Dipole-Dipole Forces are a special form of Van der Waals forces, which you already got to know in our Original. The van der Waals forces in the broadest sense are electrostatic effects caused by the interact two dipoles occurrence. In this article you will find out what the dipole-dipole forces are and how you can estimate them.
Formation of dipole molecules
Like you from the orbital theory maybe remember, form two atomic orbitals one Covalent bond out, i.e. a new one is created binding and a antibonding molecular orbital. The bonding electrons that hold the resulting molecule together are then located in the bonding orbital. As you may also remember, orbitals describe a space in which the electrons are located with a certain statistical probability.
Over time, these bonding orbitals become unequally distributed, causing the centers of gravity no longer fall on top of each other in a molecule. In this case, the chemist speaks of the fact that different partial charges appear. A dipole molecule is created temporarily or spontaneously. This case is typical for non-polar or weakly polar covalent bonds.
In the case of polar covalent bonds, from an electronegativity difference of about 0.4, permanent dipoles are formed if the centers of charge in the molecule are distributed asymmetrically, i.e. do not coincide.
Electronegativity is the measure of an atom’s ability to attract the electron cloud in a bond towards itself. The greater the electronegativity difference between two atoms in a molecule, the stronger the polar/ionic character of the bond.
In a dipole molecule there are different centers of charge, also called partial charges, which do not coincide. These can be caused by spontaneous electron displacement or are permanent due to strong electronegativity differences between the atoms in the molecule. These molecules are then spontaneous/temporary or permanent dipole called.
The word dipole comes from the Greek and literally means two ends. It is therefore a molecule with differently charged ends.
If you compare the two molecules carbon dioxide (CO2) and sulfur dioxide (SO2), you will notice that the atomic bonds are polar. The difference lies in the spatial structure of the molecules. While the CO2 has a linear structure, the SO2 has an angled structure, which is why the centers of charge do not coincide here. Consequently, sulfur dioxide is a permanent dipole, but carbon dioxide is not.
dipole moment
The dipole-dipole interaction
An important aspect to keep in mind is that these are not bindings, ie intramolecular interactions, but about intermolecular interactions. These forces are significantly weaker than the classic types of binding, which is why they deserve a separate term.
Intramolecular means «within the molecule» and intermolecular means «between molecules».
Dipole-dipole forces, or dipole-dipole interactions, always occur between two permanent dipoles that are close to each other. They are also temperature dependent. JThe stronger the polarity within the molecule, the stronger the dipole-dipole forces are for molecules of comparable size and mass.
The reason for the temperature dependence is that the kinetic energy of molecules increases with increasing temperature, making it less likely that the particles will approach each other. The dipole-dipole forces have an influence on the material properties, more on this later.
As already mentioned, permanent dipoles are based on the electronegativity difference between the involved ones bonding partners. The stronger this is, the more polar the bond is and the stronger a dipole is, as long as the centers of charge are distributed asymmetrically. You can use the electronegativity to estimate the strength of a dipole and the interactions.
If two such dipoles approach each other, they align the oppositely charged ends towards each other. This saves energy for the particles involved and is therefore a preferred state. This effect is also called directional effect. How strong this interaction is depends not only on the dipole character of the particles themselves, but also on the distance between the two molecules.
This effect is comparable to the approach of two bar magnets from physics, in which the oppositely charged poles align with each other, resulting in a stronger attraction.
The interaction energy of two particles, as described by the physicist Fritz London, decreases with the sixth power of the distance. If the distance between the molecules is doubled, the interaction is weaker by a factor of 26 = 64. The range of the dipole-dipole interaction is therefore limited. The following equation shows you the interaction energy of two different particles:
The Factors α1,2 stand for the polarizability of the particles and I1,2 for the ionization energies.
The dipole-dipole interactions are stronger than pure van der Waals forces, but weaker than hydrogen bonds. All these interactions are weaker than the primary bonds.
Special case: The hydrogen bond
The so-called hydrogen bridges are a special form of dipole-dipole forces. They only occur between molecules in which hydrogen on a strongly electronegative element, such as fluorine, oxygen or nitrogen. This is a particularly strong interaction between the particles involved.
However, you must remember that the hydrogen bond, like other interactions, is not rigid in the sense that the structure of the molecules never changes once it is established. On the contrary, these «bonds» are constantly breaking and new ones being formed, so it’s an ongoing process.
The most common example of hydrogen bridges is water itself. It is tetrahedral in structure because both are fully occupied nonbonding orbitals of oxygen but compressed. The bond angle is then only about 104.5° instead of 109.5°. Water is therefore also a permanent dipole, because all criteria are met.
Hydrogen bonds are usually arranged in a linear fashion so as to maximize the attraction between hydrogen and the electronegative element while minimizing the repulsion between the electronegative atoms. As a result, most hydrogen bonds are unsymmetrical. The hydrogen bonds are usually represented by dashed lines. Others, like water, have a symmetrical structure and are therefore much more stable.
A water molecule can be surrounded by four other water molecules at the same time and form hydrogen bonds with them. The attraction in the water is therefore very strong, which is also reflected in the material properties. With a boiling point of 100°C, this is significantly higher than that of comparable molecules. At the ammonia the boiling point is -33°C, i.e. it is already gaseous at room temperature.
One reason for this is that nitrogen (3.04) is less electronegative than oxygen (3.44), resulting in weaker hydrogen bonds. In addition, ammonia has the disadvantage that it is limited to a free pair of electrons on the nitrogen, which is why fewer hydrogen bonds can form as a result than in water. Thus, the water is inherently more stable than ammonia and requires a higher input of energy to break the hydrogen bonds.
Effects of dipole-dipole forces on material properties
As already mentioned, the influence of dipole-dipole interactions is only of real importance for small distances between the particles. In order to understand the influence of these interactions on material properties such as the boiling point or the melting point, it makes sense to compare two molecules with approximately the same size and mass in terms of their material properties. Molecular bromine (Br2) and iodine monochloride (ICl) are to be compared as an example.
While iodine chloride contains a polar bond and forms a dipole, bromine is non-polar and therefore also not a dipole, since the centers of charge coincide. Consequently, more dipole-dipole forces can be formed in iodine chloride than in bromine. Since these forces of attraction are between the molecules, it takes more energy to separate the particles from each other again.
As a consequence, the iodine chloride has both a higher melting and boiling point than bromine. While bromine is already a liquid at 0°C (melting point of -7°C) and begins to boil at 59°C, iodine chloride is still a solid at 0°C (melting point of 13.8°C at the β- form and 27.4°C for the α-form) and only begins to boil at 97.5°C.
This phenomenon is due to the intermolecular interactions present in iodine chloride, specifically the dipole-dipole forces, which provide internal stability.