溶解度 (英)

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溶解度 (英) : 鹽以及氣體溶質在液體溶劑中的溶解度

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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 right 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.