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## Intramolecular force and potential energy

Current time:0:00Total duration:6:42

# Bond length and bond energy

AP.Chem:

SAP‑3 (EU)

, SAP‑3.B (LO)

, SAP‑3.B.1 (EK)

## Video transcript

- [Instructor] If you
were to find a pure sample of hydrogen, odds are that the individual
hydrogen atoms in that sample aren't just going to be
separate atoms floating around, that many of them, and
if not most of them, would have bonded with each other, forming what's known as diatomic hydrogen, which we would write as H2. Another way to write it
is you have each hydrogen in diatomic hydrogen would have
bonded to another hydrogen, to form a diatomic molecule like this. This molecule's only made up of hydrogen, but it's two atoms of hydrogen. And this makes sense, why it's stable, because each individual hydrogen
has one valence electron if it is neutral. So that's one hydrogen there. That's another one there. And if they could share
their valence electrons, they can both feel like they
have a complete outer shell. And so this dash right over here, you can view as a pair
of electrons being shared in a covalent bond. Now, what we're going to do in this video is think about the
distance between the atoms. So just as an example, imagine
two hydrogens like this. So that's one hydrogen atom, and that is another hydrogen atom. It turns out, at standard
temperature, pressure, the distance between
the centers of the atoms that we observe, that
distance right over there, is approximately 74 picometers. And just as a refresher of
how small a picometer is, a picometer is one trillionth of a meter. So this is 74 trillionths of a meter, so we're talking about
a very small distance. But one interesting question
is why is it this distance? What would happen if we tried
to squeeze them together? What would happen if we
tried to pull them apart? And to think about that, I'm gonna make a little bit of a graph that deals with potential
energy and distance. So in the vertical axis, this is going to be potential energy, potential energy. And I won't give the units just yet. I'll just think in very
broad-brush conceptual terms, then we could think about
the units in a little bit. And then this over here is the distance, distance between the centers of the atoms. You could view it as the
distance between the nuclei. And let's give this in picometers. Now, potential energy,
when you think about it, it's all relative to something else. And so let's just arbitrarily say that at a distance of 74 picometers, our potential energy is right over here. I'm not even going to label this axis yet. Now, what's going to happen
to the potential energy if we wanted to pull
these two atoms apart? Well, this is what we
typically find them at. This is probably a low point, or this is going to be a low
point in potential energy. So if you make the distances go apart, you're going to have
to put energy into it, and that makes the
potential energy go higher. And to think about why that makes sense, imagine a spring right over here. If you want to pull it apart, if you pull on either sides of a spring, you are putting energy in, which increases the potential energy. Because if you let go, they're
just going to come back to, they're going to accelerate
back to each other. So as you pull it apart, you're adding potential energy to it. So as you have further
and further distances between the nuclei, the
potential energy goes up. And if you go really far, it's going to asymptote
towards some value, and that value's
essentially going to be the potential energy if these two
atoms were not bonded at all, if they, to some degree, weren't
associated with each other, if they weren't interacting
with each other. And so that's actually the point at which most chemists or physicists or scientists would label
zero potential energy, the energy at which they are infinitely far away from each other. And that's what this
is asymptoting towards, and so let me just draw
that line right over here. So let's call this zero right over here. And actually, let me now give units. Let's say all of this is
in kilojoules per mole. Now, once again, if
you're pulling them apart, as you pull further and
further and further apart, you're getting closer and closer to these, these two atoms not interacting. Why is that? Because as you get further
and further and further apart, the Coulomb forces between them are going to get weaker and weaker
and weaker and weaker. And so that's why they like to think about that as
zero potential energy. Now, what if we think about
it the other way around? What if we want to squeeze
these two together? Well, once again, if you
think about a spring, if you imagine a spring like this, just as you would have to add energy or increase the potential
energy of the spring if you want to pull the spring apart, you would also have to do it
to squeeze the spring more. And so to get these two atoms to be closer and closer
and closer together, you have to add energy into the system and increase the potential energy. And why, why are you having
to put more energy into it? Because the more that you squeeze
these two things together, you're going to have the positive charges of the nuclei repelling each other, so you're gonna have to
try to overcome that. That puts potential
energy into the system. And these electrons are starting to really overlap with each other, and they will also want
to repel each other. And so what we've drawn here,
just as just conceptually, is this idea of if you wanted them to really overlap with each other, you're going to have a
pretty high potential energy. And if you're going to have them very separate from each other, you're not going to have as
high of a potential energy, but this is still going to be higher than if you're at this stable point. This stable point is stable
because that is a minimum point. It is a low point in this
potential energy graph. You could view this as just right. And it turns out that
for diatomic hydrogen, this difference between zero
and where you will find it at standard temperature and pressure, this distance right over here
is 432 kilojoules per mole. So this is at the point negative
432 kilojoules per mole. And so one interesting thing to think about a diagram like this is how much energy would it take
to separate these two atoms, to completely break this bond? Well, it'd be the energy of
completely pulling them apart. And so it would be this energy. It would be this energy right over here, or 432 kilojoules. And that's what people
will call the bond energy, the energy required to separate the atoms. And we'll see in future videos, the smaller the individual atoms and the higher the order of the bonds, so from a single bond to a
double bond to a triple bond, the higher order of the bonds, the higher of a bond energy
you're going to be dealing with.

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