LONDON FORCES: CHARACTERISTICS AND EXAMPLES - CHEMISTRY - 2024

The London forces, London dispersion forces or induced-dipole dipole interactions, are the weakest type of intermolecular interactions. Its name is due to the contributions of the physicist Fritz London and his studies in the field of quantum physics.

The London forces explain how molecules interact whose structures and atoms make it impossible for him to form a permanent dipole; that is, it fundamentally applies to apolar molecules or to isolated atoms of noble gases. Unlike the other Van der Waals forces, this one requires extremely short distances.

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A good physical analogy of the London forces can be found in the operation of the Velcro closure system (image above). By pressing one side of the embroidered fabric with hooks, and the other with fibers, an attractive force is created that is proportional to the area of ​​the fabrics.

Once both faces are sealed, a force must be exerted to counteract their interactions (made by our fingers) to separate them. The same is true of molecules: the more voluminous or flat they are, the greater their intermolecular interactions at very short distances.

However, it is not always possible to bring these molecules close enough for their interactions to be noticeable.

When this is so, they require very low temperatures or very high pressures; as such is the case of gases. Likewise, these types of interactions can be present in liquid substances (such as n-hexane) and solid substances (such as iodine).

characteristics

Source: Gabriel Bolívar

What characteristics must a molecule have in order for it to interact using London forces? The answer is that anyone could do it, but when there is a permanent dipole moment, dipole-dipole interactions predominate more than scattering ones, contributing very little to the physical nature of substances.

In structures where there are no highly electronegative atoms or whose electrostatic charge distribution is homogeneous, there is no extreme or region that can be considered rich (δ-) or poor (δ +) in electrons.

In these cases, other types of forces must intervene or otherwise said compounds could only exist in the gas phase, regardless of the pressure or temperature conditions operating on them.

Homogeneous load distribution

Two isolated atoms, such as neon or argon, have a homogeneous charge distribution. This can be seen in A, top image. The white circles in the center represent the nuclei, for atoms, or the molecular skeleton, for molecules. This distribution of charge can be considered as a cloud of green electrons.

Why do noble gases comply with this homogeneity? Because they have a completely full electronic shell, so their electrons should theoretically feel the attractive charge of the nucleus equally in all orbitals.

On the other hand, for other gases, such as atomic oxygen (O), its shell is incomplete (which is observed in its electronic configuration) and forces it to form the diatomic molecule O ₂ to compensate for this deficiency.

The green circles in A can also be molecules, small or large. Its electron cloud orbits around all the atoms that make it up, especially the most electronegative ones. Around these atoms the cloud will become more concentrated and negative, while other atoms will have an electronic deficiency.

However, this cloud is not static but dynamic, so at some point brief δ- and δ + regions will form, and a phenomenon called polarization will occur.

Polarizability

In A the green colored cloud indicates a homogeneous distribution of negative charge. However, the positive attractive force exerted by the nucleus can oscillate on the electrons. This causes a deformation of the cloud thus creating regions δ-, in blue, and δ +, in yellow.

This sudden dipole moment in the atom or molecule can distort an adjacent electron cloud; in other words, it induces a sudden dipole on its neighbor (B, top image).

This is due to the fact that the δ- region disturbs the neighboring cloud, its electrons feel electrostatic repulsion and are oriented to the opposite pole, appearing δ +.

Note how the positive poles align with the negative ones, just as molecules with permanent dipole moments do. The more voluminous the electron cloud, the more difficult the nucleus will keep it homogeneous in space; and in addition, the greater the deformation of it, as can be seen in C.

Therefore, atoms and small molecules are less likely to be polarized by any particle in their environment. An example for this situation is illustrated by the small molecule of hydrogen, H ₂.

In order for it to condense, or even more to crystallize, it needs exorbitant pressures to force its molecules to physically interact.

It is inversely proportional to the distance

Although instantaneous dipoles are formed that induce others around them, they are not enough to hold the atoms or molecules together.

In B there is a distance d that separates the two clouds and their two nuclei. So that both dipoles can remain for a considered time, this distance d must be very small.

This condition, an essential characteristic of London forces (remember the Velcro closure), must be fulfilled in order for it to have a noticeable effect on the physical properties of matter.

Once d is small, the nucleus on the left in B will begin to attract the blue δ- region of the neighboring atom or molecule. This will further deform the cloud, as seen in C (the core is no longer in the center but to the right). Then, there comes a point where both clouds touch and "bounce", but slow enough to hold them together for a while.

Therefore, the London forces are inversely proportional to the distance d. In fact, the factor is equal to d ⁷, so a slight variation in the distance between the two atoms or molecules will weaken or strengthen the London scattering.

It is directly proportional to the molecular mass

How to increase the size of the clouds so that they polarize more easily? Adding electrons, and for that the nucleus must have more protons and neutrons, thus increasing the atomic mass; or, adding atoms to the backbone of the molecule, which in turn would increase its molecular mass

In this way, the nuclei or the molecular skeleton would be less likely to keep the electron cloud uniform all the time. Therefore, the larger the green circles considered in A, B and C, the more polarizable they will be and the greater will also be their interactions by London forces.

This effect is clearly observed between B and C, and could be even more so if the circles were larger in diameter. This reasoning is key to explaining the physical properties of many compounds based on their molecular masses.

Examples of London forces

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In nature

In everyday life there are countless examples of London's dispersal forces without the need to venture into the microscopic world in the first place.

One of the most common and surprising examples is found in the legs of reptiles known as geckos (top image) and in many insects (also in Spiderman).

On their legs they have pads from which thousands of small filaments protrude. In the image you can see a gecko posing on the slope of a rock. To achieve this, it makes use of the intermolecular forces between the rock and the filaments of its legs.

Each of these filaments interacts weakly with the surface on which the small reptile climbs, but since there are thousands of them, they exert a force proportional to the area of ​​its legs, strong enough so that they remain attached and can climb. Geckos are also capable of climbing smooth and perfect surfaces like glass.

Alkanes

Alkanes are saturated hydrocarbons that also interact by London forces. Their molecular structures simply consist of carbons and hydrogens joined by single bonds. Since the difference in electronegativities between C and H is very small, they are apolar compounds.

Thus, methane, CH ₄, the smallest hydrocarbon of all, boils at -161.7ºC. As C and H are added to the skeleton, other alkanes with higher molecular masses are obtained.

In this way, ethane (-88.6ºC), butane (-0.5ºC) and octane (125.7ºC) arise. Notice how their boiling points increase as the alkanes get heavier.

This is because their electronic clouds are more polarizable and their structures have a greater surface area which increases the contact between their molecules.

Octane, although an apolar compound, has a higher boiling point than water.

Halogens and gases

London forces are also present in many gaseous substances. For example, the molecules of N ₂, H ₂, CO ₂, F ₂, Cl ₂ and all the noble gases, interact through these forces, since they present homogeneous electrostatic distribution, which can undergo instantaneous dipoles and lead to polarizations.

The noble gases are He (helium), Ne (neon), Ar (argon), Kr (krypton), Xe (xenon), and Rn (radon). From left to right their boiling points increase with increasing atomic masses: -269, -246, -186, -152, -108, and -62 ºC.

Halogens also interact through these forces. Fluorine is a gas at room temperature just like chlorine. Bromine, with a higher atomic mass, is found under normal conditions as a reddish liquid, and iodine, finally, forms a purple solid that sublimates rapidly because it is heavier than the other halogens.

References

1. Whitten, Davis, Peck & Stanley. Chemistry. (8th ed.). CENGAGE Learning, p 452-455.
2. Angeles Mendez. (May 22, 2012). Dispersion forces (from London). Recovered from: quimica.laguia2000.com
3. London Dispersion Forces. Recovered from: chem.purdue.edu
4. Helmenstine, Anne Marie, Ph.D. (June 22, 2018). 3 Types of Intermolecular Forces. Recovered from: thoughtco.com
5. Ryan Ilagan & Gary L Bertrand. London Dispersion Interactions. Taken from: chem.libretexts.org
6. ChemPages Netorials. London Forces. Recovered from: chem.wisc.edu
7. Kamereon. (May 22, 2013). Geckos: The gecko and the forces of Van der waals. Recovered from: almabiologica.com