Have you ever tried to force two same-charged magnets together? Even with all your might, magnets of the same charge will always repel, even if you try to squeeze them together. However, magnets with opposite poles can snap together easily.
Magnets aren't the only things that can do this, molecules can too! Molecules with a dipole can attract or repel other molecules. These forces are very important in molecules as they can influence certain properties like boiling point.
In this article, we will be learning about dipoles and dipole moments to see how they are formed and why they are so important.
- This article is about dipole moment.
- First, we will define dipole and dipole moment.
- Next, we will learn how to determine the direction of a dipole.
- Then, we will learn about the dipole moment for water and why dipole moments don't occur in symmetrical molecules.
- Thereafter, we will cover the dipole moment formula and learn about the trends in dipole moment.
Dipole Moment Definition
Let's start by looking at the definition of dipole moment.
When two charges of opposite sign are separated by a distance, they create a dipole.
The dipole moment (μ) measures the size of a molecule's dipole.
While temporary dipoles can be formed in an atom or a molecule, it is only molecules that can have a permanent dipole (these are called polar molecules).
In a molecule, a dipole is established when there is an uneven distribution of electrons, caused by a difference in electronegativity.
Electronegativity is the tendency for an atom to pull electrons/electron density towards itself.
The most electronegative atom is fluorine, so the closer an atom is to the top right of the periodic table (i.e. the closer to fluorine), the more electronegative it is. The closer an atom is to the bottom left of the periodic table, the less electronegative it is.
When an atom is more electronegative, it will pull electron density towards itself, so that side of the molecule will be partially negative (δ-). The other side of the molecule will then be partially positive (δ+).Below is an example of a polar molecule (HF):
Fig.1-The polar molecule HF
Fluorine is more electronegative (3.98), so it has a δ- charge, while hydrogen is less electronegative (2.2), so it has a δ+ charge.There are two important things to remember when it comes to dipoles:1) Symmetrical molecules "cancel" any polar bonds, so they do not have a dipole. Essentially, the dipoles are equal and opposite, so they cancel (see Figure 4 for details).2) Atoms must have at least a 0.5 difference in electronegativity to be considered polar/have a dipole.
Dipole Moment Direction
A key fact about a dipole moment is that it is a vector, meaning it has both a magnitude and direction. The dipole moment shows the net direction of the charge.
For example, here is the dipole moment for HF:
Fig.2-The dipole moment of HF points toward fluorine
The dipole moment points towards the area of the molecule where there is the greatest amount of electron density (i.e. the more electronegative atom). The "crossed" portion of the arrow marks where there is a lack of electron density (i.e. the less electronegative atom).
In this case, the arrow is pointing towards fluorine (most EN element) and the crossed portion is at hydrogen.
Dipole Moment of Water
When a molecule has more than one bond, the net dipole moment is the total value of all the bond dipole moments. For example, here is the dipole moment of water:
Fig.3 Water has a net dipole moment facing downwards
Individual bond dipole moments are always parallel to their bonds, while the net dipole moment doesn't have to follow a bond, but will always point towards the more electronegative atom.
The magnitude of the net dipole moment is not simply the sum of the individual bond dipole moments. As we'll see later when we look at the formula, the net dipole is the sum of the dipoles in each coordinate direction (i.e. the x and y directions).
As an example, the individual dipole moment of the O-H bond is 1.5 D, while the net dipole moment of water is 1.84 D.
Dipole Moment for Symmetrical Molecules
As mentioned previously, molecules that have symmetrical/opposite dipoles don't have a dipole/dipole moment. This is because the bond polarities cancel each other, so there is no net dipole.
As an example, let's look at BeCl2:
Fig.4 The opposite bond dipole moments in BeCl2 cancel out
There is a significant difference between the electronegativity of beryllium and chlorine (it is 1.5), so the individual bonds do have a dipole moment. However, since the bond dipole moments have equal magnitude but point in exactly opposite directions, there is no net dipole moment.
This is why molecules like BeCl2 are considered non-polar, even though the difference in electronegativity is large.
Note: While a molecule like water has symmetry, it still has a dipole moment. The symmetry we are focusing on is the symmetry of the dipoles. In water, the dipoles are pointing in the same direction, while in BeCl2 they are in opposite directions, which is why they cancel.
Dipole Moment Formula
The formula for dipole moment is as follows:
$$\vec{\mu}=\Sigma_{i} q_{i} \vec{r_i}$$
Where,
\(\vec{\mu}\) is the dipole moment vector.
qi is the magnitude of the ith charge.
\(\vec{r_i}\) is the position vector of the ith charge.
This is the more complex version of the formula used for molecules that have more than two atoms. Since dipole moment is a vector quantity, you need to calculate it for each coordinate axis (x, y).
There is a simpler form of this formula that we use just for diatomic molecules:
$$\mu_{diatomic}=Q*r$$
Where,
- Q is the magnitude of the charge
- r is the bond length (distance) between charges.
The dipole moment is measured in Debye units (D), where 1 D=3.33564 x10-30 Cm (Coulomb meters).
While you won't need to use this formula for calculations at the AP level, it's helpful to know the formula, so you can understand the trends in dipole moment.
Trends in Dipole Moment
There are three factors affecting dipole moment. These are:
Distance
Polarity
Geometry of the molecule
It's important to note that these factors are working in tandem, i.e. all of these factors need to be taken into account. Just because the bond distance is greater, it doesn't automatically mean it has a greater dipole moment. For example, here is the trend in dipole moment as you move down the halogens (group 17)
| Compound | Bond length (Â) | Difference in Electronegativity | Dipole Moment (D) |
| HF | 0.92 | 1.9 | 1.82 |
| HCl | 1.27 | 0.9 | 1.08 |
| HBr | 1.41 | 0.7 | 0.82 |
| HI | 1.61 | 0.4 | 0.44 |
As you can see, even though the bond length is increasing, the polarity is decreasing. The decrease in polarity is greater than the increase in bond length, so the dipole moment decreases.
Knowing the dipole moment can be very important for determining different chemical properties. For example, here are some properties that dipole moment can affect:
Solubility
Boiling/Melting point
Stability/Reactivity
Dipole Moment - Key takeaways
- When two charges of equal magnitude and opposite signs are separated by a distance, they create a dipole.
- The dipole moment (μ) measures the size of a molecule's dipole.
- Electronegativity is the tendency for an atom to pull electrons/electron density towards itself.
- The dipole moment vector points towards the more electronegative atom.
- The net dipole moment is a sum of the individual bond dipoles.
- Symmetrical molecules don't have a dipole moment, since the individual bond dipoles cancel.
- The formula for dipole moment is:
$$\vec{\mu}=\Sigma_{i} q_{i} \vec{r_i}$$ ,where \(\vec{\mu}\) is the dipole moment vector, \(q_{i}\)is the magnitude of the ith charge and \(\vec{r_i}\) is the position vector of the ith chargeWith a simpler formula being: $$\mu_{diatomic}=Q*r$$ Where Q is the magnitude of the charge and r is the bond length (distance) between charges.
Dipole moment is affected by: distance, polarity, and molecule geometry. Dipole moment can influence properties such as boiling/melting point, solubility, and stability.
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