Naked Science Forum
Non Life Sciences => Chemistry => Topic started by: Jeffa040 on 14/01/2016 19:16:03
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I've got a question about the boiling of water. I'm a first year physics student and from the Netherlands.
I've searched alot about the boiling of water and this confused me. Everyone said something else about the cause of the boiling. Let me explain it further.
Let's say you want to cook some eggs. You put on the gas.
1. As the temperature of the water increases, the evaporation increases.
2. When the evaporation increases, the vapor pressure will increase too.
3. When the vapor pressure is equal to the external pressure, there will form a bubble.
So, my question is: is this a chain/link of causes? So the first link causes the next one? So the temperature increase causes the evaporation to increase which causes the vapor pressure to increase which causes the forming of a bubble (the actual boiling)?
I doubt if it is a link of causes (the one thing causes the other) because they happen at the same time. And in my opinion a cause happens BEFORE the consequence.
In this video they say boiling has more causes. What are these causes?
When the boiling point is reached, there happen two things:
- temperature =100 degree celcius
- vapor pressure that equals the external pressure
Are they both causes of the boiling (coming up of the bubbles)? Or is the vapor pressure that equals the external pressure the cause?
Greets, Anton
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The above description is deterministic: "if (temperature>100C) then (liquid water will become a vapor)".
However, the motion of real molecules is statistical: "as (temperatures approaches 100C) then (the percentage of liquid water that has enough energy to become a vapor approaches 100%)".
Water molecules are always jiggling due to the temperature. As the temperature increases, the average energy increases. But not all molecules have the same energy - at any instant, some have more kinetic energy than the average, while others have less than the average. The energy of any individual water molecule is always changing.
At any temperature below 100C, there are some water molecules that momentarily gain enough energy to escape from the liquid state into the vapor state. If they are near the surface, they may escape into the air in the form of water vapor. However, if they are in the body of the liquid, they have nowhere to go, immediately lose their excess energy by bumping into liquid water molecules, and rejoin the liquid state.
As the temperature approaches 100C, more and more molecules of water have enough energy to leave the water surface as a vapor.
At 100C, there are so many water molecules in the bulk of the liquid that have enough energy to become a vapor that any initial pocket of gas will be joined by other molecules, so that the bubble of vapor grows, rises to the surface, and this is seen as boiling.
Note that the water does not have to be near 100C to turn into a vapor - and does not even need to be a liquid to turn into a vapor - at low pressures and low temperatures (such as occur on Mars or at the poles of the Moon), water can turn directly from a solid to a gas.
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Thanks for the clarifying answer. The temperature reaching 100C and the vapor pressure reaching the required level to equal the external pressure happen at the same time right? Can I still say that the temperature reaching 100C is the cause of the vapor pressure to equal the external pressure and the vapor pressure is the cause of the boiling?
Because in my opinion causes happen before the consequence. But in this case the cause (temperature) happens at the same time as the consequence (vapor pressure). Although this fact, it's still the cause if I interpret you right?
Below 100C there will form tiny gas bubbles with vapor pressure lower than the external pressure, this will make them collapse when they bump into the liquid, right?
What do you mean with a pocket of gas? Is it a tiny bubble? Or is there a difference between a pocket of gas and a bubble?
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I would like to go a little beyond what has already presented, connected to bubble formation; nucleation, and nucleation centers. Nucleation centers are surfaces that ease the formation of bubbles. If we have no nucleation centers when we heat water to 100C, it is harder of the bubbles to form, resulting in a phenomena called superheating. Superheating is where a liquid can continue to exist in a metastable state above it's boiling point. Liquid water can be superheated to 240C and higher.
One way you can demonstrate superheating at home is use two glasses of water and a microwave oven (use safety glasses). One glass will contain tap water and the other glass will contain distilled water. Tap water contains impurities that will act as nucleation centers. While distilled water lacks these nucleation centers and can superheated.
In the experiment, both equal glasses are heated at the same time, until the tap water glass begins to boil. This means we have reached 100C. The distilled water glass will be calm, but superheated. Next, add a pinch of salt or sugar to the superheated water. The boiling will be begin violently as soon as these nucleation centers are introduced.
Bubble nucleation, needed for boiling, slows exponentially with the height of a free energy barrier ΔG*. This barrier comes from the free energy penalty of forming the surface of the growing nucleus.
As the liquid boils and changes phase to a gas, the growing gas bubble, in the liquid, creates surface tension between liquid and gas. There is a free energy barrier to overcome, due to the compounding surface tension. If we have an impurity, it can lower this free energy barrier, allowing the bubble to form easier. If there are no nucleation centers, we will need more than just the heat of vaporization, to overcome the free energy barrier of the growing bubble; superheat.
Extra tidbits about nucleation;
Nucleation and nucleation centers are also important to freezing. This is called supercooling.Water can remain a metastable liquid down to -48C. This version of supercooled water is called a fragile liquid. Supercooling is also connected to a free energy barrier at the liquid-solid interface. What is cool is, as we heat supercooled water, it will freeze. The heating provides the free energy needed for nucleation, so the supercooled liquid becomes solid.
There is also a version of supercooled water called a strong liquid, that can be produced from glassy amorphous ice between -123 °C and - 149 °C. It will remain a liquid in equilibrium with ice, as it is heated unit it reaches -63C. This differs from the fragile liquid in that the liquid is so cold, the molecules of water can't diffuse to form a solid, until -63C. Then the water will freeze.
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As the temperature approaches 100C, more and more molecules of water have enough energy to leave the water surface as a vapor.
At 100C, there are so many water molecules in the bulk of the liquid that have enough energy to become a vapor that any initial pocket of gas will be joined by other molecules, so that the bubble of vapor grows, rises to the surface, and this is seen as boiling.
You say two different things. First you say that before reaching 100C, the vapor molecules will leave the bulk. Later you say AT 100C there are many water molecules that have enough energy to become a vapor. Which one is right?
Do the water vapor's leave the water to join the bubble before the reaching of 100C? Because the vapor pressure is reached at 100C at the exact same time, right? So the joining of the bubble (increasing of the vapor pressure) has to occur before the 100C. Because if they would leave the bulk AT the time of reaching 100C, the vapor pressure equalling the external pressure would occur a bit later than reaching 100C.
When the temperature is below 100C, the bubbles CAN form and grow a bit, but they collapse from the external pressure around them. Right?
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in my opinion causes happen before the consequence. But in this case the cause (temperature) happens at the same time as the consequence (vapor pressure).
Trying to disentangle cause and effect for a pot of water on a gas flame:
- The rapidly jiggling molecules of the flame cause the electrons and atoms of the metal pot to jiggle more actively
- The rapidly jiggling atoms of the pot cause the molecules of water to jiggle more actively
- When we measure the temperature of the flame, the pot and the water, we are really measuring the kinetic energy of the molecules. Higher kinetic energy → higher temperature
- Conduction of the metal pot, and convection of the water spreads the jiggling through the water (approximately even temperature throughout).
- At any temperature below 100C, some water molecules near the surface may have enough energy to leave the surface as a vapor. The closer to 100C, the greater the fraction with this minimum energy to overcome surface tension, and the higher the vapor pressure.
- the greater the fraction with this minimum energy to overcome surface tension, the faster the rate of evaporation.
- At any temperature below 100C, some water molecules within the volume of the liquid may have enough energy to enter a bubble as a vapor. However, as you say, the vapor pressure is less than the water pressure, so the bubble will collapse.
- At 100C, something unusual happens - the temperature stops increasing, even though the flame continues to inject energy into the pot & water. This is called "latent heat of vaporization (http://en.wikipedia.org/wiki/Latent_heat#Table_of_latent_heats)"
- Assuming that sufficient nucleation sites exist in the body of the water (see post from puppypower above)... These effectively allow a tiny bubble to form easily (eg in a scratch in the pot).
- Water molecules with sufficient energy near the bubble will pass into vapor within the bubble. It takes energy to overcome surface tension, so this cools the liquid, keeping the temperature at 100C.
- If you keep pouring in more energy, more molecules will enter the bubble, and the bubble will grow (while keeping the water temperature at 100C)
- When the bubble gets large enough, it will detach from the nucleation site, and rise to the surface.
- The nucleation site is effectively left with a tiny bubble, to seed another bubble of water vapor.
- When there is no more water left as a liquid, the temperature of the pot will rapidly increase beyond 100C.
So I would say that the increase in temperature, and the increase in vapor pressure occur simultaneously because they derive from a common cause: The increase in average kinetic energy of the water molecules as your pour in more thermal energy from the flame.
What do you mean with a pocket of gas?
I was thinking of the nucleation sites described by puppypower.
You say two different things. First you say that (1) before reaching 100C, the vapor molecules will leave the bulk. Later you say (2) AT 100C there are many water molecules that have enough energy to become a vapor. Which one is right?
(1) is slightly misquoted: bulk-> surface, but otherwise, correct. The difference between the bulk of the liquid and the surface of the liquid is that within the bulk of the liquid below 100C, an energetic water molecule will be surrounded by other water molecules of lower energy, so the water molecule will end up back as a liquid. However, at the surface, there is only water on one side, and some of the vapor molecules will escape into the air (especially if there is a breeze to blow them away from the liquid water).
(2) is also correct. At 100C, the rate of evaporation increases, but also there are enough molecules with sufficient energy to remain as a gas within the bulk of the liquid (at nucleation sites).
I see no reason why (1) and (2) are contradictory, since they cover different temperature ranges. (1) is temperature < 100C, while (2) is at 100C.
For a superheated liquid with no nucleation sites (ie liquid water with temperature > 100C, not covered by (1) or (2)), almost all the molecules have enough energy to become a vapor; as soon as the liquid is disturbed, most of the water molecules turn into vapor explosively, bathing the unfortunate disturber in superheated steam.
add a pinch of salt or sugar to the superheated water.
I have heard of people being burnt on their face and arms while tipping coffee granules into superheated water.
Avoid doing this at home!
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Ah, sorry, I didn't notice the bulk>surface thing.
I read alot about this subject and some say that the vapor pressure is the pressure of all the vapor together, some say it's the pressure of a bubble of gas. Which one do you mean when you say that the vapor pressure rises simultaneously with the temperature?
So when the temperature increases, more water molecules will become a vapor, thus the vapor pressure increases (the general vapor pressure, of all vapor's together). When there forms a little bubble, more vapor will join the bubble so the vapor pressure equals the external pressure > this will cause a bubble to grow and rise.
Is it correct when I say that the vapor pressure of the bubble will increase a little bit after reaching 100C? Because you said that when the water is at 100C, there are many water molecules with sufficient energy to leave the bulk and join a bubble. And when they join a bubble the vapor pressure of the bubble will increase (and when it equals the external pressure, it will grow and rise).
Will the water molecules with sufficient energy to leave the surface become a vapor IN the bubble? Or do they become vapor first and then leave the surface, joining the bubble?
I am very thankful for your help, this means alot to me. I've been on this subject for a long time.
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When the average kinetic energy increases, the temperature increases. When the average kinetic energy increases, more water molecules will leave the bulk and join a bubble, this makes the vapor pressure grow. The process of leaving the bulk and joining a bubble will take some time, so the vapor pressure will increase a bit later after the kinetic energy increases, do you get what I mean?
Or did you mean more or less simultaneously in stead of at the EXACT same time?
Edit:
I read the following thing, this answers my question.
''The liquid is immediately adjacent to the bubble, and the molecules of liquid at the interface between the bubble and the liquid do not have to travel anywhere to enter the bubble.''
So when 100C is reached, this is the exact time that so many water molecules will join the adjacent bubble that it grows and rises at the same time (because of the vapor pressure reaching the external pressure). So the temperature reaching 100C is the exact time that the vapor pressure reaches the external pressure (cause of the many water molecules joining the adjacent bubble).
Doesn't the water molecules have to move at all to join the bubble? Doens't this take a very small amount of time?
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I should also point out, if it hasn't been addressed already, that you can also boil water without increasing the temperature at all (in fact it goes down). If one uses a vacuum pump, or some other means to reduce the pressure above the water to the vapor pressure of water at whatever temperature the water happens to be. The water will boil vigorously and become quite cold in short order, eventually freezing and boiling at the same time (reaching its triple point).
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Thanks for the reply. Maybe you know the answer on the following question.
Does every tiny small increasement in temperature will make a water molecule break the IMF? Let's say the temperature will go from 90.0001 C to 90.0002 C, will this also lead to a water molecule breaking the IMF and thus increasing the vapor pressure?
Will the water molecule break the IMF during the increasement or after the increasement, at the moment when the energy (temperature) reached the next level (i.e. reaching 90.0002C)?
Let's say in example: an increase from 90.0001 will break 20 water molecules from their IMF while an increase from 99.9999 to 100C will break alot more water molecules from their IMF? And thus the vapor pressure reaching the external pressure.
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There are a few seemingly simple points about this (which are actually very complicated, once you think about it for a while):
1) Temperature is a bulk property. It relates (roughly) to the average random (internal) kinetic energy of a system. By random I mean not correlated--a bullet moving 1000 mph has a lot of kinetic energy, but all the parts are moving at very nearly the same velocity (same speed and direction). A cup of hot tea can appear to be stationary, but each of the molecules within it can be moving at 1000 mph in a random direction.
But there is more to kinetic energy than just molecules moving around (translation energy). Molecules also have rotational energy, vibrational energy, and electronic energy (this last one doesn't really apply unless we get to very high temperatures).
2) There are many, many (MANY!) water molecules in a pot of water. I mean, mind-bogglingly many! Imagine a pot of water with 1.8 liters (kg) of water in it. It contains more than 6x1025 molecules of water (that's about a million times as many stars as there are in the visible universe...)
So one can easily confuse themselves by going back and forth between thinking about bulk properties (like temperature) and individual molecules, or small numbers of molecules (like less than a few trillion). The best way to avoid these confusions is to think about distributions (https://en.wikipedia.org/wiki/Boltzmann_distribution) Because we stipulate that molecules are moving randomly, there are some that are moving really fast, some hardly moving at all, and everything in between. This distribution describes the proportion of molecules in each state (speed). And because there are so freakin many molecules, even the most bizarre things happen. Think about it, if each molecule has a one in 1015 chance of doing something at any given moment, then in a collection of 1025 molecules, there are (on average) about 1010 molecules doing that very weird thing at any given moment. (think about that a while)
3) Finally, equilibrium is another bulk property. Single molecules can never be "in equilibrium" it must always be a collection (usually >1012 molecules) that can reach equilibrium. What equilibrium ultimately means is that conditions are just right such that for every molecule doing one thing, there is another molecule doing exactly the opposite (or for every group of molecules doing one thing, there is another group doing the opposite). Imagine two rooms adjoined by a single door. If the number of people going from room A to room B is the same as the number going from room B to room A, then the number of people in each room stay roughly constant. Even if only one person can fit through the door at a time--then the numbers will oscillate a little bit, straddling that equilibrium (think of it as a dynamic equilibrium).
Vapor pressure equilibrium is established when the rate of molecules going from liquid state to gas state is equal to the rate of molecules going from gas state to liquid state (and yes, it is silly to talk about a single molecule in gas or liquid state, but bear with me...). Because the rate at which molecules go from gas to liquid is dependent on the number of molecules in the gas state in a given amount of space (the density of the gas, which is essentially the pressure of the gas) there is a relationship that naturally arises in which: if there are fewer molecules in the gas state than this equilibrium pressure, then evaporation happens faster than condensation; if there are more molecules in the gas phase than the equilibrium pressure, then condensation happens faster; the system will automatically shift until it has reached equilibrium again...
So what does this all mean for your question? Will the water molecule break the IMF during the increasement or after the increasement, at the moment when the energy (temperature) reached the next level (i.e. reaching 90.0002C)?
Let's say in example: an increase from 90.0001 will break 20 water molecules from their IMF while an increase from 99.9999 to 100C will break alot more water molecules from their IMF? And thus the vapor pressure reaching the external pressure.
A tiny increase in the temperature of the water will lead to a disequilibrium, where now slightly more water molecules are leaving the liquid for the gas phase. This will continue until the pressure of gaseous water increases to the point where they are back in equilibrium (with now a slightly higher fraction of all total water molecules in gas state than when it was slightly cooler).
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There are a few seemingly simple points about this (which are actually very complicated, once you think about it for a while):
A tiny increase in the temperature of the water will lead to a disequilibrium, where now slightly more water molecules are leaving the liquid for the gas phase. This will continue until the pressure of gaseous water increases to the point where they are back in equilibrium (with now a slightly higher fraction of all total water molecules in gas state than when it was slightly cooler).
Thanks for the reply.
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I've got another question. Wikipedia says that boiling is the following'
''Boiling is the rapid vaporization of a liquid, which occurs when a liquid is heated to its boiling point, the temperature at which the vapor pressure of the liquid is equal to the pressure exerted on the liquid by the surrounding atmosphere.''
This isn't true boiling right? Doesn't the rapid vaporization of a liquid happens before the boiling point too? What i've heard you guys is that boiling is the forming, growing and rising of a bubble of gas in it.
Chemwiki says the following:
''Boiling is the process by which liquids are heated beyond their "boiling point" and undergoes the change from the liquid phase to the gaseous phase. Boiling is the converse process of condensation, in which an element or molecule in it's gaseous phase is converted to a liquid.''
The change from the liquid phase to the gaseous phase happens long before 100C, right? Like some guy here said, some water molecules DO have the energy to change to the gaseous phase before reaching 100C. And the evaporation happens too before 100C. Why is it that they say the change from liquid phase to gaseous phase is boiling?
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If you place a glass of water on warm surface, say at 40C, which is well below 100C, the water will still evaporate. However, bubbles will not continuously form. Bubble formation is separate thing from evaporation, although, both often go hand in hand at the boiling point.
Although I have seen a pot of water on a radiator, where the water is way less than 100C, with bubble on the side of the container. These bubbles are not connected to boiling, but are do the container making bubble nucleation easier. If I put water on a waxed surface, the water will bead up into a bubble without any effort. Some surfaces promote bubbles, while others making bubble formation harder.
Building a bubble, at the boiling point, is sort of like blowing up a ballon, in the sense that a force needs to be applied, against external resistance. In the case of a bubble, this resistance is called surface tension. Depending on the container, the external resistance to bubbles forming and growing, can be high or low.
In superheating, the blowing up of the bubble balloons will feel strong resistant forces. Even though superheated water may not boil in the sense of bubbles, it will still evaporate.
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