The Science of Everything Podcast - Episode 27: Intermolecular Bonds and Phase Transitions

Episode Date: December 28, 2011

An explanation of the forces and mechanisms of intermolecular bonding, including dispersion forces, dipole-dipole bonding and hydrogen bonds, with an application of these mechanisms in analysing surface tension. This is followed by a discussion of how phase transitions occur, focusing on evaporation, condensation, boiling and melting. Recommended pre-listening is Episode 15: Chemical Bonding.

You're listening to The Science of Everything podcast, episode 27, Inter-Molecular Bonds and Phase Transitions. I'm your host, James Fodor. So this episode builds on from what we talked about in episode 15 on chemical bonding. In that episode, among other things, I talked about intramolecular bonding, that is the chemical bonds that combine atoms together to form molecules. In this episode, I'll be looking at intermolecular bonds, which are the chemical bonds of a somewhat different sort,

that bind molecules together or small molecules to form bigger molecules or that bind molecules together to form larger substances like a liquid or a solid for example. So some of those include dispersion forces, dipole-dipole bonding and hydrogen bonding. And I'll also apply this knowledge of intermolecular forces to explaining how surface tension works. If you've ever heard of that phenomenon, it's a very interesting one, so we'll take a look at that. This knowledge of intermolecular bonding will then also be applied to understand phase transitions, that is, the process by which a substance goes from being a liquid to a gas, or a gas to a liquid, or a solid to a liquid, and so on. So I'll focus on evaporation, condensation, boiling, and melting, and explain the process

by which those occur with reference to knowledge of intermolecular bonding. Okay, so let's get into it, and I'll start with intermolecular bonding. So, first of all, we have to know what intermolecular forces are, that is, the forces that act between, molecules. Intermolecular forces are forces of attraction or repulsion which act between neighbouring particles. They could be ions, ions, atoms or molecules, but generally we're talking about molecules in this section. Now intermolecular forces are relatively weak compared to intramolecular forces. That is the forces acting within a molecule to keep it together. That's why substances don't sort of fall apart and the hydrogen atom in one water molecule gets

detached and attracted to another water molecule. Actually, that's a bad example because that does happen, but that's why substances stay together and don't just, unless it's an actual chemical reaction, it's because generally the forces keeping a molecule together are much stronger than the forces that keep the different molecules within that substance together. Like the intramolecular forces, intermolecular forces are really all based upon electromagnetism, that is positive and negative charges. I haven't done an episode specifically covering that topic yet, but I will do so in the future, but just for the moment, remember that like charges repel and unlike charges attract. So that's the basic idea we're applying here.

Another concept that we now need to introduce is the state or phase of matter. We have talked about this before. Just to go through them, again, we have solid, liquid, and gas are the three main stages, phases. There are some other ones too. But there are the three main ones we talking about. Now, the state or phase of matter that are given substances in is sort of can be considered to be a competition, if you like, between the thermal motion of the particles within the substance, which tends to, well, move

them around, vibrate them around, and also, and therefore sort of push them apart from each other. That's one aspect. And the second sort of side to this contest are the strength of the intermolecular bonds that act between the particles to pull them together. So if you like, the thermal energy, that's also the energy that gives an object, its temperature, as we talked about in a previous episode, is sort of pushing particles apart, but the intermolecular forces are pulling them together. The more sort of, the relatively stronger, the intermolecular forces are, or the less dispersed the particles will be, and therefore it'll be more likely that the substance will be either liquid or, if it's even stronger, a solid. Conversely, the stronger the thermal energy

forces are, well, the greater the thermal energy is, the greater the temperature of the substances, the more rapidly the particles will be moving around, the more they'll be overcoming the intermolecular bonding forces, and therefore, the more you'll tend to, the substance will tend to move from being a solid to a liquid to a gas. So gas is the most disordered state where sort of thermal energy has one. Solid is the most ordered state where intermolecular bonding forces of one and liquid's sort of in between. So those are the basic concepts you have to keep in mind when we're dealing with intermolecular forces. There's the sort of dispersing force of thermal energy and the combining or bringing together force of the intermolecular bonds. But I still haven't

explained how this drawing together force of the intermolecular bonds work. So that That's what I'm going to do now. There are basically, at least for our purposes, three different types of intermolecular bonds, intermolecular bonding forces. As I said, they all rely fundamentally on electromagnetic attraction, repulsion, the electromagnetic force, but they act in slightly different ways. So to start off with, I'll talk about dispersion forces, which are the weakest of the three types. Dispersion forces are also called London forces, but that's somewhat weird, so I prefer dispersion forces. Essentially, these arise because of the chance accumulations of electrons in a particular region of the electron cloud, which then

then induces a temporary dipole in that molecule, which in turn induces a temporary dipole in the neighbouring molecule. So let's explain what that means. If you remember the episode when we talked about polarity in molecules and electronegativity, I had talked about how in some molecules one atom attracts, has a stronger pulling power for electrons, a strong high electronegativity, than does the other atom, and therefore it tends to, in a sense, hog the electrons, or the electrons spend relatively more time near the high-electric negativity atom than they do close to the low-electron negativity atom. And therefore, within the molecule, you have a dipole established, that is a partial charge, where it's not like an ionic bond where an electron

is completely transferred from one atom to the other, but an electron has sort of been partially transferred because it's moved from, it spends relatively more time with orbiting around one atom than the other. And so there's a partial charge on two parts of the molecule. That's what a dipole is. Well, you can get dipoles, or they can form as a result of a difference in electro-ignignities within the molecule, like, for example, in water. The

oxygen atom has a higher electronegativity than the hydrogens. But you can also form a dipole, or at least a temporary dipole, as a result of random chance, or so-called dispersion forces. Basically, if you recall our episode on quantum mechanics, principles of quantum mechanics, the electrons don't literally orbit around

the nucleus of the atom. They sort of exist in a probability cloud and they have a certain probability of existing in any given place at any given time within around the atom. By the way, when we form molecules, we can also talk about the, it gets kind of complicated, but we can talk about the atoms now being in sort of shells or orbitals around the entire molecule, as opposed to just specific atoms in the molecule. That's a field called quantum chemistry that deals with that, and as you can imagine, that gets a bit complicated. So basically imagine that the, now the molecules a whole, like a water molecule or something like that has electron shells as opposed to just the single atom.

It would be easier to imagine what I'm saying if you picture it that way. So just by chance, sometimes more of the electrons will be, say, on the left-hand side of this molecule than on the right-hand side. And so just by chance at that instant, or for that short period of time, you will have a fractional charge, say, a negative charge on the left side of the molecule and a fractional positive charge on the right side because you don't have very many electrons. there. And so that is called a temporary dipole, dipole, dipole

because once again there's two charges, one on each side of the molecule, but it's temporary because it's just sort of come out by a random chance and the next nanosecond it'll be gone or changed around. However, the dipole doesn't need to last for very long or be very strong to have an effect. And if you get this temporary dipole coming to existence in one molecule, what it will do is induce a dipole in a neighbouring molecule. So suppose, for example,

we have that negative left-hand side of the water molecule, we'll say, which has relatively more electrons in it. That negative dipole will tend to repel the electrons in the neighbouring water molecule, thereby inducing a small positive dipole in its neighbouring water molecule, and correspondingly another negative dipole on the opposite side of that neighbouring water molecule. And so the initial purely random dipole that's been generated in the first water molecule has now induced a dipole to exist in the neighboring molecule, neighboring water molecule. And the positive dipole in the neighboring water molecule will then attract the negative dipole

in the original water molecule, forming a relatively weak but still important bonding force. And that force there is called the dispersion force. It's that force that exists as a result of, first of all, the chance accumulation of electrons in a particular region of the electron cloud in one molecule, and then second, that chance induced diaple. then inducing another dipole in a neighbouring molecule, then you've got a positive and a negative dipole which attract each other. Thereby causing, generating a relatively small but still noticeable attractive force between those two molecules. And of course, any individual dipole, induced

dipole interaction like this, any individual dispersion force will bond will not last very long because it's only due to random chance, but the point is they're constantly forming and breaking, so the particles, the molecules are constantly attracting each other to varying degrees. And that keeps the substance together to some degree. The dispersion forces are the only intermolecular forces that are present between non-polar molecules, at least to a first approximation that we'll be talking about now. So they're particularly important for understanding things like hydrocarbons, for example, of their non-polar, hydroparcombs like petrochemicals, for example. A lot of those are non-polar, so you need to know about dispersion forces to understand how they work. And we'll talk more about that

in a bit. Dispersion forces become stronger or larger as the size of the molecule in question increases. The reason for that is essentially because as the surface area of the molecule increases, there's greater chance for disparities in the number of electrons to occur, and so greater chance that a dipole, that a temporary dipole form. Also, there's greater ability for the temporary dipole to be larger because there's more electrons and more area for them to move about. So as the particles in question get bigger, the dispersion forces increase. Now, this principle of the relative size of the molecules,

determining the relative strength of the dispersion forces between them can be applied to understand why the byproducts of crude oil are in the given phases of matter that they are. So the different byproducts of crude oil, like refined oil, tar, wax, natural gas, these sorts of things, they're all essentially derivatives of the same base product crude oil, and they're all basically just hydrocarbons. Carbon molecules surrounded by carbon atoms,

carbon atoms surrounded by hydrogen atoms, but where the number of carbon atoms varies. So the simplest version of this, as we call them hydrocarbons, the simplest form is just one carbon atoms surrounded by four hydrogen atoms. That's called methane. That is a gas, because each methane molecule is very small, obviously, only five atoms, and therefore the dispersion forces are relatively small.

So therefore, the attractive forces, the intermolecular bonding forces, between those, between those methane molecules is relatively small, compared to the strength of the thermal motion, therefore the thermal motion wins, and at room temperature at least, methane is a gas. If you move up to a bit to something like

butane or hexane, which has, I think, four and six carbon atoms, bonded to each other in a chain and then all surrounded by hydrogens, these are larger, and so have more chance for temporary dipoles to form, and therefore have larger dispersion forces between them, these are are therefore liquids at room temperature, and these are the things that you put into your car, the type of hydrocarbons that fuel are used for liquid petrol, liquid oil.

Finally, as you increase the size of the chain even further and get up to some of the really long hydrocarbon molecules, the dispersion forces continue to increase, and so that they actually overpower the thermal motion force forces to such a degree that these substances become solids. And, of course, the long of the chains, they are the more viscous, or the sort of the more solidic the solid will be. So things like wax and tar, for example, or tar that's used for asphalt, that is essentially just really long hydrocarbon chains.

Okay, so that's an application of dispersion forces. It's also important to point out that dispersion forces are really the only, or at least, the only main intermolecular bondi force that is applicable to all types of molecules and particles. So the dispersion forces are always there, and then sometimes there's additional, forces that are added in particular circumstances, and I'll talk about what those are now.

So, we've done dispersion forces now. We're moving on to dipole-dipole bonding. So this is really, I mean, the basic idea is exactly the same as dispersion forces. It's you've got a dipole in one molecule and a dipole in a neighbouring molecule, and the positive dipole in one molecule attracts the negative dipole and the other molecule, and they form a bond in that way. The only real difference is that in dipole-dipole-bonding, the dipoles are permanent. So they'll be generated by something like a non-polar, excuse me, a polar molecule, like water, for example. The positive end of a water molecule, one of the positive ends of water molecule joins up with the negative end of a neighboring water molecule, and they form a dipole-dipole bond. Actually, that wouldn't be a dipole. Well, that might.

So the basic idea is the same, except that is the same as with dispersion forces, except, as I just said, dipole-dipal bonds are going to be permanent, or at least more permanent than dispersion forces, because the dipole itself is permanent. The dipole is a result of the internal structure of the molecule, and that doesn't change. The particular dipole-dipal bond that it has formed, whether it's with molecule A, B, or C, may change over time. But because the dipoles themselves are permanent, the bonds are going to be stronger and more permanent than they are in the case of dispersion forces. So dipole-dipal bonds do tend to be stronger than dispersion-force bonds. However, they're also rarer because they only appear in asymmetrically arranged polar molecules. So that is where one atom in the molecule is attracting, has a high electronegativity than another atom, and also when those atoms are arranged in an asymmetric relationship.

So, for example, I said that methane is a non-polar molecule. Actually, I might have said that, but it is. Even though carbon and hydrogen have different electronegativities, the reason methane is not non-polar, excuse me, is not polar, even though, as I said, the atoms within it have different electronegativities, still non-polar because it's symmetrically arranged. There's one carbon in the middle surrounded by four hydrogens. That's a symmetrical arrangement.

So sort of all of the hydrogens cancel each other out or balance each other out, might be one may be thinking about it. So there's no polar ends to it, there's no positive or negative ends to it. However, something like water, where you have only two hydrogens bonded to an oxygen, and therefore two non-bonding pairs of electrons remaining, that is not balanced out. And so that is a polar molecule. It's also possible, by the way, to have a sort of a compromise or a hybrid between the dispersion forces and the dipole-dipole bonding, where you get one molecule, that has a permanent dipole, inducing a temporary dipole in a neighbouring polar molecule.

So that's even more similar to dispersion forces in the second dipole, in say the second molecule, is induced, didn't previously exist. However, it was induced by the dipole in another molecule, which was a permanent dipole, not just a temporary one that had been caused by random chance, as in the case of dispersion forces. Okay, so that's dipole-dipole bonding. Now I'm going to move on to the last category, hydrogen bonding. Now, hydrogen bonding is really just a special case of dipole-dipal bonding.

It's not actually something different. It's just a subcategory, if you like. So it's really just dipole bonding, not dipole-dipole-bonding, plus, or dipole-dipal bonding that occurs between a hydrogen atom and a very highly electronegative atom, for example, nitrogen, oxygen, or fluorine. And we give this a special name because, well, first of all, it's common in things like organic molecules that we're interested in, and also especially bio-organic molecules. But second of all, because it's really, really strong. Hydrogen bonds can even be stronger than intramolecular bonds, depending on the situation. So

when molecules are bonded together by hydrogen bonding, that's often of particular importance. And if you remember in the, and I'll be talking more about hydrogen bonds in the episode that will be coming out soon on the structure and function of DNA, because hydrogen bonds are very important in joining together the two strands of DNA, there's two strands of nucleotides in a DNA molecule, and therefore are important to the stability of our genetic material. So the reason hydrogen bonds are so strong, just to reiterate this point, is because that in such bonds hydrogen is directly bonded to or attracted to one of the

most highly electro-negative elements, like for example nitrogen or oxygen. And therefore, the fractional positive charge or the dipole that occurs as a result of that situation is going to be very large, not just sort of a normal dipole that you occur when the electro-negativity difference is rather small, but in the case of hydrogen and oxygen or hydrogen-nitrogen, the dipole is very large, because the electro-negativity difference is very large, and therefore the fractional charge difference is going to be large, therefore the dipole-dipal bonding forces is also going to be large, because the greater of the charge difference, basically, the greater the attractive force will be between the two molecules. The effect of this hydrogen bonds is also

magnified because of the relatively low mass, or actually the very low mass of hydrogen atoms compared to other atoms. Therefore, the sort of of the bang for your buck, the bonding force that you get for a given weight, molecular weight, or atomic weight, is particularly high because of that low weight of hydrogen. Now, I mentioned before that water has dipole-dipole bonding,

which is correct, but it would be more correct to say that they have, water molecules experience hydrogen bonding, because you've got a hydrogen atom in the water molecule, which then forms a bond with the non-bonding electron pairs in neighboring water molecules. And liquid water's very high boiling point, or relatively high boiling point, is

largely due to the high number of hydrogen bonds that can be formed. So water is a relatively small molecules, just H2O, and yet it can form, each water molecule can form essentially four hydrogen bonds with its neighbours, with its two non-bonding pairs and then two hydrogens bonded to it, bonded to the oxygen. Since water is a small molecule, but it can form four very strong bonds, that is a very high bonding strength to weight ratio or to mass ratio, and therefore the intermolecular forces between water molecules are very high, so it takes a lot of thermal energy to overcome that. Therefore, the need to heat water to relatively high temperature, 100 degrees Celsius at normal atmospheric conditions, to boil it. And that doesn't perhaps sound very unusual because we're used to

that, but if you think about it, water is one of the only sort of basic substances or pure substances that we know of that is liquid at room temperature. There are some other examples like mercury, for example, and alcohol. But most of the other liquids that you actually know are just water, aqua, solutions, that is, other things dissolved in water. So water is quite unusual in being liquid at such a, what would normally be considered to be a high temperature. Okay, so now I'm going to take these principles of intermolecular bonding and apply them to the case of, to understand surface tension and how that works. Now, surface tension is a property, surface tension is a property of the surface of a liquid that allows that surface to resist an external penetrative force.

Surface tension is important because it's responsible for such phenomenon as the ability of very small, dense objects, like, for example, a pin or very small nails, to float on the surface of water, even though the objects themselves are denser than water, so by rights they should sink, and also the ability of certain insects, for example, water striders, to run over the surface of water, which you may have heard of or seen before that certain insects are actually able to walk on water in a sense. And it's not because they're floating. So this is an important concept. Surface tension is completely different to floating. Flootation, I haven't covered yet, but it will do soon in an episode. Floodation only occurs when the object as a whole is less

dense than the liquid in which it is floating. Surface tension is something different entirely so that these small objects or these water striders are able to stay above the surface of the water even though they are actually denser than the water and so are not floating on it. So how does surface tension work? It looks kind of weird when you see it and it sounds kind of counterintuitive. So I want to use these concepts of intermolecular forces to explain how it works. So remember, each of the water molecules, by the way, surface tension applies to all liquids, but I'm going to talk about it in regards to water because that's by far the most common manifestation of it that you'll see. So remember, each water molecule in the liquid water is attracted to the

neighboring water molecules because of hydrogen bonds. And hydrogen bonds, remember, are particularly strong. And especially, because these hydrogen bonds between the water molecules are particularly strong, water molecules are highly cohesive, which means they preferentially bond together with each other as opposed to bonding with other things, or at least preferentially bond with each other compared to bonding with many other materials. So you can think of it as if the water molecules are trying to pull together, are trying to get as close to each other as they can because of these intermolecular forces, especially the hydrogen

hydrogen bonds. And so what will happen is that that force will tend to lead to the surface area of the liquid being minimized. Now, if it weren't for the force of gravity and other things, that would cause the body of water to become a perfect sphere, which is what will happen to water droplets in a vacuum, and is also why dripping taps and so on look sort of roughly, the droplets look roughly spherical when they fall. It's because that minimizes the surface era, which in turn sort of minimizes the overall energy state of the water droplet, owing to the fact that it's being, all of the water molecules being pulled together by these intermolecular forces. So, given the fact that the mass of water in a sense

wants to minimize its, or tends to minimize its surface area, what will happen if you push inward in one section of the, into one section of the droplet of water or the section of water, what will happen is that the water will tend to resist being pushed in that way, resist being penetrated. And so there's an upward force that, acts upon whatever's pushing into it to push that material outwards, because the water molecules are trying to push together, to try and come together as close as possible to minimize surface

error, but this sort of penetrating object is a sort of thing to circumvent that. So there's a certain force that acts to push this penetrating object out. And that force is what generates the surface tension, because as long as the penetrating force is not too large, so that is as long as the force pushing down into the water is not too large, the force pushing outwards, the force of the surface tension will be greater than the inward pushing force, and therefore the object will not penetrate the surface of the water. So this is how a water strider, for example, can walk on the surface of water strata because the downward force on each of the water strata's feet, owing to the force of gravity acting on the water strata, is relatively low because the water strides not very heavy,

and so it's low enough to be less than the upward force acting on the water striders' length itself from the water, owing to the surface tension of the water, from essentially the water molecules pushing the water stride's legs out of the water in order to minimize the surface area and therefore keep the energy state of the water as low as possible. So this is all a consequence of the intermolecular forces that exist between the water molecules. If you didn't have those forces, there would be no surface tension because there would be nothing to act to pull the water molecules together and therefore to push sort of foreign substances out.

Also, this wouldn't occur, or at least it wouldn't be relevant if the water molecules preferentially bonded to, say, the water striers legs over. other water molecules. So it's particularly occurs in water because the hydrogen intermolecular bonds between the water molecules are quite strong compared to other potential bonds that the water molecules may form. All right, so that's it for intermolecular bonding. Now I'm going to move on and talk about phase transitions, and this might seem sort of a disconnected topic, but actually they're quite relevant because, as I said before, the temperature at which a phase transition occurs really depends upon the strength of the intermolecular forces within that substance.

So, intermolecular forces are really crucial to understanding how and why phase transitions occur. So first of all, I'm going to focus on evaporation and condensation. Evaporation is the process of the vaporization of a liquid that occurs on the surface of the liquid that is going from a liquid to a gas, and condensation is the opposite of that. Now, it's important to distinguish between evaporation and boiling. Boiling is the vaporization of a liquid that occurs inside the liquid itself, whereas evaporation only occurs on the surface of the liquid. So when a puddle dries up in the sun, that's evaporation.

When you put a pot of water on the stove and you see bumbles coming out from it, that's boiling. So it's really just where the vaporization is occurring. Is it inside the liquid, in which case it's boiling? Is it on the surface only in which case is evaporation? Okay, so how does evaporation work? Evaporation will occur when a molecule on the surface of the liquid has sufficient kinetic energy, and also is traveling in the right direction, to escape the attractive forces of surrounding water molecules,

and therefore break free of the surface and become its own free particle, in which case it's entered the gaseous phase. Condensation is obviously the opposite of this. It's when that free molecule hits the surface or sort of moves in the direction of the surface and therefore, and has low enough kinetic energy so that it's sort of captured by the attractive forces of the other liquid molecules,

and therefore joins the liquid and enters that phase. So it's really all about what the kinetic energy of a given molecule is. If it's high enough, it'll escape the attractive forces of the liquid, if it's not high enough, it will rejoin the rest of the water molecules in that liquid, in the body of liquid. Now, as we talked about in a thermodynamics episode, the kinetic energy of a molecule is directly proportional to its temperature. So if we're saying that the evaporation or condensation of a liquid depends upon its kinetic energy,

we're saying it depends upon its temperature, which of course meshes with common experience about the fact that, you know, things tend to evaporate or boil when you increase the temperature, or melt for that matter if they're a result. originally solid. One interesting thing about this process is that if you continually remove the fastest moving molecules from a substance, what you'll tend to do is you're sort of selectively removing the hottest or the highest temperature molecules, and therefore the average temperature of the molecules that are left diminishes. So that's why evaporation tends to actually

reduce the temperature of the liquid that is evaporating. Certainly reduce it lower than it would otherwise have been. This phenomenon is recalled to evaporative cooling, and it's essentially the reason why humans sweat, because as the sweat evaporates off her skin, it reduces the temperature of the skin. Now, because evaporation only occurs at the edge or surface of a body of liquid, the rate of evaporation increases with surface area, so a shallow pond or pool, puddle will evaporate much faster than a deeper one of the same volume, because there's more surface air for that to occur. therefore that essentially means more ability for the atoms to break free and leave the surface

of the liquid and become into the gaseous phase, or vice versa, it's convincing more ability for molecules to be trapped and captured by the other water molecules. Now, this process of vaporization and also of condensation constantly occur at the same time, even if it's really cold or really hot, and you say that the panel's evaporating. Actually, it's mostly evaporating, but there is still condensation going on as well. There are still some water molecules just above the surface. of the pond that are being trapped, that have low enough kinetic energy that they're being trapped by the other water molecules and therefore re-entering the liquid phase. However, if it is the case that

the puddle is in direct sunlight and it's relatively warm, then the rate of evaporation will greatly exceed the rate of recondensation, in which case the puddle will potentially completely evaporate away. However, what can occur is you reach what's called an evaporative equilibrium, which basically means that the rate of evaporation is exactly offset by the rate of, or exactly the same as the rate of condensation, and so the two processes exactly cancel each other out, and so you reach a dynamic equilibrium, where there's evaporation and condensation constantly occurring,

but no overall change in the amount of liquid oil or gas substance. This dynamic equilibrium essentially arises, because as the amount of water that has evaporated, for example, from the puddle increases, if you have an enclosed, if the entire system is enclosed, then the amounts of water that is water vapor that is dissolved in the air increases. And therefore it becomes more likely that some of those water molecules will be become trapped by the water surface again and re-enter the liquid phase. And so the rate of condensation increases.

And that process will continue to occur until the rate of condensation equals or equals out with the rate of evaporation. That process can be circumvented, obviously, if, say, the puddle or whatever is evaporating, is not in a closed system. And, for example, you've got a source of wind to blow the air away so that the already evaporated water molecules are moved on. And so then there's no or very few water molecules able to reenter the puddle or the liquid. And then so if that continues to occur, eventually, all of the water molecules evaporate.

And this is why, for example, that wind helps things to dry more quickly because they remove the water molecules that have already become vaporized, leaving more room in the air to dissolve further amounts of vapor or liquid, and also reducing the rate of return, the rate of recondensation of previously evaporated water molecules back into the liquid. Okay, so that's the process of evaporation and condensation. Now I'm going to move on to talk about boiling, which remember is when you have vaporization that occurs throughout the internal body of liquid. In a liquid, just as molecules are constantly leaving the surface and then returning into the body of water just because of thermal energy,

it's also the case that random collisions of water molecules within the liquid are constantly forming sort of little gas bubbles. You know, all it would take is a couple of water molecules to randomly hit each other in such a way that there's a sort of a small opening or a small hole, or I guess it's a bubble, inside. inside that liquid, even if it's only a few molecules across. And then you get one other water molecule that sort of randomly flies in there. It's a meeting environment, if you're like, having been cleared out by random collisions of other water molecules. And so you've essentially got one water molecule, which is sort of by itself not directly abound, and no hydrogen bonds to any surrounding water molecules.

And that's essentially a tiny little pocket of gas there. And now, these sorts of bubbles are constantly being formed, but of course, you know, a bubble that small and randomly occurring will most likely just within the next fraction of a second collapse because random motion could have caused it, but then at the next fraction of a second, the more water molecules move in and they form new hydrogen bonds and the bubble is gone. In order for boiling to occur, what you need is for those bubbles to last long enough in order to rise to the surface of the liquid through buoyant force, because essentially the bubble will be less dense than the liquid water that's surrounding.

it so they will tend to rise through essentially natural buoyancy, the force of gravity. But in order to reach the surface, they need to last long enough. So boiling a water is not so much about forming those bubbles as permitting those bubbles to last long enough in order to reach the surface, therefore sort of expelling the small amount of gas that has been accumulated into that bubble and then allowing the process to continue. So how do we increase the time that these water bubbles last for? The answer is, as you might expect, that we need to increase the temperature of the water to a sufficient level. Once we get to that crucial temperature, which is called the boiling point,

the average water molecule will have enough kinetic energy so that, in a sense, it's able to maintain that, it's able to exert enough pressure on the sides of the micro-bubbles that form to maintain the bubble. Because the bubble will be maintained, the micro-bubbles that form like that in the liquid will tend to collapse, result of the inward pressure of the water surrounding them, but they will be able to be maintained as long as the water molecules that are in the bubble are bouncing around energetically enough to offset that pressure. So the pressure pushing out from the bubble is the same as the pressure pushing in, the bubble will be stable. If the pressure pushing out is greater than the force

pushing in, by the way, the pressure pushing in, then the bubble actually grow in size, and more, and therefore you've got more water in the gaseous state. As long as the bubbles are stable, and therefore the outward is at least as same as the inward force, then the bubbles will be stable, and therefore will have sufficient time to rise through the buoyant force up to the surface of the liquid, discharging their contents, and thereby boiling away part of the liquid. And then this process continues,

potentially until all of the liquid has been boiled away. Now, this phenomenon allows us to explain why it is that pressure cookers are able to cook so much faster. The reason is that putting water under pressure increases the boiling point. So that means that the water, when water is placed under pressure, say in a pressure cooker,

you have to heat that body of water now to a high temperature in order to get it to boil than you did before. The reason is because essentially there's more pressure on the water, so the inward pushing force that tends to eliminate bubbles is greater, and so we need a correspondingly larger

outward pressure force from the water molecules within the bubble in order to offset that. And in order to get that, we need the kinetic energy of the water molecules inside the bubbles to be larger, which in turn implies a high temperature. So a higher pressure of the water when being boiled implies a higher boiling point. But when you have a high temperature of the water, that also means that you're going to cook

whatever vegetables or other materials are in that water more quickly, because that depends upon the temperature. Remember that I talked about in a previous episode that once you start having a phase transition, like once you start boiling the water, the temperature, additional energy no longer goes into increasing the temperature of water. energy goes into breaking the hydrogen bonds between the water molecules. So if you continue to heat or even turn up the heat on a boiling pot of water, for example, that would tend to increase the rate at which water is boiling, but it will not increase the temperature of the water at all.

And because the rate of cooking depends upon the temperature of the water and not how quickly it's boiling, the only way to increase the rate of cooking once you're already boiling the water is to put the water under pressure, thereby actually increasing the boiling point, and therefore increasing that temperature that you're able to maintain as the water's boiling. Okay, so now that we've covered boiling, I'm going to talk about melting and freezing. Freezing is also called solidification, is a phase change in which a liquid turns into a solid, when its temperature is lowered below the substance's freezing point. Most liquids freeze by a process called crystallization, which is basically the formation of a crystalline solid from the liquid.

A crystalline solid is basically a solid that has an orderly crystal structure, so that the molecules are arranged, or ions, or whatever they are arranged in some sort of neat, orderly pattern. And this crystallization process occurs in sort of two stages. The first is the process of nucleation, and then the second is the process of sort of expansion or growth of the crystal structure. Nucleation essentially refers to the formation of sort of little nuclei or small sort of accumulations of particles which join together in a, which joined together in a more rigid way, consistent with a solid material. So that small nucleation side, that small site with the first initial sort of solid-type bonds, intermolecular bonds that have been formed,

forms the basis upon which the growing solid crystal structure then expands. So once you've got that nucleation side, it's relatively easy for then additional previously liquid molecules, as long as they don't have too much kinetic energy, it's relatively easy for them to sort of get stuck and join that, that growing lattice structure, the growing crystal lattice structure. Once again, this will all be dependent upon the kinetic energy, and therefore the temperature of the molecules,

because the high the kinetic energy, the temperature is, the more likely that the molecules will break away, will break off from the surface of this crystal structure and rejoin the liquid. But if their kinetic energy is small enough, they'll be sort of attracted by or constrained by the growing crystalline structure, and therefore they'll remain in the solid state or enter the solid state. Once again, the rate of freezing will depend upon the surface area available, so that's why

nucleation sites are particularly important, so often if you have small impurities in a substance or you put it in a container with a large surface area of sides, it will freeze more rapidly because it has those nucleation sites to begin building the crystal on. The hardest part in a sense of freezing is forming those initial nucleation sites. So the higher the surface area you have for those to form, the faster will freeze. and on a related note, substances that are referred to as antifreeze, or at least one mechanism by which they work, is essentially just disrupting the formation of regular lattice structures or disrupting the crystallization process. So, for example, if you put salt in water, that reduces the freezing point below zero,

meaning that you now have to cool the water even more, reduce the kinetic energy of the water molecules even more in order to freeze it, because the lattice structure has been disrupted by the presence of those salt ions. The salt ions make it that much more difficult for water molecules to be attracted to the denuclation sites, to the growing crystal and lattice structure, and so they need to lose even the water molecules, need to lose even more kinetic energy before they're going to do that, therefore you need to cool the substance down even further. So in a sense, these antifreeze, or in this case the salt ions are just getting in the way of lattice formation, thereby depressing the freezing point.

melting is pretty much the exact opposite of everything I've just said. So instead of the water molecules or the liquid molecules losing kinetic energy and joining lattice structure, they gain kinetic energy and pull away from the ladder structure, and therefore the lattice structure deforms and melts. Okay, finally, I just want to very briefly talk about sublimation, which is a bit less well-known than the evaporation and melting and so on. Sublimation is the process by which a substance goes directly from being a solid to a gas, so it does not pass through the liquid stage.

Some substances sublime, for example, dry ice or solid carbon dioxide, it sublines quite readily at atmospheric temperature and pressures. So if you've ever seen blocks of dry ice, for example, they at least used to, I don't know if they still do, but at least used to use these on stage to generate smoke. Because dry ice, if it's placed in, well, just room temperature conditions, will give off this vague mist or smoke, well, sort of looks like smoke. missed is probably a better word to describe it. That is carbon dioxide gas, which is subliming, going directly from being in the solid frozen stage to the gas stage, as a result of the increase in temperature. Regular snow and ice actually do sublime as well

when their temperature is below their melting point, but they're exposed, for example, sun. So essentially the kinetic energy of the surface molecules is being increased so much that they break away from the solid ladder structure completely and just become gas molecules, but the overall temperature of the water, of the water, liquid molecules that remain, excuse me, not the liquid, the solid molecules that remain is sufficient to keep them in the lattice structure. So the whole thing doesn't melt, but you do

sort of gradually pull off or remove layers of the outermost molecules that are attached to the solid lattice. Freezer burn is an application of sublimation. Freezer burn is essentially damage or sort of oxidation marks to meat, to frozen meat, that occurs when the meat is exposed to the air or improperly packaged. It doesn't really, it doesn't make the food. It doesn't make the food. inedible, tend to make it less tasty. Reason being that if the frozen meat is exposed to the air, some of that ice, some of the frozen ice, even within the meat, will sublime, so we'll turn into a gas, thereby depriving that portion of the meat of moisture, because the water's gone, and making it become dry and shriveled and sort of look burned or blackish, and therefore

not be particularly appetizing. So that's why it's important to pack meats and other things that are going to be frozen properly so that they're not exposed to the air and therefore do not suffer this freezer burn. It's not really anything like a burn, by the way. It's just a word that's used, because the surface manifestation kind of looks like a burn, but the process is really quite different. The opposite of sublimation, by the way, is called deposition, which will be a gas going directly to the solid stage, and it has essentially a similar mechanism. Okay, so that's all the phase transitions that I want to cover, and that's the end of the episode. Hopefully you enjoyed it. If you want to send me an email with any advice or questions

or whatever. My address is Fods12 at gmail.com. Thanks for listening, and I'll talk to you next time.