A pressure cooker is basically a metal container with a lid. The lid components, vital for the pot’s function and operation, are: the rubber seal, the pressure regulator and the pressure exhaust valves. When the lid is closed, the pressure cooker creates a hermetic air and steam seal.
Since the cooking liquid is heated above boiling point, steam is created, and since such steam cannot escape the pressure sealed pot, it remains trapped and thus, producing pressure. The internal cooking temperature varies, depending on the different pressure levels created by the trapped steam.
The pressure amount is measured in kPa or in kgf/cm2. Some pressure cookers only cook at high pressure, while others have two or three pressure levels. Most pressure cookers cook at temperatures between 112° C and 118° C, which means 19° C and 25° C hotter than when boiling in a regular pot, besides considerably accelerating the cooking process.
As you probably know, a substance temperature depends on the speed with which its molecules move: a low temperature corresponds to slow molecules, high temperature are fast molecules. But, since they are an unimaginably high number of molecules in any macroscopic amount of substance (as, for example, a glass of water), that speed which determines temperature is not that of a concrete molecule, but the average speed of all the molecules together. Some of them will move at precisely this speed, others a little faster, others slightly slower, etc. The diagram can be something like this (it doesn’t intend any rigor at all, but to give you an idea of molecular velocity distribution):
In addition, the molecules are changing velocity all the time because, since they move (faster the hotter the water), they are continually colliding with each other. When they collide, they transfer velocity amongst each other – maybe one of them, after the collision, moves slower than before, and the other moves faster. Evidently, the process is so fast, the molecules so small and so many, that “from the outside” everything seems static to us, but it is not at all. Molecules gain and lose energy at each other’s expense.
As you can see, except if there is an enormous difference between both differences, there is always a molecule with enough speed to escape. The necessary energy to do so would have been “stolen” from other molecules in a collision, and the other molecule would have been left moving so slowly that it remains in the left section of the diagram, without hopes of being able to reach the gassy freedom… until it collides with another and, randomly, receives enough energy to escape.
The system’s weak point, therefore, it is the statistical velocity distribution and continual energy transference amongst some molecules in their continual collisions. And that is, by the way, something I have not seen explained often in text books, which simply say that water evaporates, period, and don’t say why. But, as you may have asked yourself, there is a catch to this aleatory escape system: the collisions amongst the molecules after the first. Even if one molecule, in the inside of the liquid, achieves the necessary velocity to get rid of the grip of those near it, it is still surrounded by a crowd of them! Yes, it will move very quickly, free as the wind for an instant… until it bumps with a nearby molecule and transfers it part of its energy, moving slower and coming back to liquid. Escaping is not so easy.
There are only two ways of preventing this. The first of them is that the molecule that has reached, even if briefly, the escaping velocity is not inside the liquid, but over the surface or very near it. If this happens, and besides molecule movement’s direction is “outwards”, then it will be able to escape the liquid before colliding with any other molecule. Since there are always molecules in the liquid’s surface (or there would not be a surface), and that there are always molecules that move with enough velocity to escape, there are always molecules escaping the liquid if it is “exposed”.
The process I just described is, therefore, one of the two ways in which the water is vaporized, i.e. turns into gas. In this case it is a slow vaporization, which is produced in the liquid’s surface, and receives the name evaporation. As you can see if you observe the diagram above, the closer the average velocity of all the molecules is to the “escape velocity”, more H2O molecules can escape the liquid for a determined period of time, with which the liquid will be able to evaporate faster. There are other products that have an influence, and we will talk about them in a moment, but I hope that the basic evaporation concept is clear.
As well you may be thinking, if every time a molecule escapes is because it has “stolen” the energy necessary from its companions, the rest of them, as a whole, will gradually lose energy, which has been taken by the escaping “traitors”. With what, if anything happens, increasingly less and less molecules would escape, and the liquid would gradually cool (since the temperature is precisely the measure of the average kinetic energy in the molecules). But ah, this doesn’t stop there.
This liquid is not alone in the Universe, so, it does begin to cool… Which means that it is colder than its surroundings: the container walls, the ground, the air, anything. And, as a consequence, the surrounding substances, warmer, transfer energy in the form of heat and in the end the liquid remains more or less as it is. Hence, when the water evaporates after a storm, the environment cools. The same happens when we sweat (in fact, that is one of the reasons why we do it), since the hottest substance to the evaporating water is our own body. And, as it is well known, a traditional drinking jug keeps the water fresher than a plastic bottle.
As we have said before, besides the temperature there are other factors that influence the evaporating velocity. Evidently, the liquid’s nature is an influence. If the forces between the molecules are very intense, it will be difficult that one of them is able to get rid of the others, for which the necessary velocity would be very big and the other way around. The amount of steam over the liquid is also an influence, since as a liquid molecule can be lucky enough to move in the right direction, one of the gas can move toward the liquid, fall in, collide with the molecules down there and be “trapped” again. The more steam there is over the liquid, more steam turns into liquid through this process, with which the net evaporation is slowed down.
More interesting yet is the effect of the liquid’s impurities, if it is not pure. On one hand, the substances dissolved in it can exercise their own “binding forces” on the liquid’s molecules that can escape with more difficulty. But, even if this doesn’t happen, any liquid with stuff dissolved in it evaporates more slowly, and the reason is quite logic: the larger the liquid’s surface the liquid can escape though is, the easier the evaporation is. But of course, if there are more things in the liquid, the entire surface is not “open field” to escape, because part of it is occupied by the dissolved substance. Therefore – to make an exaggerated example – if one of every two molecules in not the liquid’s, but the other substance’s, then the evaporation will produce two times slower than when it was pure.
But evaporation only allows the molecules close to the surface to escape. Is there any hope for the rest? Yes there is, and that is the second way the liquid can escape. If the temperature increases enough, the average molecule velocity will be high enough to allow many of them to have the necessary velocity to escape. When this happens, there is such a large number of “traitors” that they can form areas in which there are only molecules that are not bound to those around them: gas bubbles inside the liquid. This bubbles are less dense than the liquid around them, with which they eventually surface and then, at last, the molecules that form the may escape to the exterior and move freely.
This second vaporization, which is violent, is produced everywhere (not only in the surface) and requires a higher temperature, which is the boiling point, and that is what we refer to when we say that the water is boiling. The temperature in which this happens massively is precisely the liquid’s boiling temperature. So the boiling identification with evaporation is a subtle and implicit in “the water boils at 100° C”, although the phrase itself doesn’t contain it explicitly. But the cipher itself is incorrect, because there is one more factor to talk about in our silly “traitor water molecule escape” dramatization.
In the first place, we have seen what happens to the H2O molecule when it is able to escape the liquid… but not afterwards. If, as in a bad action movie, it is perfectly possible that a molecule achieves the necessary velocity to escape, and exits the liquid… to come back to it immediately. To understand why, we just need to remember that the escaping molecule – except for the rather unusual circumstances, of which we will talk about later – it doesn’t do so in vacuum. If you imagine the pot with hot water, an H2O molecule that shoots out with the necessary velocity from the liquid surface finds many other molecules out there, in the air: some of the other “traitors” that escaped before it, other O2, CO2, N2, etc.
Of course, over the surface there is a gas, with which it has a much lower density than the liquid water that the escaped molecule leaves behind: the molecules in gas are much less tight together, so the “traitor” will probably take longer to collide with another molecule than in the liquid. But surely, sooner or later, it will collide and energy transferences will produce amongst them. The consequence is that some molecules escaping the liquid, barely reach the outside air, collide with another molecule in the exterior with such bad luck that they fall back in the liquid again, they collide there with another liquid molecule transfer it part of its energy and, back again! Since they don’t have enough speed to escape.
Although I’m aware of how pathetic my “traitor” and “enslaving cooperation” example is, I continue if it helps you visualize the issue. It is as if the molecules in the air over the liquid were the last protective barrier, outside the “enslaving cooperation”: some of them hit the escaped molecules and send them back in. To be able to run away from the liquid it doesn’t only have to evade the rest of the water molecules, but it also has to evade that external barrier of several molecules that move over the liquid.
I think then it should be logical the fact that, besides the liquid temperature (at higher temperature, more molecules have enough velocity to escape), the number of “external” molecules have an influence over it: if there is too few, once they escape the liquid it is very easy to move out there without colliding and falling back in. If there is an enormous number of them, the exterior is like a can of sardines in which it is very difficult to go in without receiving a good bump from any of them and go back in the liquid.
But, what does it mean, macroscopically speaking, that there is few or many air molecules over the liquid? That is what we notice, in the atmosphere, as air pressure. The more we climb towards the top f the atmosphere, “less tight together” are the molecules, i.e. less pressure there is, and vice versa. The pressure we take as a reference is the one at sea level, a little over 100 kilopascals (kPa)
In a more technical way, the gas pressure outside the liquid has influence over the liquid’s boiling point temperature. The more the pressure, the harder it is for a molecule to escape the liquid and for this reason the higher the boiling temperature is, and the other way around. Although not everybody knows the reason behind it, this is a well known fact by those who live quite above the sea level, since it has an influence in cooking time. Yes, we are almost getting to the pressure cookers, patience!
If you want to cook macaroni on top of Mount Everest, for example, you will immediately notice that “water boils at 100° C” is a lie, because this value assumes the reference pressure of 100 kPa at sea level. On top of the Everest, the air pressure over the water pot is only 26 kPa, with which it is much easier for a molecule to escape the liquid, because it barely has air molecules over it. As a consequence, water on top of the Everest boils more or less at 70° C. The notion of boiling water being very hot is false: our intuition indicates so because it is “trained” at a determined pressure, but as we will see later, it is possible to have a glass of water cold to the touch but bubble boiling.
And this, naturally, is a problem if you wish to cook there. While the water boils, all the energy you transfer it (if you are cooking in the Everest, probably through a flame) is inverted in the change of state, and not in heating the water. So, while you are cooking your macaroni up there, the water is never going to go over the 70° C. And, in consequence, the macaroni will take much longer to be ready than if you were at sea level.
What is the solution then, if you want to cook pasta or any other thing in boiling water at the Everest? Tightening the gas molecules over the liquid, so these traitors cannot escape easily! And, what is the easiest way of achieving this? Well, as you can imagine… a pressure cooker.
Although there are safety valves and several mechanisms, ill and soon said, a pressure cooker is nothing more than a hermetic container that can endure considerable pressure differences between the “inside” and the “outside”. Lets imagine that you have one of these pots at the Everest, and you begin to heat the macaroni water.
When you close the pot and begin heating, in the beginning nothing unusual happens: the water begins rising temperature until it comes close to 70° C, so it begins turning into water faster and faster. The liquid molecules can escape without a problem, because there are very few gas molecules over them due to the scarce pressure. But, where do this free molecules go, if they are trapped in a hermetic container? Since they can’t escape the pot, there, over the liquid, more and more H2O accumulate trapped in the container.
And there is the irony of it all: the “traitor” molecules escaped from the liquid, are now the “guardians” of the ones still in the water! Gradually, the pressure inside the pot rises because the water steam accumulates… with which the boiling temperature rises and the water can be heated but without actually boiling. Besides, being water steam over the liquid, many of these “escaped” molecules go back into the liquid water and cease to be steam. When heating, more molecules reach the necessary velocity to escape the liquid, even despite the new H2O molecules that press over it… but since they join them, they increase the pressure even more and allow the boiling temperature to keep rising. This way, you can cook on the Everest with a pressure inside the pot a couple of times higher than the one on the outside, and therefore, a higher temperature, which makes your macaroni’s cooking time shorter.
But there is nothing to prevent you from doing it at sea level, and in fact we do it every time we use a pressure cooker at home. The one I have in front of me right now indicates a maximum safe pressure of 1.5 bars, i.e. 150 kPa, 50 % more pressure than sea level. Therefore, the water boiling point temperature inside mi pot, when “packed” with water steam, is over 100° C: in this case around 112° C, with which I can cook at 12 degrees more temperature than if I wasn’t using it, and those twelve degrees decrease cooking time by a lot. There are other pressure cookers that can endure quite larger pressures, of course: I have seen them up t 200 kPa, but I am convinced that there are even better ones, and can reach temperatures of 120-130° C.
Naturally, the opposite happens as well: the less the pressure on the liquid there is, the easier the H2O molecules escape from it. So, if you climb higher than the Everest summit, water boiling temperature will decrease below 69°C, 50° C, 40° C… there would be a point, at 19 km of altitude over sea level, where it would reach the value of 36.7 ° C, our body’s temperature! This point is called the Armstrong Line, named after Harry George Armstrong, the first person to describe this phenomenon. And then, experienced and brave El Tamiz reader, your body’s water exposed to the air would begin boiling as the macaroni’s. Your saliva, teardrops, any mucous exposed to the air… I don’t want to imagine the feeling. What happens to the blood and other fluids inside the body is more complicated, because the pressure inside there is higher than outside, but the dangers over the Armstrong Line are several and the situation is not pleasant at all, except if you have a pressurized suit. And farther even, if you take a look at the diagram before, you can have a glass of water at 5° C, which you’ll notice it’s cold if you place your hand inside it and, however, it would boil rabidly. I would really love to have that experience.
Of course there is a lot more to say (and I hope we do some day when we begin, at last, a series about thermodynamics), and I am sure that, if you have read about this issue, the explanations were quite more rigorous than mine. But my goal was to simply give you an intuitive idea of the difference between evaporation and boiling, and the logic reason of pressure’s influence over both. So, next time you cook in a pressure cooker, remember what you are really doing: catching those traitors to prevent them from escaping!