In chemistry the focus is often on the individual molecule its structure, name, physical properties and chemical reactivity. Molecules do not generally exist isolated and alone, however. They are immersed in the environment around them, surrounded by other molecules which may or may not be the same chemical. Either way, the molecules interact, exerting attractive forces upon one another. (This is fortunate, since if they all repelled, people, buildings, and every other object would literally fall to pieces.)
All intermolecular forces (forces between molecules) are attractive, Coulombic forces. The Coulomb is a measure of charge, and you are probably familiar with the phrase “opposite charges attract”. They do indeed, and this Coulombic force is the underlying principle for each of the forces below.
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Hydrogen Bonding
A hydrogen bond is not an actual bond. It is, however, a strong attractive force, with enough strength to hold two molecules together. The two halves of your DNA are held together by hydrogen bonds. Many of water’s unique properties (relatively high boiling point and surface tension) are a result of its ability to hydrogen bond with itself.
A hydrogen bond is formed by the attraction between a negatively charged lone pair of electrons on one molecule and a hydrogen atom that is attached to an electronegative element (like O, N, or S) on another molecule. The hydrogen atom has a partial positive charge because the electronegative element draws the electrons away from the hydrogen, leaving the proton’s positive charge unbalanced.
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Dipole-Dipole Forces
A dipole is anything that has two oppositely charged ends. This creates two poles, a positive and a negative, thus the name di- (two) pole. The hydrogen atom discussed in hydrogen bonding is an example of half of a dipole. The atom (O, N, S…) that it is attached to is the other half. Whatever atom the electron pair is on is also a dipole, with the lone pair representing the negative end. Thus hydrogen bonding is a form of dipole-dipole forces. Hydrogen bonding is far stronger than the rest, however, so it gets its own name.
A dipole results between two bonded atoms when they have significantly different electronegativities. The higher electronegativity atom becomes slightly negative, the lower, positive. An entire molecule becomes a dipole when the dipoles inside it are unbalanced. For example, carbon dioxide has two carbon-oxygen bonds, each of which is a dipole, but because they point in opposite directions, they balance out. Carbon dioxide is not a molecular dipole as a result. Water, on the other hand, has two hydrogen-oxygen dipoles, but they point out at an angle, so that a net negative (the oxygen) and positive (the side with hydrogen atoms) end are made.
When two molecular dipoles are placed near one another, they rotate so that the negative pole of one is next to the positive end of the other, and those opposite charges hold them loosely together.
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Other Dipole Forces
If a molecule is a permanent dipole, it is not only attracted to other dipoles. It will also be attracted by ions and even other molecules that lack a dipole of their own.
Any ion will attract the end of a dipole which has the opposite charge. This plays an important part in the dissolution of ionic compounds in water. Water, a dipole, is attracted to the individual ions in a crystal, and surrounds them, forming a cage of water molecules around the ion and breaking it away from the crystal. By the time the entire salt (ionic compounds are often referred to as salts) has dissolved, each ion is surrounded by water molecules that were attracted by the charge.
A molecule that has no dipole of its own will still be attracted by a dipole. When the dipole approaches the molecule, the charge at its near end causes the electrons in the second molecule to shift, either attracted or repelled by the near end (positive or negative, respectively) of the dipole. This shift of electrons creates a temporary dipole in the other molecule, known as an induced dipole, and allows it to behave as a dipole would. (These forces are weaker than the attraction between two permanent dipoles.)
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London Dispersion Forces
Strangely, this last intermolecular force includes the word “dispersion” which most people would expect to mean “spreading out”. That’s exactly not what these London forces do, however, as they are also attractive in nature, though quite weak when compared to the others.
For groups of molecules which lack a permanent dipole, there is no chance of creating even a lasting induced dipole. What does happen is that the electrons in the molecule, which are constantly in motion, create momentary charge unbalances, known as instantaneous dipoles. As their name suggests, these dipoles don’t last long at, and they bear a miniscule charge, but for that brief moment, a dipole exists in the molecule. During that moment, that dipole behaves just like a permanent dipole, attracting other dipoles, and inducing dipoles (instantaneous though they may be) in surrounding molecules. This process occurs continually, as the electrons shift, constantly creating dipoles and attractive forces. Small as those forces are, they do add up to a net cohesive force, so even non-polar molecules can stick together, albeit weakly.
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There may be days when life is hectic and you feel like you just can’t hold things together. On those days, reflect and remember the role that intermolecular forces play, and realize that without them, you’d never be able to hold anything together at all. Isn’t chemistry comforting?
