18

Vapor is a much older word alluding to dampness and it was not coined by scientists. It is in use since the 1300s. The actual meaning of meaning of vapor is "Matter in the form of a steamy or imperceptible exhalation; esp. the form into which liquids are naturally converted by the action of a sufficient degree of heat. This is the original 13th century ...


17

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 ...


17

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

The normal use distinguishes "vapour" from permanent gas At normal lab conditions there is a (fairly obvious) distinction between things that could exist as liquids and things where no liquid phase is possible. Oxygen, for example, is a permanent gas, but dichloromethane is not. But the vapour pressure of dichloromethane is pretty high and there ...


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

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{...


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 ...


10

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 ...


10

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{...


10

I'm surprised the OED has such a strict definition for gas. I could not find a strict definition in the IUPAC color books (certainly not in the gold book). Presumably these words are in such common use that their definition is assumed understood or easily found. The analytical compendium (orange book) and physical chemistry book (green book) mention vapour (...


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 ...


8

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 ...


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 ...


7

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 ...


7

Dissolution(solvation) is solvation of solute molecules by molecules of solvents. This decreases chemical potential $\mu={\left(\frac{\partial G}{\partial n}\right)}_{T,p}$ of the solute, comparing it (often hypothetically ) to the chemical potential of the solute at the same concentration in gaseous phase. Different solvents cause different chemical ...


7

If the nitrogen is truly inert, its presence will have no direct effect on the chemical potential of the water in the gas phase. Essentially, the gaseous water is "unaware" of the nitrogen. Consistent with this, in a mixture of ideal gases, the chemical potential of each gas is determined only by the temperature, and its partial pressure. It is ...


7

Chemistry is very complicated and I'm probably missing certain edge cases, but I think the answer is almost certainly 'no'. Salts lower the vapour pressure of water because intermolecular forces attract the salt ions to the water, raising the energy needed for water molecules to escape. In order for water's vapour pressure to rise, the salt would need to (on ...


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 ...


6

In college, I had a thermodynamics teacher who was awesome. He had a way of explaining things that were accurate and easy to understand. He explained this difference to us this way: A gas will not condense into a liquid with an isothermal compression. (i.e. an Ideal Gas) A vapor will form liquid when isothermally compressed. His example: When you boil water, ...


6

tl;dr– "Gas" and "vapor" aren't mutually exclusive. Generally: a gas is any material that'd fill a volume to its boundaries; and a vapor is a gas-like material that's associated with a condensed-state transition. It's a bit misleading for a state-diagram to label a region "vapor" in a manner that might imply that a vapor's ...


6

Liquid evaporation occurs at any temperature, just its rate changes with temperature. You must have seen washed clothes getting dry in open air, water evaporating at room temperature. OTOH, I guess you have not observed boiling sea, lakes or rivers to obtain clouds - putting aside geothermal activities. There are few things for you to understand: Molecules ...


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 ...


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

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 ...


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