WHAT ARE DIPOLE DIPOLE FORCES? - CHEMISTRY - 2024

The dipole dipole forces or Keesom forces are those intermolecular interactions present in molecules with permanent dipole moments. It is one of the Van der Waals forces and, although it is far from the strongest, it is a key factor that explains the physical properties of many compounds.

The term "dipole" explicitly refers to two poles: one negative and one positive. Thus, we speak of dipole molecules when they have defined regions of high and low electron densities, which is only possible if the electrons preferentially "migrate" towards certain atoms: the most electronegative.

The upper image illustrates the dipole-dipole interactions between two AB molecules with permanent dipole moments. Likewise, it can be observed how the molecules are oriented so that the interactions are efficient. Thus, the positive region δ + attracts the negative region δ-.

According to the above, it can be specified that this type of interaction is directional (unlike ionic charge-charge interactions). The molecules in their environment orient their poles in such a way that, although they are weak, the sum of all these interactions gives the compound great intermolecular stability.

This results in that compounds (organic or inorganic) capable of forming dipole-dipole interactions exhibit high boiling or melting points.

Dipole moment

The dipole moment µ of a molecule is a vector quantity. In other words: it depends on the directions where there is a polarity gradient. How and why does this gradient originate? The answer lies in the bonds and in the intrinsic nature of the atoms of the elements.

For example, in the upper image A is more electronegative than B, so in the AB bond the highest electron density is located around A.

On the other hand, B “gives up” its electron cloud and, therefore, is surrounded by an electron-poor region. This difference in electronegativities between A and B creates the polarity gradient.

As one region is rich in electrons (δ-) while the other is poor in electrons (δ +), the two poles appear, which, depending on the distances between them, originate different magnitudes of µ, which are determined for each compound.

Symmetry

If a molecule of a certain compound has µ = 0, then it is said to be an apolar molecule (even if it has polarity gradients).

To understand how symmetry - and hence molecular geometry - plays an important role in this parameter, it is necessary to consider the AB bond again.

Due to the difference in their electronegativities, there are defined regions rich and poor in electrons.

What if the links were AA or BB? In these molecules there would be no dipole moment, since both atoms attract towards them in the same way the electrons of the bond (one hundred percent covalent bond).

As can be seen in the image, neither the AA nor the BB molecule now show regions rich or poor in electrons (red and blue). Here another type of forces is responsible for holding A ₂ and B _{2 together}: induced dipole-dipole interactions, also known as London forces or dispersion forces.

On the contrary, if the molecules were of the AOA or BOB type, there would be repulsions between their poles because they have equal charges:

The δ + regions of two BOB molecules do not allow efficient dipole-dipole interaction; the same happens for the δ- regions of the two AOA molecules. Likewise, both pairs of molecules have µ = 0. The polarity gradient OA is vectorially canceled with that of the bond AO.

Consequently, dispersion forces also come into play in the AOA and BOB pair, due to the absence of effective orientation of the dipoles.

Asymmetry in nonlinear molecules

The simplest case is that of the CF ₄ molecule (or the CX ₄ type). Here, C has a tetrahedral molecular geometry and electron-rich regions are found at the vertices, specifically on the electronegative atoms of F.

The polarity gradient CF cancels in any of the directions of the tetrahedron, causing the vector sum of all these to equal 0.

Thus, although the center of the tetrahedron is very positive (δ +) and its vertices very negative (δ-), this molecule cannot form dipole-dipole interactions with other molecules.

Orientations of the dipoles

In the case of linear AB molecules, they are oriented in such a way that they form the most efficient dipole-dipole interactions (as shown in the image above). The above is applicable in the same way for other molecular geometries; for example, angular ones in the case of NO ₂ molecules.

Thus, these interactions determine whether compound AB is a gas, a liquid, or a solid at room temperature.

In the case of compounds A ₂ and B ₂ (those of the purple ellipses), it is very likely that they are gaseous. However, if its atoms are very bulky and easily polarizable (which increases the London forces), then both compounds can be either solid or liquid.

The stronger the dipole-dipole interactions, the greater the cohesion between the molecules; likewise, the higher the melting and boiling points of the compound. This is because higher temperatures are needed to "break" these interactions.

On the other hand, an increase in temperature causes the molecules to vibrate, rotate and move more frequently. This "molecular agitation" impairs the orientations of the dipoles and therefore the intermolecular forces of the compound are weakened.

Hydrogen bond interactions

In the upper image five water molecules are shown interacting by hydrogen bonds. This is a special type of dipole-dipole interactions. The electron-poor region is occupied by H; and the electron-rich region (δ-) is occupied by the highly electronegative atoms N, O, and F.

That is, molecules with N, O, and F atoms bonded to H can form hydrogen bonds.

Thus, the hydrogen bonds are OHO, NHN and FHF, OHN, NHO, etc. These molecules have permanent and very intense dipole moments, which orient them correctly to "take advantage" of these bridges.

They are energetically weaker than any covalent or ionic bond. Although, the sum of all the hydrogen bonds in the phase of a compound (solid, liquid or gaseous) makes it exhibit properties that define it as unique.

For example, such is the case of water, whose hydrogen bonds are responsible for its high boiling point and for being less dense in the ice state than liquid water; reason why icebergs float in the seas.

References

1. Dipole-Dipole Forces. Retrieved on May 30, 2018, from: chem.purdue.edu
2. Boundless Learning. Dipole-Dipole Force. Retrieved on May 30, 2018, from: courses.lumenlearning.com
3. Jennifer Roushar. (2016). Dipole-Dipole Forces. Retrieved on May 30, 2018, from: sophia.org
4. Helmenstine, Anne Marie, Ph.D. (May 3, 2018). What Are Examples of Hydrogen Bonding? Retrieved on May 30, 2018, from: thoughtco.com
5. Mathews, CK, Van Holde, KE and Ahern, KG (2002) Biochemistry. Third edition. Addison Wesley Longman, Inc., P 33.
6. Whitten, Davis, Peck & Stanley. Chemistry. (8th ed.). CENGAGE Learning, p 450-452.
7. User Qwerter. (April 16, 2011). 3D model hydrogen bonds in toilet.. Retrieved on May 30, 2018, from: commons.wikimedia.org