The video by Fogline Academy delves into the intriguing phenomenon of induced dipole attractions, or London dispersion forces. While nonpolar molecules like nitrogen and halogens appear to lack attractions due to their symmetrical electron distribution, they can still condense into liquids or solids thanks to temporary induced dipoles. These fleeting attractions occur because electrons in atoms can shift momentarily, causing temporary positive and negative regions. Larger atoms with more electrons are more easily polarizable, exhibiting stronger induced dipole interactions. This principle explains why molecules with larger electron clouds have higher boiling and melting points. The video also highlights a fun analogy of molecular velcro, illustrating how larger surface areas between molecules enhance attraction, influencing substances' physical properties like boiling points.
Highlights
Temporary induced dipoles allow nonpolar molecules to temporarily attract. 🌀
London forces explain liquid and solid forming in nonpolar molecules. ❄️
Larger electron clouds in big atoms lead to stronger attractions. 💪
Boiling points increase with molecular weight due to dipole strength. 🔥
Shape and molecular structure influence the strength of dispersion forces. 💼
Key Takeaways
Nonpolar molecules can have attractive forces due to induced dipoles. 🌟
Induced dipoles explain the liquid and solid states of nonpolar substances. 🌊
Larger atoms with bigger electron clouds exhibit stronger induced attractions. 🌌
Molecular weight correlates with polarizability and boiling points. 📈
Molecular 'velcro' illustrates how shape affects intermolecular forces. 🤹♂️
Overview
Ever wondered how nonpolar molecules like nitrogen or bromine can exist in liquid and solid states? Enter the world of induced dipole attractions—a fascinating realm where temporary electron shifts momentarily create positive and negative regions, allowing these molecules to, quite literally, stick together! 🌟
Fogline Academy shines a light on how larger molecules are more easily 'polarizable,' meaning their electron clouds shift more dramatically, hence intensifying these temporary attractions. It's a dance of electrons that makes big molecules like iodine solid at room temperature and gives bigger atoms a knack for forming liquids. 🌊
Imagine molecular velcro, where the more area you have, the stronger the bond. This analogy perfectly captures how larger molecular surfaces with more electron clouds exhibit stronger dispersion forces, ultimately affecting things like boiling points. It's why pentane boils at room temperature while dimethyl propane doesn't—it's all about those attractive forces! 🤹♀️
Chapters
00:00 - 01:00: Introduction to Induced Dipole Attractions The chapter introduces the concept of induced dipole attractions, highlighting their similarity to ionic bonds in terms of involving attractions between positive and negative partial charges on molecules. It poses a question about the nature of nonpolar molecules like nitrogen in this context.
01:00 - 03:00: London Forces and Dispersion Forces This chapter discusses the existence of halogen molecules like bromine and iodine in liquid or solid states despite being nonpolar. It highlights the question of how nonpolar diatomic molecules with identical atoms sharing electrons equally can exist in condensed states such as liquids and solids at room temperature (notably iodine, which is a solid).
04:00 - 06:00: Strength of Induced Dipole Attractions The chapter discusses the concept of induced dipole attractions, also known as London forces or dispersion forces. It emphasizes that although certain molecules lack distinct positive or negative sides, there are still attractive forces at play. These forces allow the molecules to adhere together, facilitating the formation of condensed liquids or solids.
06:00 - 09:00: Boiling Points of Halogens and Hydrocarbons This chapter explains the boiling points of halogens and hydrocarbons, emphasizing the role of intermolecular attractions. It discusses how even in neutral and nonpolar atoms or molecules, like a helium atom—a noble gas that doesn't bond—the electron distribution can vary. This is attributed to the electron's nature of being like a cloud, leading to instantaneous polarities and attractions despite an overall neutral charge.
09:00 - 12:00: Molecular Velcro and Effect of Shape This chapter discusses the concept of electron distribution within atoms and molecules, specifically focusing on moments when the electron density is asymmetric. Such asymmetry can lead to a concentration of electrons on one side of the atom. This distribution results in a temporary negative charge that can influence adjacent electron clouds by repelling them. The chapter explores the implications of this fluctuating electrical nature.
12:00 - 15:00: Ranking Boiling Points of Different Molecules The chapter discusses how a temporary dipole can be induced in molecules. When a cloud shifts a neighboring electron cloud away, it creates an asymmetrical distribution of electron density. This causes one side of the atom to be slightly positive and the other to be slightly negative, leading to an attraction between these induced dipoles.
00:00 - 00:30 In this video, we're gonna talk briefly about
what are known as induced dipole attractions. So now that we've established that intermolecular
forces are in some ways analogous to ionic bonds, that is they involve attractions between positive
and negative partial charges on molecules, brings up an interesting question, which is how
is it that nonpolar molecules, such as nitrogen or
00:30 - 01:00 the halogens, bromine, iodine, and so forth, how
is it that these molecules can exist as liquids, or in some cases even solids, iodine at room
temperature is a solid, how is it that these things can exist in these condensed states if
they're completely nonpolar because, of course, by definition these diatomic molecules have two
identical atoms that share electrons equally,
01:00 - 01:30 so there is no positive or negative side of
such molecules, and yet there must be some sort of attractive force between these molecules that
allows them to stick together and form a condensed liquid or solid. So, these types of attractions
go by various names, often called induced dipole attractions, London forces, dispersion forces,
essentially the concept of these induced dipole
01:30 - 02:00 attractions is that in atoms or molecules that
are completely neutral and nonpolar, for example, let's take a helium atomm remember helium is a
noble gas so it doesn't bond to anything, and in theory at any given instant the electron should
be uniformly distributed around the atom. But since electrons we think of as being this sort of
cloud or amorphous moving material, it's possible
02:00 - 02:30 that the electron distribution, or the electron
density can for at least a moment be asymmetric call in an atom or molecule. And if, for example,
we have electrons concentrated on one side of the atom in the cloud at some instant, that will then
create a slight excess of negative charge that may repel the electrons in the neighboring cloud.
And if that happens, that fluctuating electrical
02:30 - 03:00 cloud can temporarily induce a dipole because
as it shifts the neighboring electron cloud away from itself it creates an asymmetry in the two
clouds. And as soon as that asymmetry occurs, it means that one side of one of the atoms
is slightly positive, while the other side is slightly negative, and has more electron density.
And as a result there's an attraction between the
03:00 - 03:30 slight positive of one atom and the slight
negative of the neighboring atom. And so we would say that we have these mutually reinforcing
induced dipoles. Now of course. in a gas or even in a liquid because these are temporary dipoles
that are just caused by fluctuations in electrical clouds, these induced dipoles and these temporary
polarities will disappear almost as soon as
03:30 - 04:00 they're formed. So, in one moment we'll have this
induced dipole in this attraction, in the next instant these electron clouds will fluctuate back
in the other direction. However, if you have a large number of atoms or molecules, if you sort of
time averaged over all of the atoms and molecules in the sample there is enough of these temporary
fluctuating attractions to allow these materials,
04:00 - 04:30 if it's cool enough, to condense into a liquid
or solid state. And so once again these temporary fluctuating dipoles are called dispersion forces,
or induced dipole, induced dipole attractions. Now, the strength of these induced dipole
attractions turns out to be very important in understanding the trends in all sorts of physical
properties of lots of different substances. And so
04:30 - 05:00 we want to talk a little bit about the idea of
what causes these induced dipole or dispersion forces to be stronger in some substances than
others. So basically, the concept here is if we were to compare two different atoms, one that's
relatively small and one that's relatively large, and think about the size of that electron cloud
and what happens as we shift the electron cloud
05:00 - 05:30 and distort it. We recognize that in a larger
atom with a large electron cloud, that's not held as tightly, that electron cloud can shift
more dramatically than it can in a small atom. And as a result, it creates a larger asymmetry,
where there's more positive and negative, in that large atom, than there would be in a smaller atom
that shifts. And so we would say that these larger
05:30 - 06:00 atoms are more polarizable, meaning it's easier to
induce this temporary dipole and polarity in these large atoms. Some evidence to show that that's
the right way to think about things is to look, for example, some substances that are very similar
but have different sized electron clouds. And so,
06:00 - 06:30 for example, we could look at me halogens:
fluorine, chlorine, bromine, iodine, that are all in the same column, they're all diatomic
nonpolar molecules. And if we look at the boiling points and melting points of these substances,
we recognize that, of course, for fluorine, the melting and boiling point is lower than it is
for chlorine. As we move from chlorine to bromine,
06:30 - 07:00 it goes up even more. In fact, fluorine and
chlorine are both gases at room temperature, while bromine is a liquid, so it's boiling
point is above room temperature, and then as we move from bromine to iodine it goes up even
more, that is the iodine molecules are even more strongly attracted and are in fact solid at
room temperature. Now, since the size of electron clouds goes up with the number of electrons, and
since the number of electrons goes up with the
07:00 - 07:30 number of protons and also neutrons, there is
a strong correlation between molecular weight and polarizability. So, in essence molecules
that have a larger molecular weight will tend to have larger atoms with larger electron clouds
that are more easily distorted and so they will tend to have higher melting and boiling points.
Now, a similar but slightly different version of
07:30 - 08:00 that argument can be seen here, where we have a
graph of boiling point versus molecular weight, for a series of hydrocarbons: pentane,
hexane, heptane, octane, nonane. Now, at first one thing that might be confusing about
this is that in every one of these molecules
08:00 - 08:30 we're talking about the same sized atoms. We
have carbon atoms and hydrogen atoms, and so regardless of which molecule we're talking about
the polarizability of a carbon atom is roughly the same in all of them same for the hydrogen atom,
and yet as the chain of the molecule gets longer and longer. We can see that the boiling point goes
up, and so obviously there must be stronger and stronger attraction between molecules. And one
common way to think about this is in terms of
08:30 - 09:00 what they often call molecular velcro, as being
a nice analogy to describe dispersion forces, and simply the concept here is that like having
pieces of velcro that are, say, attached to wood, if you have two relatively small pieces of wood
that have velcro on them and are stuck together, they will, those two boards or pieces of wood will
be easier to pull apart, then if say you had two
09:00 - 09:30 longer boards that were covered with velcro
and stuck together. It would take more effort to separate those two boards. And so essentially
even though per square inch there's no difference in the attraction of the velcro, the fact that you
have more square inches or more area of velcro on the long molecules means that will be harder
to separate those molecules and they'll have higher boiling points. And that way of thinking
is reinforced even further in thinking about the
09:30 - 10:00 effect of shape on these dispersion forces. So,
for example, let's take two different isomers of c5h12 on the left we have pentane. five carbons in
a row, and on the right we have dimethyl propane, an isomer of that right but arranged in the
slightly different way, a structural isomer.
10:00 - 10:30 And if we think about the molecular velcro on
these molecules we consider that in the case of pentane, where the carbons are all in what
we would call a straight chain arrangement, there's a larger area for interaction between a
larger number of those carbon and hydrogen atoms, then there is in the dimethyl propane, which is
sort of a more globular kind of shaped molecule. And not as many atoms can interact with each other
on adjacent molecules smaller area for interaction
10:30 - 11:00 less attractive force. And, of course, this means
that the boiling points for the two substances are different, in fact, quite dramatically different
so much, so that pentane on the left has a boiling point that's above room temperature, that is
pentane is a liquid, whereas dimethyl propane, on the right, has a boiling point that's below
room temperature. And so it's a gas at room
11:00 - 11:30 temperature, even though the two substances have
exactly the same molecular weight. So finally, we can take those concepts and apply them you can
challenge yourself to rank the boiling points of the four substances listed here. So, I would
encourage you to pause the video and to think
11:30 - 12:00 about what we just talked about and try to rank
the boiling points of these four substances. Of course, one of the first things we need to
think about is what are the molecular weights and structures of these different substances. Of
course, we know that oxygen, o2, is a diatomic molecule with two oxygen atoms, molar mass of
32. Nitrogen is another diatomic molecule, n2,
12:00 - 12:30 with a molar mass of 28. We remember butane from
our study of organic chemistry, is CH3CH2CH2CH3, whereas methyl propane is CH3CHCH3, with a methyl
group sticking off the middle carbon. Of course,
12:30 - 13:00 they're both c4h10, so they both have a
molar mass of 58, but they're two different structural isomers. And of course, we know
that based on what we just talked about, that the smallest boiling point or the lowest
boiling point of these should be for nitrogen,
13:00 - 13:30 since it has the lowest molecular weight. Then,
second should be oxygen, 32. And then, the highest boiling point should be the two isomers
of butane. And of course, the straighter version butane should have a higher boiling point than
the methyl propane, which because of this shape does not stick together as well. So, if we look
at the actual boiling points of these substances, which are listed here, we recognize that is, in
fact, the case nitrogen has the lowest of these.
13:30 - 14:00 And you'll note that I put these in both Kelvin
and Celsius, since the celsius temperatures are mostly negative numbers makes a little harder
to compare. If we look instead at the absolute boiling point, and Kelvin we can see that, in
fact, that nitrogen is the lowest number, 77, and it goes up next oxygen, then methyl propane, and
finally the highest of these butane has expected.