woelen
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Ampouling of SO2Cl2
I have some thionyl chloride and sulphuryl chloride. I already ampouled some of the thionyl chloride in glass ampoules, with melted down opening, such
that it is hermetically closed. This ampouling makes the storage of the thionyl chloride much more pleasant. No more nasty corrosive fumes.
I want to do the same with sulphuryl chloride (ampouling in 20 ml ampoules), but I read that this compound slowly decomposes to Cl2 and SO2. Is it
safe to ampoule this stuff, or can I expect the ampoules to explode after some time of storage, due to pressure buildup of the Cl2 and SO2?
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not_important
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I believe sulphuryl chloride is fairly stable up to its boiling point. The reaction is reversible, so even slow decomposition happens it would seem
that it would be self limiting.
So search turned up homework questions involving the decomposition, but all were gas phase elevated temperatures. Even the the pressure was less than
a bar. Didn't find anything regarding the commercial packaging, which would face similar problems if they exist.
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len1
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My search yielded the same exam questions not important must have been talking about but we can use them
Kc at 50C is 0.12 = pSO2 x pCl2
so pT=2 sqrt(0.12) ~ 0.7atm, thats the actual pressure on the glass, which is substantial, you need at least 3mm I would say If the SO2Cl2 contains
impurities then this Kc is not applicable. If it containes water even more so.
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12AX7
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Any idea what pressure your ampoules can tolerate at room temperature? If it does decompose, it would be useful to know how much danger there is.
Tim
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microcosmicus
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@12AX7
An easy way to find out would be to seal some water in an ampoule and see how
hot one can heat it before it breaks, then use the vapor pressure of water to convert
that temperature into a pressure. To be sure, the failure will be rather catastrophic,
assuming the form of a steam explosion (there is a reason that old-time sealed
tubes were referred to as "bombs"), so you should take some reasonable safety
precaution, such as enclosing the water-filled ampoule in an iron pipe before heating it.
In particular, at 150C, the pressure will be 4.7 atm, so if one of your ampoules filled
with water can withstand being heated to that temperature without exploding, it
should be up to the task of storing contents at twice atmospheric pressure.. For instance, seal an
ampoule, enclose it in a cast-iron nipple with endcaps (leaving somewhere for the
steam to escape, maybe some holes in the pipe), then put it in the oven at 150C
and see what happens..
If you are worried about doing the test at room temperature, you could
try butane instead since its vapor pressure is 2 atm at 20C. Since it
boils at -0.5C, you could put the ampoule in an ice bath, fill it, seal it,
then let it warm up inside a plastic jar and see if it breaks. If not, you might
jack up the temperature to 50C, at which point the vapor pressure is 5 atm.
@len1 and woelen:
Returning to the original question, according to von Richter,
(Text-book of Inorganic Chemistry, p. 195)
sulphuryl chloride boils at 69.1C and only decomposes into
chlorine and sulphur dioxide significantly above 130C.
Thus it sounds like there should not be a problem with
pressure if you are going to store your ampoule at room temperature.
Also, shouldn't that equilibrium equation read
0.12 pSOCl2 = pSO2 x pCl2? Then, we would have pSO2 = pCl2 =
0.36 sqrt (pSOCl2). Since it is in equilibrium with a liquid phase,. pSO2Cl2
equals its vapor pressure. Hence, pT = SO2 + pCl2 + pSOCl2 =
pSOCl2 + 0.72 sqrt pSOCL2. Since 50C is less than the boiling point
of 70C. this vapor pressure will be less than 1 atm, so the total
pressure will be less than 1.7 atm. That is at 50C, but if the ampoule
is going to be stored room temperature, the vapor pressure should be
significantly smaller --- according to the MSDS. the vapor pressure
at 20C is 0.12 atm. Using your equilibrium constant, this gives
pT = 0.4 atm. Hence, it sounds like ampouling sulphuryl chloride
should not be all that differrent than ampouling other liquids
with low boiling points. (also, the MSDS makes no reference to
problems due to pressure, only corrosivity.)
sciencelab.com/xMSDS-Sulfuryl_chloride-9925154
By the way, is the problem set the two of you are referring to the following:
web.clark.edu/nfattaleh/classes/132/Sp05/Homework/132Sp05Ch16Hwk.pdf
[Edited on 13-3-2008 by microcosmicus]
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len1
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If a material participating in a gaseous reaction is a liquid at the temperature at which the equilibrium constant for the reaction is quoted, then
its concentration does not appear in the expression for the equilibrium constant. Reason being that its concentration is then a constant which is
incorporated into Kc. So expression for Kc=0.12 is such. Only point where substantial variation of such a Kc can be expected is near a phase
transition.
Similar reasoning applies to solubility products. Ksp=[Ca][SO4] because the concentration of [CaSO4] is a constant.
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woelen
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I will try one ampoule of 10 ml (I have them in 10 and 20 ml formats) with a small amount of SO2Cl2 in it (1 ml or so) and then I'll drop it in a
beaker, filled with water of 50 C and leave it in there for a few days (every now and then replacing the cold water with hot water again) and put the
beaker somewhere outside. If the ampoule does not crack, then I consider it safe to ampoule all of my SO2Cl2. The ampoules have thin glass (< 1
mm).
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len1
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Good idea. It has also occured to me that if before sealing the SO2Cl2 is let to attain equilibrium, then SO2 and Cl2 being heavier than air will be
near their equilibrium partial pressures above the liquid - with air making up the difference of about 0.3atm to atmospheric pressure. If one seals
such an ampule, it should store indefinitely with the pressure inside varying from 1atm only as the Kc and partial pressure of SO2Cl2 changes with
temperature.
[Edited on 13-3-2008 by len1]
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microcosmicus
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The reason I thought of the equilibrium constant as I did was that I thought in
terms of the reactions
SO2Cl2(l) <---> SO2Cl2(g)
SO2Cl2(g) <---> SO2(g) + Cl2(g)
rather than an overall reaction. The analogy with CaSO4 doesn't work that
well because it is ionic but SO2Cl2 is covalent --- whilst there will be
molecules of SO2Cl2 in the gaseous phase which can condense into the liquid,
there are not molecules of CaSO4 in the solution, but the Ca++ and SO4--
ions individually. The vapor pressure puts an upper bound on pSO2Cl2,
but the same equilibrium would hold for a gas whose pressure was less
than the vapor pressure.
However, from the standpoint of thermodynamics, either convention is fine, so now
I understand why pSO2 = pCl2 = 0.35 atm at 50C. At room temperature,
this pressure will be less.
[Edited on 13-3-2008 by microcosmicus]
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len1
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No-no, there are molecules of CaSO4 in solution, the equilibrium has to be a local thing. The equilbria constant the way they are constructed,
althoug a lot of chemists dont understand this, actually reflect the probability of species meeting and interacting - Ksp for CaSO4 refelcts the
concentrations of Ca and SO4 needed so the amount of combination = amount of recombination. If there was not an amount of CaSO4 in solution there
will only be combination and eventually all the ions will end up as solid CaSO4 at the bottom of the beaker. The solid percipitate of CaSO4 at the
bottom maintains a constant concetration of CaSO4 in solution by
CaSO4(s) -> CaSO4(aq)
In this way the situation is totally analogous to liquid SO2CL2 maintaing a constant concentration of gaseous SO2Cl2 above it.
I wish my chem lecturer understood this when I was a student, instead of trying to bullshit his was out of it - saying the precipitate is a pure
material and therefore its concentration is constant, something they write in books. Its only after years as a physicist that I understood this
properly. The correct answer is that the solution is saturates with CaSO4, which means its concetration is constant.
[Edited on 14-3-2008 by len1]
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microcosmicus
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You're right, there will be a minute number of CaSO4 molecules in a solution --- I think of them
a "ionic nanocrystals" which happen to consist of 2 ions, also there will be nanocrystals
of several ions forming and unforming. Likewise, in the gas, there will be "microdroplets"
forming and unforming as one or more SO2Cl2 molecules get attracted to each other.
What I had in mind was that this quantity is negligible so that we can safely ignore the CaSO4
molecules whilst the concentration of undissociated SO2Cl2 is too large to ignore. This is
one sense in which I was wary of the analogy. The other was that, whilst SO2 and Cl2 first
join to form SOCl2 molecules which attract each other by van der Waals forces to make a
liquid, there are no individual CaSO4 molecules to be found in the crystal.
Alright, all this discussion has gotten me eager to burn a sulphur candle, whip up some Cl2
as per your practical guide, shine a light on the gas mixture, then condense the product
and store it in an ampoule, then later maybe measure these darned equilibrium constants
we have been talking about 
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len1
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Yeah but negligible compared to what? [CaSO4] is not negligible compared to the amount of Ca2+ or SO4- in solution which is also small, in fact its
just right to maintain equilibrium, one cant neglect it. The same occurs for SO2Cl2 which is just right to maintain equilibrium with SO2 and Cl2,
except that we are talking about a liquid/gas instead of a solid/liquid heterogeneous equilibrium - but both gases and liquids are fluids. The main
point is that [CaSO4] and pSO2Cl2 are both constants.
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microcosmicus
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Sure, one cannot neglect the tiny quantities in order to understand the equilibrium.
However, exactly because they are so small and unform rather quickly, one can
safely ignore the partial osmotic pressure of undissociated salt as negligible in
comparison to the pressure of the ions. By contrast, the contribution of undissociated
SO2Cl2 to the pressure on woelen's ampoule is not negligible.
A better analogy to SO2Cl2 would be HgCl2 since, when dissolved
in water, most of the molecules are undissociated. Like SO2Cl2,
the chlorines of HgCl2 are covalently bonded and, in the solid phase,
one only finds the covalently bonded molecules, not the ions.
[Edited on 14-3-2008 by microcosmicus]
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Nicodem
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The reagent bottle of sulfuryl chloride at my workplace looks like 20 years old and it has never blown up. However, it is in a Fluka bottle for which
the producer claims it can contain >1 bar overpressure. Yet, I think even a thin ampoule would do since, like it was already said, the
SO2Cl2<=>SO2+Cl2 reaction is an equilibrium and even normal ampoules were in the old times used as sealed vessels for heating reactions above
the boiling point (though they have to be properly small!).
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woelen
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I found this old thread by accident while searching something else. Eight years ago I decided to take the risk and ampoule all of my SO2Cl2 and store
the ampoules in a big plastic container, filled with old paper in which the ampoules are wrapped.
I can tell that it is safe to ampoule SO2Cl2. I have these ampoules lying around for 8 years now and they have felt cycling of cold winter (-10 C) and
hot summer (30+ C) and none of them cracked or leaked. I store them in my garage, not in the house, but I think that safety measure is not really
necessary.
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aga
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8 years !
What a wonderful follow-up.
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