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Heat of vaporization of water and ethanol

Heat of vaporization of water and ethanol.

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Video transcript

- [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.