>> host: Welcome back to
"Chemistry Matters",
and our unit on solutions,
acids, and bases.
Chemists often create
and work with solutions.
Some solutions have
a lot of solute in them.
A qualitative way of
describing such a solution
is to say it's concentrated,
like this concentrated
orange juice.
A qualitative way of
describing a solution
that does not have
a lot of solute in it
is to say that it's dilute,
like the orange juice
that I'm making.
However, it's not practical
or safe to always use these
qualitative descriptions
of solutions.
We need measured
or quantitative descriptions
of our solutions.
There are many ways
that we can describe
a solution's concentration,
but in this unit,
we're going to focus on three.
Mass percent, molarity,
and molality.
So let's head back
to our classroom,
where our students are exploring
different ways to express
solution concentration.
>> professor: Let's look at
three ways of describing
a solution's concentration,
starting with mass percent.
Mass percent is a simple way
of expressing how concentrated
a solution is.
It's found by taking
the mass of the solute
in the solution
and dividing it by the total
mass of the solution,
and then multiplying it by 100.
One thing to keep in mind
when calculating mass percent,
is that it's the mass
of the entire solution,
not just the solvent
in the denominator.
So, to determine
the denominator,
you have to add
the mass of the solute
to the mass of the solvent
before you can find
the mass of the solution.
It'll help you understand
if we practice.
What would be the mass percent
of an aqueous solution
of salt water
if you placed 1 gram of salt
into 99 grams of water?
Use the formula to
figure out the answer.
>> student: It would be 1%?
>> professor: That's correct.
How did you get that?
>> student: Well, since the
mass of the solute is 1 gram,
that would go into
the numerator.
So 100 grams would have to
go into the denominator,
since that's the total mass
of the solution.
Then, divide, and you get .01.
So I just multiplied .01
times 100 to convert it
to a percentage,
and that is 1%.
>> professor: Very good.
And to create a solution
with a specific mass percent,
the only equipment
you really need
is a balance of some sort,
and, of course,
the solute and the solvent
that you plan to use.
Can anyone think of an example
of something like this?
>> student: My little brother
had poison ivy
a few weeks ago, and he used
hydrocortisone cream on it.
The cream had 1% on the label.
Is that what you mean?
>> professor: That's exactly
what I mean.
So the active ingredient,
hydrocortisone, was only
1% percent of the cream.
Good example.
Now, the next way to express
how concentrated a solution is,
and, by far, the gold standard
for chemists around the world,
is known as molarity.
Molarity is a ratio of moles
of solute to the volume
of the solution in liters.
It's calculated by dividing
the moles of solute
in the numerator
by the volume of the solution
in liters in the denominator.
Molarity is abbreviated
with a capital 'M'.
A key point to notice here
is the use of liters
in the denominator.
Since we want to measure the
volume of the entire solution,
special glassware is used
to create solutions
of known molarities.
A volumetric flask,
like the one I'm holding,
will allow you to put
the number of moles of the
solute in the flask first,
then add distilled water
up to the line.
So now what I'd like
for you guys to do
is to create your own
solutions of a known molarity.
Team one should create
500 milliliters of a 1 molar
copper(II) nitrate solution.
Team two will make
1 liter of a 1 molar
copper(II) nitrate solution.
Come up with
your procedures first,
and let's discuss them.
Once I approve your procedures,
you can make your solutions.
Okay, everyone,
who's ready to present the plan
that they're gonna use
to create their solution?
>> student: We'll give it a try.
We have to make 500 milliliters
of a 1 molar
copper(II) nitrate solution,
which means we actually
need to have .5 moles
of copper(II) nitrate
dissolved in a total of
500 milliliters of solution,
right?
>> student: Since molarity
is the ratio of moles
to 1 liter, we figured that
to make a half liter
of 1 molar solution,
we would need half of a mole
of copper(II) nitrate.
>> professor: That's right.
>> student: Looking at
the molar masses
from the periodic table,
copper(II) nitrate has copper
with a molar mass
of 63.55 grams.
We multiply that by 1
because there's only
1 copper in the formula.
Nitrogen has a molar mass
of 14.01 grams, and because
copper(II) nitrate has
2 nitrogens in the formula,
we multiply that by 2.
And oxygen's molar mass
is 16.00 grams.
And we multiply that by 6,
because copper(II) nitrate
contains 6 oxygens
in the formula.
After multiplying all of those
and adding them up,
1 mole of copper(II) nitrate
would have a mass
of 187.57 grams.
So half a mole
would be 93.79 grams.
>> student: So now,
we just add that to the flask
and add about 500 milliliters
of water in there?
>> student: Wait.
Actually, I think we just
put the copper(II) nitrate in,
and then add water
until it comes up to
the 500 milliliter mark.
>> student: Oh, right.
The denominator has liters
of solution, not just solvent.
>> student: Yeah.
The copper(II) nitrate
is going to take up space,
which is why we have to
use this flask, I guess.
>> professor: Great catch,
and you're exactly right.
That's why we have to
create these solutions with
specific types of glassware.
We have to be able to account
for the volume or the space
that the solute is
going to take up.
Your team is good to go.
Okay, team two, are you ready
to tell me with your plan?
>> student: I'm pretty sure
we are.
We need to make 1 molar
solution of copper(II) nitrate,
which is 1 mole of
copper(II) nitrate dissolved.
>> student: Yeah,
and according to the masses
on the periodic table,
the molar mass of
copper(II) nitrate
is 187.57 grams.
>> student: Exactly.
So we just need to mass
that out and place it into
the volumetric flask, right?
>> student: And then
we'll add water up to
the 1 liter mark,
to make sure it's 1 mole
per 1 liter of solution.
>> professor: That sounds
like a great plan.
Everyone, put on
your safety goggles,
and you can all go ahead
and start creating
your solutions.
[indistinct chatter]
>> student: Careful.
Keep it going.
[indistinct chatter]
>> student: Yeah,
I don't think we will.
>> student: Careful, careful,
careful, careful.
>> host: Our students
already described their
very detailed plans for you,
making your job
a little bit easier.
But you've still got
the activity itself to do.
By the way,
depending on how you
round off your atomic mass,
your exact numbers
could be a little different
from those of our students.
We'll take a break here.
When you're done with
the molarity lab,
you can compare your results
with the results
of our students.
We'll take a look at those
at the top of the next video
on the playlist.
I'll see you there.