There are two main kinds of inter-molecular forces - dipole interactions, which is when
polar molecules are attracted to each other, and dispersion forces, which are caused by
the motion of electrons.
Collectively, these intermolecular forces are often called van der Waals interactions.
Sometimes, especially in Biology, people get casual in their language, and they call a
specific KIND of intermolecular force a van der Waals interaction.
No big deal.
We'll talk about all of these forces individually.
First, let's look at the various kinds of dipole interactions.
Ion-Dipole forces occur when ions are attracted to a polar molecule.
Positive ions are attracted to the negative end of a polar molecule, and negative ions
are attracted to the positive end.
For example, when NaCl is in solution in water, the negative end of the polar water molecules
are attracted to the positive Na+ ions.
So you can see the water molecules are oriented so the oxygens, with their partial negative
charges, are pointed towards the Na+ ions, and the Hydrogens are pointed away.
Meanwhile, the positive end of the water molecule is oriented towards the negative Chloride
ion.
There are also cases of ion-induced dipole interactions.
This is when the PRESENCE of the ion induces a temporary dipole in another molecule that
usually does not have poles.
It's sitting there, being nonpolar, no partial positive or negative charges,
until an ion approaches.
The charge of the ion causes a temporary distortion of the electron cloud of the nonpolar molecule,
INDUCING a dipole, which is then attracted to the ion.
For example, as this Chloride ion, Cl -, approaches hexane (C6H14), a nonpolar molecule, the negative
charge on the ion slightly distorts the electron cloud - pushes the electrons away,
just a little.
As a result, there is a slight partial positive charge on this side of the hexane, which is
attracted to the negatively charged ion.
Dipole-Dipole interactions occur when two polar molecules are very close to each other,
and the positive end of one polar molecule is near the negative end of the other molecule.
The strength of the attraction increases with increasing polarity.
Here, for example, you can see two molecules of formaldehyde, CH2O.
Each of these individual molecules contains a polar covalent bond.
The molecules have positive and negative ends, or poles.
If two formaldehyde molecules come close together, the partial positive end of one molecule will
be attracted to the partial negative end of the other molecule.
A special case of dipole-dipole interaction is the Hydrogen Bond.
This is when the hydrogen atom in a very polar molecule is attracted to the negatively charged
atom in another polar molecule.
The hydrogen is most often involved in a bond with a fluorine, oxygen, or nitrogen.
Fluorine, oxygen, and nitrogen are so electronegative, they are typically involved in
extremely polar bonds.
These electronegative atoms pull electrons closer to themselves, resulting in a
partial negative charge.
The hydrogen has a partial positive charge, so the two atoms in the different molecules
are attracted to each other.
Now, here's a caveat/ an exception/ something to keep in mind.
Not every hydrogen bond is a dipole-dipole interaction, or even an intermolecular bond.
There are cases of hydrogen bonds within a single molecule.
For instance, if you study biochemistry, you'll learn about how hydrogen bonds are formed
within a single polypeptide chain (that is, a string of amino acids), as it folds up to
form a protein.
So that's a case of hydrogen bonds forming between different parts of the same molecule.
But in chemistry, you're almost always talking about intermolecular hydrogen bonds.
Now, we've done a survey of dipole interactions.
Let's move on to dispersion forces.
You can think of the dispersion force as kind of the default interaction
between chemical species.
It involves the motion of electrons.
Since absolutely every atom and every molecule and every compound has electrons, every single
kind of chemical species is going to experience dispersion forces.
Thinking about the word "dispersion" is the key to remembering this force.
How are the electrons dispersed, or distributed, in the cloud around the chemical species?
Are they absolutely evenly distributed?
Yes, they could be, some of the time.
But a lot of the time, in fact, almost all of the time, there will be an uneven distribution
of electrons on the very, very surface of any given molecule.
Think of it like patchiness.
There will, for a brief moment, be a patch that is more negative than another patch.
One patch will be more positive than another patch.
In the next moment, because the electrons are constantly in motion, the patch is gone.
There will be a different kind of patchiness.
Because this is true of all chemical species - every molecule, every compound, every atom
- there will be some degree of attraction due to this kind of patchiness.
For the briefest of moments, the negative patch on one molecule will line up with a
positive patch on another molecule.
Then in the next instant, these patches have disappeared, or rather redistributed.
Imagine these bonds blinking on and off at a dizzying speed.
There are ranges of strengths of dispersion forces.
The larger the molecule, the bigger the electron cloud, the farther the electrons are from
nuclei and so they are held less tightly, and so they can move around more easily and
thus more easily create these sorts of patches of charge.
We sometimes call these "instantaneous dipoles."
This kind of interaction is the weakest of all the intermolecular bonds.
Careful you don't confuse INTRAmolecular bonds with INTERmolecular bonds.
In previous videos, we discussed Intramolecular Bonds - the strong bonds that hold compounds
together: Ionic bonds and Covalent bonds.
Intra refers to the bonds WITHIN ONE molecule or ONE ionic compound.
INTER refers to the bonds BETWEEN two different chemical species.
Intra-molecular forces, the forces that hold one compound together, are much stronger than
inter-molecular forces - the forces between different molecules.
Intermolecular forces are all, essentially, electrostatic interactions - attractions between
positive and negative charges - just far weaker than the strong electrostatic attractions
involved in ionic bonds.
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