Sciencemadness Discussion Board - SO3 from pyrosulfate and H2SO4

SO3 from pyrosulfate and H2SO4

ave369 - 10-9-2015 at 09:57

In my topic "A safer procedure of mineral acid synthesis", one of the users suggested that heating a pyrosulfate with conc. sulfuric acid can yield SO3 at a much lower temperature then with pyrolysis.

Today, I tried to check it and boiled a small amount of sodium pyrosulfate with 80% sulfuric acid (I didn't have any stronger at hand). I think I managed to produce some SO3, but only some. Here is the detailed description of my experiment.

A retort was mounted on a tripod. The retort's nose was put into a small flask containing 80% sulfuric acid, to make any evolving gas bubble through this acid. In the retort was pyrosulfate and more 80% sulfuric acid. I started to heat the retort with a gas burner.

Early on, some kind of gas started to bubble from the retort's nose, and make white smoke on contact with outside air. I think it was SO3. Soon, the evolution of smoke from the receiving flask ceased, but that same smoke started to form inside the retort. Soon, the inside of the retort turned milky white. This milky white fog continued to bubble through the receiving flask, and made no further smoke with outside air. I think it was H2SO4: I've driven all possible SO3 from the liquid inside the retort, all of the other SO3 combined with water present in the 80% sulfuric and start to form H2SO4 mist. If the acid in the retort was stronger to begin with, I would get more SO3.

Am I right?

gdflp - 10-9-2015 at 10:02

Garagechemist has experimented with this method quite a bit, you can see his results here. He mentions that :

Quote:

Any water is very detrimental to the yield

So yes, your explanation sounds quite plausible. It might be worth trying again, but instead use the acid you produced from this experiment as the catalytic acid, instead of the 80% acid.

ave369 - 10-9-2015 at 10:13

The tiny amount of SO3 I've made only managed to increase the concentration by a percent or two, if even that.

clearly_not_atara - 10-9-2015 at 10:58

The way to concentrate H2SO4 is to simply boil it long enough to remove all of the water. According to Wikipedia (shut up, you use it too) 98.3% concentration forms at the boiling point (due to partial thermal decomposition). Sulfuric acid boils at 337º C, so there won't be an azeotrope.

Magpie - 10-9-2015 at 13:22

This may be useful:

http://www.sciencemadness.org/talk/viewthread.php?tid=10332

ave369 - 10-9-2015 at 13:33

Quote: Originally posted by clearly_not_atara

The way to concentrate H2SO4 is to simply boil it long enough to remove all of the water. According to Wikipedia (shut up, you use it too) 98.3% concentration forms at the boiling point (due to partial thermal decomposition). Sulfuric acid boils at 337º C, so there won't be an azeotrope.

337º C is where the azeotrope boils. The azeotrope is 98,3%. But that's not what I mean.

What I mean is:

Have you really tried to boil battery acid? I mean, ever? Are you aware of the "white mist threshold" at 70-80%? Have you ever caught a whiff of that mist, which, even rarefied to the point of invisibility, will burn your lungs? Have you ever lost a piece of glassware to thermal cracking, because the difference between 337º C and whatever's outside is too large? Have you ever been frustrated at the fact that the goddamn white mist carries away as much acid as water, and the concentration does not want to grow beyond 90%, despite the World War I level gas attack around your boiling vessel?

I've boiled more battery acid down to the 80% threshold than brewed coffee in the morning. But very rarely I was fortunate enough to reach 90%. THAT's why I'm looking for other ways to boost it up to 98% and beyond.

[Edited on 10-9-2015 by ave369]

DISTILL!!!!DISTILL!!!!DISTILL!!!!

Deathunter88 - 10-9-2015 at 20:00

No doubt you were excited when you distilled your first liquid, either water or ethanol. (or something else) Now apply that skill to sulfuric acid.
To get 98% sulfuric acid, distilling is both cost efficient, easy, safe(relatively) and feasible. Take your 80% acid and add 300ml to a 500ml RBF. Use a ring stand and metal clamp to hold it 5mm above the top of the hotplate/radiator (air-bath), then attach a a 300mm vigreux column as an air condenser. I think you can also use a actual air condenser or a Liebig condenser without water. Turn up the hotplate/radiator to full power and wait for the acid to boil. This is when you have to experiment with the power setting of the hotplate/radiator. Not high enough and you get very little distillate, too high and it overpowers the condenser and you get the WW1 fumes from the sulfuric acid vapors touching the air. Let it distill into a beaker and every 2 minutes hold a paper towel above the receiving beaker and let a drop of the distillate fall on it. At the start nothing will happen. But after a while the distillate will burn through the tissue. At this point change the receiving beaker and start collect acid at around 80%. Let this distill until you get a further 50-100ml of acid.(After a while you will get good at determining the exact amount here) Then change the beaker again to get your 98% acid. I distilled 1500ml of 96% drain cleaner in 3 hours. Got 1250ml of 98% acid.

[Edited on 11-9-2015 by Deathunter88]

ave369 - 10-9-2015 at 23:48

I would be very glad to try this, but I'm a bit equipment-deficient. My distillation setup is a retort, and I'm unsure if it can survive distillation of sulfuric acid without thermal cracking. I don't have another retort, after all.

Next time I order something on AntrazoXrom, however, I'll try to add some simple modern distillation setup to the list.

[Edited on 11-9-2015 by ave369]

AJKOER - 11-9-2015 at 05:50

There may be safer ways to obtain SO3, at least in very small amounts.

For example, start with a naturally occurring process, the formation of acid rain. To quote Wikipedia ( https://en.m.wikipedia.org/wiki/Acid_Rain ) on acid rain:

"In the gas phase sulfur dioxide is oxidized by reaction with the hydroxyl radical via an intermolecular reaction:[5]

SO2 + OH· → HOSO2·

which is followed by:

HOSO2· + O2 → HO2· + SO3

In the presence of water, sulfur trioxide (SO3) is converted rapidly to sulfuric acid:

SO3 (g) + H2O (l) → H2SO4 (aq)"

Now, there is a convenient path to the hydroxyl radical via the action of either sunlight or pulse radiation on N2O in the presence of water, for example, as I have documented previously on SM. Or even possibly, a damp nitrate or nitrite for a path to hydroxyl radical generation via the aqueous photochemical reaction:

NO3-(aq) + hv → •NO2 + •OH (see http://pubs.acs.org/doi/abs/10.1021/ja073609 )

Also, NO2-(aq) + hv → NO + •OH (actually, more complex, see, for example, discussion on page 2 at
https://www.google.com/url?url=http://scholar.google.com/sch... )

An alternate path in place of the hydroxyl radical, is via the contact of SO2 with fumes of decomposing strong HNO2 as, per Atomistry (link: http://nitrogen.atomistry.com/nitrous_acid.html ) on the oxidizing ability of Nitrous acid, to quote:

"Stannous chloride is converted into stannic chloride, sulphuretted hydrogen into sulphur, sulphur dioxide into sulphur trioxide. Iodine is liberated from potassium iodide"

together with the comments:

"The decomposition of a cold dilute solution follows the reaction

3HNO2 ⇔ HNO3 + 2NO + H2O,

whereas stronger solutions at higher temperatures decompose according to the equations

2HNO2 ⇔ N2O3 + H2O ⇔ NO + NO2 + H2O"

And, the following gas phase reaction:

SO2 (g) + NO2 (g) = SO3 (g) + NO (g) (see discussion on the 5th page at this link address https://www.google.com/url?sa=t&source=web&rct=j&... )

And:

N2O3 + SO2 = N2O4 + SO3

Source, see for example, https://books.google.com/books?id=EvsbBQAAQBAJ&pg=PA209&... which also notes:

.HO2 + .NO → .NO2 + OH· (for another source, see http://www2.nau.edu/~doetqp-p/courses/env440/env440_2/lectur... )

where the HO2 radical formed from HOSO2· + O2 → HO2· + SO3 above, can help address the presence of Nitric oxide.

[Edited on 11-9-2015 by AJKOER]

ave369 - 11-9-2015 at 08:42

Instead of unnecessary fiddling to make nitrous acid, I can just toss a copper coin into nitric acid and direct the resulting foxtail into a stream of SO2. But I doubt that this will be useful in producing highly concentrated sulfuric acid: the industrialized version of this process produced chamber grade acid, which was weaker than the result of my boiling experiments.

But so far, making metric craptons of pyrosulfate and distilling it with conc. sulfuric appears to be the most promising method to me.

[Edited on 11-9-2015 by ave369]

CuReUS - 11-9-2015 at 08:54

recently I learnt that SO₃ can be made by heating ferric sulphate

ave369 - 11-9-2015 at 09:24

Quote: Originally posted by CuReUS

recently I learnt that SO₃ can be made by heating ferric sulphate

I know that. But by dissolving nails in battery acid, you get ferrous sulfate. Where do I get ferric sulfate?

Pumukli - 11-9-2015 at 10:25

Ferric sulfate is used in my vicinity in waste water treatment plants to make Fe-hydroxy-flakes to help in the water clearing process. They use it by the ton, so it should be a relativley cheap (and perhaps available) industrial chemical, even where you live. I don't know your circumstances but if you have friends around such plants it would worth asking.

On the other hand when I opened some old ferrous sulfate bags I always found it partly oxidized (yellow). So O₂ from air surely can do the oxidation, speed is in question though. :-)

Maybe a ferrous sulfate solution, an aquarium pump, bubble stone, sort of heating setup, and maybe a catalitic ammount of something (redox catalysator, Cu²⁺, Mn²⁺, etc) in an acidic solution would do the trick.

I've never tried such a reaction just thinking aloud. And assuming that getting ferrous sulfate is not nearly as problematic as getting ferric of course.

[Edited on 11-9-2015 by Pumukli]

ave369 - 11-9-2015 at 12:34

I live in a little town where all industry died out back in the days of the fall of the Soviet Union. Thank you very much, capitalism. Ferrous sulfate and copper sulfate, they are easy to get because they are used in agriculture.

AJKOER - 11-9-2015 at 13:32

Quote: Originally posted by ave369

Instead of unnecessary fiddling to make nitrous acid, I can just toss a copper coin into nitric acid and direct the resulting foxtail into a stream of SO2. But I doubt that this will be useful in producing highly concentrated sulfuric acid: the industrialized version of this process produced chamber grade acid, which was weaker than the result of my boiling experiments.

But so far, making metric craptons of pyrosulfate and distilling it with conc. sulfuric appears to be the most promising method to me.

[Edited on 11-9-2015 by ave369]

Please note, my citations of SO3 formation is particular to a gas (not aqueous) phase reaction at low pH.

So, heating a dry nitrate salt (as a path to NO2 or N2O3) in the presence of SO2 is one possible route to test (akin to the Lead Chamber process absence water). Adding CaSO3 to nearly boiling strong HNO2 (which I would also test in strong sunlight as the Nitrous acid can produce NO and OH radicals) may form a mist containing SO3, or just some aqueous H2SO4.

My favorite (which I may get around to testing) is the simply safe photolysis of water vapor, SO2, N2O and O2 in a closed vessel allowing uv exposure and see if any solid SO3 is created.

Another experiment (which I plan on performing) is saturating a solution with some SO2 and lots of N2O, and then treating the mix in the presence of added O2 to photolysis or pulse radiation from a microwave. The expected product could include the aqueous HSO5- anion, or H2SO5 (assuming it has not otherwised reacted, see https://books.google.com/books?id=vVvrCAAAQBAJ&pg=PA143&... ), which can be used to create persulfate salts. Reference, see discussion on the sulfite radical on page 6 and the middle of page 7 at https://www.google.com/url?sa=t&source=web&rct=j&... and also http://pubs.acs.org/doi/abs/10.1021/jp011255h .

[Edited on 11-9-2015 by AJKOER]

CuReUS - 13-9-2015 at 01:25

Quote: Originally posted by ave369

I know that. But by dissolving nails in battery acid, you get ferrous sulfate. Where do I get ferric sulfate?

like everything in life,you will get it if you search for it

like pumukli said,it is used in waste water treatment.It is also used as a pickling bath for Al and steel.

Quote: Originally posted by Pumukli

Maybe a ferrous sulfate solution, an aquarium pump, bubble stone, sort of heating setup, and maybe a catalitic ammount of something (redox catalysator, Cu²⁺, Mn²⁺, etc) in an acidic solution would do the trick.

the wiki method mentions heating an acidic solution of FeSO₄ with an oxidising agent to get ferric sulphate.
you could also make it using electrochemistry
https://books.google.co.in/books?id=vSsBCAAAQBAJ&pg=PA13... (pg-137,138)

[Edited on 13-9-2015 by CuReUS]

Marvin - 13-9-2015 at 06:17

Len has some cool stuff about Oleum from Sodium Pyrosulphate in his book.

Starting with distilling Sodium Hydrogen Sulphate, at the point it reaches 500C 90% of the water has left but only 10% of the potential SO₃. So starting to collect at 500C produces a 90% yield of Oleum containing 86% free SO₃.

Starting to collect at 580C produces a 65% yeild of Oleum containing 94% free SO₃. Almost pure SO₃ coming off.

His end point is 820C which leaves the neutral sulphate as residue and he used a Quartz distilling flask with side arm. This is disappointing. Prior to this and in previous discussions I'd rather hoped the reaction would be over by 500C and it really hasn't even started.

A rather cool and whip smart observation he makes is that adding neutral Sodium Sulphate to Sulphuric Acid effectively breaks the Water - Sulphur Trioxide azeotrope.

I'm still looking for evidence that cooling the acid mixture precipitates neutral Sodium Sulphate, which might potentially lead to much lower temperatures for production. Realistically even precipitating the Acid Sulphate would result in a concentration. I can't currently try the reaction myself.

ave369 - 13-9-2015 at 14:25

I've found a very interesting tidbit from a list of long-obsolete industrial technologies:

Na2S2O7+MgSO4=Na2Mg(SO4)2+SO3

An entry in an early XX century encyclopedia claims that adding magnesium sulfate to sodium pyrosulfate significantly lowers the temperature of its decomposition.

Anybody knows more about this reaction? I've got a big fat jar of magnesium sulfate and will probably try this ASAP, but it would be good if anyone knows any hidden problems.

[Edited on 13-9-2015 by ave369]

AJKOER - 14-9-2015 at 05:13

Quote: Originally posted by ave369

In my topic "A safer procedure of mineral acid synthesis", one of the users suggested that heating a pyrosulfate with conc. sulfuric acid can yield SO3 at a much lower temperature then with pyrolysis.

Today, I tried to check it and boiled a small amount of sodium pyrosulfate with 80% sulfuric acid (I didn't have any stronger at hand). I think I managed to produce some SO3, but only some. Here is the detailed description of my experiment....

Actually, just discovered that heating Ammonium sulfate can, at one stage in its thermal decomposition, forms the pyrosulfate. Here is a reference: http://onlinelibrary.wiley.com/doi/10.1002/jctb.5010200408/a... . To quote the abstract:

"The thermal decomposition and vaporisation of ammonium sulphate, (NH4)2SO4, is shown to take place via two distinct sets of reactions. In the first, ammonium pyrosulphate, (NH4)2S2O7, is the primary condensed phase product: 2(NH4)2SO4 =(NH4)2S2O7+2NH3+H2O

The second stage concerns the decomposition of the pyrosulphate. Ammonia, sulphur dioxide, nitrogen and water are the major products, the dominant reaction being 3(NH4)2S2O7 =2NH3+6SO2+2N2+9H2O"

Apparently, the formation of the pyrosulfate has been employed as a basis for a US patent 2,004,023,441to manufacture (or, possibly concentrate) Sulfuric acid. See, "Chemical and thermal decomposition of ammonium sulphate into ammonia and sulphuric acid", US 20040234441 A1. Here is the abstract:

"A method is described for manufacturing ammonia and sulphuric acid by decomposition of ammonium sulphate by a chemical and thermal handling process The method concerns mixing of ammonium sulphate with concentrated sulphuric acid, where the mixture is heated to a temperature above 235° C., whereby ammonium sulphate is melted, and thereafter by heating the mixture to above 280° C. but below the boiling point of concentrated sulphuric acid. It is hereby achieved to decompose ammonium sulphate into ammonia gas and sulphuric acid liquid simultaneously, where the produced sulphuric acid during the reaction time is mixed ideally with the concentrated sulphuric acid from the initial basis mixture."

IrC - 14-9-2015 at 05:33

Simpler to just post the patent so the process can be studied, plus a few more that appear to be useful.

Attachment: US20040234441A1.pdf (18kB)
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Attachment: US20080056983A1.pdf (159kB)
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Also informative:

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AJKOER - 14-9-2015 at 07:18

Here is a reference I just noticed, in my cited patent, that states the formation of oleum, in particular, requires pyrolyzing in the presence of added O2, to quote:

"U.S. Pat. No. 4,490,347 (Gelblum) describes the production of oleum (sulfuric acid containing sulfur trioxide) from a mixture of ammonium bisulfate and sulfuric acid, by pyrolyzing the mixture in the presence of oxygen. The production of oleum is not desirable in the recovery of ammonia and sulfuric acid from ammonium sulfate."

where some sources claim the formation of ammonium bisulfate (see, for example, http://pubs.acs.org/doi/abs/10.1021/i260036a001 ) as one of the initial products on heating ammonium sulfate.

Said Patent (link http://www.google.com/patents/US4490347 ) from a safety perspective (explosion hazard), further notes:

"A first aspect to achieving the above objects is to feed oxygen-enriched air, i.e., air that is enriched with oxygen, into the furnace. Since the ratio of oxygen to inerts in oxygen-enriched air is higher, less inerts are introduced into the process per mole of oxygen consumed. Thus less fuel is needed to heat gases in the furnace. The degree of enrichment of the air with oxygen has practical limitations. As the concentration of oxygen is increased, the burner temperature increases, flame profiles change, and more sophisticated burners and controls are needed to attain efficient combustion, long equipment life, and to minimize the potential for explosion. Oxygen-Enriched air containing 22 to 40% oxygen by volume can be used in conventional sulfuric acid regeneration (SAR) furnaces with a reasonable amount of control instrumentation."

[Edited on 14-9-2015 by AJKOER]

aga - 14-9-2015 at 13:55

Quote: Originally posted by ave369

I've got a big fat jar of magnesium sulfate and will probably try this ASAP, but it would be good if anyone knows any hidden problems

None of us know дерьмо about it really.

Googling and quoting is never as exciting as Doing and Finding Out. удача !

ave369 - 15-9-2015 at 07:09

Today I tried this synthesis but failed it on the stage of making pyrosulfate. I chose a bad vessel for calcination and fouled the entire yield.

However, it seems that I made an interesting discovery. I'm still unsure if I understand correctly what happened, but I'll describe.

I had a small flask filled with a mixture of H2SO4 and HCl standing on a shelf in my lab for a long time. It was made by bubbling hydrogen chloride gas through 80% sulfuric acid. Today I tried to purify this acid by distilling HCl out of it.

I did it, and here's what I found: an addition of HCl completely suppresses the formation of the infamous white mist! Instead of it, azeotropic HCl vapors come out, which are easily absorbed by poking the retort's nose in water or even just condensed and dripped down, if the speed of distillation is slow enough. When all HCl was distilled away from the mixture, in the retort was 95% sulfuric acid! When I opened the retort's tubulus while it was still hot, it started fuming profusely and the retort filled itself with white mist, but while closed, no SOx at all emerged!

[Edited on 15-9-2015 by ave369]

UC235 - 15-9-2015 at 08:04

Quote: Originally posted by ave369

When I opened the retort's tubulus while it was still hot, it started fuming profusely and the retort filled itself with white mist, but while closed, no SOx at all emerged!

You do know that the white mist is SO3 vapor absorbing water from the air to form a mist of sulfuric acid (a known carcinogen, by the way, if inhaled). If the vessel is closed, there is no moisture to do this even if you are forming SO3 vapor (which you were)

How do you know that the product was 95%? It sounds like you're just pulling a number out of your ass.

With regards to the original post, you did not produce any quantifiable SO3. If there was water present, there isn't any SO3 except traces in equilibrium (no more than the fumes from regular hot sulfuric acid). The bubbles produced are air being forced out of the apparatus as it's heated and the air expands. If you want to test the other poster's suggestion, you need to distill 98% sulfuric with the pyrosulfate in it and see if you can condense any sulfur trioxide.

ave369 - 15-9-2015 at 08:51

Quote: Originally posted by UC235

If the vessel is closed, there is no moisture to do this even if you are forming SO3 vapor (which you were)

How do you know that the product was 95%? It sounds like you're just pulling a number out of your ass.

When I did this before, the mist did form even in a closed retort. Because when sulfuric acid boils, it evaporates its own water along with SO3 AND the water that it was diluted with (that's the very point of boiling it down. If there wasn't more water than SO3 in the vapors, boiling down wouldn't work at all). There were times when the inside of the retort turned milky white.

I determined the concentration of the acid by measuring its density. However, my scale isn't particularly precise, so I used another method to confirm that: by measuring the time required for charring wooden sticks of different size and composition (such as matchstick, toothpick, long thin chip of firewood) with this acid.

Quote: Originally posted by UC235

The bubbles produced are air being forced out of the apparatus as it's heated and the air expands.

I've been living in this world for a long time and have seen many wonders. But air that fumes on contact with air - that's a miracle and a mystery!

BTW, thanks for warning about carcinogens, but I'm already a chain smoker. A little SOxes won't do any more harm.

[Edited on 15-9-2015 by ave369]

papaya - 15-9-2015 at 11:36

Quote: Originally posted by ave369

Today I tried this synthesis but failed it on the stage of making pyrosulfate. I chose a bad vessel for calcination and fouled the entire yield.

However, it seems that I made an interesting discovery. I'm still unsure if I understand correctly what happened, but I'll describe.

I had a small flask filled with a mixture of H2SO4 and HCl standing on a shelf in my lab for a long time. It was made by bubbling hydrogen chloride gas through 80% sulfuric acid. Today I tried to purify this acid by distilling HCl out of it.

I did it, and here's what I found: an addition of HCl completely suppresses the formation of the infamous white mist! Instead of it, azeotropic HCl vapors come out, which are easily absorbed by poking the retort's nose in water or even just condensed and dripped down, if the speed of distillation is slow enough. When all HCl was distilled away from the mixture, in the retort was 95% sulfuric acid! When I opened the retort's tubulus while it was still hot, it started fuming profusely and the retort filled itself with white mist, but while closed, no SOx at all emerged!
[Edited on 15-9-2015 by ave369]

Hmm, this is quite interesting, if your observations are right (getting higher concentration H2SO4 than initiall was there) then it seems HCL can pull water from conc. H2SO4 and put it into gas phase. Something similar happens in industry when benzene is used to get absolute alcohol (benzene forming azeotrope with water and leaving the dry alcohol). Now, what if you try using the same very dry benzene or even toluene to pull water from H2SO4 (add it to your 80% H2SO4 and distill together)? I doubt it though, but who knows?

[Edited on 15-9-2015 by papaya]

ave369 - 15-9-2015 at 14:31

Obtaining toluene here in my town is a mess. A full day's work extracting it from an OTC solvent. Benzene is right out.

However, in my experiment, the concentration of sulfuric acid definitely increased. Maybe I was wrong about the number 95%, but not by a very wide margin: it's at least 94%. There can be no mistake. It has the density. And it does everything real concentrated H2SO4 does and 80% doesn't: it quickly chars wood, water dropped in it goes pop-sizzle. My normal boildowns never go that far: I usually stop at 80%, or at most 90% if I do the boiling down by remote hotplate control in an unmanned bath house, so the SOx vapors are less an issue. This boildown with HCl, this one I did at home, with absolutely none choking gas attacks.

[Edited on 15-9-2015 by ave369]

gdflp - 15-9-2015 at 15:01

Quote: Originally posted by papaya

Now, what if you try using the same very dry benzene or even toluene to pull water from H2SO4 (add it to your 80% H2SO4 and distill together)? I doubt it though, but who knows?

That won't work. Under those conditions, both benzene and toluene will readily sulphonate and will simply contaminate, and add water to, the sulfuric acid.

gdflp - 15-9-2015 at 15:05

Quote: Originally posted by ave369

I'm curious as to whether the concentrated acid still contains traces of chloride, do you have any silver nitrate to test it?

papaya - 15-9-2015 at 15:08

At lease benzene should not sulphonate? If I remember correctly an old procedure called to purify benzene from thiophene (very close physical properties to C6H6, cannot be separated easyly) is boiling technical grade benzene with conc. H2SO4, that destrois thiophene. I'm not sure on toluene, at least 80% sulfuric is not like 96% sulfuric, hardly it can do anything to toluene at that concentration, still needs to be tested.

Magpie - 15-9-2015 at 15:12

ave369 it seems that you have made a very useful discovery for those wishing to obtain 98% H2SO4 (bp = 338°C) starting with weaker acid.

The HCl-water maximum boiling azeotrope boils at 108.6°C at a concentration of 20.2%. So if you add HCl to, say, 80% H2SO4, boiling should remove water in the distillate until the distillate temperature reaches ~108.6°C. Then it will constant temp/constant concentration (20.2%) boil until the HCl has all been evaporated.

If, when the concentration of H2SO4 reaches 98% and there is still HCl left, I imagine it would evaporate off very quickly.

gdflp - 15-9-2015 at 15:15

I believe that the standard procedure for removing thiophenes from benzene is by shaking with room temperature sulfuric acid, at least that's what Vogel calls for. When purifying toluene, both the toluene and sulfuric acid need to be chilled to prevent partial sulfonation of the toluene, I've done this before. And while the initial 80% sulfuric acid may not sulfonate the toluene or benzene, the more concentrated acid desired as the product will at boiling.

papaya - 15-9-2015 at 15:25

I don't insist on anything, but still it's known that aromatic sulfonations are reversible, meaning that when benzene sulfonate is starting to form, there should be only a little water left, that's why oleum is preferred (higher yields), isn't it? Benzene at least can be re-used if this scheme works at all, hope organics people will correct me if wrong. And HCl find is a good find ave, congrats!

[Edited on 15-9-2015 by papaya]

gdflp - 15-9-2015 at 15:43

The issue I am seeing is that, even though the reaction is indeed reversible, I believe that it will inevitably contaminate the sulfuric acid with aromatic sulphonates. Here's an example : Benzene is added to 80% sulfuric acid and the mixture is refluxed with a Dean-Stark, or simply subjected to distillation. Initially, the less concentrated sulfuric acid may or may not be strong enough to sulphonate the benzene, so the benzene/water azeotrope will begin to be collected. As water is removed from the reaction mixture, the equilibrium will begin to shift towards the aromatic sulphonate side due to water being lost, though the loss of benzene mitigates this somewhat. As the aromatic sulphonate forms, the amount of volatile benzene drops and, after a certain point, disappears, however some sulphonate still remains. Seeing as this is dissolved in now nearly concentrated sulfuric acid, there is no way to shift the equilibrium back towards the benzene side to remove this, so I don't see any easy way to remove it. Do you agree, or am I missing something?

ave369 - 15-9-2015 at 22:58

Quote: Originally posted by gdflp

I'm curious as to whether the concentrated acid still contains traces of chloride, do you have any silver nitrate to test it?

Yes, I do, and already did that during the distillation. The still bottoms do not contain trace chloride; I was distilling HCl out until the distilled liquid gave a very faint cloudiness with AgNO3 (and a profuse precipitation with Ba (OH)2, which meant it has much more H2SO4 than HCl).

Upd: repeated the experiment today. The effect was observed, but to a lesser degree (I concentrated sulfuric acid from 78 to 86%). Probably messed up the required proportion between H2SO4 and HCl: I've got that particular mixture by pure random chance. I'll conduct further experiments to come up with a definite methodics.

[Edited on 16-9-2015 by ave369]

[Edited on 16-9-2015 by ave369]

chloric1 - 3-1-2016 at 15:27

Yes the novel HCl trick is a great find. I am so used to seeing the reverse. Adding concentrated sulfuric to concentrated HCl solution to get HCl gas. I was thrilled when I learned calcium chloride could release HCl gas from the acid.

What about adding dehydrated magnesium sulfate to 80% sulfuric acid and distilling? That would be simular to the magnesium nitrate concentration of nitric acid.

byko3y - 4-1-2016 at 05:09

chloric1, what HCl trick are you talkin bout? Dehydration of acidic chloride sources with either CaCl2 or H2SO4 or some other salt is a common way of HCl gas preparation.
Sulfuric acid can be dehydrated to oleum (SO3) with Na2SO4-MgSO4, but you need a high temperature (like 400 °C) and there's a problem with cleaning the device from moltent salt left after the reaction.

chloric1 - 4-1-2016 at 16:10

Quote: Originally posted by byko3y

chloric1, what HCl trick are you talkin bout? Dehydration of acidic chloride sources with either CaCl2 or H2SO4 or some other salt is a common way of HCl gas preparation.
Sulfuric acid can be dehydrated to oleum (SO3) with Na2SO4-MgSO4, but you need a high temperature (like 400 °C) and there's a problem with cleaning the device from moltent salt left after the reaction.

I was talking about ave369 adding HCl to 80% sulfuric acid and distilling off azeotropic hydrochloric acid to leave 93 -95% sulfuric acid behind. It is directly opposite of adding concentrated sulfuric acid to hydrochloric acid to get HCl gas. If Na2SO4 MgSO4 mix will liberate SO3 from sulfuric acid, it would surely concentrate 80% sulfuric acid to 93% or more.

byko3y - 5-1-2016 at 05:47

You don't need to add anything to make 93-95% sulfuric acid, simple distillation o dillute H2SO4 is enough.
No, sulfates can't simply dehydrate sulfuric acid. There's only few compounds that can do this, like polyphosphoric acid.

clearly_not_atara - 6-1-2016 at 15:04

It makes sense to me: the affinity of HCl for H2O is huge, comparable to H2SO4 itself.

Quote:

Upd: repeated the experiment today. The effect was observed, but to a lesser degree (I concentrated sulfuric acid from 78 to 86%). Probably messed up the required proportion between H2SO4 and HCl: I've got that particular mixture by pure random chance. I'll conduct further experiments to come up with a definite methodics.

I'm guessing it simply gets better the more HCl you use. It probably won't be possible to use too much HCl, as it has relatively little love for H2SO4; any "extra" will simply boil away (an unpleasant gas, but not as bad as H2SO4). The caveat is that you will, obviously, need to both generate dry HCl and trap the outgas.

The SO3 can be trapped as a complex; it forms complexes most famously with pyridine but also dioxane and DMF. Any of these is easier to handle than oleum...

byko3y - 7-1-2016 at 05:32

Are you kidding me? The affinity of H2SO4 for water is so huge it can dehydrate organic like sugar into carbon at room temperature.

Praxichys - 7-1-2016 at 08:03

A common lab practice to generate HCl gas is to drip aqueous HCl onto concentrated sulfuric acid. The sulfuric acid strips the water from the HCl very effectively. Azeotropes exist because of affinity between two polar compounds. When another compound with any degree of affinity is introduced, the azeotrope shifts, and in some cases is destroyed.

Certainly the HCl you distill from the mixture cannot be its 20.2% azeotrope at atmospheric pressure, because if the affinity H2SO4 has for water. However, HCl is still acting as a volitile entrainer with some affinity for water. I suspect that for a given temperature, heating H2SO4/HCl/water mix is more effective at dehydrating the sulfuric acid than just heating H2SO4/water alone. However, the caveat is that it will reach an asymptotic equilibrium of concentration where the affinity for water is the same for both acids at that particular concentration, and no further concentration can be achieved. This will likely be substantially less than the 98% achievable by boiling.

Theoretically, one needs to add HCl of a concentration higher than the azeotrope that can be distilled from the desired concentration of sulfuric acid, to achieve that concentration. The theoretical maximum concentration of H2SO4 achievable with this method would be obtained by bubbling pure, gaseous HCl through aqueous sulfuric acid at some arbitrary temperature, preferentially less than that required to drive the water off using the old boiling method, but higher than the boiling point of the resulting HCl/water mixture. In this way I think it could be concentrated further than its 98% azeotrope, although it would take a long time unless the bath was simultaneously heated on the same terms as the boiling method.

In short, 31% hardware store HCl will only do so much. I fear that it may only serve to be a grossly inefficient way to achieve a concentration of H2SO4 much less than 98%, when boiling it could achieve an equal concentration at a reasonable temperature with far less apparatus. Concentrating beyond that would require the generation of HCl gas without conc. H2SO4 which involves more apparatus and reagents. At that rate, you may as well invest in a throwaway hotplate from WalMart and a $4 beaker, and just boil it down outside. It's free, and everything is totally reusable as long as you don't do anything stupid with the hot glass.

Besides, running all that HCl through the H2SO4 will concentrate metallic impurities that are common in hardware-store HCl and leave them in the sulfuric acid, while at the same time introducing traces of chloride which are very hard to remove.

[Edited on 7-1-2016 by Praxichys]