Dipole Chemistry

Up until now, you have probably heard that water has many cool properties such as being polar, having cohesive and adhesive forces, and being a great solvent! But, what you ever heard about water being a dipole and wondered what exactly that means? If your answer is yes, you came to the right place!

• First, we will talk about the definition of a dipole and how dipoles are formed.
• Then, we will dive into the different types of dipoles in chemistry and give some examples.

Dipole Definition in Chemistry

Dipoles occur when electrons are shared unequally between atoms in the same molecule due to a high difference in the electronegativity of the atoms involved.

Determination and Formation of a Dipole

The formation of a dipole depends on the polarity of a bond, which is determined by the difference in electronegativity between the two atoms involved in the bond.

Types of Bonds

The three types of bonds you should be familiar with are non-polar covalent bonds, polar covalent bonds, and ionic bonds.

In non-polar covalent bonds, the electrons are equally shared between atoms. In polar covalent bonds, the electrons are shared unequally between atoms. In ionic bonds, the electrons are transferred.

• In ionic bonds, there are no dipoles.
• In polar covalent bonds, dipoles are always present.
• Non-polar covalent bonds do have dipoles but they cancel out due to symmetry.

Predicting Bond Polarity

To determine whether a bond is nonpolar covalent, polar covalent, or ionic, we need to look at the electronegativity values of the atoms involved and calculate the difference between them.

• If the difference in electronegativity is less than 0.4 → non-polar covalent bond
• If the difference in electronegativity falls between 0.4 and 1.7 → polar covalent bond
• If the difference in electronegativity is greater than 1.7 → ionic bond

The electronegativity values are given by Pauling's scale of electronegativity. In the periodic table below, we can see the electronegativity values for each element. Notice the trend here: electronegativity increases from left to right and decreases down a group.

_{Fig.1-Periodic table showing Pauling's scale of electronegativity}

Let's look at an example!

Dipole Moment in Chemistry

To show the dipole moment, we use arrows pointing toward the more electronegative element. For example, in the figure below we can see an HCl and an SO₃ molecule.

• In HCl, chlorine has a higher electronegativity value compared to hydrogen. So, the chlorine will have a partial negative charge and the hydrogen will have a partial positive charge. Since chlorine is more electronegative, the dipole arrow will point towards chlorine.
• In SO₃, the oxygen atom has an electronegativity value higher than that of the sulfur atoms. So, the sulfur atom will have a partial positive charge and the oxygen atoms will have a partial negative charge. In this molecule, the symmetry causes the dipoles to cancel each other out. So, SO₃ has no dipole moment.

Dipole moment of a bond can be calculated by using the following equation: $\mu =Q*\overrightarrow{r}$ where $Q$ is the magnitude of the partial charges δ⁺ and δ^- , and $r$ is the distance vector between the two charges. You can think of the distance vector as an arrow with pointing to the more electron-negative element from the less electron negative one. Dipole moment is measured in Debye units (D). The bigger the dipole moment of the bond, the more polar the bond is.

A dipole moment of a molecule is the sum of the dipole moments of the bonds. This is why it is important that we are using vectors. Vectors have a property called directionality, meaning they point from somewhere to somewhere. You see if two vectors are equally long and point in the opposite direction ( + and -) their sum will be zero. So in theory, if the molecule is perfectly symmetric, meaning all vectors will add up to 0 the dipole moment of the whole molecule will be zero. Okay, let's take a look at an example.

Types of Dipole in Chemistry

The three types of dipole interactions you might encounter are called ion-dipole, dipole-dipole, and induced-dipole induced-dipole (London dispersion forces).

Ion-Dipole

An ion-dipole interaction occurs between an ion and a polar (dipole) molecule. The higher the ion charge, the stronger the ion-dipole attractive force is. An example of ion-dipole is sodium ion in water.

Fig.3-Ion-dipole forces holding sodium ion and water

Another type of interaction involving ions is ion-induced dipole force. This interaction occurs when a charged ion induces a temporary dipole in a non-polar molecule. For example, Fe³⁺ can induce a temporary dipole in O₂, giving rise to an ion-induced dipole interaction!

Dipole-Dipole

When two polar molecules possessing permanent dipoles are near each other, attractive forces called dipole-dipole interactions hold the molecules together. Dipole-dipole interactions are attractive forces that occur between the positive end of a polar molecular and the negative end of another polar molecule. A common example of dipole-dipole forces is seen between HCl molecules. In HCl, the partial positive H atoms get attracted to the partial negative Cl atoms of another molecule.

Fig.4-Dipole-dipole forces between HCl molecules

Hydrogen Bonding

A special type of dipole-dipole interaction is hydrogen bonding. Hydrogen bonding is an intermolecular force that occurs between the hydrogen atom covalently bonded to an N, O, or F and another molecule containing N, O, or F. For example, in water (H₂O), the H atom covalently bonded to oxygen gets attracted to the oxygen of another water molecule, creating hydrogen bonding.

Fig.5-Hydrogen bonding between water molecules

Dipole-induced Dipole forces

Dipole-induced dipole forces arise when a polar molecule with a permanent dipole induces a temporary dipole in a non-polar molecule. For example, dipole-induced dipole forces can hold molecules of HCl and He atoms together.

London dispersion forces

Dipoles can be a little intimidating, but once you get the hang of it you will find it simple!