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November 6

filling an oxygen in the bottles from three others continuously

I've noticed when the First-Aid-responders fill their small tanks, they fill the oxygen by three big tanks. It says, they take one small tank (2 liter approximately) and connect it to one big tank and open the tap, after a minute they close the tap they pass the small tank to another tank big one for a minute, and eventually they do the same processes with another big one tank (third time...) I'm volunteer student, and I would like to understand what happen here... 213.57.99.33 (talk) 00:08, 6 November 2014 (UTC)[reply]

My supposition is that this arrangement results in a higher final pressure for all the small tanks than if the three big tanks were simply connected in parallel, or used individually. The first big tank loses the most pressure, but the pre-pressurisation of the small tank to whatever its pressure is results in the second big tank losing much less and retaining a higher pressure, and the repetition on the second tank preserves most of the pressure in the final topping-up tank. —Quondum 00:25, 6 November 2014 (UTC)[reply]

(edit conflict) Can you point to any images or documents that show this behavior? The process you describe is probably specific to the equipment being used, so I doubt we'll be able to give any kind of a definite answer without more context. Dragons flight (talk) 00:27, 6 November 2014 (UTC)[reply]

Here's a nice article from Praxair (one of the largest commercial suppliers of gas products in the United States): An Inside Look. I know from personal experience: it is really complicated to fill a tank with cryogenic liquid/gas like Liquid Oxygen! There are a lot of safety procedures and a lot of practical problems. For example, do you tap from the top or the bottom of the tank? The answer depends on what you are doing. As you fill, the temperatures and pressures in the supply- and the destination vessel are changing - gas always flows from high pressure to low pressure. If there's a gas pressure head, it will drive the liquid to flow in the same direction - even working against gravity! If you aren't carefully watching the pressure, you might be using the small tank to fill the big tank! Be sure to follow safe procedures when you are working with oxygen (... it is a strong oxidizer)! Nimur (talk) 00:39, 6 November 2014 (UTC)[reply]

(EC) It sounds like you're referring to a Cascade storage system used for transfilling. Note that images e.g. [1] you can see may be a bit confusing (and I didn't find the explanation in our article very good and most search results are just vendors). It looks like all the tanks are connected together and there is only one output valve you connect to, which may suggest you're using them all at once. However as can be seen in this video [2], you only open one of the large tanks at a time, moving from low pressure to high pressure which seems to be similar to what you were describing (well you didn't know/mention pressure differentials).

The reasons appear similar to what Quondum said, and as described in the video and also to some extent here [3]. You basically start with a low pressure tank, fill the bottle to equilbrium, and then close it and move on to the next higher pressure one until you reach the desired pressure. That way you can end up with a relatively high pressure in the last larger bottle (which you may not use at all, if your earlier ones are high enough).

Once the last/highest pressure larger bottle is below the minimum pressure needed for you final fill or alternatively if the first/lowest pressure is so low it's rarely useful (like if the tanks you're filling would generally be at a higher pressure), you can just remove the lowest pressure bottle (to be an empty to be filled. And replace it with a full bottle, now the new last tank. You renumber the bottles rather than moving them around, I presume because it's safer and easier.

So you don't end up with a lot of low pressure bottles that are below the minimum needed but still have a fair amount of oxygen (or whatever), in other words, you get more from each large bottle for a given pressure. I'm not sure but there may also be an advantage in possibly reduced temperatures changes, a lower pressure differential and easier monitoring for safety reasons.

The disadvantage may be [4], although I'm not sure that using a single bottle will help in any way. It may be the alternative the commentators in that link are thinking of is professionally/externally filling the small bottles rather than using a cascading storage system to fill the small bottles by someone for who it isn't their primary job.

BTW, it's somewhat unclear to me if in your example the larger bottles are all connected together (although I note you seem to have added the same video below). As mentioned, this appears to be the norm in such systems (based on the images and vendors) and there would seem to be advantages (namely you aren't moving and reattaching the smaller bottle so much which would seem to be an avenue for mistakes) but I guess it's possible for cost or other reasons, in your example they aren't. I presume they otherwise work similar.

P.S. In terms of numbers, if you only ever use 3 in this set-up before you need to replace the lowest pressure one, then I guess there's probably no advantage to adding 4 more to get 7. Unless for some reason you prefer to keep the "empty" ones in situ and simply unnumber or whatever and don't use them until they are replaced in bulk. Which is possible, since they are probably quite heavy and there are likely additional safety precautions. Nil Einne (talk) 01:28, 6 November 2014 (UTC)[reply]

It's kind of what we see in this video (but we have only three big tanks, here there is 7 I think). I must to say, that the three big tanks that I'm talking about they are the same thing: oxigen. (it's written on) 213.57.99.33 (talk) 01:02, 6 November 2014 (UTC)[reply]

That cascade storage system article makes a lot of sense. Basically, I assume that the site wouldn't want to return every oxygen container to be refilled the moment its pressure drops below whatever they're filling the little containers up to. Not when you can still use the depleted containers to fill little containers that are nearly empty... the rest follows logically. To me this seems very similar in concept to a countercurrent exchange. Wnt (talk) 02:52, 6 November 2014 (UTC)[reply]

This is a really neat analogy. With enough bottles in the cascade, one would expect that energy put into the compressed gas in the small cylinders is transferred from the cascade with efficiency approaching 100% (asymptotically); also, the cooling due to decompression will be asymptotically zero. —Quondum 14:08, 6 November 2014 (UTC)[reply]

The relation between Pulmonary embolism and birth

From time to time we hear about women who get PE (Pulmonary Embolism) as a result of birth. My question is what is the relation between the two. And why does the birth raise the risk of getting it? 213.57.99.33 (talk) 01:13, 6 November 2014 (UTC)[reply]

In this freely-available, high-quality article describing the risk factors for venous thromboembolism (e.g. PE), the authors note that the risk starts rising in the first trimester, and is primarily attributable to changes in the blood levels of clotting (and natural anticoagulant) factors. They argue teleologically that the shift to increased tendency for clotting might protect against hemorrhage, a major cause of death during pregnancy. -- Scray (talk) 02:02, 6 November 2014 (UTC)[reply]

Thank you deeply! now it makes sense. 213.57.99.33 (talk) 02:14, 6 November 2014 (UTC)[reply]

Air permeability of glass

Hi. the article on Spontaneous Generation has this picture of a glass flask sealed by Louis Pasteur in the 1860s. Let's assume for the sake of argument that there are no microfractures from the sealing process, and the glass is all uniformly unblemished. Is the air inside it the same air (in other words, all the same atoms) that were sealed in 150 years ago? My Dad and I had a disagreement; he felt that the majority of it would have cycled through, my impression was that, even if it's not "all" the same air it's surely mostly the same. Looking for data with Google gave me inconclusive info, but it seems like glass is (to some greater or lesser degree) permeable by air. Does anyone have any info on just "how" permeable it is? Anyone care to make an estimate of what percentage of the atoms currently in the flask (stored carefully but not exhaustively so in a lab and then a museum archive) would have been in there in the 1860s? Obviously, there would be some degree of temperature and thus pressure fluctuation driving air in and out (assuming glass is to some degree permeable). Clearly it's not very permeable by water because there's still broth in it (though whether the amount of broth has changed in the last 150 years is a separate question), but H2O would seem (to this non-chemist) to be a much larger molecule than H or N or O that make up the majority of air. 75.140.88.172 (talk) 08:07, 6 November 2014 (UTC)[reply]

I find it HIGHLY unlikely. The internuclear distance between the silicon and oxygen atoms in SiO₂ (glass, quartz, etc) is about 1.67 angstroms (0.167 nanometers).[5]. Air is basically a mixture of O₂ and N₂. The internuclear distance between oxygen and nitrogen atoms in molecules of O₂ and N₂ is 0.120 nm and 0.110 nm respectively: [6]. So, IF you had a sheet of glass that consisted of a SINGLE atomic layer of silicon dioxide the pore size would JUST be big enough to let a molecule of air slip through OCCASIONALLY, if it hit the EXACT hole at the EXACT right angle, maybe. In glass as thick as that flask, you've got a network many trillions and trillions of atoms thick; there just aren't any pores large enough to let any molecules of the air through, given the tolerances we're talking about here. Ain't no way. Assuming there isn't any other way for the air to leak out, it certainly isn't coming through the glass itself. --Jayron32 12:21, 6 November 2014 (UTC)[reply]

Diffusion does occur in solids. Water, in fact, is particularly fast as it actually reacts with silica up to about 1% of its normal density! At high temperatures, the effective diffusion coefficient is high enough that the glass is quite permeable indeed; I get a 6-meter scale length for the diffusion profile after 150 a. However, at room temperature the Arrhenius equation predicts so much slower a process that the scale length is just 3 μm. The diffusion coefficients for air are theoretically larger at high temperatures (above 1500 K, which (coincidentally?) is the glass transition temperature for fused quartz), but because it cannot diffuse chemically, the activation energy is almost twice as large and the room temperature rates are much smaller. It looks like room-temperature nitrogen can cross only about one atomic layer in the glass even over 150 a. --Tardis (talk) 14:53, 6 November 2014 (UTC)[reply]

Nice explanations, I think the OP has a good answer now. Related: does anyone think this would be worth adding to our article on Long-term_experiments? On the one hand, we could say the experiment is over, and helped disprove spontaneous generation. On the other hand, we could say that it has only disproven spontaneous generation on a ~150 yr time scale! I see this jar as similar to the Oxford_Electric_Bell (which is also included in the article). The Oxford Bell is similarly not really an active experiment, it's just a neat old science artifact that is still around. SemanticMantis (talk) 16:36, 6 November 2014 (UTC)[reply]

Isn't the much larger electron cloud of Si going to be more important in stopping diffusion than the relative internuclear distances, nucleii being point particles for the purposes of this exercise? Also, N2 and O2 are much bigger molecules than H2O, aren't they? I imagine trying to sneak a pair of apples through a wall of bricks a mile thick. Of course, if the internal surface of the glass oxidizes, could not a charge differential eventually lead to the loss of a few O2 molecules or SiO2 molecules from the outside. Maybe over the length of time the flask will simply shrink as (X)O2 is lost from its outer skin? μηδείς (talk) 18:55, 6 November 2014 (UTC)[reply]
Glass contains no O2 molecules, nor strictly, does it contain any SiO2 molecules. Silicon dioxide (quartz, glass, etc.) is a network solid and does not contain any discrete molecules, as such. The formula SiO2 is merely the ratio of Silicon to Oxygen in lowest terms and does not imply the existence of discrete molecules as such. The actual composition is basically this pattern of atoms repeated billions and billions of times in all dimensions. --Jayron32 02:46, 7 November 2014 (UTC)[reply]

I am entirely aware of that. My point is that on the inner and outer surfaces we should expect places where there are imperfections. Should oxygen react with the inside surface, the change of charge might induce the release of oxygen from the outside of the bottle, perhaps at a hugely slow rate. Obviously if the oxygen can't react with the inside surface of the bottle you can drop my suggestion. I don't think you've said that. The idea is the same as reducing rust in iron by attaching a sacrificial zinc plate to it. μηδείς (talk) 20:19, 8 November 2014 (UTC)[reply]

Right, and in case anyone doesn't know, the structure in the bottle walls is not precisely that pattern, because glass is an amorphous solid. SemanticMantis (talk) 14:17, 7 November 2014 (UTC)[reply]

Sideeffects of Madras Eye

Recently, people who suffered Madras Eye(Conjuctivitus)are found having some sideeffects in addition. The skin of their whole back is reddish with itches. Is this the sideeffect of Madras Eye.117.193.119.71 (talk) 16:31, 6 November 2014 (UTC)[reply]

See what Pinkeye has to say about it. Anecdote: I've had it in the past, and my back was unaffected. ←Baseball Bugs ^{What's up, Doc?} carrots→ 16:58, 6 November 2014 (UTC)[reply]

Conjunctivitis is a type of disease caused by many different pathogen. The OP should look at Madras Eye and if that doesn't help, google ("Madras Eye" and "Back Rash") if this is a question of curiosity or, if he has a rash, ask a doctor about it. μηδείς (talk) 17:02, 6 November 2014 (UTC)[reply]

Also see Hives for the back itches. Ariel. (talk) 12:03, 7 November 2014 (UTC)[reply]

Oils and the Trachea

Can some edible large-chain Oils, and especially Mineral-oils "find there way" into the Trachea and "Settle" there? while food is swallowed?, if yes, can it be problematic to Humans&Animals consuming them?, thanks. Ben-Natan (talk) — Preceding undated comment added 17:40, 6 November 2014 (UTC)[reply]

This sounds like some sort of bizarre plot to kill Batman in the 1960 TV series. Not a direct answer, but see cough, and aspiration pneumonia. μηδείς (talk) 18:58, 6 November 2014 (UTC)[reply]

If I understand correctly, Whenever oils get into the Trachea, the Cough-able Organism would Cough till they come out? (As clued, I don't know if I understand correctly). Ben-Natan (talk) 21:09, 7 November 2014 (UTC)[reply]

You might also want to look thru Lipid pneumonia. Mihaister (talk) 23:04, 7 November 2014 (UTC)[reply]

It still not clear to me if oils consumed by Eating can in some cases, accidentally go into the Trachea\Lungs? Ben-Natan (talk) 03:11, 8 November 2014 (UTC)[reply]

Absolutely, I am sure you can accidentally choke on some oily food. You will most likely cough. But if your lungs are unhealthy or you cannot cough (say, an elderly stroke patient with aspiration pneumonia,) lipids in the food could be a major problem.

Or are you perhaps talking about like sewer pipes getting clogged with a waxy build-up where restaurants dump their cooking oil down the drain? That's been illegal for years in much of the US for just the point you suggest. μηδείς (talk) 05:35, 8 November 2014 (UTC)[reply]

I didn't know about the US restaurants doing that and I'm not a US resident of any kind... Gladly, My lungs are just well and I can cough in such a choking case case - It is good to get this data from you and I'll keep it in mind and hope others do. It's good to know that indeed. Ben-Natan (talk) 13:07, 8 November 2014 (UTC)[reply]

Ugh. I just looked up mineral oil aspiration on PubMed and found this and a whole bunch of mostly shorter term studies with similar results noted - see the search for more. Wnt (talk) 09:44, 8 November 2014 (UTC)[reply]

T

Kindle region restriction

I've taken the liberty of moving this to WP:Reference desk/Computing#Kindle region restriction. Wnt (talk) 20:42, 6 November 2014 (UTC)[reply]

How much oxygen liquid will be in the same tank

If I have 2.4 liter oxigen tank in the room temprature,how much oxygen liquid will be in the same tank in condition of the boiling point? How can I calculate it?5.28.172.180 (talk) 20:52, 6 November 2014 (UTC)[reply]

Are you asking how much volume 2.4 liters of liquid oxygen would take up once it has boiled and been brought to room temperature at standard pressure? -- ToE 22:27, 6 November 2014 (UTC)[reply]

I think the OP is asking "If we cool the tank from room temperature to the boiling point of oxygen, what will be the volume of the liquid?" To answer this, we need to know the pressure of the oxygen tank at room temperature. From this, we can work out the mass of the oxygen using the ideal gas equation - we know p, V, T and R, and the molar mass of oxygen is 32 g/mol (it's actually 31.9988, but 32 will do for this calculation). The density of liquid oxygen is 1.141 g/cm³ (from our liquid oxygen article), which enables us to calculate the volume from the mass. Tevildo (talk) 22:45, 6 November 2014 (UTC)[reply]

Sorry for my abstruseness. I mean to ask: how much oxygen liquid could enter in the same tank of 2.4 liter oxygen (gas)? 149.78.27.187 (talk) 03:13, 7 November 2014 (UTC)[reply]

2.4 litres of liquid oxygen have a mass of 2.738 kg (this just uses the density). It wouldn't be possible to use the oxygen tank to store liquid oxygen, as the tank is not insulated, and the oxygen will boil off. If you want to know how much liquid oxygen is needed to fill the tank with oxygen gas at room temperature, we still need to know the (maximum) pressure. 2.74 kg of oxygen in a 2.4 litre tank at room temperature would require a pressure of about 900 bar, which is much higher than the maximum pressure of a normal oxygen tank. Tevildo (talk) 11:40, 7 November 2014 (UTC)[reply]

Thank you, bu it's my mistake. I've found that the 2.4 liter is the volume of the tank (not weight), but the oxygen needs to be more less of this. So the question is to know how much is it, if it was changed to liquid. In simple words, if you take the tank of 2.4 liter (volume) and you change the state of matter, how much liquid it will be. 149.78.27.187 (talk) 00:19, 8 November 2014 (UTC)[reply]

You may be interested in the concept of molar volume, which allows you to do quick calculations while avoiding explicit use of the ideal gas law. The molar volume for 1 atmosphere of pressure is 24.465 L/mol at 25 °C. One mole of O₂ masses 32 g, so if you divide by the density, 1.141 g/cm³, you find that one mole of liquid oxygen takes up only 28.05 cm³. If you divide that into the 24,465 cm³ molar volume you can show that when liquid oxygen boils off and reaches 25 °C at 1 atmosphere of pressure it increases in volume by a factor of 872. Thus you would need a 2094 L balloon to hold the result of boiling off 2.4 L of liquid oxygen. But I think that you are asking about storing that O₂ in a tank instead of a balloon, so as Tevildo says, you need to specify the pressure of that tank, as the higher the pressure the smaller the volume is needed to contain the same amount of gas. From our oxygen tank article: "Oxygen is rarely held at pressures higher than 200 bar / 3000 psi due to the risks of fire triggered by high temperatures caused by adiabatic heating when the gas changes pressure when moving from one vessel to another." 1 atm (standard atmosphere) of pressure is just over one percent more than 1 bar. If we compress our 2094 L of O₂ to 200 atm (2393 psi -- note that 3000 psi is the working pressure of the AL80 scuba tanks commonly used in the United States), then the volume will decrease by a factor of 200 to 10.47 L. (See Boyle's law, one element of the ideal gas law.) Thus you need a 10.47 L, 200 atm working pressure tank to hold the same oxygen as held by a 2.4 L Dewar flask filled with liquid oxygen.

So the take away here is that O₂ at 200 atm takes up 4.36 times as much volume as when liquefied.

( 1 atm / 200 atm ) * (24.465 L/mol) / [ (32 g/mol) / ( (1.141 g/cm³) * (1000 cm³ / 1 L) ) ]

In practice, most medical O₂ tanks I have seen are pressurized to 2200 psi which is only 150 atm, but searching online I see that 230 bar (3336 psi or 227 atm) high pressure medical oxygen cylinders are available, so the actual multiple you want will depend on the working pressure of your tank.

Is this what you were looking for? -- ToE 13:53, 7 November 2014 (UTC)[reply]

Looking at his question, and using the numbers given above, I think the answer he wants is a 200atm, 2.4L tank of oxygen is equivalent to approximately 550ml of liquid oxygen. 75.140.88.172 (talk) 01:13, 8 November 2014 (UTC)[reply]

Thanks!!! ToE, this is just what I looked for! But I would like to see the formula again in a normal writing because I don't understand this form of presentation. May you put it into the LaTex? for example (I'm not sure it's accurate):

{1atm \over 200atm}*{24.465 \over 32}etc

5.28.179.11 (talk) 03:02, 8 November 2014 (UTC)[reply]

OK, so Tevildo's original interpretation of your question was correct, but that's no problem because our two interpretations are simply reciprocals of one another. Note that Tevildo's use of the ideal gas law is a good way to solve this. I suggested using molar volume because I though it an easier way to conceptualize the problem as opposed to just plugging some numbers into a mysterious equation.

Prettifying(?) my fractions (per your request) we have the expansion ratio for liquid oxygen boiling off at 1 atm and reaching 25 °C room temperature as

{\frac {24.465{\text{L/mol}}}{\frac {32{\text{g/mol}}}{\left(1.141{\text{g/cm3}}\right)\cdot \left(1000{\text{cm3/L}}\right)}}}=872

, and if you then compress that to 200 atm you get the ratio

{\frac {1{\text{atm}}}{200{\text{atm}}}}\cdot {\frac {24.465{\text{L/mol}}}{\frac {32{\text{g/mol}}}{\left(1.141{\text{g/cm3}}\right)\cdot \left(1000{\text{cm3/L}}\right)}}}=4.36

.

So if you take 2.4 liter of oxygen at 200 atm and 25 °C and then liquefy it, the final volume will be

{\frac {2.4{\text{L}}}{4.36}}=0.550{\text{L}}=550{\text{mL}}

, the answer given above by 75.140.88.172.

Now I'm concerned that this may just look like a big confusing fraction, but it is very understandable. The 872 is just a ratio, with the molar volume at 25 °C & 1 atm on top and the volume per mole of liquid oxygen on the bottom, that calculated by dividing the molar mass by the density (after converting the density from grams per cm³ to grams per L). Note that the concept of molar volume works because one mole of any gas (ideally) has the same volume (at any given temperature and pressure). Once liquefied, different compounds have different molar densities, so the bottom half of the ratio is specific to liquid oxygen. -- ToE 13:02, 8 November 2014 (UTC)[reply]

If you are up to the math, you should try to rework the problem using the ideal gas law (as suggested by Tevildo) and confirming that the results are the same. I'd suggest using R = 0.08206 L·atm·mol⁻¹·K⁻¹, as that matches most of the units we have been working with here, leaving you only to convert 25 °C from Celsius to Kelvin. -- ToE 13:10, 8 November 2014 (UTC)[reply]

Thank you for the explanation of the formula. But now I noticed it's not 200 atm but it's 200 Bar. Is it the same for you? Could I put the atm's value into the formula too?5.28.179.11 (talk) 01:27, 9 November 2014 (UTC)[reply]

You have several options. 1 bar is 100,000 Pa (Pascal, the SI unit of pressure equal to 1 Newton per square meter; note that 100,000 Pa would typically be written 100 kPa) while 1 atm is 101,325 Pa. Since the two units differ by only 1.325%, you could just keep the previous answer of 550 ml, figuring that it is close enough. After all, do you really know the volume of the nominally 2.4 L pressure cylinder to an accuracy of 1%? Alternately, you could plug the correct pressure of 200 bar = 197.4 atm into the equations above in place of the 200 atm. You will get an expansion ratio (after boiling off and compressing) of 4.42 (the previous value times 1.01325) and a final volume of liquid oxygen resulting from taking 2.4 L of O2 at 200 bar and 25 °C and liquefying it of 543 ml (the previous value divided by 1.01325). But if you know from the start that you are working with bar, you might as well use appropriate molar volume of 24.789 l/mol for ideal gases at 1 bar and 25 °C. (Note that this value is 1.01325 times the atm based molar volume we used earlier.) This would give us an expansion ratio of

{\frac {1{\text{bar}}}{200{\text{bar}}}}\cdot {\frac {24.789{\text{L/mol}}}{\frac {32{\text{g/mol}}}{\left(1.141{\text{g/cm3}}\right)\cdot \left(1000{\text{cm3/L}}\right)}}}=4.42

, and

{\frac {2.4{\text{L}}}{4.42}}=0.543{\text{L}}=543{\text{mL}}

of LOX.

If you are up to the math (and it is as easy or easier than what we did here), I still suggest that you confirm this number using the ideal gas law, but now that you are using bar you should use R = 0.08314 L·bar·mol⁻¹·K⁻¹. -- ToE 07:10, 9 November 2014 (UTC) Dare I link to real gas?[reply]

ToE, I thank you deeply for the help. Now it's very clear! 5.28.177.33 (talk) 04:44, 12 November 2014 (UTC)[reply]