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.
A dipole is a molecule or covalent bond that has a separation of charges.
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.
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!
Predict the type of bond polarity between the following atoms:
a) H and Br
H has an EN value of 2.20 and Br has an EN of 2.96. The electronegativity difference between these atoms is 0.76 so it would have a polar covalent bond.
b) Li and F
Li has an EN value of 0.98 and F has an EN of 3.98. The electronegativity difference is 3.00 so it would have an ionic bond.
c) I and I
I has an EN value of 2.66. The electronegativity difference is 0.00 so it would have a non-polar covalent bond.
Dipole Moment in Chemistry
To measure the separation of charges in a molecule we use dipole moment. Dipole moments are present in polar molecules that have asymmetric shapes because, in asymmetric shapes, the dipoles do not cancel out.
Dipole moment is referred to as a measurement of the magnitude of a dipole.
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: where is the magnitude of the partial charges δ+ and δ- , and 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.
You can learn more about the different molecular shapes by reading "Valence Shell Electron Pair Repulsion (VSEPR) Theory.
Which of the following compounds has a dipole moment? PCl3 or PCl5?
First, we need to take a look at their lewis structures. If the structure is symmetrical, then the dipoles will cancel out and the compound will not have a dipole.
In PCl3, the bond is polar because of the difference in electronegativity between P and Cl atoms, and the presence of a lone pair of electrons gives PCl3 a tetrahedral structure.
On the other hand, PCl5 is considered non-polar because its symmetrical shape, which is trigonal bipyramidal, cancels the dipoles out.
Fig.2-Lewis diagrams of phosphorus trichloride and phosphorus pentachloride
If you need to go back and learn how to draw Lewis structures, check out "Lewis Diagrams".
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!
So what does it mean to induce a dipole? If you put an ion near a non-polar molecule, you can start affecting its electrons. For example, a positive ion will attract these electrons to the side on which the ion is. This will create a larger concentration of ions there and lead to a dipole forming on the originally non-polar molecule.
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
Induced-dipole Induced-dipole interactions are also known as London dispersion forces. This type of interaction is present in all molecules, but it is most important when dealing with non-polar molecules. London dispersion forces occur because of the random movement of electrons in the cloud of electrons. This movement produces a weak, temporary dipole moment! For example, London dispersion forces are the only type of attractive force holding F2 molecules together.
Examples of Dipoles in Chemistry
Now that you have a better understanding of what dipoles are, let's look at more examples! If the figure below you can see the structure of acetone. Acetone, C3H6O, is a polar molecular with a bond dipole.
Fig.6-Dipoles in Acetone
Another common example of a molecule containing dipoles is carbon tetrachloride, CCl4. Carbon tetrachloride is a non-polar molecule that contains polar bonds, and therefore, has dipoles present. However, the net dipole is zero due to its tetrahedral structure, where the bond dipoles directly oppose each other.
Fig.7-Structure of Carbon Tetrachloride
Let's look at one last example!
What is the net dipole moment in CO2?
CO2 is a linear molecule that has two C=O bond dipoles equal in magnitude but pointing in opposite directions. Therefore, the net dipole moment is zero.
Fig.8-Dipoles in Carbon Dioxide
Dipoles can be a little intimidating, but once you get the hang of it you will find it simple!
Dipoles - Key takeaways
- Dipoles occur when electrons are shared unequally between atoms due to a high difference in the electronegativity of the atoms involved.
- A dipole moment is referred to as a measurement of the magnitude of a dipole.
- Dipole moments are present in polar molecules that have asymmetric shapes because, in asymmetric shapes, the dipoles do not cancel out.
- Types of dipoles include ion-dipole, dipole-dipole, and induced-dipole induced-dipole (London dispersion forces).
Saunders, N. (2020). Supersimple Chemistry: The Ultimate Bitesize Study Guide. London: Dorling Kindersley.
Timberlake, K. C. (2019). Chemistry: An introduction to general, organic, and Biological Chemistry. New York, NY: Pearson.
Malone, L. J., Dolter, T. O., & Gentemann, S. (2013). Basic concepts of Chemistry (8th ed.). Hoboken, NJ: John Wiley & Sons.
Brown, T. L., LeMay, H. E., Bursten, B. E., Murphy, C. J., Woodward, P. M., Stoltzfus, M., & Lufaso, M. W. (2018). Chemistry: The central science (13th ed.). Harlow, United Kingdom: Pearson.
References
- Fig.1-Periodic table showing Pauling's scale of electronegativity (https://upload.wikimedia.org/wikipedia/commons/thumb/4/42/Electronegative.jpg/640px-Electronegative.jpg) by ad blocker on wikimedia commons licensed by CC By-SA 3.0 (https://creativecommons.org/licenses/by-sa/3.0/)
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