Heat of vaporization of water and ethanol | Biology | Khan Academy

– [Voiceover] So we have two
different substances here and just for the sake of an argument, let’s assume that they
are in their liquid state. Well you probably already recognize this substance right here, each molecule has one oxygen atom and two hydrogen atoms, this is
water and we have drawn all neat hydrogen bonds right over there. Now this substance, at least right now, might be a little less familiar to you, you might recognize you have an O-H group, and then you have a carbon chain, this tells you that this is an alcohol, and what type of alcohol? Well you have two carbons here, so this is ethyl alcohol
or known as ethanol. So this right over here,
let me write that down. This is ethanol, which is
the primary constituent in the alcohol that people drink,
it’s also an additive into car fuel, but what I
wanna think about here, is if we assume that both of these are in their liquid state and let’s say they’re hanging out in a cup and we’re just at sea level so it’s just a standard
pressure conditions. Which one is going to
be easier to vaporize or which one is going to have more of it’s molecules turning into vapor, or I guess you could say
turning into vapor more easily? Well you immediately see that
they both have hydrogen bonds, you have this hydrogen bond between the partially negative end and
the partial positive ends, hydrogen bond between
the partial negative end and the partial positive ends. The other thing that you notice is that, I guess you could think of
it on a per molecule basis, on average you have fewer hydrogen bonds on the ethanol than you have on the water. Ethanol– Oxygen is more electronegative, we already know it’s more
electronegative than hydrogen, it’s also more
electronegative than carbon, but it’s a lot more
electronegative than hydrogen. So you have this imbalance here and then on top of that, this carbon, you have a lot more atoms here in which to distribute a partial charge. There could be a very weak partial charge distributed here amongst the carbons but you have a stronger
partial charge on the hydrogen but it’s not gonna be
strong as what you have here because, once again, you
have a larger molecule to distribute especially
around this carbon to help dissipate charging. So you’re gonna have
weaker partial charges here and they’re occurring in fewer places so you have less hydrogen
bonding on the ethanol than you have on the water. Let me write that, you
have less hydrogen bonding. As we’ve already talked about, in the liquid state and frankly,
in the solid state as well, the hydrogen bonding is what is keeping these things together,
that’s what’s keeping the water together, flowing
next to each other. This is what’s keeping
the ethanol together. So if you have less hydrogen–
Let me write this down, less hydrogen bonding, it
actually has more hydrogen atoms per molecule, but if you
have less hydrogen bonding, it’s gonna take less energy
to break these things free. Before I even talk about
breaking things free and these molecules turning into vapors
entering their gas state, let’s just think about how that happens. When we talk about the
temperature of a system, we’re really just talking about
the average kinetic energy. Each molecule, remember
they’re all bouncing around in all different ways, this
one might have, for example, a much higher kinetic
energy than this one. They’re all moving in
different directions, this one might have a little bit higher, and maybe this one all of a sudden has a really high kinetic energy
because it’s just been knocked in just the exact right ways and it’s enough to overcome
both these hydrogen bonds over here and the pressure
from the air above it. Remember this isn’t happening
in a vacuum, you have air up here, air molecules,
I’ll just draw the generic, you have different types of things, nitrogen, carbon dioxide,
etcetera etcetera. But if I just draw generic air molecules, there’s also some pressure from
these things bouncing around but this one might have enough,
this particular molecule might have enough kinetic
energy to overcome the hydrogen bonds and overcome the pressure
from the molecules above it to essentially vaporize,
to turn into its gas state. The same thing might be true over here, maybe this is the molecule that has the super high kinetic energy
to be able to break free. In that case, it is going to
turn into its gaseous state. The hydrogen bonds are gonna break apart, and it’s gonna be so far from
any of its sibling molecules, I guess you could say, from
the other ethanol molecules that it won’t be able to
form new hydrogen bonds. Same thing with this
one, once it vaporizes, it’s out in gaseous state, it’s
much further from any other water molecules, it’s not going to be able to form those hydrogen bonds with them. Because there’s more
hydrogen bonds here to break, than here, you can imagine
it would take, on average, more heat to vaporize this thing
than to vaporize this thing and that is indeed the case. The term for how much heat do you need to vaporize a certain mass of a
substance, you can imagine, is called the heat of vaporization,
let me write that down, heat of vaporization and you can imagine, it is higher for water
than it is for ethanol and I will give you the numbers here, at least ones that I’ve
been able to look up. I found slightly different numbers, depending on which resource
I looked at but what I found for water, the heat of vaporization
is 2260 joules per gram or instead of using joules,
remember joules is a unit of energy it could be a unit of
heat, instead of joules if you wanna think of it in terms of calories, that’s equivalent to 541
calories per gram while the heat of vaporization for
ethanol is a good bit lower. The heat of vaporization for
ethanol–let me make this clear this right over here is
water, that’s for water. The same thing for ethanol. The heat of vaporization for ethanol is, based on what I looked
up, is 841 joules per gram or if we wanna write them as
calories, 201 calories per gram which means it would require, roughly, 201 calories to evaporate,
to fully vaporize a gram of ethanol at standard temperature, keeping the temperature constant. We could talk more about
that in other videos, but the big thing that
we’re talking about here is, look, it requires less
energy to vaporize this thing and you can run the experiment,
take a glass of water, equivalent glasses, fill them
up the same amount of time, a glass of water and a glass of ethanol and then see how long it takes. You can put a heat lamp on top of them or you could just put them outside where they’re experiencing the same atmospheric conditions,
the same sun’s rays and see what’s the difference–
how much more energy, how much more time does it take for the water to evaporate than the ethanol. There’s a similar idea here
which is boiling point. We’ve all boiled things, boiling point is the point at which the vapor
pressure from the substance has become equal to and starts
to overcome the pressure from just a regular atmospheric pressure. And so you can imagine that water has a higher temperature
at which it starts to boil than ethanol and
that is indeed the case. Water’s boiling point is
exactly 100° Celsius, in fact, water’s boiling point was
an important data point for even establishing the Celsius
scale, so by definition, it’s 100° Celsius, while
ethanol’s boiling point is approximately 78° Celsius. So it boils at a much lower temperature an that’s because there’s just fewer hydrogen bonds to actually break.