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In this article, we will be discussing London dispersion forces, the weakest of the forces. We will learn about how these forces work, what properties they have, and what factors affect their strength
- This article covers the topic of London dispersion forces.
- First, we will define London dispersion forces.
- Next, we will look at diagrams to see what is happening at the molecular level.
- Then we will learn about the properties of dispersion forces, and what factors influence them.
- Lastly, we will walk through some examples to solidify our understanding of the topic.
London dispersion forces definition
London dispersion forces are a temporary attraction between two adjacent atoms. One atom's electrons are unsymmetrical, which creates a temporary dipole. This dipole causes an induced dipole in the other atom, which leads to the attraction between the two.
When a molecule has a dipole, its electrons are unevenly distributed, so it has a slightly positive (δ+) and slightly negative (δ-) end. A temporary dipole is caused by the movement of electrons. An induced dipole is when a dipole is formed in response to a nearby dipole.
Attractive forces that exist between neutral molecules are of three types: hydrogen bonding, dipole-dipole forces and London dispersion forces. In particular, London dispersion forces and dipole-dipole forces are types of intermolecular forces which are both included under the general term of van der Waals forces.
Table 1: Types of Intermolecular Interactions:
| Type of Interaction: Intermolecular | Energy Range (kJ/mol) |
| van der Waals (London, dipole-dipole) | 0.1 - 10 |
| Hydrogen Bonding | 10 - 40 |
Hydrogen Bond - attractive force between a strongly electronegative atom, X, bonded to a hydrogen atom, H, and a lone pair of electrons on another small, electronegative atom, Y. Hydrogen bonds are weaker (range: 10 kJ/mol - 40 kJ/mol) than covalent bonds (range: 209 kJ/mol - 1080 kJ/mol) and ionic bonds (range: lattice energy - 600 kJ/mol to 10,000 kJ/mol) but stronger than intermolecular interactions. This type of bond is represented by:
—X—H…Y—
where, the solid dashes, —, represent covalent bonds, and the dots, …, represent a hydrogen bond.
Dipole-dipole Force - an attractive intermolecular force that causes molecules that contain permanent dipoles to align end-to-end, so that the positive end of a given dipole on one molecule interacts with the negative end of a dipole on an adjacent molecule.
Covalent Bond - a chemical bond in which electrons are shared between atoms.
Electronegativity - a measure of the ability of a given atom to attract electrons to itself.
To better understand these definitions, let's look at some diagrams.
London dispersion forces diagram
London dispersion forces are due to two types of dipoles: temporary and induced.
Let's start by looking at what happens when a temporary dipole is formed.
Fig. 2: The movement of electrons leads to a temporary dipole. StudySmarter Original.
Electrons in an atom are constantly in motion. On the left, electrons are evenly/symmetrically distributed. As the electrons move, they will occasionally be asymmetrical, which leads to a dipole. The side with more electrons will have a slightly negative charge, while the side with fewer electrons will have a slightly positive charge. This is considered a temporary dipole, since the motion of electrons leads to a constant shift between symmetrical and asymmetrical distributions, so the dipole won't last long.
Now onto the induced dipole:
Fig. 3: The temporary dipole causes an induced dipole in a neutral molecule. StudySmarter Original.
The temporary dipole approaches another atom/molecule that has an even distribution of electrons. The electrons in that neutral atom/molecule will be drawn toward the slightly positive end of the dipole. This movement of electrons causes an induced dipole.
An induced dipole is technically the same as a temporary dipole, except one is "induced" by another dipole, hence the name. This induced dipole is also temporary, since moving the particles away from each other will make it disappear, since the attraction isn't strong enough.
London dispersion forces properties
London dispersion forces have three main properties:
- Weak (The weakest of all the forces between molecules)
- Caused by temporary electron imbalances
- Present in all molecules (polar or non-polar)
London dispersion forces factors
There are three factors that affect the strength of these forces:
- Size of the molecules
- Shape of the molecules
- Distance between the molecules
The size of a molecule is related to its polarizability.
Polarizability describes how easily electron distribution can be disturbed within a molecule.
Isomers are molecules that have the same chemical formula, but different molecular geometry.
Fig. 4: Neopentane is less "accessible" so it's a gas, while n-pentane is more accessible, so it is a liquid. StudySmarter Original.
London dispersion forces examples
Now that we've learned all about London dispersion forces, it's time to work on some example problems!
Which of the following will have the strongest dispersion forces?
a) He
b) Ne
c) Kr
d) Xe
The main factor here is size. Xenon (Xe) is the largest of these elements, so it will have the strongest forces.
For comparison, their boiling points (in order) are -269 °C, -246 °C, -153° C, -108° C. As the elements get bigger, their forces are stronger, so they are closer to being liquids than those that are smaller.
Between the two isomers, which has the stronger dispersion forces?
Fig. 5: C6H12 isomers. StudySmarter Original.
Since these are isomers, we need to focus on their shape. If we were to put an atom at each of their points-of-contact, it would look like this:
Fig. 6: Cyclohexane has more points of contact. StudySmarter Original.
Based on this, we can see that cyclohexane has more points of contact. This means that it has the stronger dispersion forces.
For reference, cyclohexane has a boiling point of 80.8 °C, while 4-methyl-1-pentene has a boiling point of 54 °C. This lower boiling point suggests it is weaker, since it more likely to go into the gas phase than cyclohexane.
London Dispersion Forces - Key takeaways
- London dispersion forces are a temporary attraction between two adjacent atoms. One atom's electrons are unsymmetrical, which creates a temporary dipole. This dipole causes an induced dipole in the other atom, which leads to attraction between the two.
- When a molecule has a dipole, its electrons are unevenly distributed, so it has a slightly positive (δ+) and slightly negative (δ-) end. A temporary dipole is caused by the movement of electrons. An induced dipole is when a dipole is formed in response to a nearby dipole.
- Dispersion forces are weak and present in all molecules
- Polarizability describes how easily electron distribution can be disturbed within a molecule.
- Isomers are molecules that have the same chemical formula, but a different orientation.
- Molecules that are larger and/or have more points of contact have stronger dispersion forces.
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Frequently Asked Questions about London Dispersion Forces
What are london dispersion forces?
London dispersion forces are a temporary attraction between two adjacent atoms. One atom's electrons are unsymmetrical, which creates a temporary dipole. This dipole causes an induced dipole in the other atom, which leads to attraction between the two.
What does London dispersion force depend on?
London Dispersion Forces depend on the weight and shape of molecules.
Why is London dispersion the weakest force?
They are the weakest because for a very brief second they are dipoles, which means that, there is a partially positive element interacting with a partially negative element, making it easy to disrupt them.
Which has the strongest London dispersion force?
Iodine molecules
How do you know if a molecule has London dispersion forces?
ALL molecules have it
What are london dispersion forces?
A temporary attraction between two adjacent atoms. One atom's electrons are unsymmetrical, which creates a temporary dipole. This dipole causes an induced dipole in the other atom, which leads to the attraction between the two.
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