Solubility and intermolecular forces | Chemistry | Khan Academy
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Solubility and intermolecular forces | Chemistry | Khan Academy

Now that we know what a solution
is, let’s think a little bit about what it takes
to get a molecule to be soluble into a solution
or into a solvent. So let’s say I start off with
a salt, and I’ll do a little side here, because in chemistry,
you’ll hear the word salt all the time. Let me write it down: salt. And in our everyday language,
salt is table salt. It makes food salty,
or sodium chloride. And this indeed is both a salt
from the Food Channel point of view and from the chemistry
point of view, although the chemistry point of view does not
care about what it does to season your food. The chemistry point of view, the
reason why it’s called a salt is because it’s
a neutral compound that’s made with ions. So we all know that this is
made when you take sodium. Sodium wants to lose its one
electron in its valence shell. Chloride really wants to
take it, so it does. Chloride becomes a negative ion
and sodium is a positive ion, and they stick to each
other really strongly because this guy’s positive now, and
this guy’s negative after he took away his electron. Imagine your house is too small,
so you have to give away your dog to someone who has
room for the dog, but now you have to hang out at that
person’s house all the time because they have the
dog you love. I don’t know if that analogy
was at all appropriate. But I think you get the idea. A salt is just any compound
that’s neutral. The other common ones, potassium
chloride, you could do calcium bromide, or I could
do a bunch of them, but these are all salts. And what we want to think about
is what happens when you try to essentially dissolve
these salts in water. So we know what water is
doing, liquid water. So let me draw some
liquid water. So if that’s the oxygen and then
you have two hydrogens that are kind of lumping off of
it, I’ll draw it like that. I’ll draw a couple of them. And then, of course, you have
another oxygen here. Maybe the hydrogens are in this
orientation because the hydrogen ends are attracted
through hydrogen bonds– we’ve learned this– to the oxygen
ends because this has a slight negative charge here, a slight
positive charge here. These are the hydrogen
bonds that we’ve talked so much about. And maybe you have another
oxygen here and it’s got its hydrogens there and there. You have some hydrogen
bonds there. I could do another oxygen here,
and you can kind of see the structure that forms,
although what I’m drawing, this is actually more of a– if
you were in a solid state, this would be kind of
rigid and they would just vibrate in place. In the liquid state, they’re
all moving around. They’re rubbing up against
each other, but they’re staying very close. Actually, the liquid state for
water is actually the most compact state for water. Now, when you’re dealing with
stuff like this– these are moving around, maybe this guy’s
moving that way, that guy’s moving that way– and you
want to dissolve something like sodium chloride. Sodium chloride’s actually
quite a large molecule. If you look at the Periodic
Table up here, oxygen is a Period 2 element. Hydrogen is very small. We know when it gets into a
hydrogen bond with oxygen, it’s really just a proton
sitting out there because all the electrons like to hang out
with the oxygen, while, say, sodium and chloride, they’re
considerably larger. I won’t go into the exact
molecular sizes, but maybe sodium– let’s do sodium–
which actually, just as a review, which is larger. We know that it becomes smaller
as you go to the right of the Periodic Table, so sodium
is quite a large atom, while chloride is a good bit
smaller, but they’re both bigger than oxygen and a lot
bigger than hydrogen. So let me draw that. So sodium– I’ll do sodium
as a positive. It’s pretty big. Maybe it looks like this. Sodium is positive and then
you have the chloride. The chloride I’ll
do in purple. They’re still pretty big. The chloride, it’ll
look like this. And what happens when you
put it into water, it disassociates. Even though these guys in
a solid state, they’re jam-packed to each other. When you put it into water,
the positive cations are attracted to the negative
partial charges on the oxygen side of the water, and the
negative anions are attracted to the positive sides
of the hydrogen. But in order to get, for
example, this sodium ion into the water, it has
to fit in there. So, for example, I drew this as
a liquid initially, but if this was a solid and you had
this structure, it would be extremely difficult. In fact, it would be next to
impossible to squeeze these huge sodium ions in place
to make it soluble into, say, solid ice. And as even cold water, these
bonds are still going to be pretty strong and they’re
going to be just kind of barely moving past each other
because there’s not a lot of kinetic energy. So what you need to do is, the
warmer the water you have– I mean, you can fit it into cold
water, because at least cold water has some give, but the
warmer the better, because you have some kinetic energy, and
that essentially gives space. Or it makes room for this sodium
ion that’s entering in to kind of bump its way into
a configuration that’s reasonably stable. And a reasonably stable
configuration would look something like this. Sodium would look– and then
you’d have a bunch of– sodium is positive. It would be attracted to the
negative end of the water molecules, so the oxygen end. So it looks like that, the
oxygen end, and then the hydrogen ends are going
to be pointing in the other direction. The hydrogen ends are going
to be on the other side. And, of course, the chlorine
atom is going to be very attracted to that other side, so
the chlorine atom might be right over here. So the chlorine atom might want
to hang out right here. In order to get as much of the
sodium chloride into your water sample, you want
to heat up the water as much as possible. Because what that does is it
allows these bonds to not be taken as seriously and these
relatively huge atoms to kind of bump their way in. So, in general, if you think
about solubility of a solute in water– or especially if you
think of a solid solute, which is sodium chloride– into
a liquid solvent, then the higher the temperature while
you’re in the liquid state, the more of the solid
you’re going to be able to get into the liquid, or you’re going
to raise solubility. So temperature goes up,
solubility goes up. For example, if you were to take
some table salt, and you could experiment with this. It doesn’t seem too dangerous
and not too expensive because salt is reasonably cheap. Keep putting it into
a glass, and at some point it’ll dissolve. You could shake it a little
bit, just to make sure. You could think about what’s
happening at the molecular level while you shake it and why
does that help to shake or stir things? But at some point, you’re going
to end up with– if this is your glass of water, the salt
will keep going in there, but at some point, you’ll have
salt crystals at the bottom of your glass. At that point, your water is
saturated with salt at the temperature that you’re trying
to deal with it. Now, right when you start seeing
that, if you were to put it in the microwave or if
you were to heat it up, you would see that even these guys
are able to be absorbed in the water, and that’s because the
extra kinetic energy from the temperature is making it more
likely that these guys are going to be able to bump out of
configuration for just long enough for these guys
to bump in. And just a little side note,
when you take these salts, which are just ionic compounds
that are neutral, they’re made of ions, but they cancel
each other out. When you put them in water,
these compounds by themselves aren’t normally– when they’re
in the solid state, they don’t normally conduct electricity. Even though they’re charged,
they’re very closely stuck to each other, so there’s
not a lot of room for movement of charge. But once you disassociate them
in water or dissolve them in water, now, all of a sudden,
you have these floating charges in the water, and this
does conduct electricity, so it becomes quite a reasonable
conductor of electricity. So the general rule of thumb is,
if you’re dealing with a solid in a liquid solvent,
lowering the temperature will decrease the solubility, because
it’s harder to jam the molecules in there, and
increasing the temperature will increase the solubility. But what about a gas? What if you make some soda and
you want to dissolve some carbon dioxide into, let’s
say, water again? So here, the way to think about
it when we did it with salts, these are ionic
compounds. They had some natural attraction
to the different polar ends of the
water molecule. But gases, for the most
part, do not have strong attractive forces. That’s why they’re gases,
especially at room temperature. They like to be free. A gas, they have a good bit of
kinetic energy, but more important, the bonds between
them, for example, in ideal gases we talked about it, they
just have their London dispersion forces. They have very weak bonds, and
that’s why at, say, the same temperature and pressure that
water would be a liquid, a lot of these gases are gases. They jump away from each other
because they don’t want to touch each other. Now, when you put this in
liquid, and this is at least my intuition, so let’s just say
this is a bunch of water molecules here. If you were to dissolve– let’s
say it’s carbon dioxide. You can ignore this
stuff up here. If you were to dissolve carbon
dioxide in water– so if you were to dissolve this in water,
so those are some carbon dioxide molecules. I’m just drawing the whole
molecule as a circle. What do these molecules
want to do? It’s natural state is a gas and
it is a gas at let’s say the standard pressure, so it
really wants to escape from this water, but it just can’t
do it that easily because there’s water molecules
all around it, right? This guy right here, he might
want to bump out, but he’s surrounded by water molecules. So what would help
him bump out? Well, if you raise the average
kinetic energy of the system, if you made all of these guys,
that these guys were moving faster, and especially if the
carbon dioxide molecules themselves had more kinetic
energy, then maybe they could break out. And as you have from personal
experience with Coke bottles, you could also shake the system,
because if you shake the system, it just moves
everything around enough that these guys can escape. So when you’re dissolving
a gas inside of a liquid solvent, when the solute is
a gas, it actually has the opposite effect, that
rising temperature. So when temperature goes up,
solubility goes down because these guys want to escape. They want to be free. They want to be away from other
molecules and they want to bounce around in open– I
shouldn’t use the word air– in open space. And so anything that lets the
system move around more, they’re going to go up. And likewise, if temperature
goes down, solubility goes up. The other factor, and it’s not
as big of a factor when you talk about a solid solute, but
when you talk about a liquid solute– let me just
do it again. So those are the carbon dioxide
molecules and then you have a bunch of water
molecules– they should all be the same size– that
it’s dissolved in. I think you get the idea. Pressure is also a big factor. I already said that these guys,
their natural state is to roam free. They want to get out. They want to somehow bounce
out of the water. But if you have a really high
pressure up here– just the atmosphere up here has just
tons of molecules bouncing really hard down on the surface
of our solution– so if there’s just tons of
molecules bouncing really hard off the surface, it’ll
be harder for anything to escape upwards. And that’s why, when you have
pressure going up, or at least this is the intuition, when
pressure goes up, solubility of a gas also goes up. And this is for a gas. So just the interesting thing
to remember is that when you think about solubility, solids
do the inverse of gas. Temperature is good for solid
solubility, right? We said when you put salt or
sugar in water, it’s good to increase the temperature. You’ll be able put
more in there. On the other hand, with a
gas, it’s the opposite. You want colder temperatures
to put more gas into the solution, or you want higher
pressure to keep it– at least in the way my mind works– from
escaping out the top. Anyway, hope you found
that useful.

About James Carlton

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100 thoughts on “Solubility and intermolecular forces | Chemistry | Khan Academy

  1. I wish he would put adds on his videos, so he could have a full job by helping hundreds upon thousands of students with their work. I really appreciate this, KhanAcademy, thank you so much! ^^

  2. where is the Like Button, to just Like ever and any Video you Post..I mean, you Teach so well, Easy to Follow and calming…BRAVO for all Videos I have watched, and those I have yet to get to…Thank You Khan!

  3. Once Bill Gates found him, this became his Full time Job. Bill gates pays him to make these for us…Awesome

  4. Sal in an ionic compound the anion gets really big and the cation shrinks so chloride would actually be bigger than sodium. in NaCl.

  5. And how is this relevant to my question? Maybe the smalleness of your organ can be relevant, in order to probe atomic sizes of the ions and also, i guess, for enforcing the unofficial statistics that the size of your penis is directly proportional to the size of your brain.

  6. It's both of the ions acting within the solution. Na+ and Cl- are what is dissolved in a solution, and to have a neutral charge on the original compound you need BOTH the cation and the anion together. so a salt is a combination of a cation and an anion. the ionic nature of the bond within a salt is what allows the Na+ and the Cl- to dissociate into their respective ions…

  7. if non-chemically you are asking what makes salt taste salty, its that your taste buds have receptors for NaCl/ other ionic compounds(salts), not the ions themselves.

  8. "He's surrounded by Waler Mocules" @ 10:45 I always laugh when you stumble man, you say some funny stuff! Thanks BTW!!!!

  9. FYI: Guys he's wrong about the sizes of Na+ and Cl-. He made a video about it, but basically Na+ is bigger than Cl- because when Na loses an electron to become positive, the 2nd shell is exposed and thus the atom becomes smaller.

  10. So.. Solid solvent -> (NACL solid solvent example:) solubility goes up, temp. goes up.

    And Liquid solvent -> (Putting CO2 gas in liquid solvent- water:) solubility goes up, temp. goes down. right?

    What about Gas solvent?

  11. They both have the same number of energy levels but Chlorine has more protons so it is more electronegative……I think……:-)

  12. I want to ask, why is Na bigger than Cl? Cl has more number of electrons revolving and more have protons in nucleus, shouldn't it be a bigger atom than Na?

  13. The Cl atom has a smaller radius BECAUSE it has more protons and therefore it pulls the electrons it has closer to it's nucleus. The number of electrons in this case does not matter because they both (Na and Cl) have the same number of energy shells.

  14. Correction: Not the case with Na+ and Cl- ions!! Sal explains this in video number 38 in the Chemistry playlist called "Mini-Video on Ion Size".

  15. More electrons and more protons attracting each other making the element smaller. The ones on the left have a "looser grip" letting the orbital be bigger than the ones on the right.

  16. just saying the Na + should be smaller than the Cl – coz the Na has no outer shell electrons so is a shell bhind the chlorine

  17. Thats exactly not correct. The further right you go the more electronegative you get. The more electronegative the molecule is the closer it holds its electrons.

    Chlorine does have both more electrons and an additional shell, but its so electronegative that atomically it is smaller than sodium

  18. No he's right actually. When the sodium looses one elctron the ion that is formed has almost half the radius in fact.
    2 reasons:
    1. There is one less electron, so the nucleus attracts the remaining electrons more strongly.
    2. One energy level has been removed, so we have a noble gas arrangement of the sodium.

    And for Chlorine it's vice versa. So the sodium ion becomes bigger than the chlorine ion (as the ions are formed)

  19. to the right of the periodic table, the more electrons you have on the last layer, and the more electrons you have, the more they will be pulled towards the nucleus witch is positive, making the atom smaller

  20. Is that right? in the gas part? the more pressure the more solubility? I just Imagine in a coke.. that if it is sealed the soda or carbon dioxide is trap in the bottle.. so which mean the high pressure it get the less solubility it get.. PLEASE answer me am a little bit confuse.. thank you..

  21. Learning this stuff is so awesome and interesting, but try explaining it to someone who doesn't do chemistry and they will be like wat..?

  22. Acually, becuase the choride has an electron, it is bigger, because the net electon force of the nucleus on the electron cloud is less that the charge of the cloud. On the other hand,because positive attracts negative, the nucleus of the sodium should attract the electron cloud more, it should be smaller.

  23. In 4:25 he states that Sodium ion is bigger then Chlorine ion and that's incorrect. it's correct only for Sodium atoms and Chlorine atoms, because as an ion the element has the same electron configuration as the nearest noble gas, Sodium, atomic number 11 , nearest noble gas is Neon, atomic number 10. as we know the electron orbitals determine the atomic radius so now the Sodium ion has the radius of Neon but a bit smaller because is has more protons so it wants the electrons to be closer to the nucleus. and the same with Chlorine but that it will become to be like Argon because Chlorine is 17 and Argon is 18, but a bit bigger because it has less protons so compared to Argon, Chlorine ion wants the electrons near him less. so from all I've explained above , the Sodium is a bit bigger then Neon and Chlorine is a bit smaller then Argon so we can ignore it, so now the ratio between Sodium ion to Chlorine ion is the same ratio of Neon atom to Argon atom and argon is much bigger.

  24. Im really enjoying my first year in college . With these courses and class on How to shspe a Big Booty is very interesting. Dont ya think? Leave your final thoughhts on the ending of 2014 Happy New Year

  25. 1:20 story of your life? xD JK Lovin the videos

  26. So… When you blood is warmer your blood will carry more O2 and when your blood is colder your blood will carry less O2?

  27. I LOVE Khan Academy, but i wish there are videos about general chemistry that are not as basic. Nothing too advanced, but lets say something like Calc or Organic Chem videos

  28. omg I am gonna die tommorrow is my practical and i haven;t understood a single one, what a dumb mind of mine………………………………#####**

  29. Great vid, but one thing: Doesn't increased pressure mean increased temperature as well? So wouldn't increased pressure mean decreased solubility for gases?

  30. 4:29 isn't the diagram for sodium supposed to be smaller due to the fact that it is a cation, whereas the CL should be bigger since its an anion?

  31. Love the way you stretch imagination to the outmost (as feynmann said), but it's time to disagree (maybe wrongly). The radius of chlorine and sodium ions doesn't make sense: sodium gets electronic configuration (EC) of neon and an extra proton, so it should be smaller than neon; chlorine EC-argon-like but substracting a proton, it should be larger. Then, as argon is bigger than neon: Cl^->Ar>Ne>Na^+. Even so, there is lot of discussion.

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