Chem102EffectsofTempandPressureonSolubility

How temperature and pressure affect the solubility of a solute in a solution. Henry's Law.

>> This module is
about how pressure

and temperature affect
the solubility

of a solute in a solution.

So pressure.

First of all, if it's, if the
solute is a liquid or a solid,

then the pressure of gas

above the solution has
no appreciable effect

on the solubility of that liquid
or solid solute, so that's easy.

However, if the solute is
a gas, then the pressure

above the solution does matter.

In particular, what matters
is the partial pressure

of the gas that's dissolved in
the solution above the solution.

And it ends up that the
solubility of a gas that's

in a solution, dissolved
in a solution,

is directly proportional to the
partial pressure of that gas

above the solution in
certain situations.

This is called Henry's Law.

You should memorize
it -- C equals kP.

Here C is the concentration
of the dissolved gas.

The units of the concentration,
by the way, depend upon this.

This is called Henry's
Law constant, and you can,

you'll find Henry's Law
constant in different units.

We're going to use moles
per liter atmosphere here.

And P is the partial
pressure of the gas that's

in the solution above
the solution.

So k depends upon the,
that particular solution,

that particular combination
of solute, gas, and solvent.

Now, Henry's Law does not apply,

there is not this
linear relationship

when it's not a dilute solution

and when the gas
that's dissolved

in the solution either
reacts with a solvent

or dissociates in the solvent.

And if that's the case, then
Henry's Law's out the window.

For example, hydrogen
chloride by itself is a gas,

and if you were to dissolve it
in water, then the hydrogen,

hydrochloric acid that forms
dissociates immediately

into hydrogen ion
and chloride ion.

Hydrochloric acid is
a strong electrolyte,

which means that it completely
dissociates into the ions,

and that dissociation will
affect the concentration

of hydrogen chloride gases
dissolved in the solution.

So Henry's Law does
not work for that.

But there's a lot of instances
when we can use Henry's Law.

For example, let's suppose we
have a bottle of beer, right.

And when the, when it was
bottled, they pressurized

that bottle with carbon
dioxide gas to a pressure

of 4.8 atmospheres at
22 degrees Celsius.

We can use Henry's Law here
to calculate the concentration

of carbon dioxide
that's dissolved

in that beer before it's opened
and then after it's opened.

And now, other things we have
to know are, first of all,

what Henry's Law constant is
for carbon dioxide and water,

and it's this number right here.

And we also have to know
what the partial pressure

of carbon dioxide is in
just the normal atmosphere.

Now, let's assume it's
about 1 atmosphere.

The atmospheric pressure's
about 1 atmosphere.

And then, it ends up, because
there's only a small fraction

of the atmosphere
that's actually composed

of carbon dioxide, it ends
up the partial pressure

of carbon dioxide's about
this, .00040 atmospheres, okay.

Knowing that, all we're going
to do is plug into Henry's Law,

find the concentration
before we open the bottle

and after we open the bottle.

So before we open the bottle,
we take Henry's Law constant --

this is the k -- times
the partial pressure

of the carbon dioxide above
it, which is 4.8 atmospheres,

multiply them together, and
we get .148 moles per liter.

Two sig figs to get
.15 moles per liter.

After we open it, opens up.

Carbon dioxide gas escapes.

And all we have is essentially
the, well, is the atmosphere

with this many atmospheres
is the partial pressure

of carbon dioxide above it.

Plugging into Henry's Law again,

this is k. This is
the partial pressure

of carbon dioxide
in the atmosphere.

We get this number,
124 time, 1.24 times 10

to the minus 5th molar.

Two sig figs, 1.2 times
10 to the minus 5th molar.

And you can see here
that it's about 4 orders

of magnitude roughly difference
between before we opened it

and after we opened it.

So where does all that gas go?

It's, you know, some of the
bubbles that you see coming

out the, that's where
it's going.

That's why come out, okay.

Now, temperature.

What effect does temperature
have on the solubility?

Well, if it's a solid, then
increasing the temperature

for most solids increases their
solubility, but not all of them.

However, if it's a gas that's a
solute, then, across the board,

if you increase the
temperature of that solution,

you decrease the
solubility of the gas.

So these graphs here just give
the solubility of some solids

as a function of temperature.

And most them certainly
increased --

some more than others.

Potassium nitrate has a
real sharp curve here.

Notice potassium chloride, it's
pretty shallow curve right here.

But some of them, the
solubility decreases

as you increase the
temperature --

sodium sulfate and
cerium sulfate.

We talked a little
bit about that before.

These are gases, okay.

Solubility times 10 to the
minus 3 moles per liter.

And as you increase
the temperature, they,

their solubilities all decrease.

And that's that.