The intermolecular forces are a set of interactions occurring between the molecules a single chemical compound or not, and which hold them together. Thanks to these forces the molecules manage to group together and define a solid, liquid, or gas; that is, they are responsible for the physical properties of substances.
Intermolecular forces can be of two types:
electrostatic or Van der Waals. The electrostatic ones are all those where we have ions, which are electrically charged atoms or molecules (+ or -). Meanwhile, those of Van der Waals involve neutral atoms or molecules, which due to fluctuations in their electronic distribution become polarized and attract or repel each other.
Intermolecular forces can be compared to the bond between different pieces of Legos. Depending on their shapes and orientations, their joints become weaker or stronger, as do the design and final dimensions of the construction.
The stronger these forces, the more difficult it will be to separate the Legos pieces or, in the case of chemistry, the molecules. This also means that there will be more compaction and therefore the molecules will define more viscous liquids, heavier gases, or denser solids.
Understanding intermolecular forces are essential to describing many of the chemical, physical, and biological concepts.
Electrostatic forces (intermolecular forces)
Oppositely charged ions attract with great force, which obeys Coulomb’s law, and decreases rapidly the further the ions move away.
For example, the Ca 2+ ion is positively charged, a cation, while CO 3 2- is a negatively charged ion, an anion. As their charges are opposite, Ca 2+ and CO 3 2- attract each other, showing that the closer they are, the more together they will remain.
On the contrary, ions with the same charge, whether positive or negative, repel each other with the same force but in the opposite direction. Thus, Ca 2+ would repel ions like Na +, Mg 2+, K +, etc., unless there are other anions that decrease the repulsion.
Of all the intermolecular forces, the ion-ion type is the simplest; but at the same time, the strongest of all. It takes a lot of energy or heat to separate the ions.
Ions can also interact with neutral atoms or molecules, such as those in water. This is especially true when it comes to the molecules of a solvent, which is solvate, that is, it surrounds the ions in its liquid.
For example, sodium chloride, NaCl, is made up of Na + and Cl – ions. When its crystals are thrown into the water, the water molecules, H 2 O or HOH, hydrate or surround the NaCl ions; but the way they do it varies depending on the ion in question.
Thus, we see in the upper image that the Na + ion is hydrated by the oxygen atoms of H 2 O. Meanwhile, the Cl – ion is hydrated by the hydrogen atoms.
Because water is a polar substance, which means that it has a dipole: one pole with a positive partial charge and another with a negative partial charge.
The oxygen atom, being more electronegative, attracts electrons towards itself, therefore, electrons are located more frequently around that atom; This is not the case with hydrogen atoms, less electronegative. Oxygen, being more negative, is oriented towards Na +; while hydrogens, being more positive, are oriented towards Cl –, since opposite charges attract each other.
Not only polar or dipole molecules can interact with ions. For neutral atoms or molecules, even if they do not have dipoles, the distribution of their electrons is susceptible to the inductive effects of nearby ions; that is, the ions cause a momentary and brief polarization, enough for there to be a remarkable interaction.
Suppose, for example, the interaction between the OH – ion and the CO 2 molecule. CO 2 is a neutral compound, whose molecule (in purple) lacks a dipole. However, as it approaches OH –, its negative charge repels electrons from the oxygen atoms of CO 2.
The closer the OH – and CO 2 are, the stronger the repulsion. Consequently, an OH-induced dipole begins to settle over CO 2 . A positive pole δ + then appears due to the repelled electrons “migrating” to the other end of the molecule.
Thus, OH – and CO 2 stay together long enough for them to react to each other. This is the reason why CO 2 is particularly soluble in alkaline solutions.
Van der Waals forces (intermolecular forces)
Van der Waals forces, in principle, refer only to those that exist between neutral atoms or molecules, without ionic charges.
Opposite charges attract, like charges repel. The same is true for dipoles: opposite poles (δ + and δ-) attract, while like poles repel. The positive poles or δ + are represented above with the color blue, whereas the negative poles or δ-, are represented with the red color.
Notice how the molecules above are oriented and arranged in such a way that opposite poles meet, pushing equal poles apart in the process. This ordering is what is known as the dipole-dipole interactions or forces (Keesom forces), and they are the most important intermolecular forces between polar molecules.
For example, the molecules of H 2 O, HCl, HF, and CO, among others, are arranged in similar ways. The more polar they are, the stronger their dipole-dipole forces will be; and therefore, the more difficult it will be to separate their molecules.
HF is more polar than HCl, so the boiling point of HF is higher (19.5 ºC) than that of HCl (-85.05 ºC). The effect of intermolecular forces on physical properties is amazing. The smallest variation and the substance will behave totally different than expected.
Induced dipole-dipole (Debye)
Dipoles, like ions, can also fluctuate or affect the electronic distribution of neutral atoms or molecules. So we see, in the image above, that a dipole suddenly polarizes a neutral molecule. It is therefore said that it is an induced dipole-dipole force since the second dipole is momentary, not permanent.
For example, water is capable of dissolving a small amount of O 2, but enough for marine fauna to breathe. If H 2 O could not induce a dipole in O 2, all the oxygen in the seas would escape to the surface, as there would be no interactions between the two molecules.
Instantaneous dipole-induced dipole (London)
Neutral atoms or molecules do not need neighboring ions or dipoles to suffer by themselves fluctuations in the distribution of their internal charges. The electrons are not stationary but move throughout the molecule. At some point, an instantaneous dipole will take place, which if it is very close to a neutral molecule or atom, will induce a dipole in them (see above).
The larger and more asymmetric these molecules are, the more likely instantaneous dipoles will occur. That is why compounds with higher atomic masses have stronger instantaneous dipole-induced dipole forces. This type of Van der Waals force is better known as London forces.
For example, noble gases (He, Ne, Ar, etc.) are held together by London forces, as are gases in air (O 2, N 2, CO 2, etc.). Likewise, this force is the most predominant among hydrocarbons such as methane, CH 4, and propane, CH3 CH2 CH3.
Hydrophobic forces (intermolecular forces)
Finally, we have the hydrophobic forces, which are a special result of the London forces and of the repulsions between substances with different polarities. The fats have little affinity to the water, the reason why it is insoluble in this.
The fat molecules seek to group together in such a way that their interactions with water are as minimal as possible, and this is achieved by forming 3D structures like those shown above.
Thus, fat molecules, such as phospholipids, unite to form lipid bilayers, micelles, and liposomes.
If observed, these molecules have a white head that represents a polar portion, akin to water; as well as an apolar tail, which repels water molecules. The tails seek to be located inside these supramolecular structures so that they do not interact with water.
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