>> host: Welcome back
to "Chemistry Matters", Unit 7,
Solutions, Acids, and Bases.
By now you should have
completed the solubility lab
with salts.
As you saw near the end
of the last video
on our playlist,
our students were discovering
that not all substances
dissolve in water,
even when the students
stirred the mixture
and increased the temperature.
And if you performed
this lab correctly,
you should have gotten
similar results.
Chemists have a word that
describes this phenomenon:
insoluble.
Insoluble means that
a solid, liquid, or gas
will not dissolve in
a particular solvent.
Solubility depends on
the attractions between
the solute and the solvent.
For example,
when sodium chloride,
or common table salt,
is added to water,
the sodium and chloride ions
are more attracted to the water
than they are to each other,
but, as our students saw
in the solubility lab,
calcium and sulfate ions
are more attracted to each other
than they are to water.
Some other examples
of insoluble substances
are salt in gasoline
or copper in water.
If copper could dissolve
in water,
the Statue of Liberty
would've disappeared into
New York Harbor long ago!
Soluble, on the other hand,
means that a solid, liquid,
or gas can be dissolved
in a particular solvent.
Solubility is a term
that describes the maximum
amount of a substance
that can be dissolved in
a given quantity of solvent
at a given temperature
to produce a saturated solution.
Now we just used
a new term there:
saturated solution, so let's
look at what that means.
A saturated solution
is a solution in which
the maximum amount of solute
has been dissolved in
a given amount of a solvent
at a particular temperature.
For example, let's say you're
trying to make some sweet tea,
and you've already brewed
the tea by pouring
very hot water over
the tea bags to get
your un-sweet tea.
Now you want to add sugar.
So you start adding sugar
one spoonful at a time.
At first you'll notice
that most of the sugar
dissolves almost immediately.
That's because the solution
can hold more sugar
at that temperature.
But as you continue
to add sugar,
you'll eventually reach
a point where you can no longer
dissolve any more sugar into
the tea at that temperature,
regardless of how much you stir.
This is called
the saturation point,
which is where you can
no longer dissolve
any more solute in that
solution at that temperature.
With that in mind, let's
head back to our classroom
where our students are in
the process of investigating
how temperature can affect
the solubility of
three different salts.
>> professor: Okay, let's
look at your data.
Tell me what you found.
>> student: Well, we found
different results for each
of the salts.
>> student: Yeah,
the calcium sulfate didn't
dissolve at any temperature.
>> student: The potassium
nitrate dissolved completely
at 30 degrees Celsius,
but only some of it
dissolved at 5 degrees.
>> professor: How about
your team?
Did you get similar results
for the calcium sulfate
and potassium nitrate?
>> student: Yep, pretty much
the same results.
>> professor: What about
the sodium chloride?
How soluble was it?
>> student: The sodium chloride
completely dissolved even
at the lowest temperature.
>> professor: Very good.
And I like the terminology
you're using to describe
your data.
For soluble substances,
like potassium nitrate
and sodium chloride,
temperature does indeed
play a role.
As a matter of fact,
we say that solubility
is temperature dependent,
because in most cases,
the higher the temperature
of the solvent,
the greater the amount
of solute can be held.
Take a look at these graphs.
They're called
solubility curves.
And they show the temperature
dependence on the solubility
of both the potassium nitrate
and the sodium chloride.
Team 1, as you interpret
this graph, can you explain
your observations with respect
to the sodium chloride?
>> student: Yeah I think I can.
According to that graph,
it looks like we could have
dissolved almost 40 grams
of sodium chloride
in our beaker that was
at 5 degrees Celsius.
I guess that's why all
25 grams completely dissolved
even at that low temperature.
>> student: It looks like
the solubility of
sodium chloride only goes up
a little bit though as
the temperature is increased.
>> professor: Very good,
both of you!
Team 2, what about
the potassium nitrate
you investigated?
Does your data support
what's displayed here?
>> student: I believe so.
We noticed when we tried
to dissolve the 25 grams
of potassium nitrate
at 5 degrees Celsius,
it looked like only
a little bit of it dissolved.
>> student: Yeah, and looking
at that graph I can see why.
If we have 100 grams of water
at 5 degrees Celsius,
it looks like we can only
dissolve around 18 grams
of potassium nitrate.
No wonder some of it
wouldn't dissolve!
>> student: Right, but look
at how much the temperature
affects the solubility
of potassium nitrate!
I can see why all of it
dissolved when the water
was 30 degrees Celsius.
It looks like we could have
dissolved around 45 grams
of potassium nitrate
at that temperature!
>> professor: Exactly right.
You've all done a very nice job
of interpreting the data
on this graph, and now you can
apply this type of thinking
to other graphs too.
>> host: It looks like our
students are getting a handle
on how to interpret
solubility curves.
But let's explore solubility
curves a little more closely.
A solubility curve shows
the dependence of solubility
on temperature.
On this graph from
our classroom,
which depicts the solubility
of sodium chloride
and potassium nitrate,
the y-axis shows grams of solute
per 100 grams of water.
And the X-axis
displays temperature
in degrees Celsius.
As you know, sodium chloride
and potassium nitrate
are solids, but other solids
dissolve in water too.
Like sodium nitrate, NaNO3,
potassium chloride, KCl,
ammonium chloride, NH4Cl,
and potassium iodide, KI.
One thing you'll notice
is that for solids,
which are represented
on the graph as solid lines,
as temperature increases,
greater and greater amounts
of solute can be dissolved
in that 100 grams of water.
But gases, such as ammonia, NH3,
will also dissolve in water.
So will sulfur dioxide, SO2,
and hydrogen chloride, HCl.
For gases, which are
indicated with dashed lines,
the greater the temperature,
the less soluble they are.
This pattern, which is
a great example of our patterns
cross cutting concept,
is the opposite of what
we see with solids.
And that's why most people
like their soft drinks
to be cold.
They're more bubbly that way!
You'll notice that
calcium sulfate is not
on the graph.
You might be wondering
why that is.
Let's go back to our classroom
and find out what
our students think.
>> student: I have a question
about the calcium sulfate.
I know it didn't dissolve,
but does that mean it doesn't
have a solubility curve?
>> professor: That's
an outstanding question.
We can generate a solubility
curve for calcium sulfate,
but because it's
essentially insoluble,
we'd need to drastically change
our units on our axes
to be able to visualize it.
Only about .21 grams
of calcium sulfate
will dissolve in 100 grams
of water when its temperature
is about 20 degrees Celsius.
>> student: Whoa, and we
tried to dissolve 25 grams.
No wonder it just floated
around at the top
of the beaker.
>> professor: Yes, it doesn't
take very much calcium sulfate
to dissolve at all
for the solution
to become saturated.
Occasionally if
a saturated solution
is allowed to slowly cool,
you can get what's called
a supersaturated solution.
This is a solution
that's holding more
dissolved solute than
it would normally hold
at that temperature.
The reason I've asked you
to wear your safety glasses
is because, in these flasks,
I have aqueous solutions
of sodium acetate trihydrate
that were heated
above 70 degrees Celsius
and saturated
at that temperature.
I then allowed them
to cool overnight.
When we arrived here today,
they're considered
a supersaturated solution
and is very unstable.
In fact, all it would take
for all of that solute
to come out of the solution
and to crystallize is a single
sodium acetate crystal.
So what I've done is placed
a few sodium acetate crystals
on the watch glass
in front of you,
and I'm going to ask you
to slowly pour
your supersaturated solutions
onto the watch glass.
Go ahead.
>> student: Whoa!
>> student: Hey, that's crazy.
>> student: That's awesome!
>> student: Man, this
so awesome!
>> student: It's snow!
>> student: It's-- yeah!
>> student: Can it get taller?
>> student: All the ones
are sitting on top.
It's going to completely fall.
It looks like it's gonna fall!
>> student: It's like snow.
>> students: This is so cool.
>> student: This is so awesome.
>> student: That gained height
really quickly.
>> student: I think ours
is taller.
>> student: There we go.
>> host: Wow, that was
a pretty clever demonstration
of a solution that was
holding too much solute
at a particular temperature.
The process we just observed
is called crystallization.
Lots of supersaturated solutions
will crystallize
when their conditions change,
but they won't always do it
in such a dramatic fashion.
Not only can solutes
be thought of as making
a solution saturated
or even supersaturated,
they can also be classified
as nonelectrolytes,
weak electrolytes,
or strong electrolytes.
I'm sure you've seen commercials
for sports drinks touting
their ability to replenish
electrolytes lost
while sweating, but what
exactly are electrolytes?
Well, in the simplest sense,
electrolytes are ions,
or charged particles.
Some solutes, when dissolved,
actually produce
a lot of ions.
Solutes like ionic compounds
or even strong acids are
known as strong electrolytes
for their ability to produce
these ions in solutions.
Solutes, such as weak acids,
produce very few ions
in solutions and are therefore
known as weak electrolytes.
And covalently bonded
substances, such as table sugar,
don't produce any ions at all
when they dissolve.
They're known as
nonelectrolytes.
Why the term
electrolyte though?
Let's head back to
the classroom for another
"shocking" demonstration.
>> professor: What I have here
is a light bulb
with one electrode
connected to this wire,
and the other electrode
connected to the other.
For safety reasons,
I'm only going to plug it in
once this is in the solution.
When the circuit is completed
and electrons can flow,
electricity will pass from
one electrode to the other
and the light bulb
will light up.
For the electricity,
the moving electrons,
to make it from one electrode
to the other, there must
be charged particles
in the solution.
It's kind of like a frog
trying to make it from one end
of the pond to another
jumping across logs
and lily pads.
The electrons must move
from ion to ion
throughout the solution.
There have to be ions
in the solution
for the current to flow.
The more ions,
the more current can flow.
The solvent, here,
is distilled water,
which has no ions in it.
See? No light.
Now we're gonna test
some solutions.
First I'm gonna dissolve
some sugar...
into the distilled water.
And nothing happens,
no light because
there are no ions.
We can characterize sugar
as a nonelectrolyte.
Now I'm gonna place
some sodium chloride,
or common table salt,
and then stir it up
to dissolve it.
Voila!
>> student: Oh!
So ions must have formed.
That makes sense
because sodium chloride
is an ionic compound.
>> student: Yeah, and I can see
bubbling in there too.
>> professor: Yes, what you're
seeing is electrolysis,
or the decomposition of water,
just like we saw in
Unit 5 when we studied
chemical reactions.
The fact that the electrons
can make it from one electrode
to the other indicates
that there are ions
in the salt water solution
and that electricity
is causing the water
to decompose into
hydrogen and oxygen gases.
That's why sodium chloride,
along with most ionic compounds,
is known as a strong
electrolyte.
When dissolved,
strong electrolytes allow
a current to be passed
through them.
Lastly, I have some
acetic acid here,
which is known as
a weak electrolyte.
You'll notice that there's
only going to be a faint
glow of the light
because there's a relatively
small amount of ions
in this solution.
>> student: You said
'most' ionic compounds are
strong electrolytes, but what
about the calcium sulfate
we looked at earlier?
That wouldn't be
a strong electrolyte, right?
I mean, it didn't really
dissolve at all.
>> professor: That's a great
question and observation,
and you're exactly right.
Calcium sulfate won't
fit our definition
of a strong electrolyte
since it didn't produce
a large number of ions
in the solution.
>> student: So how come
certain solutes dissolve,
while others don't?
>> professor: Well, there are
several factors involved,
but one that we've studied
back in unit 4 has to do
with the shape and polarity
of the solute and solvent
and their attraction
to each other.
Remember, polar substances
have areas of slightly positive
and slightly negative charges
where the electrons are
unequally shared, and there
are attractive forces
between these slightly positive
and negative charged areas
and the particles with
opposite charges.
When solutes and solvents
have similar
intermolecular force,
the forces between particles,
the solute is much more likely
to be soluble in the solvent.
Therefore, a polar,
or ionic solute, is usually
soluble in a polar solvent
and a nonpolar solute
would be soluble
in a nonpolar solvent.
But you asked about
calcium sulfate, which is
ionic, and is not soluble
in water, which is polar.
Remember that soluble
ionic compounds separate
into ions when they dissolve.
They do that because
the ions are more attracted
to the solvent than they
are to each other.
In the case of calcium sulfate
in water, most of the
calcium ions and sulfate ions
are more attracted
to each other than they are
to the water.
That's why they don't dissolve.
>> host: Our teacher was just
making a very important point,
and it's related to a saying
you might have heard:
"like dissolves like."
Have you ever gotten oil
or something greasy
on your hands
and you couldn't get it off
with just water?
Water is polar, while grease
is nonpolar.
Soap, or detergent, like I have
in this urban meyer flask,
is a great example
of like dissolving like.
Soap is both polar, like water,
and nonpolar, like grease.
The soap molecule has
a polar "head"
with a long, nonpolar "tail."
The nonpolar part of
the soap molecule sticks
to the grease and dirt,
while the polar head
of the molecule sticks
to the water.
So when you wash your hands
with soap, the dirt and grease
are removed from your skin
and go down the sink
with the water
when you rinse your hands.
Now that we have
a fundamental understanding
of what solutions are
and how they form,
we're ready to more closely
explore the composition
of solutions.
So next we're going to learn
how to express the concentration
of solutions, which is
another way of describing
how much solute is present
in the solution.
And we'll do that at
the top of the next video
on the playlist.
I'll see you there!