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Picric-A
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You know what Sauron? i dot realy care...
sorry for the confusion, i did mention drain cleaner as it is a fairly cheap source of impure conc sulphuric.
if the product is SO2, why use pure, lab grade H2SO4 when drain opener grade works just as well?
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chloric1
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| Quote: | Originally posted by Picric-A
You know what Sauron? i dot realy care...
sorry for the confusion, i did mention drain cleaner as it is a fairly cheap source of impure conc sulphuric.
if the product is SO2, why use pure, lab grade H2SO4 when drain opener grade works just as well? |
Yeh I second that. Reagent or lab grade should be saved for important synthesis not menial task such a SO2 production. If the drain cleaner sulfuric
acid is readily available in your neighborhood then why not use it for SO2, HCl, or drying gases instead of paying hazmat charges.
Doing such does not make one a punk. Actually, your derogatory tone diminishes your professional image in my eyes.
Fellow molecular manipulator
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Klute
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This whole second page is uselessly unconstructive... Could we go back to topic and avoid such worthless bitterness? Please?
..Could a mod clean this up?...
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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ScienceSquirrel
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| Quote: | Originally posted by Sauron
The entire forum is being polluted by the likes of Picric-A and is becoming well nigh uninhabitable.
Why shy away from the truth? That is not courtesy, it's cowardice. |
It is not that bad.
Posters come and go.
I view posting here as a night down the pub.
If you do not get on with someone then ignore them. Move on and chat to someone else.
If someone gets offensive then the mods will tell them to 'drink up' 
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Sauron
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I just requested quotation on lecture bottles, stainless control valves and CGA110 PTFE sealing washers. The control valves are knob type with 1/4"
hose barbs.
I'll report what they quote. Doubtless it is a lot less in USA with S-A Singapore and their agent in Thailand out of the loop.
[Edited on 31-8-2008 by Sauron]
Sic gorgeamus a los subjectatus nunc.
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chloric1
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Ok anywho- Just as a helpfull tip, I heated dry sodium metabisulfite in a test tube a year ago because an Advanced Inorganic Chemistry text had a list
of solids that evolved dry gases. Sodium Metabisufite was recommended for SO2 generation. What I found was part of the sulfite reduced itself and I
got polysulfides and the test tube cracked by being exposed to a strongly alkaline melt. IRC the Bisulfite can exist in more than one configuration
with one of those having a hydrogen bonded directly to the sulfur. I realize that metabisulfite is supposed to be the pyrosulfte but I am considering
partial hydration here. Now my metabisulfite was technical grade and it might of had sulfite as an impurity.
This is just something to consider. One might try mixing metabisulfite with and inert filler, like aluminum oxide. had heat dry. Another thought is
heating sodium metabisulfite with a monobasic phosphate salt and removing moisture.
Fellow molecular manipulator
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Sauron
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A quick look at the stoichiometry of such reactions the other day led me to conclude that they were not the best way to get the most SO2 out of a
given amount of substrate, even without considering unexpected side reactions like yours. Add to that the energy input, and it's just not a good
tradeoff for the sole advantage of not having to dry the SO2 stresm, which is after all, easy.
According to Ullmann's SO2 is somewhat soluble in H2SO2, varying with concentration of the acid, minima being at 85%. I will have to look up this
reference to see how much SO2 gets lost this way. It can't be much as this is almost universally the drying agent of choice for SO2.
Chart is attached. At ordinary temperatures and concentrations of H2SO4, 20-40 g SO2 dissolve in a Kg of acid. SO2 is 64 g/mol. However, as the
solubility drops way off at elevated temperatures, dissolved SO2 can be recovered at the end of a prep run by warming the H2SO4. Adjusting the acid to
85% first would be helpful.
[Edited on 31-8-2008 by Sauron]
[Edited on 31-8-2008 by Sauron]
Sic gorgeamus a los subjectatus nunc.
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panziandi
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looking back over this thread it appears there are two options: sulphites + acid, or reducing the sulphuric acid with copper... you seem to have
dismissed the sulphite method on grounds of not a good trade-off substrate:SO2 and energy input. But sulphites are cheap so you are saving there.
Reducing sulphuric acid with copper is the other suggestion, Cu(s) + 2H2SO4(aq) → CuSO4(aq) + SO2(g) + 2H2O(l), looking at that and seeing as
you want to fill several lecture bottles with SO2 and not being bothered to do the math, I'd say it'll cost a lot more in copper metal (expensive at
the moment) and also cost more in acid consumption and you will have a large volume of copper (II) sulphate to dispose of, because I can't imagine you
wanting to keep a simple lab made chemical like that. Large quantities of copper (II) compounds constitute an environmental hazard and would require
professional disposal, although having the Chao Phraya running cupric-blue may be interesting! 
p.s not to mention the copper method requires hot concentrated H2SO4 certainly not the nicest but you also have energy input there for heating a large
volume of acid.
[Edited on 31-8-2008 by panziandi]
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Sauron
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Come now, I said no such thing. I dismissed only the thermal decomposition of sulfites, not their reaction with moderately dilute H2SO4 (like 50%) to
generate the SO2, preferably dropping the sulfite as saturated soln into the acid. Then, passing the SO2 through a wash bottle or two of conc H2SO4
before passing it into a large Dewar condenser with integral collecting flask (1 liter) and taking that liquid SO2 off to fill lecture bottles 300 ml
at a time.
I accept Klute's bad experience with the copper method, and therefore will not consider it further. The only one who recommended copper was someone I
would not believe if he told me the sun was shining and I had a window open. In the category of BURN BEFORE READING.
[Edited on 31-8-2008 by Sauron]
Sic gorgeamus a los subjectatus nunc.
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panziandi
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Ah I see! Well I certainly think sulphite method is the best. In fact I may try this next week if I get a chance I'm thinking along similar lines of
preparation then drying through H2SO4 (only one wash bottle as the quantities will be minimal) then through a coldfinger and I will attempt to store
the liq SO2 in a glass pressure vial closed with a HiVac teflon tap (basically it looks like a schlenk solvent tube but with a constriction so you can
ampuole off the bottom if required) ... but I have molecular genetics exams this week so chemistry is put on the back burner for now!)
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chloric1
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What about using a drying tube filled with pumice stones moistened with H2SO4. That way less H2SO4 s used and less of your SO2 will be dissolved on
drying.
My unexpected results in regards to the pyrolisis of metabisulfite are really a reminder that what you think should be straight forward in chemistry
is often much more complex and involved. I still might someday try another dry pyrolisis of metabisulfite but I will try different conditions and
additives so that SO2 can come off at lower temperatures to limit side reactions.
Fellow molecular manipulator
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Sauron
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Also, as we will be generating the SO2 from 50% H2SO4, which will gradually become more dilute as the satd sulfite solution is added, significant
amounts of SO2 will remain dissolved in the generator flask and will have to be driven off by heating at the end, just as with the wash bottle acid.
The amounts of dissolved SO2 in the wash bottles are relatively trivial. The amount of dissolved SO2 in the generator is not. It is well worthwhile to
replenish the acid in the generator to maintain it at the maximum practical concentration (probably more than 50% but not more than 85%) to minimize
this. Alternatively, or additionally, the generator can be externally heated throughout the run to a temperature between ambient and 100 C, consistent
with minimizing production of water vapor. 60-70 C perhaps. Keeping the temperature up and the acid concentration up will keep SO2 solubility at or
below 20 g/Kg acid.
Panziandi, you just restated Murphy's Law without attribution.
Sauron's Law: Murphy was an optimist.
Now let's talk about practical matters so far unmentioned.
Let's suppose we want to use 1 Kg Na2SO3 to generate SO2.
MW 126, solubility 230 g/L water ambient. Very roughly therefore 4.3 L and concentration a little under 2 M.
Stoichiometry:
Na2SO3 + H2SO4 -> SO2 + Na2SO4 + H2O
So that Kg of sulfite gets me 8 mols SO2, slightly over 500 g
For the reaction I need 8 mols H2SO4 = 800 g 100%, 640 g 95% acid, d 1.85 so how many mls?
Dissolve in 800 ml water. (USUAL CAUTIONS).
So we are going to react 4.3 L satd sodium sulfite soln by dripping it into 50% H2SO4 about 1300 ml so at end, reactants volume is 5.6 L. If we let
this alone the resulting H2SO4 concentration at the end will be low (actually if we stick to stoichiometry, nil, because all the acid will have
reacted and now will be sodium sulfate.) So clearly we need to add conc H2SO4 slowly through the reaction or intermittently anyway, in amounts
calculated to keep the acid concentration where we want it (50%) and preferably at the end we want to hike that to 85% and drive off all dissolved
SO2. External heating may not be required because the heat of dissolution of the conc H2SO4 will be considerable.
The product 8 mols SO2 is exactly what I want to load into a standard LB.
[Edited on 1-9-2008 by Sauron]
[Edited on 1-9-2008 by Sauron]
Sic gorgeamus a los subjectatus nunc.
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Klute
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I like the idea of adding concentrated acid to keep the concentration up. That way you can start with a minimal amount of water in the acid at the
beggining.
P2O5 might be a good option for drying the acid too, perhaps after a small H2SO4 washbottle to avoid consuming the P2O5 quickly: very little SO2 in
the wash bottle, efficient drying and no absorption from the P2O5. The used P2O5 coudl then be neutralized with water to form conc. H3PO4 with
possible traces of SO2, which could be used to generated more SO2 from bisulfite, instead of H2SO4?
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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panziandi
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By adding H2SO4 towards the end would generate enough heat perhaps to avoid requiring external heat?
I expect you will actually get NaHSO4 and not Na2SO4 so perhaps just bare that in mind and adjust accordingly if you are aiming for 500g SO2 adjust to
generate an excess to ensure you'll have plenty.
I would use as high a % of H2SO4 as possible, in fact I'd prefer to use 98% neat onto the sulphite and in a larger flask (so as foaming is not an
issue) just to keep waste minimal, and flash the larger flask with air or etc at the end to ensure all SO2 is flushed into the Dewar and collected.
With an excess you'll be sure to have enough to fill a LB and any not bottled can be discarded or used.
I like the idea of H2SO4 on pumice actually - never has occured to me as a drying wash agent more of a dessicator dryer but would work well I'd
imagine! P2O5 maybe overkill? I imagine two washes with H2SO4 would be plenty but mild steel does rust so perhaps P2O5 to remove trace moisture before
bottling. But will you simply be decanting the liq SO2 into a LB or will u vac out the LB and attach via adapter to the base of Dewar etc? if you are
decanting forget P2O5 as the SO2 could attract trace moisture whilst decanting. I'd only use P2O5 if the system was closed. Does that make sense or
too waffly?
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Sauron
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I have two large dewar condensers as cold traps with integral 1 L receiving flasks and bottom drains. I was planning to run a tuve from that drain
into the dry, unclosed, chilled LB and after filling, quickly screw on the control valve with ptfe sealing washer - valve open, then once the valve is
well started in the thread, close the valve and tighten thread (it takes a wrench.) Then the closed LB can come out of the cold bath and slowly come
to ambient, I can use some soap soln to check for leaks.
If reacting reagent conc H2SO4 and dry sulfite (anhydrous as it usually is these days) the SO2 would not need drying. I agree P2O5 is overkill. Woelen
just quoted Vanino's book saying CaCl2 is a good drying agent for SO2, as well. (See parallel thread.) P2O5 is flocculent and a mess. I only handle it
in a glove box, no breezes and no moisture. The bottle once opened stays in the dry box. Otherwise it turns into expensive H3PO4.
------------------
OK so we worked out the stoichiometry for sodium sulfite.
Na2SO3 + H2SO4 -> SO2 + Na2SO4 + H2O
How about sodium bisulfite?
NaHSO3 + H2SO4 -> SO2 + NaHSO4 + H2O
MW 104 so almost 10 mols/Kg and solubility 300 g/L water at ambient. Conclusion: sodium bisulfite is better choice than sodium sulfite, and should
produce >600 g SO2 per Kg.
Next up, sodium meta-bisulfite.
Na2S2O5 + H2SO4 -> 2 SO2 + Na2SO4 + H2O
2 mols SO2 per mol metabisulfite but the MW is almost 2X that of bisulfite so yield SO2 per Kg is no better (c. 600 g/Kg).
Solubility:
I have not been able to find this quantified, most describe this salt as "freely soluble" but surely there's a limiting amount?
[Edited on 1-9-2008 by Sauron]
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panziandi
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OK! Well bisulphite looks the best candidate then. All the sulphites dissolve in cold water BUT I only speak from the experience of Na2S2O5 as it's
cheapest for me to get, it dissolves in cold water but you get small lumps which remain stubborn and need stirring or breaking up to dissolve. On the
solubility front:
Na2SO3:
Solubility in water:
23 g/100 mL (20°C)
Solubility in other solvents:
soluble in glycerol
practically insoluble in alcohol
NaHSO3:
Solubility in water:
300 g/l water
Solubility in other solvents:
SOLUBLE IN 2 PARTS BOILING WATER
SOLUBLE IN ABOUT 70 PARTS ALCOHOL
Na2S2O5:
Solubility in water:
540 g/L (20°C)
Solubility in other solvents:
Freely soluble in glycerol
Slightly soluble in alcohol
So maybe metabisulphite would be better due to its higher solubility?
[Edited on 1-9-2008 by panziandi]
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chloric1
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The metabisulfite is indeed highly soluble. On two occassions at least I remember making a 40% solution. Once
to clean up permanganate stains and secondly for react with ketones.
Never tried 98% H2SO4 directly on a sulfite. Concentrated sulfuric acid is a strong oxidizer. You might get a mess. It might oxidize your sulfite
to dithionate while itself be reduced to SO2. You would still get SO2 but at what yields? What would be the stoichometry of this redox? I feel there
would be several species involved here much like 98% added to potassium iodide crystals.
Fellow molecular manipulator
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Sauron
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Oleum is a strong oxidizer. Conc H2SO4 is a strong dehydrating agent but not a strong oxidizer. The Vanino procedure cited by woelen employs conc
H2SO4 with calcium sulfite admixed with calcium sulfate as diluent.
Now that's his procedure not mine but while I think acid somewhere in the 50% to 85% range is best, I do not think any oddball redox reactions are
going to happen with conc H2SO4 (95-98%).
Sounds like metabisulfite gets the mix down to <2 L fot the metanisulfite soln containing >5 mols (1 Kg, MW 190) and only 750-800 ml or so of
50% H2SO4. So very compact, well under 3 L total and a 5L flask is fine. Very nice as the metabisulfite is also cheapest of the three salts. I'd
still add the solution to the acid to minimize foaming.
[Edited on 2-9-2008 by Sauron]
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ScienceSquirrel
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I don't suppose you have any home brew shops out in Thailand?
The stuff is used as a sterilant / final wash for equipment so a few pounds buys 500g in the UK.
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Klute
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I'd think it's even cheaper as a technical industrial product, 2E/kilo... Hard to get lower.
Sauron, in order to insure your SO2 satys well anhydrous, you ould add a little (non-agressive) dessicant in the receiver? CaCl2?) Do you plan on
distilling the SO2 from the first Dewar to the LB? Or just transfering?
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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ScienceSquirrel
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| Quote: | Originally posted by Klute
I'd think it's even cheaper as a technical industrial product, 2E/kilo... Hard to get lower.
Sauron, in order to insure your SO2 satys well anhydrous, you ould add a little (non-agressive) dessicant in the receiver? CaCl2?) Do you plan on
distilling the SO2 from the first Dewar to the LB? Or just transfering? |
My local home brew shop sells food grade 97% for a third of the price of Sigma Aldrich's 97%. It does not come in a nice Sigma Aldrich tub but it
probably comes from the same factory in China!
Of course if Sauron is buying a 25kg drum he is going to get it cheaper...
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Sauron
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No home brew shops (illegal but popular.)
Anyway plenty industrial chemical suppliers and lab chemical suppliers.
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Klute
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That's what's called selling a name...... 
\"You can battle with a demon, you can embrace a demon; what the hell can you do with a fucking spiritual computer?\"
-Alice Parr
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Sauron
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I doubt I will buy a 25 Kg drum, at least initially. The 2.5 Kg packing looks attractive. Food grade is really overkill for this application.
The dewar will be terminated with a drying tube. I was just going to drain the receiver into the LB, distilling the SO2 ought not to be necessary,
IMO.
Remember this is short term storage only. The SO2 will get turned into SOCl2 or SO2Cl2.
Sic gorgeamus a los subjectatus nunc.
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Sauron
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I have now been quoted c.$150 for a mild steel 440 ml lecture bottle. I'd guess this is $75 in USA, the rest being shipping, duty, VAT and middleman
greed.
A 900 ml Sure/Pac cylinder is half that price. But rated 240 psi vs 1800 psi for the LB.
The real shock however is the quote on the control valve. A simple SS316 control valve, nothing fancy, not a regulator, no gauges, is $275. So
probably $140 in USA.
I think this is ridiculous. An empty LB and a simple control valve >$400??
NUTS!
I'm not going to put up with this. I will generate the SO2, pass it through drying train and use it directly. If I need intermediate storage I will
put it in a dewar flask.
Sic gorgeamus a los subjectatus nunc.
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