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## Bond energies

Current time:0:00Total duration:5:21

# Worked example: Interpreting potential energy curves of diatomic molecules

AP.Chem:

SAP‑3 (EU)

, SAP‑3.B (LO)

, SAP‑3.B.1 (EK)

, SAP‑3.B.2 (EK)

## Video transcript

- [Instructor] In a previous video, we began to think about
potential energy as a function of internuclear distance
for diatomic molecules. What do I mean by diatomic molecules? Well, we looked at
molecular hydrogen, or H2, which is just two hydrogens
covalently bonded to each other. And at standard temperature and pressure, there, they would naturally, the distance between the two nuclei would be based on where there is the lowest potential energy. And if you were to squeeze them together, you would have to put
energy into the system and have a higher potential energy. Or if you were to pull them apart, you would have to put
energy into the system and have a higher potential energy. What I want to do in this video is do a little bit of a worked example. Over here, I have three potential energies as a function of
internuclear distance graphs. And what I'm going to tell you is one of these is molecular hydrogen, one of these is molecular
nitrogen or diatomic nitrogen, N2, and one of these is diatomic oxygen. And what I want you to think
about, pause this video, is which graph is the potential energy as a function of internuclear distance for each of these diatomic molecules. And I'll give you a hint. Look at the low point in potential energy. The low point in potential energy is what you would typically observe that diatomic molecule's
internuclear distance to be at standard
temperature and pressure. And this distance right over here is going to be a function of two things. It's going to be a function of how small the atoms actually are, how small their radii are. So smaller atoms are, in general, going to have a shorter
stable internuclear distance. But the other thing to think
about is the bond order between these atoms, and I'll give you a little bit of a hint. Diatomic hydrogen, you just
have a single covalent bond. For diatomic nitrogen,
it is a triple bond. And for diatomic oxygen,
it is a double bond. So the higher order the bond, that will also bring the
two atoms closer together, and it also makes it have
a higher bond energy, the energy required to separate the atoms. Remember, we talked about
it in the previous video. This right over here is the bond energy. And so with that said, pause the video, and try to figure it out. Which of these is the graphs of H2, which is N2, and which is O2? So let's first just think about
it in terms of bond energy. If you look at it, the single bond, double
bond, triple bond here, you would expect the
highest order bond here to have the highest bond energy, and the highest bond energy is this salmon-colored
one right over here. So just based on that, I would say that this is
a good candidate for N2. So this one right over here, this looks like diatomic nitrogen to me. Then the next highest bond energy, if you look at it carefully, it looks like this purple
one right over here. And so just based on bond order, I would say this is a
good candidate for O2. And then the lowest bond energy is this one right over here. And so just based on the bond order here, it's just a single covalent bond, this looks like a good
candidate for diatomic hydrogen. But let's also think about
the radii of these atoms. If we get a periodic
table of elements here, we can see that hydrogen
only has one electron in that first shell, and so it's going to be the smallest. So that makes sense over
here, that your distance, where you have the
lowest potential energy, is shortest for the diatomic molecule that's made up of the smallest atoms. But then when you look at the other two, something interesting happens. Remember, your radius
for an atom increases as you go down a column. But as you go to the right on
a row, your radius decreases. 'Cause you're adding
more and more electrons to the same shell, but the
Coulomb forces are increasing between that outermost
shell and your nucleus. And so if you just look at that trend, as you go from nitrogen to oxygen, you would actually
expect your atomic radius to get a little bit smaller. They're right next to each other. They might be close, but
you say, okay, oxygen, you have one extra electron
in that same second shell, maybe it's going to be
a little bit smaller. So if you were to base
things just on that, you'd say, all right, well,
the internuclear distance for this salmon-colored one
is a little bit shorter, maybe that one is oxygen, and
maybe this one is nitrogen. But they would be close,
and I would say, in general, the bond order would trump things. And the bond order, because
you see this high bond energy, that's the biggest
giveaway that this is going to be the higher bond order
diatomic molecule or N2. They're close in atomic radius, but this is what makes
all of the difference. And we'll take those two nitrogen atoms and squeeze them together
just a little bit more, even though they might
be a little bit bigger. And so I feel pretty
good with this labeling.