Molarity vs. molality | Lab values and concentrations | Health & Medicine | Khan Academy

Let’s talk about the
difference between two words molarity and molality. So the first one has obviously,
a little r in it, and the other has an l in it. And in fact, sometimes
when people say it quickly, it’s hard to even hear
which one they just said, so just make sure
you listen carefully because there is a slight
difference, actually, two little differences
that we’ll talk about. So molarity, let’s
start there, is really talking about moles of
something, some particle, over 1 liter of solution. And I haven’t been writing
out liter of solution, but that is what
is meant, so one liter of solution,
that is molarity. Now, molality is
slightly different, so let me do it in
a different color. Molality is actually
the moles, again– so that part is the same– over
1 kilogram– so this is now actually looking at mass– 1
kilogram of solvent, so not solution, but solvent. And so let me make
that very clear. So this is 1 liter of solution. This is 1 kilogram of solvent. So both are looking
at the same numerator, but the denominator
is different. So let me actually
draw out an example of what this might look like. Let’s say we have our
solution down here, and let’s say it’s
mostly water, so I’m going to fill this
in with water, and let’s fill it in more
quickly using a little air brush. So let’s say this is
our solution of water. I’m going to make
it nice and even. You can see the line
that it goes up to. So this is carefully
measured out, and I’m going to put
1 mole of something, and we can decide whatever
that something is going to be. And in this case,
I don’t know, let’s say we decide to put
some urea in there, little molecules of urea. And we know the
urea is something our body uses to get rid
of nitrogen, so something that we throw into
urine and actually even sounds like urine. So I’m going to put some
molecules of urea in here, and we’re going to make
our urea, I don’t know, let’s say kind of a pink color. This is our urea, and
this is 1 mole of it. So I’m only going to draw a
few of the little molecules, but you know that
whenever I say there’s a mole that must mean
that there’s 6.02 times 10 to the 23rd of these molecules
in here, so lots and lots of molecules. So 1 mole, this refers
to 6.02 times 10 to the 23rd, so lots of
molecules hanging out in our solution now. And so let me actually then
now cut and paste this. We’re going to cut and paste
this over to the other side, so that’s over here, and we’ll
move this underneath here. So you know this is
exactly the same, right? So this is just the
exact same solution, and I still have my
little pink urea. And now on this side, let’s
say I measure this out and check the level here. Let’s say this level right
here is exactly 1 liter. Well then, I would say this
solution has 1 mole of urea in 1 liter of solution, so we
have 1 molar solution of urea. So that’s our molarity. Right? So so far so good. And these are little
ureas just to make sure that we’re clear about that. So this is our
molarity, but what about the other
side, our molality? So for that, I actually
need to use a little eraser. So imagine now that I
actually removed all the urea because I don’t
want the solution, I just want the solvent. I just want the
part that is water. I don’t care about
the molecules. I just want to first get
1 kilogram of the solvent. So to do that, I’ve got
to get rid of all my urea. So I take out all the
little molecules of urea, and immediately
you can’t imagine that the water is
actually going to allow those little holes
to be like that. They’re going to fill in
those holes immediately. So right away, those holes are
going to get filled in, right? So let’s fill them
in with water, so the water rushes into
those holes and fills them in. But in doing so, in filling in
these little holes, of course, the level falls. Right? You actually have a
little bit less water. So you actually drop the
level of water a little bit, so let me erase some water
up here because the water level falls just a
little bit to fill in all those holes of
solvent that I took away or all the holes of
urea that I pulled out. So now my level is fallen, and
so if I was to measure this, let’s say this is
less than 1 liter. Let’s say it’s 0.99 liters. It’s going to be
very close, but it’s going be slightly less, right? so let’s say about 0.99 liters. So it’s a little bit
less than a liter. And remember, 1 liter
equals 1 kilogram. So for water and
most temperatures, 1 liter for a water, 1
liter equals 1 kilogram. So I guess I have
to ask the question, does this equal 1 kilogram? Well, the answer is no, right? It actually equals
about 0.99 kilograms. It’s actually slightly
less than a liter. So that’s going to weigh
less than 1 kilogram, just 0.99 kilograms. So now I have really
1 mole of urea. Thinking back to how
much I dumped in, I had put in 1 mole of urea. I’m just going to
let it hover here because this is
where it was right before it fell into my water. 1 mole of urea was going in to
only 0.99 kilograms of water. And so if that’s the
case, then my molality is actually going to
be slightly different. It’s going to be 1 mole of urea
over 0.99 kilograms of solvent, of water. And so 1 divided by a
number slightly less than 1 will be a little
bit more than 1, so my molality will
actually be maybe let’s say 1.01 or thereabout. It will be just
slightly upwards of 1, and that will be the molality. So they’re very similar, right? Like 1 molarity,
in this case was, was going to equal just a
little bit higher molality and that’s because we know
that the molecules of urea take up a little bit
of volume and that makes the overall volume of
the solvent a little bit less. So that’s the key difference,
and if you think about this really when you’re talking
about blood and things that are dissolved into
blood, most clinicians will jump back to
molarity because it’s just easier to work with
and you don’t actually have to figure out the
exact amount of solvent. You can just think
about the solution. So most of the clinicians
or doctors and nurses will think in terms of
molarity, but most of the time when you’re working
in a lab setting and you can be more precise,
people think about molality.

54 comments

  1. Molarity is usually slightly LOWER than molality when dealing with water based solutions or solvents (like blood).

  2. I think it would be useful and easier to understand the difference if you gave a situation where it makes more sense to use molarity and one where it makes sense to use molality.

  3. thanx alot.. before watching this video i had no idea about molality and molarity 🙂
    keep up the good work (Y)

  4. When first learning about molarity, I accidentally did the Molality side of my worksheet. Why is this concept such a troll???

  5. C'est génial ! En France on a pas ça, du coup pour réviser mes cours je regarde tes vidéos et je pense pas que je trouverais mieux 😀

  6. so is Molality ALWAYS greater than molarity in aqueous solutions, or are there factors such as density that will change that ? 

  7. It might be easier if molarity is described as adding the urea first then filling up to the 1L mark, and molality as filling the beaker up to 1L and then adding the urea. Either way, molarity in this case is more concentrated.

  8. Don't pay attention to the example, it's not clear (and I think one part is wrong). But, all you need to understand is that Molarity is moles over 1 liter of solution, and Molality is moles over 1 kilogram of solvent.

  9. This is a truly terrible example. Why on earth would you attempt to explain it this way? Remove the solute and then add it back? Bizarre. Offer something that is a clear contrast, not the exact same solution.

  10. The point he wanted to make is that osmolality is no affected by change in volume of solution or by temperature

  11. The important difference between Molarity and Molality is this :
    While finding Molarity you have to look at all the solution , while finding Molality you have to look at only the solvent !

  12. But if you have added solute in second beaker then how come it will decrease the volume of solvent in it?

  13. Using water as an example is bad. It wouldn't make sense if you use the same solvent that was used to define what a liter is. if you used a different solvent, say something that has 1 Kg of mass for every 2 liters in volume, it would become much clearer.

    Assume you have a 1 liter of a solvent, and you decide to measure it's mass, and you find that it's only 0.5 kg. Then you decide to add to that liter about 1 mole of urea. Assuming that 1 mole of urea, when added to that liter of solvent, would increase its volume by 0.1 liters, you would end up with a 1.1 liter of a solution that contains 1 mole of urea. If you were to measure it's molarity it would be 1/1.2 = ~0.9 mol/L.

    At the same time, if you wanted to measure the molality of that same thing, it would be fairly different. First, you would ignore the effect in mass of adding 1 mole of urea to that liter of solvent, since we're only concerned with the mass of the solvent itself. Then, you would just do the calculations straight forward. It would be 1 mole of urea in a 0.5 kg of solvent, that means 1/0.5 = 2 mol/kg.

    The reason the guy in the video used water as an example is because blood is a water-based solution, and in water-based solution, at room temperature, molarity and molality are usually close.

  14. My gen chem is a bit rusty. So would it be, for example to get the molarity of 1L of saltwater, you'd look for moles of NaCl existing in this liter of already combined saltwater? But for molality, could you say it would be the moles of NaCl that could be added to 1kg of plain water to quantitatively get 1L of saltwater?
    Yup, makes more sense in chemistry lab settings xD
    *After reading I may have gotten the two switched. Lol

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