving nonpolar substances.

While molecules can be described as "polar covalent" or "ionic", it must be noted that this is often a relative term, with one molecule simply being more polar or less polar than another. However, the following properties are typical of such molecules. Polar molecules tend to:

• have higher melting points than nonpolar molecules
• have higher boiling points than nonpolar molecules
• be more soluble in water (dissolve better) than nonpolar molecules
• have lower vapor pressures than nonpolar molecules

Example $$\PageIndex{2}$$:

Label each of the following as polar or nonpolar.

1. Water, H2O:
2. Methanol, CH3OH:
3. Hydrogen Cyanide, HCN:
4. Oxygen, O2:
5. Propane, C3H8:

Solution

1. Water is polar. Any molecule with lone pairs of electrons around the central atom is polar.
2. Methanol is polar. This is not a symmetric molecule. The $$\ce{-OH}$$ side is different from the other 3 $$\ce{-H}$$ sides.
3. Hydrogen cyanide is polar. The molecule is not symmetric. The nitrogen and hydrogen have different electronegativities, creating an uneven pull on the electrons.
4. Oxygen is nonpolar. The molecule is symmetric. The two oxygen atoms pull on the electrons by exactly the same amount.
5. Propane is nonpolar, because it is symmetric, with $$\ce{H}$$ atoms bonded to every side around the central atoms and no unshared pairs of electrons.

Exercise $$\PageIndex{2}$$

Label each of the following as polar or nonpolar.

a. SO3

b. NH3

non polar