BerthelotOnCrack
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SO3 preparation : help about excessive smoking
Hello,
I recently prepared SO3 using the NaHSO4 pyrolisis method in a quartz glass flask.
From the moment pyrosulfate was formed and especially when the sulfur tryoxide finally came down the condenser tho, a *lot* of white mist come out of
the apparatus, (see pic is the smoke produced in something like 5 seconds and fills the 500mL bottle already)
I was expecting some fuming but not as much, considering all of the gases pass throught the water-cooled condenser. It's like half the SO3 refused to
condense, even when the cooling water is at 15°C.
Assuming this is sulfuric acid mist, I tried but was *completely* unable to condense it, even using a cold water trap not shown here, which is
supposed to be too hygroscopic for SO3 fumes..
Any advice on how to efficiently deal with the very high volume of mist ? Is it even H2SO4, it does not seem hygroscopic at all.
While we are here, if anyone tries to dissociate NaHSO4, my extensive record of failed attemps shows that only a nichrome heating wire will transfer
enough heat to the solution, even a bunsen burner will fail at providing enough energy to liberate the SO3 at a reasonable rate.
[Edited on 31-7-2024 by BerthelotOnCrack]
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Rainwater
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Had this question show up on a test recently.
If I recall correctly the reaction of the trioxide with water is so exothermic it has two major effects
1) the solubility of the gas in water decreases at higher temperatures. Slowing down the reaction
2) the kinetics of this reaction slow down with increasing temperature.
But basicly the atmosphere is so saturated with vapor you get fog, like on a rainy day. Without physical forces acting on the cloud to force it to
condence, it just hangs around.
Edit: it should be noted that i did not ace that class.
[Edited on 31-7-2024 by Rainwater]
"You can't do that" - challenge accepted
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BerthelotOnCrack
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Quote: Originally posted by Rainwater  | Had this question show up on a test recently.
If I recall correctly the reaction of the trioxide with water is so exothermic it has two major effects
1) the solubility of the gas in water decreases at higher temperatures. Slowing down the reaction
2) the kinetics of this reaction slow down with increasing temperature.
But basicly the atmosphere is so saturated with vapor you get fog, like on a rainy day. Without physical forces acting on the cloud to force it to
condence, it just hangs around.
Edit: it should be noted that i did not ace that class.
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Thank you, so i'm making H2SO4 clouds now ? Ok I see, but how do they not dissolve when I literally bubble them through water ? That SO3 had no hard
time condensing on my throat
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Twospoons
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The particles are so fine they remain suspended in the bubbles.
This might be a good place to try an electrostatic precipitator. You would need a voltage of 5-10kV connected to a needle to create an ion stream that
would charge the SO3 droplets. The charged droplets would then be attracted to the nearest surface at ground potential, which could be a metal bar
inside your bottle.
Helicopter: "helico" -> spiral, "pter" -> with wings
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Keras
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Okay, I tried this reaction too and here is my advice:
1. A suggested in the book Small Scale Preparation of Laboratory Reagents, you should use a cut. I’m sure most of the mist you get is from the
simultaneous release of both water and sulphur trioxide, which happens mainly under 300/350 °C. So: preheat your bisulphate at 350 °C, leaving your
flask open. You’ll get a lot of fumes, but these are mainly water, very few sulphur trioxide. Once the emission of mist stalls, remove the heat
(presumably a heating mantle), let it cool, attach the flask to the distillation apparatus and resume heating. Sulphur trioxide gas is colourless when
pure.
2. For heating, I used two blow torches, one on each side of my quartz RBF. Three might be better. Considering they’re worth pennies, consider using
two or three to get the necessary heating power.
3. Condense the trioxide in a RBF immersed into cold water, using a vacuum adapter at the end of your condenser. Connect the output of the vacuum
adapter to a wash bottle full of concentrated sulphuric acid.
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BerthelotOnCrack
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| Quote: Originally posted by Twospoons | The particles are so fine they remain suspended in the bubbles.
This might be a good place to try an electrostatic precipitator. You would need a voltage of 5-10kV connected to a needle to create an ion stream that
would charge the SO3 droplets. The charged droplets would then be attracted to the nearest surface at ground potential, which could be a metal bar
inside your bottle. |
Wouldn't a metal bar choke and die instantly in SO3 ?
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BerthelotOnCrack
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| Quote: Originally posted by Keras | Okay, I tried this reaction too and here is my advice:
1. A suggested in the book Small Scale Preparation of Laboratory Reagents, you should use a cut. I’m sure most of the mist you get is from the
simultaneous release of both water and sulphur trioxide, which happens mainly under 300/350 °C. So: preheat your bisulphate at 350 °C, leaving your
flask open. You’ll get a lot of fumes, but these are mainly water, very few sulphur trioxide. Once the emission of mist stalls, remove the heat
(presumably a heating mantle), let it cool, attach the flask to the distillation apparatus and resume heating. Sulphur trioxide gas is colourless when
pure.
2. For heating, I used two blow torches, one on each side of my quartz RBF. Three might be better. Considering they’re worth pennies, consider using
two or three to get the necessary heating power.
3. Condense the trioxide in a RBF immersed into cold water, using a vacuum adapter at the end of your condenser. Connect the output of the vacuum
adapter to a wash bottle full of concentrated sulphuric acid.
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Thank you, but shouldnt the trioxide dissolve better in cold water than in a sulfuric acid trap ?
The fumes would be produced art 350°C but the SO3 is already coming out, how could the temperature be so low, bisulfate is supposed to dissociate a
mt a reasonable rate at ~700°C
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Keras
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| Quote: Originally posted by BerthelotOnCrack |
Thank you, but shouldnt the trioxide dissolve better in cold water than in a sulfuric acid trap ?
The fumes would be produced art 350°C but the SO3 is already coming out, how could the temperature be so low, bisulfate is supposed to dissociate a
mt a reasonable rate at ~700°C |
No, don’t ever try dissolving sulphur trioxide in water, even cold. The reaction is so exothermic you will end up with a mess. On the other hand,
SO₃ can be easily dissolved in concentrated sulphuric acid. You obtain oleum, that you can then dilute with water to get 100% sulphuric acid again.
If you have a look at the book I mentioned, you’ll see that about 20% of the SO₃ passes between 250 and 350 °C (IIRC), while 85 to 90% of the
water is vaporised in this range. This means that, in the range 400 to 700 °C, you get 80% trioxide and nearly no water.
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BerthelotOnCrack
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| Quote: Originally posted by Keras | | Quote: Originally posted by BerthelotOnCrack |
Thank you, but shouldnt the trioxide dissolve better in cold water than in a sulfuric acid trap ?
The fumes would be produced art 350°C but the SO3 is already coming out, how could the temperature be so low, bisulfate is supposed to dissociate a
mt a reasonable rate at ~700°C |
No, don’t ever try dissolving sulphur trioxide in water, even cold. The reaction is so exothermic you will end up with a mess. On the other hand,
SO₃ can be easily dissolved in concentrated sulphuric acid. You obtain oleum, that you can then dilute with water to get 100% sulphuric acid again.
If you have a look at the book I mentioned, you’ll see that about 20% of the SO₃ passes between 250 and 350 °C (IIRC), while 85 to 90% of the
water is vaporised in this range. This means that, in the range 400 to 700 °C, you get 80% trioxide and nearly no water. |
Thanks a lot for the book, I was looking at a less precise paper from 1969, I'll set up a sulfuric acid trap, will there be a suckback like with NO2 ?
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Keras
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Yes, especially at the end when you remove the heat source,
so use two washing bottles to avoid any possible suck-back.
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phlogiston
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I've never played with SO3, so for what it is worth, but passing an aerosol through a plug of glass wool in a tube is very effective at trapping tiny
aerosol droplets, much more so than wash bottles.
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
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Rainwater
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sorry, I take safety too serious
First I have to address this. putting it mildly, The dangers of this compound are extreme. If you can smell it, or get any burning sensation, you are
not using the proper PPE and containment procedures
You can die a $<"(ing god awful painful death. This stuff has been used in war crimes and is as unforgiving as a bullet! you will not get many
second chances.
Glad your ok.
keeping an open bottle of concentrated ammonia near your apparatus will quickly reveal any leaks and convert the deadly SO3 into the
irritating ammonium sulfamate
next is a bubblier. the larger the bubbles, the more ineffective the device. I like to stuff a ton of loosely packed glass wool or sand into mine to
decrease the bubble size. 100% absorption means that the bubbles go in, but don't make it to the surface.
as well as a secondary containment vessel in case my bubblier blows its top and turns into a guider, geezer, ----- old faithful ----.. in case my
bubblier starts spraying stuff everywhere.
"You can't do that" - challenge accepted
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Keras
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SO₃ is odourless (contrarily to SO₂, probably because the trioxide is never found free in the atmosphere, therefore no one ever evolved a need to
identify it – nor would it have been possible anyway), but it is visible. Don’t dismiss its harmfulness, but don’t go overboard either. Operate
outside, and you’ll be fine. I did my experiment in a barn with enough ventilation and all went well.
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chornedsnorkack
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| Quote: Originally posted by Rainwater | Had this question show up on a test recently.
If I recall correctly the reaction of the trioxide with water is so exothermic it has two major effects
1) the solubility of the gas in water decreases at higher temperatures. Slowing down the reaction
2) the kinetics of this reaction slow down with increasing temperature.
But basicly the atmosphere is so saturated with vapor you get fog, like on a rainy day. Without physical forces acting on the cloud to force it to
condence, it just hangs around.
Edit: it should be noted that i did not ace that class.
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I wouldn´t think it acing either. Was heating the reason offered?
The reasoning I´ve seen starts as follows:
Sulphuric acid is very refractory compared to its components. The boiling point of azeotropic (98,4%) sulphuric acid is 338 degrees, compared to 100
degrees of water and 45 degrees of SO3.
Therefore if gaseous water and SO3 meet, they condense even at very low concentrations - quite dry air. And they condense into fine mist.
Now, these fine mist particles undergo very much slower thermal movement/Brownian movement/diffusion compared to gaseous SO3 molecules.
SO3 gas could be readily absorbed by concentrated sulphuric acid surface because it freely diffuses there and is caught by the surface. But
on water surface, the water surface gives off water molecules, SO3 molecules meet them midway, condense as mist - and then cannot move the
final inch.
On the other hand, since the droplets are so fine, they are not (entirely) caught by sieves, bubbling through etc. - they get through wherever air
gets through.
As for the heat of hydration, you might dispose of it by employing solid water. But solid water still evaporates just as well as liquid water on
freezing point - which means that sulphur trioxide and solid water still form a fine mist off the surface of water. I suspect (due to the fineness of
mist) that sulphuric acid mist might get through a filter of snow?
Another similarly refractory compound of two gases is ammonium chloride - evaporates at +340 like sulphuric acid, but in this case the components
boil at -85 and -33. Does ammonium chloride also form hard to condense mist?
About smell: SO3 is not common in nature to evolve a response to, but considering how fast it hydrates, does a proton smell? For example,
do HCl, HBr and HI have any difference in smell, or is it just protons that are perceived?
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BerthelotOnCrack
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| Quote: Originally posted by Keras | Yes, especially at the end when you remove the heat source,
so use two washing bottles to avoid any possible suck-back.
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Two bubblers actually prevent suckback ? Can I use this for HNO3 synthesis too ?
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BerthelotOnCrack
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| Quote: Originally posted by phlogiston | | I've never played with SO3, so for what it is worth, but passing an aerosol through a plug of glass wool in a tube is very effective at trapping tiny
aerosol droplets, much more so than wash bottles. |
Good idea, would work like an aquarium bubbler
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BerthelotOnCrack
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| Quote: Originally posted by Rainwater |
sorry, I take safety too serious
First I have to address this. putting it mildly, The dangers of this compound are extreme. |
Thanks a lot for the advice regarding the bubbler, but on that I respectfully disagree
I made 400mL of it and aside from the irritating fumes it just felt like hot sulfuric acid for me. I guess if I breathe all the fumes for 2 mn my
lungs would give up, but I know no toxicity mechanism other than the burning effect.
I am curious of your opinion about NOx gases for instance, that kill you on one breath and about which I am much more paranoid about bubbling when i
distill HNO3
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Rainwater
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My opinion of NOx is about the same.
For both SO3 and NOx, it is not really the toxicity that is concerning but the corrosive effects I worry about. In the case of
inhealing trace amounts of the compounds damage can take days or weeks to show itself, by the time you feel any different, you have forgotten the
moment of exposure and associate it with just being sick.
This can easily lead to repeated exposure, under the presence that "it didnt hurt me last time i smelt it".
At first your lungs undergo abnormal osmosis, to balance the ph, this leads to fluid buildup, diminished lung capacity, and your bodies physical
immune system underperforms. Dont forget your lungs process more environmental material than any other part of your body, it is a complex filter and
liquid/gas interface membrane covering about 100 square meters.
In my profession, inhalation exposure is the cause of over 90% of occupational related illnesses. I take it seriously, not just so my coworkers can
die of old age, but because I hate training the green horns.
In short, if you can smell it, you have a containment breach.
Have a plan in place, preferably a way to instantly neturalize any hazardous compounds.
For example, producing HNO3 under a slight vacuum generated by a water asporator, containing an basic solution such as potassium carbonate.
If accounting for the KNO3 produced, you will quickly see yields around 100% due to the collection of decomposition products normally
vented to atmosphere.
Due to SO3 violent reaction with water, a water asporator would be a risky method. A cold trap, with some kind of particulate filter seams
like a better option.
"You can't do that" - challenge accepted
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Twospoons
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The world is full of people who wish they had heeded safety advice. Don't be one of them.
Helicopter: "helico" -> spiral, "pter" -> with wings
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