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Course: Biology library > Unit 2
Lesson 3: Chemical bonds and reactionsIntermolecular forces
Different types of intermolecular forces (forces between molecules). Created by Jay.
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- If you have a large hydrocarbon molecule, would it be possible to have all three intermolecular forces acting between the molecules? Or is it just hydrogen bonding because it is the strongest? How do you determine what forces act when you have big and diverse molecule like an anhydride, e.g. acetic anhydride: Would here be dipole-dipole interactions between the O's and C's as well as hydrogen bonding between the H's and O's? Or just one of the two? What about the london dispersion forces?(13 votes)
- You can have all kinds of intermolecular forces acting simultaneously. Usually you consider only the strongest force, because it swamps all the others.
When you are looking at a large molecule like acetic anhydride, you look at your list of intermolecular forces, arranged in order of decreasing strength. Then you go down the list and stop at the first one that fits your molecule.
We can omit all the forces involving ions. That leaves
H-bonding
Dipole-dipole
Dipole - induced dipole
London Dispersion
What is the strongest .intermolecular force in acetic anhydride?
H-Bonding ? No, because there are no O-H, N-H, or F-H bonds in the molecule.
Dipole-dipole ? Yes, because the molecule has polar C=O groups.
Dipole-induced dipole and London Dispersion forces are also present, but they are small in comparison to the dipole-dipole forces.(37 votes)
- Hi Sal,
I am a 60 year old woman who has been taking Geometry, Chemistry and Algebra in order to help my 15 year old niece. I wondered about something, and there seemed to be no definitive answer when I Googled it. The question is: do thoughts have mass? I think that they must have an infinitesimal, at this time, immeasurable mass.
Thanks,
Sue
P.S. I think that I am addicted to Khan Academy!(20 votes)- thoughts do not have mass. the reason is because a thought merely triggers a response of ionic movement (i.e. Na+, K+ ) these ions already exist in the neuron, so the correct thing to say is that a neuron has mass, the thought is the "coding" or "frequency" of these ionic movements. so a thought does not have mass. i like the question though :)(25 votes)
- Can someone explain why does water evaporate at room temperature; having its boiling point at 100ºC?(16 votes)
- In water at room temperature, the molecules have a certain average kinetic energy.
But some of them will be moving faster, and some of them will be slower.
Some of the molecules at the surface of the water will have enough kinetic energy to escape into the atmosphere.
The water will slowly evaporate, even though it is at room temperature.(28 votes)
- At, he says that the boiling point for methane is around -164 degrees celsius. So, can methane be in a liquid form, if it is colder than -164 degrees celsius? And if so, is it present in a liquid form naturally anywhere? 7:40(8 votes)
- Yes. Titan, Saturn's largest moon, has clouds, rain, rivers and lakes of liquid methane.(33 votes)
- how can a molecule having a permanent dipole moment induce some temporary dipole moment in a neighbouring molecule(3 votes)
- Suppose you're in a big room full of people wandering around. If I bring a smelly skunk into the room from one of the doors, a lot of people are probably going to move to the other side of the room. When the skunk leaves, though, the people will return to their more even spread-out state. Conversely, if I brought a bunch of cupcakes there might be a rush for my side of the room, though people would spread out again once the cupcakes were gone.
In this terrible metaphor, the people walking around the room are like the electrons evenly spread around a molecule. If you bring something negatively charged near the molecule (like the negative end of a permanent dipole), then the negatively charged electrons will be repelled and will concentrate on the far side of the molecule, making the near side slightly positively charged and the far side slightly negatively charged (ie you've made an induced dipole!). When the permanent dipole goes away, though, it is more stable for the electrons to spread out again. Conversely, if you bring the positive side of a permanent dipole near a molecule, the electrons in the molecule will rush towards the positive side, leaving the far side of the molecule with fewer electrons and thus a temporary positive charge.(36 votes)
- In the notes before this video they said dipole dipole interactions are the strongest form of inter-molecular bonding and in the video he said hydrogen bonding is the strongest. can you please clarify if you can. Thanks(5 votes)
- Hydrogen bonding is also a dipole-dipole interaction, but it is such a strong form of dipole-dipole bonding that it gets its own name to distinguish it from the others.(9 votes)
- if hydrogen bond is one of the strongest inter molecular force why is ammonia a gas and hydrogen fluoride (most electronegative of the FON elements) has a boiling point of 19.5 degree C and water 100 degree C?..... quite a wide variation in boiling point and state of matter for compounds sharing similar inter-molecular force(6 votes)
- Good question!
From your question I assume you've worked out why HF has a higher boiling point than NH₃.
To see why H₂O has a much higher boiling point, try drawing each molecule and all possible hydrogen bonds they can make. What do you see?
Does this help you answer your question?(3 votes)
- At, he says "double bond." I know where he gets that term from but what is exactly a double bond? 1:27(3 votes)
- A double bond is a chemical bond in which two pairs of electrons are shared between two atoms.(5 votes)
- At, a textbox shows up on the screen clarifying that the partial positive charge is on the carbonyl carbon rather than on the CH3 carbon. Why is this? Is this because the net ionic charge on the carbonyl carbon is more positive (would I need to actually make a dot structure diagram to figure that out if I theoretically wanted to)? 1:26(3 votes)
- Two reasons why.
First is the electronegativity difference between carbon and oxygen. Oxygen is more electronegative than carbon so the electrons in their covalent bond would rather spend more of their time around the oxygen as opposed to the carbon. Electrons are negatively charged so the oxygen with more electrons around it at any given time means it'll have a partial negative charge as a result. And carbon which has a lack of electrons around it because of this will experience a partial positive charge because of the lose of electron density around it.
Second is the contribution of resonance. If you draw a resonance structure of the acetone molecule inyou'll find you can push pi electrons of the double bond of the carbonyl into the oxygen so it has three lone pairs around it now instead of just two. If you add up for the formal charges on this second resonance structure you'll find the oxygen has a 1- charge and the carbonyl carbon has a 1+ charge. Now the actual structure of acetone is a hybrid of combination of these two resonance structures, what we call the resonance hybrid. The resonance structure with the formal charges also helps add some partial negative charge to the oxygen and partial positive charge to the central carbon. 1:26
So it's a combination of both these reasons as to why we get those partial charges on acetone.
Hope that helps.(4 votes)
- Why does oxygen need to take 4 electrons from Carbon? It only needs 2 electrons to fulfill it's octet so shouldn't there just be one line that is drawn?(3 votes)
- So for acetone's Lewis dot structure, if you just kept a single line (or single bond) between the central carbon and the oxygen you would indeed give oxygen an octet since it already has the 6 electrons from the lone pairs. However, this results in two things we have to consider. First with only a single bond to the oxygen, the central carbon now lacks an octet since it will only have 3 single bonds (1 to the oxygen and 2 to the other carbons). This means the central carbon only has 6 electrons from the three single bonds. The second thing we have to consider are the formal charges around the oxygen and the central carbon atom. The central carbon with only three single bonds will have a formal charge of +1. The oxygen with 3 lone pairs and a single bond will have a formal charge of -1. These formal charges of +1 and -1 still add to an overall charge of 0 for the molecule which makes sense since it's neutral. So what this means together is that you can draw the Lewis structures of acetone with a C=O bond or a C-O and both are valid structures. When we can draw multiple Lewis structures for the same molecule we call these resonance structures. Resonance structures individually don't represent how the molecule actually looks, rather the molecule looks like combination of the two resonance structures called a resonance hybrid. However the resonance structure with a C=O bond has less charge separation (no + or - formal charges on any of the atoms) so we say that it is the dominant resonance structure of the two. So conventionally we draw acetone with with a C=O bond more so when drawing its Lewis structure for simplicity sake. So drawing acetone your way isn't wrong, it's just less conventional. Long explanation, but I hope that helps.(2 votes)
Video transcript
In the video on
electronegativity, we learned how to determine
whether a covalent bond is polar or nonpolar. In this video, we're going
to see how we figure out whether molecules
are polar or nonpolar and also how to apply
that polarity to what we call intermolecular forces. Intermolecular
forces are the forces that are between molecules. And so that's different from
an intramolecular force, which is the force within a molecule. So a force within
a molecule would be something like
the covalent bond. And an intermolecular
force would be the force that are
between molecules. And so let's look at the
first intermolecular force. It's called a
dipole-dipole interaction. And let's analyze
why it has that name. If I look at one of these
molecules of acetone here and I focus in on the
carbon that's double bonded to the oxygen,
I know that oxygen is more electronegative
than carbon. And so we have four
electrons in this double bond between the carbon
and the oxygen. So I'll try to highlight
them right here. And since oxygen is
more electronegative, oxygen is going to pull
those electrons closer to it, therefore giving oxygen a
partial negative charge. Those electrons in yellow are
moving away from this carbon. So the carbon's losing a
little bit of electron density, and this carbon is becoming
partially positive like that. And so for this
molecule, we're going to get a separation of charge, a
positive and a negative charge. So we have a polarized
double bond situation here. We also have a
polarized molecule. And so there's two
different poles, a negative and a positive pole here. And so we say that this
is a polar molecule. So acetone is a
relatively polar molecule. The same thing happens to this
acetone molecule down here. So we get a partial negative,
and we get a partial positive. So this is a polar
molecule as well. It has two poles. So we call this a dipole. So each molecule
has a dipole moment. And because each
molecule is polar and has a separation of
positive and negative charge, in organic chemistry we know
that opposite charges attract, right? So this negatively
charged oxygen is going to be attracted to
this positively charged carbon. And so there's going to be
an electrostatic attraction between those two molecules. And that's what's going to hold
these two molecules together. And you would
therefore need energy if you were to try
to pull them apart. And so the boiling
point of acetone turns out to be approximately
56 degrees Celsius. And since room temperature
is between 20 and 25, at room temperature
we have not reached the boiling point of acetone. And therefore, acetone
is still a liquid. So at room temperature and
pressure, acetone is a liquid. And it has to do with
the intermolecular force of dipole-dipole
interactions holding those
molecules together. And the intermolecular
force, in turn, depends on the
electronegativity. Let's look at another
intermolecular force, and this one's called
hydrogen bonding. So here we have two
water molecules. And once again, if I think
about these electrons here, which are between the
oxygen and the hydrogen, I know oxygen's more
electronegative than hydrogen. So oxygen's going to pull
those electrons closer to it, giving the oxygen a partial
negative charge like that. The hydrogen is losing a
little bit of electron density, therefore becoming
partially positive. The same situation exists in
the water molecule down here. So we have a partial negative,
and we have a partial positive. And so like the
last example, we can see there's going
to be some sort of electrostatic attraction
between those opposite charges, between the negatively
partially charged oxygen, and the partially positive
hydrogen like that. And so this is a polar molecule. Of course, water is
a polar molecule. And so you would
think that this would be an example of
dipole-dipole interaction. And it is, except
in this case it's an even stronger version of
dipole-dipole interaction that we call hydrogen bonding. So at one time it
was thought that it was possible for hydrogen
to form an extra bond. And that's where the term
originally comes from. But of course, it's not an
actual intramolecular force. We're talking about an
intermolecular force. But it is the strongest
intermolecular force. The way to recognize when
hydrogen bonding is present as opposed to just
dipole-dipole is to see what the hydrogen is bonded to. And so in this case, we have
a very electronegative atom, hydrogen, bonded-- oxygen,
I should say-- bonded to hydrogen. And then that hydrogen
is interacting with another electronegative
atom like that. So we have a partial negative,
and we have a partial positive, and then we have another
partial negative over here. And this is the
situation that you need to have when you
have hydrogen bonding. Here's your hydrogen showing
intermolecular force here. And what some students forget
is that this hydrogen actually has to be bonded to another
electronegative atom in order for there to be a big enough
difference in electronegativity for there to be a little
bit extra attraction. And so the three
electronegative elements that you should remember
for hydrogen bonding are fluorine,
oxygen, and nitrogen. And so the mnemonics
that students use is FON. So if you remember FON as the
electronegative atoms that can participate in
hydrogen bonding, you should be able to remember
this intermolecular force. The boiling point of water is,
of course, about 100 degrees Celsius, so higher than
what we saw for acetone. And this just is due to the
fact that hydrogen bonding is a stronger version of
dipole-dipole interaction, and therefore, it takes
more energy or more heat to pull these water
molecules apart in order to turn
them into a gas. And so, of course, water is
a liquid at room temperature. All right. Let's look at another
intermolecular force. And this one is called
London dispersion forces. So these are the weakest
intermolecular forces, and they have to do with the
electrons that are always moving around in orbitals. And even though the
methane molecule here, if we look at it,
we have a carbon surrounded by four
hydrogens for methane. And it's hard to tell in how
I've drawn the structure here, but if you go back and
you look at the video for the tetrahedral
bond angle proof, you can see that in
three dimensions, these hydrogens are
coming off of the carbon, and they're equivalent
in all directions. And there's a very
small difference in electronegativity between
the carbon and the hydrogen. And that small difference
is canceled out in three dimensions. So the methane molecule becomes
nonpolar as a result of that. So this one's nonpolar, and,
of course, this one's nonpolar. And so there's no
dipole-dipole interaction. There's no hydrogen bonding. The only intermolecular
force that's holding two methane
molecules together would be London
dispersion forces. And so once again, you could
think about the electrons that are in these bonds
moving in those orbitals. And let's say for the
molecule on the left, if for a brief
transient moment in time you get a little bit
of negative charge on this side of the molecule,
so it might turn out to be those electrons have a net
negative charge on this side. And then for this
molecule, the electrons could be moving the
opposite direction, giving this a partial positive. And so there could be
a very, very small bit of attraction between these
two methane molecules. It's very weak, which is why
London dispersion forces are the weakest
intermolecular forces. But it is there. And that's the only thing that's
holding together these methane molecules. And since it's weak, we would
expect the boiling point for methane to be extremely low. And, of course, it is. So the boiling point for methane
is somewhere around negative 164 degrees Celsius. And so since room temperature
is somewhere around 20 to 25, obviously methane
has already boiled, if you will, and
turned into a gas. So methane is obviously a gas at
room temperature and pressure. Now, if you increase
the number of carbons, you're going to increase the
number of attractive forces that are possible. And if you do that,
you can actually increase the boiling point
of other hydrocarbons dramatically. And so even though
London dispersion forces are the weakest, if you
have larger molecules and you sum up all
those extra forces, it can actually turn out to be
rather significant when you're working with larger molecules. And so this is just
a quick summary of some of the
intermolecular forces to show you the application
of electronegativity and how important it is.