Electronegativity
electronegativity and intermolecular forces
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- 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 can figure out whether molecules are polar or non polar,
- and also how to apply that polarity to what we call "inter-molecular forces.
- Intermolecular forces are the forces that are between molecules.
- And so thats different than an intra-molecular force, which is the force within a molecule.
- So the force within a molecule would be something like a covalent bond.
- An intermolecular force would be the force that is between molecules.
- So lets look at the first intermolecular force.
- Its called a dipole-dipole interaction
- and lets analyze why it has that name.
- If i look at one of these molecules of acetone here,
- and I focus in on the carbon thats double bonded to the oxygen,
- I know that oxygen is more electronegative than carbon,
- and so we have 4 electrons in this double bond between the carbon and the oxygen.
- I'll try to highlight them right here.
- 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 this carbon is losing a little bit of electron density
- and this carbon is becoming partially positive, like that.
- For this molecule, we're going to get a separation of charge.
- A positive and a negative charge.
- 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.
- 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,
- and it has two poles. We call this a dipole.
- 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.
- So this negatively charged oxygen is going to be attracted to this positively charged carbon.
- There's going to be an electrostatic attraction between those two molecules.
- And thats what's going to hold these two molecules together, and you would therefore need energy
- if you were to try to pull them apart. The boiling point of acetone turns out to be
- approximately 56 degrees celsius. 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.
- At room temperature and pressure, acetone is a liquid. And it has to do with the intermolecular force
- of dipole-dipole interaction holding those two molecules together.
- And the intermolecular force, in turn, depends on the electronegativity.
- Lets look at another intermolecular force, and this one's called hydrogen bonding.
- 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 is more electronegative
- than hydrogen, so oxygen is 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.
- 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.
- This is a polar molecule, of course water is a polar molecule.
- You would think that this would be an example of dipole-dipole interaction, and it is,
- except that in this case its an even stronger version of dipole-dipole interaction that we call "hydrogen bonding".
- At one time it was thought that it was possible for hydrogen to form an extra bond
- and thats where the term originally comes from.
- But of course its not an actual intra-molecular force, its an inter-molecular 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. In this case, we have a very electronegative atom, oxygen, bonded to hydrogen.
- That hydrogen is interacting with another electronegative atom like that.
- We have a partial negative, and a partial positive, and we have another partial negative over here.
- 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.
- The mnemonic that students use is "FON". So if you remember "FON" as the
- electronegative atoms that could 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.
- Higher than what we saw for acetone.
- This is due to the fact that hydrogen bonding is a stronger version of dipole-dipole interaction, and therefore
- takes more energy or more heat to pull these water molecules apart in order to turn them into a gas.
- So, of course, water is a liquid at room temperature.
- Alright, lets look at another intermolecular force. This one is called "london dispersion forces".
- These are the weakest intermolecular forces, and they have to do with the electrons that are always moving around
- in orbitals. Even though the methane molecule here, if we look at it, and we have a carbon surrounded by
- hydrogens for methane. And its hard to tell from how i've drawn the structure here, but if you go back and
- see the video for the tetrahedral bond angle proof, you can see that in three dimensions, these hydrogens are
- coming off the carbon and they're equivalent in all directions.
- There's a very small difference in electronegativity between the carbon and the hydrogen.
- That small difference is cancelled out in three dimensions, and so the methane molecule becomes
- non-polar as a result of that. So this one is non polar and this one is non polar.
- There's no dipole-dipole interaction, there's no hydrogen bonding.
- The only intermolecular force thats holding two methane molecules together would be london dispersion forces.
- Once again you can think about the electrons that are in these bonds moving in those orbitals.
- Lets say for the molecule on the left, 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.
- There could be a very small bit of attraction between these two methane molecules.
- Its very weak, which is why london dispersion forces are the weakest intermolecular force, but
- it is there. And thats the only thing thats holding together these methane molecules.
- Since its 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.
- Since room temperature is somewhere around 20-25, obviously methane has already boiled, if you will,
- and turned into a gas. Methane is obviously a gas at room temperature and pressure.
- Methane is a gas .
- 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. 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. This is just a quick summary of some of the intermolecular forces to show you
- the application of electronegativity, and how important it is.
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