Intermolecular Forces
Transcript
Now let's look at intermolecular forces. Take note, this is an important section on the MCAT because it can be used to predict trends and things like boiling point and the melting point. So make sure that you have this down pat. So the first of these intermolecular forces we'll look at are the London dispersion forces.
And these involved are called induced dipoles. For example, if we have a long, straight alcane like this, and for an instant, let's suppose that all the electrons are on one side of the molecule. Once this happens, these electrons here will repel the electrons on this next molecule and induce a dipole here. So now we have a temporary induced dipole.
We saw earlier that non-polar compounds had a much lower boiling point than polar compounds. Well by inducing a dipole, we've made this a quasi polar compound. And so this will slightly increase the boiling point. Now London dispersion forces are size dependent. Meaning the longer this chain is, or essentially the more electrons that exist in the compound, the easier it is to polarize the compound, which makes it so that the London dispersion forces are stronger.
If you don't have very many electrons, then you can't polarize it. Now the next intermolecular force are the dipole-dipole interactions. These occur between polar molecules. An example of a polar molecule would be CH3Cl or chloromethane. Chloromethane is a carbon, with three hydrogens and a chlorine attached to it. Now the carbon-hydrogen bonds are essentially non-polar.
Cuz the carbon-chlorine bond is polar. Which results in a slightly negative charge on the chlorine, because it is more electronegative and will suck up those electrons to hold on to it. And the carbon which is slightly positive. This results in a much greater attraction between the molecules for each other in the solid and the liquid.
In the vapor, there is still this attraction if they come close to each other, but since vapor molecules tend to be so far apart, this force isn't as strong. The result of this dipole-dipole interaction though, is that molecules wanna stay as the liquid for as long as possible, because there's intermolecular force holding them together.
Now you might think since CH3Cl is slightly polar, if we added a whole bunch of chlorines that way we can make it even more polar. Let's do carbon tetrachloride, CCl4. Now while it's true, that the carbon won't have a slight positive charge, and these a slight negative, each chlorine. In effect they cancel each other out, so to determine if a molecule is polar or not, only must it have a polar covalent bond in it, it can't be cancelled out by others in symmetry.
So carbon tetrachloride is in fact non-polar, even though it contains polar bonds. So when you see a compound that has electronegativity difference as greater than 0.4, you can't just assume that it's going to be polar, unless you know the bond connectivity. And the last of the intermolecular forces is hydrogen bonding.
This results from hydrogen being bonded to nitrogen, oxygen, and fluorine. So for example. Water has hydrogen bonded to oxygen. Well this hydrogen is strongly attracted to the oxygen. And the resulting intermolecular force significantly raises the boiling point of water.
Again nitrogen to hydrogen, as an amines. Oxygen to hydrogen, as in alcohols, and other compounds. These bondings will significantly increase the boiling point over something that doesn't have a hydrogen bonding. Now fluorine, we have FH which is just hydrogen fluoride. Once fluorine is attached to this hydrogen, that's it.
If you try to attach an alcohol group and the fluorine, you'd notice that this violates the octet rule because fluorine only needs one extra electron, which is already supplied by the hydrogen. Now which of these three forces is the strongest? Hydrogen bonding, dipole-dipole, London dispersion? Hydrogen bonding is far and away the strongest.
London dispersion forces, which are just induced dipoles and not permanent dipoles, are the weakest. The dipole-dipole interactions are in between the two. Now remember, while London dispersion forces are the weakest and hydrogen bonding the strongest, if the chain, or something like an alkane group starts to get very long.
All of a sudden the hydrogen bonding is a much smaller portion of the total molecule. And becomes less significant than the London dispersion forces. This doesn't mean that hydrogen bonding is weaker, just that the molecule is so big at this point, hydrogen bonding can't overcome the rest of the molecule. So let's use what we've learned to determine a few boiling point trends.
Which of these has the highest boiling point? Methane, methane propane, or normal butane? In this case, we're looking purely at London dispersion forces. These are greatest when the size is largest, so we would pick D. Normal butane would have the highest boiling point. Now the N, for normal butane, just indicates that it's just a straight alkane, with one, two, three, four, carbons.
The other choice would be the isobutane. One, two, three, with the fourth up here. So you've got one, two, three, four, and they're not in a straight linear chain. Let's do another set. Which of these has the highest boiling point? Water, hydrogen sulfide, ammonia, or phosphine?
Now while hydrogen sulfide and phosphine both have hydrogen attached something, they don't actually hydrogen bond. Remember, hydrogen bonding is only between O, N, and F with hydrogen. Now what about water versus ammonia? Well hydrogen bonding is typically stronger between oxygen and hydrogen compounds, than it is between nitrogen and hydrogen compounds.
So we can cross off the ammonia, we're left with A, water. Water will have the greatest hydrogen bonding, resulting in the highest boiling point. Okay, let's do one more. Which of these has the highest boiling point? Well, we scan through these very quick and we see, look.
There is water, hydrogen bonding, this is the answer. But wait a minute, when we're doing the MCAT we need to make sure we don't just pick things that have hydrogen bonding, we need to look through all the answers. Let's confirm that we're right. CH3Cl, this is a dipole. Dipole-dipole interactions less than H2O.
Normal pentane, okay lung dispersion. NACl, hold up for a second, sodium chloride is an ionic compound. Ionic compounds have much higher boiling points than any of the covalent compounds. Don't forget that even though hydrogen bonding leads to higher boiling points for covalent compounds, it will never top the ionic compounds. So ionic compounds will have far and away higher melting points and higher boiling points.
This is something that's often thrown into questions. So make sure that if you see an ionic compound, it has the highest boiling point or melting point, even higher than something that's hydrogen bonding.
Read full transcriptAnd these involved are called induced dipoles. For example, if we have a long, straight alcane like this, and for an instant, let's suppose that all the electrons are on one side of the molecule. Once this happens, these electrons here will repel the electrons on this next molecule and induce a dipole here. So now we have a temporary induced dipole.
We saw earlier that non-polar compounds had a much lower boiling point than polar compounds. Well by inducing a dipole, we've made this a quasi polar compound. And so this will slightly increase the boiling point. Now London dispersion forces are size dependent. Meaning the longer this chain is, or essentially the more electrons that exist in the compound, the easier it is to polarize the compound, which makes it so that the London dispersion forces are stronger.
If you don't have very many electrons, then you can't polarize it. Now the next intermolecular force are the dipole-dipole interactions. These occur between polar molecules. An example of a polar molecule would be CH3Cl or chloromethane. Chloromethane is a carbon, with three hydrogens and a chlorine attached to it. Now the carbon-hydrogen bonds are essentially non-polar.
Cuz the carbon-chlorine bond is polar. Which results in a slightly negative charge on the chlorine, because it is more electronegative and will suck up those electrons to hold on to it. And the carbon which is slightly positive. This results in a much greater attraction between the molecules for each other in the solid and the liquid.
In the vapor, there is still this attraction if they come close to each other, but since vapor molecules tend to be so far apart, this force isn't as strong. The result of this dipole-dipole interaction though, is that molecules wanna stay as the liquid for as long as possible, because there's intermolecular force holding them together.
Now you might think since CH3Cl is slightly polar, if we added a whole bunch of chlorines that way we can make it even more polar. Let's do carbon tetrachloride, CCl4. Now while it's true, that the carbon won't have a slight positive charge, and these a slight negative, each chlorine. In effect they cancel each other out, so to determine if a molecule is polar or not, only must it have a polar covalent bond in it, it can't be cancelled out by others in symmetry.
So carbon tetrachloride is in fact non-polar, even though it contains polar bonds. So when you see a compound that has electronegativity difference as greater than 0.4, you can't just assume that it's going to be polar, unless you know the bond connectivity. And the last of the intermolecular forces is hydrogen bonding.
This results from hydrogen being bonded to nitrogen, oxygen, and fluorine. So for example. Water has hydrogen bonded to oxygen. Well this hydrogen is strongly attracted to the oxygen. And the resulting intermolecular force significantly raises the boiling point of water.
Again nitrogen to hydrogen, as an amines. Oxygen to hydrogen, as in alcohols, and other compounds. These bondings will significantly increase the boiling point over something that doesn't have a hydrogen bonding. Now fluorine, we have FH which is just hydrogen fluoride. Once fluorine is attached to this hydrogen, that's it.
If you try to attach an alcohol group and the fluorine, you'd notice that this violates the octet rule because fluorine only needs one extra electron, which is already supplied by the hydrogen. Now which of these three forces is the strongest? Hydrogen bonding, dipole-dipole, London dispersion? Hydrogen bonding is far and away the strongest.
London dispersion forces, which are just induced dipoles and not permanent dipoles, are the weakest. The dipole-dipole interactions are in between the two. Now remember, while London dispersion forces are the weakest and hydrogen bonding the strongest, if the chain, or something like an alkane group starts to get very long.
All of a sudden the hydrogen bonding is a much smaller portion of the total molecule. And becomes less significant than the London dispersion forces. This doesn't mean that hydrogen bonding is weaker, just that the molecule is so big at this point, hydrogen bonding can't overcome the rest of the molecule. So let's use what we've learned to determine a few boiling point trends.
Which of these has the highest boiling point? Methane, methane propane, or normal butane? In this case, we're looking purely at London dispersion forces. These are greatest when the size is largest, so we would pick D. Normal butane would have the highest boiling point. Now the N, for normal butane, just indicates that it's just a straight alkane, with one, two, three, four, carbons.
The other choice would be the isobutane. One, two, three, with the fourth up here. So you've got one, two, three, four, and they're not in a straight linear chain. Let's do another set. Which of these has the highest boiling point? Water, hydrogen sulfide, ammonia, or phosphine?
Now while hydrogen sulfide and phosphine both have hydrogen attached something, they don't actually hydrogen bond. Remember, hydrogen bonding is only between O, N, and F with hydrogen. Now what about water versus ammonia? Well hydrogen bonding is typically stronger between oxygen and hydrogen compounds, than it is between nitrogen and hydrogen compounds.
So we can cross off the ammonia, we're left with A, water. Water will have the greatest hydrogen bonding, resulting in the highest boiling point. Okay, let's do one more. Which of these has the highest boiling point? Well, we scan through these very quick and we see, look.
There is water, hydrogen bonding, this is the answer. But wait a minute, when we're doing the MCAT we need to make sure we don't just pick things that have hydrogen bonding, we need to look through all the answers. Let's confirm that we're right. CH3Cl, this is a dipole. Dipole-dipole interactions less than H2O.
Normal pentane, okay lung dispersion. NACl, hold up for a second, sodium chloride is an ionic compound. Ionic compounds have much higher boiling points than any of the covalent compounds. Don't forget that even though hydrogen bonding leads to higher boiling points for covalent compounds, it will never top the ionic compounds. So ionic compounds will have far and away higher melting points and higher boiling points.
This is something that's often thrown into questions. So make sure that if you see an ionic compound, it has the highest boiling point or melting point, even higher than something that's hydrogen bonding.
