L26 IMFs P3. Dipole Attractions

Summary

This video by Fogline Academy delves into the fascinating world of dipole-dipole attractions in polar molecules like iodine chloride and HCl. It discusses how these attractions impact boiling points and polarity, using various examples to illustrate. A close examination of concept dipole moments sheds light on the relationship between molecular structure and boiling points, as well as the quantitative measure of polarity. The video emphasizes how dipole moments and molecular weight interplay to determine the boiling point of substances with similar molecular weights, offering insights into the ranking of boiling points based on dipole moments.

Highlights

• Dipole-dipole attractions are like magnets for polar molecules, drawing them closer thanks to partial positive and negative sides. 🧲
• Higher boiling points are a typical sign of strong dipole-dipole interactions between polar molecules. 🔥
• Iodine chloride and HCl are classic examples illustrating these powerful dipole attractions. 🌌
• Molecular weight plays a crucial role; heavier non-polar molecules might have higher boiling points than lighter polar ones due to dispersion forces. ⚖️
• The dipole moment is key in understanding molecule polarity; it's both about charge separation and the distance between charges. 📏

Key Takeaways

• Permanent dipoles in polar molecules lead to stronger attractions and higher boiling points. 🔥
• Dipole-dipole attractions are affected by molecular weight and polarity. ⚖️
• Comparing substances with similar molecular weights reveals how polarity influences boiling points. 🌡️
• Understanding dipole moments helps predict boiling point trends in molecules. 📈
• Some unexpected trends in boiling points are explained by examining other forces beyond just dipole moments. 🤯

Overview

In this enlightening video on dipole attractions from Fogline Academy, viewers are introduced to the concept of dipole-dipole attractions in polar molecules. These attractions occur due to the asymmetric charge distribution in molecules like HCl and iodine chloride, leading them to exhibit stronger interactive forces and subsequently higher boiling points than non-polar counterparts.

The discussion takes a deeper dive into why molecules like carbon monoxide, even being polar, have lower boiling points compared to non-polar bromine. It turns out that induced dipole attractions play a significant role too, especially when larger molecular weights come into play to overshadow the influence of dipole-dipole interactions between molecules.

The video thoroughly examines dipole moments, breaking it down into its core components of charge and distance. With examples like iodine chloride to highlight significant differences, viewers gain insights into how molecule geometry and bond lengths affect these moments and overall polarity, further affecting boiling points. Concluding with organic molecules, it stresses how similar molecular weights reveal polarity's guiding role in boiling point trends.

Chapters

• 00:00 - 00:30: Introduction to Dipole-Dipole Attractions In the chapter titled 'Introduction to Dipole-Dipole Attractions,' the focus is on understanding the forces between permanent dipoles, specifically in polar molecules, which are commonly referred to as dipole-dipole attractions. These attractions occur in any polar molecule, such as iodine chloride, HCl, or water. The essential aspect discussed is the presence of an asymmetric charge distribution within these molecules, leading to partial positive and partial negative charges on different sides of the molecule.
• 00:30 - 03:00: Relation Between Polarity and Boiling Points The chapter focuses on the relationship between the polarity of molecules and their boiling points. It explains that polar molecules exhibit stronger attractions due to positive and negative attractions, leading to higher boiling points. This concept is illustrated through a comparison between nitrogen (N2) and carbon monoxide (CO).
• 03:00 - 10:00: Role of Dipole Moment The chapter discusses the role of the dipole moment in influencing the boiling points of substances. It highlights that carbon monoxide, despite having the same molecular weight as nitrogen, has a slightly higher boiling point due to its dipole moment. Similarly, differences in polarity explain the boiling points of bromine and iodine chloride.
• 10:00 - 15:30: Dipole Moment and Boiling Point Trends in Organic Molecules The chapter explores the relationship between dipole moments and boiling point trends in organic molecules. It starts by examining a case where a polar molecule with a dipole moment has a higher boiling point than a non-polar molecule, despite having the same molecular weight. The example given compares a polar substance with a boiling point of 97°C to bromine with a boiling point of 59°C. The chapter suggests that understanding these subtleties requires deeper analysis, such as comparing carbon monoxide (polar) to bromine (non-polar) and examining their boiling points to understand the influence of polarity on boiling point.
• 15:30 - 20:00: Ranking Boiling Points of Compounds The chapter discusses the boiling points of different compounds and the factors affecting them. It highlights that bromine has a much higher boiling point than carbon monoxide despite being nonpolar. This is attributed to the presence of dispersion forces, or induced dipole attractions, which are significant even in nonpolar molecules.
• 20:00 - 20:00: Anomaly in NH3 Boiling Point The chapter discusses the concept of molecular attractions, specifically focusing on polar molecules and their interactions. It highlights how polar molecules not only have dispersion forces or induced dipole attractions but also have additional attractions due to their permanent dipoles. However, the chapter emphasizes that the additional attraction from permanent dipoles is insufficient to compensate for significant differences in molecular weight when comparing two substances. The chapter likely aims to explain why NH3 (ammonia) has an anomalously high boiling point relative to its molecular weight.

L26 IMFs P3. Dipole Attractions Transcription

• 00:00 - 00:30 So in this video, we'll talk a little bit about  the attractions between permanent dipoles,   that is polar molecules, what we would call  dipole-dipole attractions. So, these are the type   of attractions that you would find in any polar  molecule. For example, iodine chloride, or HCl,   or water, for example, where once again, we're  talking about a molecule that has an asymmetric   charge distribution, and so a partial positive  and partial negative side to the molecule that
• 00:30 - 01:00 are attracted to each other. Now, of course, it's  a little more obvious here the fact that these   molecules are attracted to each other based  on positive, negative attractions, and it's   no surprise that we should expect that polar  molecules that have these type of attractions   are going to be more strongly attracted and,  therefore have higher boiling points. And in fact,   that is the case so if we compare saying  nitrogen, N2, to carbon monoxide, CO, it's
• 01:00 - 01:30 no surprise that carbon monoxide has a slightly  higher boiling point, about 4 degrees higher,   then does the nitrogen even though they have  exactly the same molecular weight. And similarly,   if we look at say bromine, Br2, that's completely  nonpolar, it's no surprise that iodine chloride,
• 01:30 - 02:00 which is polar, has a higher boiling point,  97, as compared to bromine, 59, even though   they essentially have the same molecular weight.  Now, there's a couple of subtleties here that we   need to figure out. Now, one of those is if  we compare carbon monoxide to bromine. So,   in that case we know that carbon monoxide is polar  and bromine is non-polar, but yet it's clear that
• 02:00 - 02:30 bromine has a much higher molecular, or sorry  a much higher boiling point, then does carbon   monoxide even though the bromine is nonpolar.  Now the reason for that is because it turns out   that those dispersion forces, those induced dipole  attractions, exist not only in nonpolar molecules,
• 02:30 - 03:00 but really in all molecules. And so, a better  way to think about the attractions between   polar molecules is that polar molecules have these  dispersion, or induced dipole attractions, plus a   little bit extra due to its permanent dipole. But  that little bit extra due to the permanent dipole   is not enough to make up for a large difference  in molecular weight. So, if we're comparing two   substances, where one has a much higher molecular  weight than the other that substance is going to
• 03:00 - 03:30 have a higher boiling point, even though it might  be nonpolar, and the other substances polar. So,   you can really only use polarity in order  to give it to rank a substance with a higher   boiling point, if they have comparable molecular  weights. Now, the other thing to notice here is
• 03:30 - 04:00 to look at the difference in the boiling points  in the two cases. In other words, when we look at   bromine as compared to iodine chloride we see  that when, by being polar the iodine chloride   has a dramatically higher boiling point than the  bromine does, roughly 38 degrees Celsius higher as   a result of being polar. Whereas when we look at  the carbon monoxide in comparison to the nitrogen,
• 04:00 - 04:30 we only find a four degree difference in boiling  points. So, the polarity doesn't really seem to   add a whole lot in this particular case. So we  might wonder why is polarity so important in one   case and not in the other case? And that brings  up another issue that we need to go back and think   about a little bit more, and that is the concept  of dipole moment. So, if we remember back in
• 04:30 - 05:00 Chapter nine, the textbook discussed when we were  discussing the concept of polarity and molecules,   the idea that if you have some molecules, such as  N2, whose Lewis structure is shown here versus say   carbon monoxide, with the Lewis structure shown  here, that the way that we figure out whether
• 05:00 - 05:30 a molecule is polar or not is by looking at  electronegativities. So, in the case of nitrogen   both atoms are identical. And so, we would say  that nitrogen is nonpolar because both atoms   pull equally. But in the case of carbon monoxide,  where they have different electronegativities,
• 05:30 - 06:00 we know that the oxygen pulls more strongly on the  electrons than does the carbon, and we would say   that carbon monoxide is polar. And of course,  we can extrapolate that to more complicated   molecules, such as carbon dioxide or water, where  we not only need to take into account the polarity
• 06:00 - 06:30 of the individual bonds, but also what the net  effect is in a molecule where you have multiple   bond polarities and how those bond polarities add  up. So, in the case of carbon dioxide the fact   that they cancel out creating a nonpolar molecule,  and in the case, of water where they do not cancel   out and you end up with some net polarity. But the  other thing that was brought up in Chapter 9, is
• 06:30 - 07:00 it's not only a case of being polar or nonpolar,  but sometimes we want an actual quantitative   measure of the amount of polarity of a molecule.  And that quantitative measure is called dipole   moment, usually represented by the Greek letter  mu, and having units of Debye, usually abbreviated   with a capital D. And as your book discussed, the  dipole moment is not only about the fact that the
• 07:00 - 07:30 electrons are shifted, but also about the distance  over which they are shifted. So, it turns out that   dipole moment is equal to Q times R, where Q here  represents the amount of positive and negative   charge that's separated, and R represents the  distance over which those charges are separated.
• 07:30 - 08:00 So, when we're talking about a polar molecule  what we're really talking about here is that Q   positive and negative represent these fractional  or partial charges that are created by shifting   the electron cloud, R, the distance really has  to do with the bond length in the molecule, so   how far apart are those atoms. So, the reason that  turns out to be important is if we're to go back
• 08:00 - 08:30 and look at carbon monoxide, we recognize that  carbon and oxygen are both relatively small atoms,   and more importantly we see that in the Lewis  structure there is a triple bond. And so the   distance over which the charges are separated in  carbon monoxide is quite small because we remember   that a triple bond is shorter than a double bond,  a double bond is shorter than a single bond,
• 08:30 - 09:00 in other words, that bond length is quite small  and carbon monoxide, Whereas if we were to look at   something say like iodine chloride, where even  though the difference in electronegativities   between iodine and chlorine is actually not as  large as difference between carbon and oxygen,   iodine chloride has a single bond and both of  the atoms are much, much larger than carbon and
• 09:00 - 09:30 oxygen. And so even though, in some sense, the  charge separation in terms of electronegativities   is not as extreme in iodine chloride, the  distance over which the charges are separated   is much more dramatic. And as a result, the dipole  moment for carbon monoxide is quite small, only   0.1 Debye, to by whereas the dipole moment for  iodine chloride is much much larger, 1.6 Debye,
• 09:30 - 10:00 due to this much larger distance over, which the  charges are separated. And so in fact, iodine   chloride is in a sense much more polar. Now, it  turns out that dipole moments can, for relatively
• 10:00 - 10:30 simple molecules, be calculated using physics and  concepts about electron charge distribution in   molecules, and so forth, but in a lot of ways it's  easier and more accurate to measure dipole moments   of molecules, which is done using by putting  molecules in the gas phase in electrical fields,   often fluctuating electrical fields, and measuring  how the molecules interact with those electrical
• 10:30 - 11:00 fields. And so as a result, there are measured  dipole moments for a large number of different   molecules that you can look up. And really  then, what we need to know when trying to make   predictions about the strength of IMF's between  molecules is how large is that dipole moment for   a particular molecule that we're interested. So,  if we know now go back to our previous discussion,   we can compare the dipole moments of the molecules  that we were talking about, and we can see that
• 11:00 - 11:30 when you compare carbon monoxide to nitrogen see  the difference, in dipole moment is fairly small,   and so it's no surprise that there's a relatively  small friends and boiling points. Whereas when we   compare iodine chlorine and bromine, we see a much  larger difference in dipole moments, and therefore   a much larger difference in boiling points. And to  reinforce that, we can bring up one other example,
• 11:30 - 12:00 HCN where here the molecular weight is pretty  similar to that of nitrogen and carbon monoxide   however, because now, we are we have three atoms  and so we can stretch that charge separation   over a much larger distance, we end up with a  molecule that has a very large dipole moment,   three Debye, and as a result the boiling point  of HCN is dramatically higher than the other two
• 12:00 - 12:30 molecules of similar molecular weight. Although  once again, it's still not enough to make up for   the difference in molecular weight with these two  larger molecules, so even though it is quite polar   it still has a lower boiling point than bromine  and iodine chloride because of their much larger   molecular weights, and therefore much larger  induced dipole attractions. Now, we can take that
• 12:30 - 13:00 concept of dipole moments and dipole-dipole forces  and apply it to some of our organic molecule   examples that we've been thinking about recently.  So, here we have a number of different organic   molecules that all have very similar molecular  weights. They're all in sort of the low to mid   40 of molecular weight. So, we have propane,  dimethyl ether, ethylene oxide, acetaldehyde,
• 13:00 - 13:30 and acetonitrile. And along the x-axis, we see  that we have plotted the dipole moment for each   of these different substances. And, no surprise,  since they have similar molecular weight the ones   with the stronger dipole moment will tend to have  stronger IMF attractions between molecules and so   will tend to have a higher boiling point. And  it's worth taking a moment just to focus in on
• 13:30 - 14:00 three of these, that is propane, dimethyl ether,  and acetaldehyde, because these are all functional   groups that we've looked at previously, so  we have a regular straight chain alkane that   is essentially nonpolar, we see that it's a  dipole moment is close to zero, because there's   very little difference in electronegativity  between carbon and hydrogen and the molecule is
• 14:00 - 14:30 essentially symmetric. When we look at the ether,  here we have a slightly different situation. Here   we have remember a molecule where there's an  oxygen in between two carbon atoms, and as a   result of that, we have polar bonds between the  oxygen and the carbon atoms and because it's bent   around that oxygen those polarities do not cancel  out and we have some net polarity for an ether.
• 14:30 - 15:00 And of course we see that it does have a dipole  moment, but that dipole moment is not extremely   large mostly because of the fact that a lot of  that carbon-oxygen bond polarity gets canceled out   because of the shape of molecule. If on the other  hand we look at acetaldehyde, and remember that an   aldehyde similar to a ketone has a C double-bonded  O, with the geometry as shown here, and of course,
• 15:00 - 15:30 here once again we have this carbon oxygen  bond polarity that points towards the oxygen,   and in this case, there's nothing cancelling out  that bond polarity. And so the dipole moment is   quite a bit larger. And so that's a lesson we can  keep in the back of our mind, that when we talk   about aldehydes and ketones with that carbonyl  group we're going to expect to see higher dipole
• 15:30 - 16:00 moments than we would in the case of an ether,  where although they're somewhat polar the dipole   moments are smaller. So finally, we can now take  this information and apply it to rank the boiling   points of the five substances shown here. Once  again, I would encourage you to pause the video   and then think through, the problem-solving  here, and come up with your own ranking.
• 16:00 - 16:30 The things that you're gonna do first is to figure  out what's the molecular weight of each of these   substances. So, armed with the periodic table you  can quickly come up with these molar masses for   these substances. And you'll notice what right  away that the first two in the list have nearly   the same molecular weight. And the other three  have roughly double that amount. Next of course,
• 16:30 - 17:00 we would need to draw the lewis structure  for each of these substances and think   about whether the molecules are polar or not.  And as we just discussed in addition to that,   we need to think about the dipole moment, that  is the charge separation. And to save us some   time I'm just going to display those  dipole moments here. Now, hopefully,
• 17:00 - 17:30 in drawing Lewis structures you were able  to figure out pretty quickly that CH4,   methane, and SiH4, which has the same lewis  structure, are both completely nonpolar,   so they have zero dipole moment. And you  should been able to conclude that NH3 and PH3,   which remember because of that lone  pair on the nitrogen or phosphorus
• 17:30 - 18:00 has that trigonal pyramid shape, and so  is polar is gonna have a dipole moment,   and H2S, to remember has a little structure  very similar that of water, is bent and of   course it's polar. As well now with the actual  dipole moments here we can now finally make   some conclusions. First, as we talked about the  primary guiding factor is molecular weight. So,
• 18:00 - 18:30 we would expect the three bottom ones in the  list to have the higher boiling points. And we   should expect the two at the top of the list  to have lower boiling points because of that   difference in polarizability or induced dipole  attractions. And then within each of those groups,   of course, the more polar that is the higher the  dipole moment the higher the boiling point. So,   if we look at the actual boiling points, we see  that that is essentially the case. That CH4,
• 18:30 - 19:00 which is the smallest molecule and is nonpolar has  the lowest boiling point. And then H2S, which has   the highest molecular weight and is the most polar  of those three substances has the highest boiling   point of the substances listed here. However,  the one surprise is NH3. So, it turns out that
• 19:00 - 19:30 although the boiling point of NH3 we would predict  it to be in between CH4 and SiH4, that is it's   more polar than methane so it should be higher  than that, but because of its small size we expect   it to be lower than SiH4. But in fact, it's higher  and not only is it higher than that, it's higher
• 19:30 - 20:00 than all the other ones in the list. So, the  question, of course, then is why does NH3 have   a much higher boiling point than expected based  on the trends? Which brings up our next video.