16
Boiling is not about the upper surface at all. What happens there is but of little interest to us. Boiling means that bubbles are forming at the bottom and probably within the bulk of the liquid. Guess what is inside these bubbles? It is pure water vapor. That's where the equilibrium between vapor and water actually comes into play.
As for the second ...
15
Yes they will boil all right. Sure, there might be some kinetic impediment to it if you let the liquids to settle in layers, but if you stir them so as to expose their surfaces, they will boil†. This is what steam distillation is all about.
As for the first law, it will hold just fine. You burn your firewood, you get the heat, but it is not for free: ...
14
How are the vapors transported?
The additional question from the comments:
Why [is] the amt. of vapours taken up proportional to vapour press. Why directly proportional? Why not nothing else?
First things first, let's get a definition (Wikipedia, italics added):
Vapor pressure [...] is defined as the pressure exerted by a vapor in thermodynamic ...
13
I don't believe there is an equation that you can use for melting points of a general substance as a function of pressure (since the melting phase transition has a lot to do with the geometry of the molecule and the structure of the solid), but there is one for the boiling point of any pure substance when you are not near the critical point.
The liquid-...
12
This is not just some vapor pressure. This is the equilibrium vapor pressure. Thermodynamics is all about equilibrium, you know. And equilibrium, roughly speaking, is what takes place in a closed container after a billion years.
Immiscible as they are, the liquids still have some solubility in each other (maybe extremely low, but anyway). Over the course of ...
11
First some general comments about Raoults law before discussing the particular solutions in question.
Raoult’s law suggests that the partial pressure of each substance above a solution is proportional to its mole fraction x, thus $p=p^ox$ where $p^o$ is the vapour pressure of the pure substance.
Experimentally there are deviations from Raoult’s law and ...
11
Yes, this is indeed the case. The reasoning behind it is using chemical equilibria to their fullest.
If you have a Brønsted acid and a Brønsted base in the same vessel, you will always have an equilibrium of the following kind:
$$\ce{HA + B- <=> A- + HB}$$
It depends on the nature of the acid and the base — i.e. their $\mathrm{p}K_\mathrm{a}/\mathrm{...
9
There is a liquid state for carbon dioxide. Borrowing the $\ce{CO2}$ phase diagram from Wikipedia, we can see that $\ce{CO2}$ will condense at a few atmospheres, dependent on temperature. At still higher pressures, the liquid will solidify. Below the triple point temperature, the gas will transition directly to solid.
9
I think that the second explanation is the correct one. A simple experiment can prove it: if you place a porous cover over the pure solvent, its vapour pressure will not change. In alternative, you can place a few corks and observe the same result (no variation in the vapour pressure). In both cases surface sites are blocked, but the vapor pressure will ...
9
Nothing special would happen, immiscible liquids would just form layers. As for the expression, $$p_T=\Sigma p^o_i$$
I suggest you read this answer. Quoting Ivan Neretin:
This is not just some vapor pressure. This is the equilibrium vapor pressure. Thermodynamics is all about equilibrium, you know. And equilibrium, roughly speaking, is what takes place in a ...
8
For most substances, higher pressure (or air pressure, in your case) will cause the melting temperature to go up. To think about it intuitively, imagine that you have a certain solid. Melting it would increase the volume of that substance because liquids take more space than solids. If you increase the pressure, it becomes harder for that transformation to ...
8
Although "paradox" is not quite the right term, what you have discussed is actually a simple, yet interesting and important phenomenon.
Given the ideal situation as you have presented, your thoughts on what would happen are correct. If the system were to achieve $\pu{100\%}$ humidity with respect to the pure water, that would always be slightly over $\pu{...
7
Look up the term “phase diagram”. At different temperatures and pressures different phases are favored. Think about water. At 0 °C and 1 atm water and ice are equally favorable. There can be freezing or melting. At other temperatures and pressures there can be sublimation and deposition.
The sublimation of water is involved in freeze drying and freezer burn ...
6
As Philipp stated in the comments it is Le Chatelier's principle in play.
The principle in layman's language states that
If you try to bring about any physical change in a system the system will try to but not necessarily succeed in cancelling the change.
The reaction you have given is an equilibrium reaction where all the compounds are in gaseous ...
6
Is there any extra force holding the two components together?
Correct. Namely, hydrogen bonding. As a quick refresher, hydrogen bonding is an electrostatic attraction between an electropositive hydrogen atom on one compound being (weakly) attracted to a highly electronegative atom on another, nearby molecule, typically oxygen, ...
6
It almost certainly was due to thermal expansion of water and the air above it. Consider that the thermal expansion coefficient of water near room temperature is ~$0.000214$ K$^{-1}$. [Though it is not linear over a broad range, assume for the small temperature rise in your car that it is.]
If the water was cold (e.g. ~$277$ K, maximum density) and it ...
6
There is a book The Design of Controlled-atmosphere Chambers for the Study of Oxygen Toxicity, so I suggest "contolled-atmosphere chamber".
6
"Sealed in a vacuum" is an oxymoron, a contradiction in terms, much like "frozen with fire". If you seal a liquid in a flask containing nothing else but vacuum, then a part of the liquid will quickly evaporate and fill the flask with vapor, so it would no longer be a vacuum.
The said vapor will exert some pressure, depending on the temperature and the ...
5
Let's reword it:
Traces of oil and grease are distilled off at temperatures below their boiling points.
The temperature in the frying pan is above 100 °C.
The food does contain significant amounts of water.
This sounds like the conditions for steam distillation, except that you're collecting the mixture of water and water-insoluble material everywhere in ...
answered Dec 19 '14 at 7:48
Klaus-Dieter Warzecha
41.3k7 gold badges85 silver badges148 bronze badges
5
There can be no scientific explanation, because the argument itself is incorrect.
For the most part air is nitrogen. To be specific, about 80% of dry air is made up from nitrogen. In fact both nitrogen and oxygen, the second major components of air that takes up almost all the other 20%, behave extremely like ideal gas. Pressures of dry air and pure ...
5
Since the room temperature butane didn't instantly cool when it hit the container, it was still above the boiling point.
If you had a large freezer that was roughly at equilibrium at a temperature below $-1^\circ C$, and you put both the pressurized container and the empty container in there and let them reach equilibrium, then when you opened the ...
5
You don't really need to do anything to get LPG vapour. It is a gas a room temperature and pressure so as long as you just want the clean gas not mixed with air (which would be dangerous as the mixture will be flammable and potentially explosive), all you need to do is allow the output of the canister to vent into a balloon.
LPG is mostly propane and butane ...
5
Let's have a look at the definitions of dew:
dew (the free dictionary)
Water droplets condensed from the air, usually at night, onto cool surfaces.
Something moist, fresh, pure, or renewing: "The timely dew of sleep / ... inclines / Our eye-lids" (John Milton).
Moisture, as in the form of tears or perspiration, that appears in small drops.
...
5
It is true that "non-volatile" is a relative term and that virtually all compounds have some, possibly infinitesimal, vapor pressure, but the idea is that this is negligible. And you are also correct that "...the molecules of the solute that are at the surface between the solution and the vapor and should not be held by the strong intermolecular bonds of the ...
5
Step 0) Realize that sometimes the answer given for a homework problem is wrong.
In this case it appears that the correct answer is $\pu{6.11 kg}$ of water, not $\pu{6.00 kg}$ of water.
The approach to the problem given below is to simply calculate the total water mass in the house before and after cooling, then subtract the latter from the former to ...
5
If water molecules go into the gas phase from the liquid then they are in a different phase. But water molecules go into the gas phase one at a time. If there are the same number of water molecules leaving the liquid as coming into it then the atmosphere above the liquid would have 100% humidity and be in equilibrium with the liquid.
Now if some small ...
5
Is it possible to boil a liquid by just mixing many immiscible liquids together?
No*, boiling is when the vapor pressure of a phase is greater than the ambient pressure. You might create a total pressure of all of the partial pressures in excess of ambient, but it wont cause boiling since the partial pressure of each of the phases is less than ambient, no ...
4
It is not that surprising that you got this answer: the existence of azeotropes is due to non-ideal behaviour of mixtures (see for instance the Wikipedia page and the associated diagram). It is not surprising that when you impose an ideal behaviour of the solvent that follows Raoults law, you get some unexpected results. Basically, what your calculation ...
Only top voted, non community-wiki answers of a minimum length are eligible
