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NaHSO4

axehandle - 28-1-2004 at 22:40

Sorry for cross-posting, but this merits attention:

"NaHSO4 mixed with alcohol gives a solution of H2SO4 in alcohol and Na2SO4 as a precipitate.
Boil, or better, distil off the alcohol and get the acid."

2NaHSO4 + C2H5OH --> H2SO4 + Na2SO4 + C2H5OH

LOOKS valid, and if it worked it would be kind of a H2SO4 holy grail (NaHSO4) can be
bought by the tons at pool supplier's.

Anyone ever tried this? Have I missed something (sometimes I can't even hit a barn with a BB-gun, sometimes I'm so right you could cry... =)

BromicAcid - 28-1-2004 at 23:05

Immediately I would wonder about the ability of the reaction to go foreword. If you dropped sodium hydrogensulfate into alcohol and it is not soluble to begin with then the reaction would proceed incredibly slow. Even if it is soluble in alcohol then you have the problem of an insoluble solid forming on the outside of the sodium hydrogensulfate and preventing further reaction.

You would probably have to dissolve the sodium hydrogensulfate in something like water then add and watch the precipitate but the water is going to be able to dissolve appreciable amount of the sodium sulfate formed, and the alcohol is going to distill off before the water so this reaction would get all mucked up.

In addition if you managed to get pure sulfuric acid and ethanol then if you try and boil it off you'll produce ethane and diethyl ether due to the anhydrous acid concentration continuously rising and the temperature being elevated. But, all this is off the top of my head, I should look up some solubility's but just giving my two cents for now.

axehandle - 28-1-2004 at 23:15

Here are my 3 cents:
http://www.jackspcs.com/chemdesc.htm

"SODIUM BISULFATE
Other names: Sodium acid sulfate; sodium hydrogen sulfate; sodium bisulphate; acid sulfate; sodium pyrosulfate; niter cake (impure).
Description: Sodium bisulfate is available as colorless, free-flowing crystals or as white, fused lumps; One part of the anhydrous bisulfate is soluble in 2 parts water or 1 part boiling water, but alcohol decomposes it into sodium sulfate and sulfuric acid.
Precautions: Keep container tightly closed. Use with adequate ventilation. Do not inhale dust as contact with sodium bisulfate causes burns. Avoid contact with eyes, skin, and clothing. Always wear safety glasses or goggles and rubber or similar impervious gloves. Do not swallow.
First Aid: For eye contact irrigate the eyes thoroughly with plenty of water, then contact a physician or an eye specialist. Thoroughly flush contacted skin with plenty of water. Remove contaminated clothing at once and wash clothing before re-use. If swallowed, wash out mouth thoroughly with plenty of water; give much water to drink, then give milk of magnesia. Call a physician at once.
Photographic Formulas: Iron Green/Blue Toner."

The_Davster - 28-1-2004 at 23:22

To make sure little/no diethyl ether is formed couldent you just mix the reactants in stoichiometric ammounts? or am I missing something really obvious here?

axehandle - 28-1-2004 at 23:36

Where would the diethylether come from?

axehandle - 28-1-2004 at 23:38

Ah, sorry BromicAcid, I see your point. I'll just have to make damned sure my fume hood is working properly.

chemoleo - 29-1-2004 at 02:09

Well, if this works it's great!
For one thing, you *will* get diethylether, and temporarily ethylsulphate, but the good thing is alcohol is continuously shuffled out of the system by distilling it off, as ether.
What I don't see, however, is, if only H2SO4 and Na2SO4 are present (the alcohol is now all boiled off as ether), what stops them from recombining to NaHSO4? In other words, what we DON'T want is to destill off the C2H5OH, as it maintains the equilibrium being shifted to the right side!
So, the question then poses - how are you going to purify the H2SO4 if you can't get rid of the Ethanol by destillation?
Conversely, you need to find a way to separate the H2SO4 from the Na2SO4 - and then you are all set

axehandle - 29-1-2004 at 03:30

Na2SO4 is precipitated as a solid. I'll just pass
the ethanol/H2SO4 solution through a filtering
paper. Na2SO4 == history. =)

chemoleo - 29-1-2004 at 03:45

did you try it? If you can easily separate the two, then i guess all your problems are solved! Congratulations, you just found a way to make H2SO4 from NaHSO4, and ether as a byproduct!
Anyway, let us know if you get it to work well.

PS by the way, filtering paper is almost certainly not good, as you will turn it into carbon with the conc H2SO4! Borosilicate filter might be better, or just let it settle.

[Edited on 29-1-2004 by chemoleo]

axehandle - 29-1-2004 at 03:52

chemeleo: Thanks for the tip. To think that I
was going to use a coffee filter *shudder*. What was I thinking?

I'll just let it settle, then pour off as much as
I can. Without the Na2SO4 there should be
no ether, right? In that case, a slow boil on
the stove's lowest setting might do just the
trick to remove the ethanol.

And in fact, I'm gonna try the method today!
I found NaHSO4 for appr. $20/7kg.

Then I'm gonna buy 4 liters of 95% ethanol.

Wish me luck, and I'll report my success (or failure) later.

[Edited on 2004-1-29 by axehandle]

chemoleo - 29-1-2004 at 04:45

forgot to say, maybe it might be wise to use an *excess* of ethanol - the reason being that it will be easier to filter off any precipitated Na2SO4 or unreacted NaHSO4.
I would try to do this quantiatively, for instance take 10 g of NaHSO4, and calculate the theoretical amount of H2SO4 you'd get. Then calculate how much Na2CO3 you'd need to neutralise the carbonate to get Na2SO4. If you manage to dissolve the carbonate with bubbles appearing to the very end (CO2), then you know that the yield was very good. If there are no more bubbles after you added half of the theoretical amount necessary, then you know the yield is at 50% and so on.
Also, for starters, i would grind the NaHSO4 as fine as possible, and also heat the solution up to 60 deg C or so (not over a Bunsen flame, but heat plate!!!! ). This with continuous stirring. Do it for an hour or so. Actually, you might want to reflux it - that way you will keep the ether.
Good luck

Ramiel - 29-1-2004 at 04:54

I may be wrong too, but here's my 2 peneth worth (that's about 1/4 cents!

).
If there is little water, and the H2SO4 is more ionic than the NaHSO4, then won't the sulphuric displace the less ionic salt from solution, and favour the forward reaction?

something sounds fishy about all this, but I just wonder if this helps the mechanism of the reaction (to do what one might not expect!).

notagod - 29-1-2004 at 05:20

axehandle! Where did you get the NaHSO4 from. (which company, city). You can PM me if you don't want to reply in thread.

axehandle - 29-1-2004 at 08:25

Any swimming pool chemical supplier should carry it.

17:15 I've just begun the first experiment, at room temp. The reaction is going slowly, so I'm heating the reaction vessel with a hair drier.

17:25 Definitively getting an acid response. Litmus paper turned red almost instantly.

We'll see how it goes. Don't get too excited and blow all your money on NaHSO4 now, mates.... I'll fill you all in on the developments.

axehandle - 29-1-2004 at 08:44

17:40 Added an aqarium air pump as an ad-hoc stirring device. Temp. of alcohol approaching 50 degrees C.

axehandle - 29-1-2004 at 08:45

17:44 Stupidly tasted some of the liquid. Shock. Tish, tish, definitily acidic.

axehandle - 29-1-2004 at 08:49

17:48 Decided to go and wash my hands to avoid further acidic nailbiting incidents.

axehandle - 29-1-2004 at 09:35

This rection takes forever, albeit *something*
is definitely happening. I should have put the NaHSO4 prills (~2mm dia) in my ball mil first.

Trouble is, NaHSO4 is toxic as hell and can be absorbed through the skin as well as the lungs. Wouldn'r really want to inhale the stuff when it's airfloat.....

Have I finally found the holy grail?

axehandle - 29-1-2004 at 09:57

I dissolved 1 part NaHSO4 in 1 part (by mass)
cold water. (Credits to BromicAcid for the suggestion.) Then I added a generous amount of ethanol. A precipitate immediately formed,
making the liquid milky until it sank to the bottom.

Litmus paper test indicated strong acidity. Turned bright red instantly.

I'm going to perform a larger test in a 2 liter
PET bottle.

Does anyone know if PET is resistant to BOTH ethanol AND sulfuric acid?

[Edited on 2004-1-29 by axehandle]

BromicAcid - 29-1-2004 at 10:15

Just a note, sodium bisulfate itself is strongly acidic in solution.

NaHSO4 ----> Na(+) + H3O(+) + SO4(2-)

So just testing for acid really isn't enough. You would probably need a titration to get the most accurate reading of how much H2SO4 you produce.

axehandle - 29-1-2004 at 10:21

Yes, I know. From breathing the stuff firsthand. Very nasty stuff this.

Since I've forgotten all my college chemistry, could you please tell me of a kitchen-method way of detecting sulfuric acid?

axehandle - 29-1-2004 at 10:26

Whooo haaa!!!!!! Look at THAT precipitate!!!
It's INSTANT.

(Just added about 1 liter of ethanol to one liter of water+NaHSO4 solution.)

Now for the slow boil until (hopefully) white fumes appear.....

axehandle - 29-1-2004 at 10:44

If this really works, I'll have to build a condenser for the ethanol. I just spent 5 liters
for a measly 0.5kg of NaHSO4, and that makes
about $18 just for the ethanol.

Don't get ahead of yourself

Polverone - 29-1-2004 at 12:44

NaHSO4 itself is probably much less soluble in ethanol than it is in water, just like most other ionic solids. If you made a strong solution of NaCN in water and then added ethanol, you would get a precipitate also, but it doesn't mean the NaCN, water, or alcohol have changed. The solubility of the solvent system is just different.

Further, ethanol that contains even a little bit of water could possibly dissolve enough NaHSO4 to give an acid reaction on test paper, without any significant formation of ether, sulfuric acid, or ethyl hydrogen sulfate.

Finally, sodium pyrosulfate is (contrary to that information above) distinct from sodium bisulfate, and I would expect it to react with alcohol more than I would expect sodium bisulfate to. Sodium bisulfate can be transformed to sodium pyrosulfate by heating at 150-200 C for a couple of days (!) or by heating somewhat higher than that for a few minutes. If it behaves like potassium bisulfate, it will melt, give off some vapor (water), then solidify again, requiring much higher application of heat to melt once more. This higher-melting solid should be mostly sodium pyrosulfate, and I think it could be worth trying to prepare and use some if your experiments with sodium bisulfate end in failure.

I don't want to rain on your parade, just warn that initial encouraging results can be misleading. Also, no reaction conditions were given for the reaction of NaHSO4, so it may be that you'll get satisfactory results from it but only if you're willing to wait a long time or perform the reaction with, say, hot ethanol vapor run over heated NaHSO4.

guaguanco - 29-1-2004 at 13:08

Quote:

Originally posted by Polverone
perform the reaction with, say, hot ethanol vapor run over heated NaHSO4.

Off the top of my head, running EtOH vapor over hot (presumably anhydrous) NaHSO4 sounds like a good way to produce either ethylene or diethylether. Hot NaHSO4 is an excellent cracking catalyst.

the pyrosulfates

Polverone - 30-1-2004 at 17:00

I've been chatting with axehandle a bit via U2U. Here's some information that others might benefit from too:

Quote:

The pyrosulphates of the alkali metals are obtained most easily by heating the corresponding acid sulphates to 300-350 C:
2KHSO4 = K2S2O7 + H2O.
This dehydration, however, does not proceed to completion even under reduced pressure, and in order to obtain the pure salts it is preferable to allow the finely divided anhydrous normal sulphate to interact with sulfur trioxide. ... When heated strongly, for instance above 450 C, the pyrosulphates undergo decomposition into sulphur trioxide and the corresponding normal sulphate. The sodium salt, which melts at 400.9 C, exhibits appreciable dissociation at about 460 C. In contact with moist air or water the pyrosulphates become converted into the corresponding acid sulphates:
K2S2O7 + H2O = 2KHSO4

Friend (editor), A Text-book of Inorganic Chemistry

axehandle - 30-1-2004 at 17:55

So, NaHSO4x1H2O into paint can. Hole with cork and glass pipe into lid. Heat the little sucker in my propane furnace until no water is formed anymore, continue heating when SO3 is produced, down into bubbler with water until reaction stops.

I have lots of empty paint cans, corks and glass pipes. I also have a determination. I will NOT yield to the stupid ban on acids this fucking government has impelled on me. I SHALL have my sulfuric acid, I'll die trying!

Hmmm, btw, isn't SO3 a solid at room temp.?

[Edited on 2004-1-31 by axehandle]

rikkitikkitavi - 31-1-2004 at 09:56

Polverone, actually , Na2SO4 shows very little solubility in water/EtOH solutions up to 30-40 % water content.

I will see if I can find my copies of gmelin about the Na2SO-NaHSO4-H2SO4 /H2O/EtOH /pyrosulfates that I have somewhere. It is only in german though, but I can scan it. Very informative.
There is a few patents regarding decomposing NH4HSO4 by EtOH/MeOH since this is easier obtainable , but there is probably old patents regarding NaHSO4 to, since this was a byproduct with little use in manufacturing high conc HNO3 in late 19th.

axehandle, forget any metal in contact with molten hydrogensulfates. It eats it up in notime, actually molten pyrosulfate/hydrogensulfate is used for dissolving in chemical analysis.

[Edited on 31-1-2004 by rikkitikkitavi]

[Edited on 31-1-2004 by rikkitikkitavi]

axehandle - 31-1-2004 at 19:38

Thanks. You just saved me a furnace contaminated with NaHSO4.

I'm right now dry distilling anhydrous FeSO4,
using empty paint can with cork and copper
tubing. Seems to work, albeit slowly.

ehem

Organikum - 1-2-2004 at 12:00

Quote:

Hole with cork

axehandle? SO3? Cork? Uiiiiiiiiiii.........

Ceramics, clay, pottery!!!!!
Wasserglas.....
Flowerpot and kruzifix nochamal!

Herrgottdaskannsdochnichtgeben.......

chemoleo - 1-2-2004 at 12:10

Lol Organikum - at times of need you just have to use cork

- or paper plugs!
By the way, all he will have is a little carbon impurity

Are you bavarian?

axehandle - 3-2-2004 at 21:46

Orkanikum: Kork bist entwas für die chemische reaktion ich habe lust für.

Unt ich brugt nicht wasserglass, ich brugt Duranglas.

Ein Reich, eine Folke. ein Fuhrer!

I'm for the 1st: Joking, and for the 2nd: I don't know German, althouch my girlfriends father is an immigrated German.....

axehandle - 3-2-2004 at 21:49

Organikum: I actually understood all of your post, learning German as I am. I choose to ignore it.

Tacho - 4-2-2004 at 02:11

I have found by experience that NaHSO4 can substitute H2SO4 in many reactions. Every now and then I use it instead of H2SO4 for simple experiments. I have made ethyl bromide using it.

I first got interested in it when I read somewhere that you could do destructive distilation (with silica?) to obtain H2SO4 the same way you do with Iron sulfate. Never tested this.

Very easy to get, any pool shop should have it.

Also, when heated with common salt, produces a very acid gas (test paper goes red very quickly) which has to be HCl.

Anhydrous NaHSO4 should produce anhydrous HCl.

[Edited on 4-2-2004 by Tacho]

[Edited on 4-2-2004 by Tacho]

Polverone - 4-2-2004 at 09:12

I've wondered recently if NaBr could be used the same way to produce HBr without any oxidation like you'd get with H2SO4. I do recall reading in some ancient manual of industrial chemistry that KHSO4 or NaHSO4 could be heated with sodium formate to give formic acid, without the oxidation that H2SO4 would cause.

rikkitikkitavi - 4-2-2004 at 11:33

HSO4 pKa 2
HCOOH Pka 3,75

thus the reaction

Na-Fo + NaHSO4 => HFo + Na2SO4 has a pk of 1,75 , i e shifted to the right

Since HFo is volatile, it can be evaporated easily , thus drives reaction towards completion.

NaHSO4 reacts with NaCl to produce HCl , but at high temperatures, the hydrosulfate needs to be molten.

/rickard

[Edited on 4-2-2004 by rikkitikkitavi]

Now I've had it!

axehandle - 5-2-2004 at 13:38

After filling a steel pipe with anhydrous FeSO4 and putting it in my propane furnace in a (futile) attempt to produce H2SO4 (through bubbling), with the result being stinking down my entire flat with SO2 (there must have been a leak somewhere, or perhaps the FeSO4 didn't like 1600 C of temperature and decomposed into FeO2 + SO2), I've decided to go the whole way and build myself a nice, old-fashioned lead chamber.

You heard me: A lead chamber!

The contact method is out, mostly due to the fact that sulfur poisons the platinum, and I can't find a catalyst that would work without extreme temperatures and be very hard to apply to a suitable substrate.

Now, I'm going to to try to find an enormous quantity of lead sheet for my chamber. The damned thing will probably weigh half a tonne, but I'll be damned if I let these setbacks stop me!!!

So, unless someone can suggest a very easy to prepare and use catalyst for the contact method, lead chamber it is. Who knows, it might even impress girls. (-Whats that? -Oh, it's just my lead chamber)....

:mad:

Mumbles - 5-2-2004 at 15:32

1600C? I can tell you your problem without a doubt. The SO3 decomposed to SO2 and O2. I think it decomposes around 400 or 500. if you can get it right near the decomposition point of the Sulfate the SO3 could probably cool enough before it decomposed to a great extent.

I do assume by the bubbling method you mean into Sulfuric Acid. Into water could do more than fill your flat with SO2 fumes.

axehandle - 5-2-2004 at 15:44

My suspicions confirmed are, somewhat. Thanks.

wait a second

Polverone - 5-2-2004 at 17:09

1600 C?! No way! Your steel pipe would have been converted to an incadescent puddle of metal. How are you estimating the temperature achieved? Also, I believe that mixing fine silica with your sulfate may help it to decompose in the fashion you desire, since I recall the addition of silica from descriptions of old techniques for producing H2SO4 from ferrous sulfate.

[Edited on 2-6-2004 by Polverone]

axehandle - 5-2-2004 at 17:21

Estimated maximum of the furnace. The inside is lined with a special mixture of bentonite, perlite and cement, which make a very goog heat insulator.

The pipe had, however, a yellow glow, so I would estimate 1000C.

Sorry for the confusion.

to fast

Organikum - 5-2-2004 at 19:19

Using cork but cant stand a joke..... tz tz

No, seriously:
You heated the FeSO4 to fast and to high as it sounds, you have to give it some time to dehydrate first, the water doesnt jump out of the molecule at a certain temperature! This takes a while. So heat it to above the dehydration temperature and hold it there for some hours - then power it up to make the H2SO4 etc. At this point your setup should be completely enclosed - and really, water-glass aka sodium silicate is favorable for lining the steeltube insides and for sealing - I dont tease you, I try to help.

Lead is used in high-power electricity cables shielding - the kind of buried cables. About 100gram/10cm. Contact your local industrial scrapyard for this. It is a fuck to get it out but its the best source for lead I know.

But you might consider to try the FeSO4 decomposition again - outsides so possible - and slowly as advised. Lining the steel with the silicate is a good idea - believe me, not without reason writes the MERCK that water-glass is used for lining vessels used for acid concentration.

If you wont trust me,
trust the MERCK.

But hey! I got one more:

Der Führer war ein armes Schwein, er hatte keinen Führerschein!

Chemoleo:
I? Barbarian? I? - eh, perhaps....

[Edited on 6-2-2004 by Organikum]

axehandle - 5-2-2004 at 22:20

>to fast

>Using cork but cant stand a joke..... tz tz

I know that, but it was the only sealing material I had at hand, and the pipe _was_ sticking about 15cm out of the furnace... for
tubing I used a soft-annealed copper pipe.

>No, seriously:
>You heated the FeSO4 to fast and to high >as it sounds, you have to give it some
>time to dehydrate first, the water doesnt
>jump out of the molecule at a certain
>temperature! This takes a while. So heat it >to above the dehydration temperature
>and hold it there for some hours - then
>power it up to make

I did. I put a generous amount in the kitchen oven and dried it at 300C for about 3 hours. Should have done the trick, it changed colour and the density went down considerably. It was pure Hell to break up the lump and get them through the small cork-hole in the pipe though...

>the H2SO4 etc. At this point your setup
>should be completely enclosed - and
>really, water-glass aka sodium silicate is
>favorable for lining the steeltube insides
> and for sealing - I dont tease you, I try to
> help.

I know that, your tone doesn't suggest otherwise. But I'm not planning to dry distill FeSO4 ad nauseam, it was a one-time shot and I missed. I won't do it again, my furnace basically has two settings: 1: Off, 2: Full blast (I've built the propane burner myself and it has a lowest setting that is not low enough...).

>Lead is used in high-power electricity
>cables shielding - the kind of buried cables.
>About 100gram/10cm. Contact your local
>industrial scrapyard for this. It is a fuck to
>get it out but its the best source for lead I >know.

I know a better one: Used X-ray shield plates. About 1.5mm thick, 200mm wide and anywhere from a piece to 2 meters in length. My local scrapyard deals alot in them They are basically for free, but heavy as lead to carry, pun intended.

>But you might consider to try the FeSO4
>decomposition again - outsides so possible
> - and slowly as advised. Lining the steel
>with the silicate is a good idea - believe
>me, not without reason writes the MERCK
>that water-glass is used for lining vessels
>used for acid concentration.

I won't try it again, I've contacted a jeweller about some platinum. Next time it's either contact method, or most likely, lead-chamber.

I only really need a small amount of acid, to
make a small amount of nitric to make aqua regia to make a catalyst on a kieselguhr substrate, etcetera. Catch-22.

Do you perchance know why many texts quote that the maximum concentration attainable in a lead-chamber is 78%? Shouldn't it be possible to boil the water off, or has it got asomething to do with SO2(aq)?

>If you wont trust me, trust the MERCK.

I trust you both.

>But hey! I got one more:

>Der Führer war ein armes Schwein, er
>hatte keinen Führerschein!

The Führer was a poor pig, he had no foreskin (direct translation (and I think the _correct_ word for foreskin is Führschein, nicht war?)).

no

Organikum - 6-2-2004 at 03:17

driver-license is the right translation....

axehandle - 6-2-2004 at 09:38

oops

Freudian slip, it seems...... heh

But I think it was close enough. What's the consolation prize?

[Edited on 2004-2-6 by axehandle]

There, just drilled a 10mm hole through my apartment balcony wall for an exhaust from a fume fan. I wonder if my landlord will notice? *shudder*

[Edited on 2004-2-6 by axehandle]

surely you mean 100mm!

Magpie - 6-2-2004 at 20:27

axehandle - 8-2-2004 at 19:25

No, 10mm is exactly what I meant. I was a bit unclear though, it's not really for a fume hood exhaust, just for a trickle of air + NO from my almost-finished Birkeland-Eyde reactor. The air flow comes from two humongous aquarium pumps, but 10mm is more than enough.

rikkitikkitavi - 15-2-2004 at 12:18

axehandle., FeSO4 decomposes into SO2.

You need Fe2(SO4)3 whcih decomposes into SO2 and SO3 when heated, and at lower temperature so the SO3 doesnt decompose.

/rickard

Friedrich Wöhler - 15-2-2004 at 13:31

That's true, that pure FeSO4, when heated, give a mix of SO2 and SO3.

But: [Hofmann/Rüdorff: Anorganische Chemie]

Centuries ago, they let FeSO4 ("iron vitriol"

with moisture oxidize by air to an Iron(III)-hydroxide-sulfate Fe(OH)SO4.
This, when heated, gives that "fuming sulfuric acid":
2 Fe(OH)SO4 ===(heat)===> Fe2O3 + H2S2O7 (thus H2SO4 + SO3).

You really can get it by blowing air through an FeSO4-solution, this Fe(OH)SO4 is unsoluble, you can it filter and dry.
How easily to oxidize an FeSO4-solution is, you can see when you dilute pure FeSO4 in hot water and in few minutes in contact with air it will change its color ang get cloudily because of Fe(III)- building.

Just I don't know that decomposition temperature but I can remember it was very low. (Even must not glow...)

fritz - 16-2-2004 at 09:57

Instead of Pt you can use following catalysts for oxidation of SO2:

V2O5-catalyst:
pumice and V2O5 in the ratio of 2:1 are wetted with water. Dried in vacuum and heated 30mins at 120°C
OR:
To pumice NH4VO3-sln (saturated) is added as much as the pumice can absorb. You let dry and make it glow carefully.

best use at 400-500°C
reaction starts at about 200°C

Cr2O3-catalyst:
Cr2O3 is heated strongly and brought onto glass-wool.
OR:
(NH4)2Cr2O7 is heated carefully and added to glass-wool.

best use at 200-300°C

Fe2O3-catalyst:
saturated FeSO4-sln. is absorbed on pumice or pieces of brickstone. After drying
heat it to 800°C until no more white smoke appears.
OR:
Fe2O3 is brought on pumice or brickstone

best use at 500-600°C

Fe2O3/CuO/Cr2O3-catalyst:
a mixture of Fe2O3; Cr2O3 and CuO (ratio:???) is dried at 120°C and brought onto glass-wool.

CuO/Cr2O3-catalyst:
CuSO4 and CrCl3 solution (both saturated) is absorbed on pumice. heat to 800°C until no more white smoke appears.

best use at 300-400°C

hope this will help you.
My source suggests asbestos as absorbing material but this seems not to be a good idea...

rikkitikkitavi - 16-2-2004 at 09:59

hence the name "vitriol oil" ?

Interesting , that thing about FeSO4 becomes usefull. I once found an old, non opened plastic can with FeSO4 (used as a biocide) wich
to much surprise wasnt full of light green crystals (FeSO4) but a brown mush of what must have been Fe(OH)SO4 then. Fe2+ easily oxidizes to Fe3+

/rickard

axehandle - 18-2-2004 at 08:01

Pumice seems hard to get. Does anyone think perlite would work as a substitute? Or even those little clay balls used in the bottom of flower pots ("lekakulor" in Swedish, can't remember the English name)?

[Edited on 2004-2-18 by axehandle]

Mr. Wizard - 18-2-2004 at 08:14

Is Vermiculite the pumice substitute you were thinking about? It is made by heating a volcanic rock which then expands into little worms (vermi?).

axehandle - 18-2-2004 at 08:17

No, perlite originally has a density close to water and is technically a glass. It is made into a very porous material by soaking it in water, then heating it. The steam expands the material to about 10 times its original volume.

Edit: It's white, btw. Can be bought for next to nothing in any gardening store, is very porous, can absorb lots of water (and catalyst...?).

Edit2: Also, it's used in refractory comps because of its porosity and resistance to heat (in the 2000C range!). + I have about 100 liters of it in a huge plastic bag. Weighs about 3kg.

[Edited on 2004-2-18 by axehandle]

[Edited on 2004-2-18 by axehandle]

axehandle - 19-2-2004 at 06:43

On my eternal quest for H2SO4, I'm now manufacturing large volumes of Fe(OH)SO4 using FeSO4, a glass pipe, a stainless but stained steel bowl, an aquarium pump and hot water... the reaction is very fast, I couldn't believe it. We're talking 100g of FeSO4 converted to Fe(OH)SO4 in 15 seconds here. This is the good part.

Quote:

That's true, that pure FeSO4, when heated, give a mix of SO2 and SO3.

But: [Hofmann/Rüdorff: Anorganische Chemie]

Centuries ago, they let FeSO4 ("iron vitriol"

Does anyone have access to a textbook that would cite the temperature required? I don't want the flat filled with SO2 thing to happen again.... If it really is very low (e.g. lower than the boiling point of H2SO4), then it could be made on the stove.

Otherwise I'll have to build an improvised Liebig cooler.....

And I'll have to buy a propane Bunsen burner..... My furnace doesn't allow for much control, it's basically two settings: Off, and full blast, respectively.

Edit: Slowly drying the Fe(OH)SO4 on the stove. The brown sludge bears a striking resemblance to what I produced the last time I had a stomach illness.

[Edited on 2004-2-19 by axehandle]

Friedrich Wöhler - 19-2-2004 at 20:50

Quote:

Does anyone have access to a textbook that would cite the temperature required?

Let me search at the weekend.

Quote:

The brown sludge bears a striking resemblance to what I produced the last time I had a stomach illness.

Don't confound them, the last one is not suitable to produce sulfuric acid!

axehandle - 20-2-2004 at 02:50

Thank you.

Quote:

Don't confound them, the last one is not suitable to produce sulfuric acid!

I'm not so sure about that.