All the bonding we've talked about thus far are the attractions of one atom to another within a molecule. We call them covalent bonds but in a bigger sense, these attractions within a molecule are called intramolecular forces which means forces within a molecule.
THIS section is about another very important set of forces that occur between molecules and these are collectively known at intermolecular forces (IMFs). Intermolecular forces are the forces of attraction that pulls molecules together so that there can be properties of matter for condensed states. A condensed state is the opposite of gas state. A condensed phase is either a solid or a liquid and the molecules are all held close together by IMFs. The strength of the IMFs will determine a multitude of properties such as melting point, boiling point, viscosity, and vapor pressure. Those are all physical properties. Physical properties of matter are governed by the intermolecular forces within the molecules making up that particular matter. There are three major classes of IMFs.
If you have a polar molecule, you have a permanent dipole moment for your molecule. This means that there is a distinct permanent separation of net positive charge and net negative charge within the molecule. The separation is ever so slight though, meaning you do not get a full +1 and –1 of charge like you do with ionic compounds. Here, you are getting a much smaller amount of charge. We refer to this small amount as partial charge and we use a lowercase delta symbol to show it. So partial positive is shown as \(\delta^+\) while partial negative is shown as \(\delta^-\). This was pointed out when polarity was discussed in the previous chapter.
Once you have the polar molecules, you will have attractions between the oppositely charged regions within the molecules. Coulomb's law is still in effect but the forces are much weaker now. IMFs are much much weaker than chemical bonds - about 10× weaker. But that is strong enough to make molecules stick to each other and form condensed phases.
REMEMBER: only polar molecules have dipole-dipole interactions
This is a special case of dipole-dipole. So much so, it gets its own designation. H-bonding is typically the strongest version of dipole-dipole interaction. It occurs when a covalently bound hydrogen is attached to a very electronegative element. To be specific, H-bonding occurs anytime a H is covalently bonded to a N, O, or F atom. Read that again. The H HAS to be directly bonded to a N, O, or F. When this happens, the electrons are drawn in close to the N, O, or F leaving the proton of hydrogen's nucleus very exposed and with a high relative partial positive charge. Also, the N, O, and F in this setup end up with a just as high negative partial charge. So this leads to sticking points that are basically best case scenarios for IMF attractions.
REMEMBER: only when H is directly bound to N, O, or F do you have H-bonding
So how do completely non-polar molecules stick together? No permanent dipole. What makes them stick?
ALL molecules have fluctuations in their electron clouds.. There are brief moments of slight charge separations that lead to momentary dipoles. These are also known as temporary dipoles. They ARE real and they change rapidly. But on an atomic time scale this is an eternity and the molecules are able to line up and allow the attractions between momentary partial positive and partial negative to take hold. ALL molecules have these little temporary fluctuations of charge distribution and therefore ALL molecules have some degree of dispersion forces.
One-on-one, a temporary dipole would lose to a permanent one. The attraction just isn't as much. However, because these fluctuations occur all over the surface of the molecule, there is a bigger and bigger chance of attractions the bigger and bigger the molecule gets. Truly small molecules and atoms have very low dispersion forces. But as you build bigger and bigger molecules, you have more and more surface area and more and more sticking points which means that dispersion forces scale with molecular size. Big molecules have much bigger dispersion forces than small ones. The forces are even big enough that the overall effect is bigger than the other two IMFs. This is what we mean when we say there is strength in numbers. Thousands of tiny little interactions add up quick.
A classic example of this in the real (macroscopic) world is Velcro. Velcro is a hook-and-loop fastener where one surface is thousands of little plastic hooks while the other surface is thousands of little fiber loops. A single hook and fiber loop is not very strong. But put thousands of those together and you've got a pretty good fastener. Behold the power of Velcro! So too is the power of dispersion forces.
REMEMBER: ALL molecules have dispersion forces. The forces scale with molecular size.
Below is a bit of a quickie video of Dr. McCord explaining the three intermolecular forces.