Sciencemadness Discussion Board - Preparation of Dilute H2SO4 from FeSO4

Method 2. By reacting with HCl (or other strong acid):

FeSO4 + 2 HCl ---> FeCl2 + H2SO4

*sigh*
If only it were that easy. Do you have an enthalpy of reaction value for that reaction?

Have no fear, I'll calculate it for you: -46kJ/mol (give or take).

On paper it seems like it would work, but in practice, why would you want to swap HCl for sulfuric acid?

Your FeSO4 + 2 NaCl ---> Na2SO4 + FeCl2 is interesting to say the least. I doubt that reaction would take place smoothly in practice, even though I believe the ∆H works out to be negative (I didn't double check it so don't take my word for it).

weiming1998 - 22-1-2012 at 17:10

If I remember correctly, the dry distillation of iron sulfate actually makes fuming sulfuric, or oleum.

But is there some reason that you're going through all these contortions to make H2SO4?

Is it academic interest? Or what? I hate to use the T word, but your posts kinds of have that aroma.

Some chemicals are just not worth making. I make hydrochloric acid just because the hardware stores no longer carry it and I kind of enjoy making chemicals.

But making sulfuric acid just seems to be so 17th century. If you don't know where to buy it, I can tell you.

I don't. Tried drain cleaner. No sulfuric acid drain cleaners. Tried asking low fume pool acid at Sigma Chemicals. They only come in 15L drums, and even if I want to buy it, I probably can't because the workers were already beginning to ask a frenzy of questions (like what's your pool's condition?, Do you have a sample of your pool water with you? etc. Buying a whole car battery, then draining the acids is too expensive and I can't find anywhere that sells car battery electrolytes.

entropy51 - 22-1-2012 at 18:59

I can't believe Oz is so backward that they don't have sulfuric drain openers or battery acid. Somewhere.

I can believe that people don't want to sell them to 13 year olds, however.

It took me several years of chemistry discussions to earn the trust of the local pharmacists, but once I did, they would order anything I wanted.

Did you know that you can make HNO3 by dry distillation of a mixture of alum, CuSO4, and NaNO3? Of course you can't do this until you find a source of nitrates.

Less whining and more looking is my advice. Good luck.

weiming1998 - 22-1-2012 at 20:00

I can't believe Oz is so backward that they don't have sulfuric drain openers or battery acid. Somewhere.

I can believe that people don't want to sell them to 13 year olds, however.

It took me several years of chemistry discussions to earn the trust of the local pharmacists, but once I did, they would order anything I wanted.

Did you know that you can make HNO3 by dry distillation of a mixture of alum, CuSO4, and NaNO3? Of course you can't do this until you find a source of nitrates.

Less whining and more looking is my advice. Good luck.

Oh well, just asking. If you people won't help, then I'll look for them on my own like I've always did. Simple as that. And no, I'm not complaining.

weiming1998 - 22-1-2012 at 20:07

Also, if I tried to talk about those pool workers about how I am doing home chemistry and need sulfuric acid from the low- fume pool acids, then they will either:
.Sell nothing to me
.Sell things to me and act normal, then ring the police.
.Sell nothing to me AND kick me out of the shop.
.Sell nothing to me, then ring the police.
. Actually having a discussion, and selling me SOME of the things.
. Lecture about how these things are dangerous and telling me to go play/get a girlfriend (like you people have done)

Aqua_Fortis_100% - 22-1-2012 at 20:16

H2SO4 only loses to water on chemical use. So it cant be an easy task to completely hide it from you.

In my country, you (commom people, not firm/enterprise) could only buy max. 2lt of 98% H2SO4 per month from chemical supplier, and with your civil/adress information. Since I dont usually use large amounts of H2SO4 in short times, and dont make anything illegal with it, thats suits my experimental needs very well.

However, before 18 (actually before 16), I have boiled battery acid to obtain my own "98%" sulfuric (which was at most 85-95%). This is, IMHO, the only real way of "making" large amounts of concentrated H2SO4 by your own. The greatest drawback is that is an extremely dangerous operation, I and many members cant stress enough how dangerous is that.. Cold concentrated sulfuric can blind you instantly, so you can imagine what >300°C concentrated could do to your skin/face.. This without mention the SOx/H2SO4 nasty fumes.. Today Im much older and think that I dont need to expose myself with this kind of danger.

Stick with entropy51 idea. It dont worth the effort to make your own H2SO4, except only if you really cant find it (which is probably impossible.. Otherwise, you havent looked hard enough). Try to convince your parents or older friends to buy it for you (remember, if you have on hands a bottle of H2SO4 , treat it with greatest care.. It will not be nice if have a chance to hit you).

For me, the idea of FeSO4 pyrolisis being useful is just one: oleum. Oleum, unlike normal H2SO4 is very hard to come across and is a very valuable chemical, and anhydrous ferrous sulfate could be put in use to make it instead of plain sulfuric.. (even though, better options for SO3 generation are up to amateur reach and are already discussed in the forum), just bubbling SO3 in conc. H2SO4.

[Edited on 23-1-2012 by Aqua_Fortis_100%]

weiming1998 - 22-1-2012 at 20:37

Quote: Originally posted by Aqua_Fortis_100%

H2SO4 only loses to water on chemical use. So it cant be an easy task to completely hide it from you.

In my country, you (commom people, not firm/enterprise) could only buy max. 2lt of 98% H2SO4 per month from chemical supplier, and with your civil/adress information. Since I dont usually use large amounts of H2SO4 in short times, and dont make anything illegal with it, thats suits my experimental needs very well.

However, before 18 (actually before 16), I have boiled battery acid to obtain my own "98%" sulfuric (which was at most 85-95%). This is, IMHO, the only real way of "making" large amounts of concentrated H2SO4 by your own. The greatest drawback is that is an extremely dangerous operation, I and many members cant stress enough how dangerous is that.. Cold concentrated sulfuric can blind you instantly, so you can imagine what >300°C concentrated could do to your skin/face.. This without mention the SOx/H2SO4 nasty fumes.. Today Im much older and think that I dont need to expose myself with this kind of danger.

Stick with entropy51 idea. It dont worth the effort to make your own H2SO4, except only if you really cant find it (which is probably impossible.. Otherwise, you havent looked hard enough). Try to convince your parents or older friends to buy it for you (remember, if you have on hands a bottle of H2SO4 , treat it with greatest care.. It will not be nice if have a chance to hit you).

For me, the idea of FeSO4 pyrolisis being useful is just one: oleum. Oleum, unlike normal H2SO4 is very hard to come across and is a very valuable chemical, and anhydrous ferrous sulfate could be put in use to make it instead of plain sulfuric.. (even though, better options for SO3 generation are up to amateur reach and are already discussed in the forum), just bubbling SO3 in conc. H2SO4.

[Edited on 23-1-2012 by Aqua_Fortis_100%]

Actually, my parents are considering to buy things off eBay for me. So maybe in the not-so-distant future, I will get it. You are right about hot, concentrated sulfuric acid being dangerous, and I will take all the precautions necessary. SO3 is even more dangerous than simply H2SO4, and can dehydrate/carbonize your skin instantly. Also, would an international shipping of something hazardous, like sulfuric acid, be extremely expensive? Chemical suppliers don't cater to individuals here.

entropy51 - 22-1-2012 at 21:22

Oh well, just asking. If you people won't help, then I'll look for them on my own like I've always did. Simple as that. And no, I'm not complaining.

I don't know what you mean by "You people won't help." We can't buy chemicals for you. If you were in the US, we could refer you to all kinds of suppliers. But you aren't, and we can't. I have offered all the advice that I can. And you do sound like a lot of the kewls who visit here and bitch about not being able to obtain chemicals even though you are "entitled" to .

To tell the truth, the more you reveal about your self, the more I doubt that you could safely handle nitric acid.

weiming1998 - 22-1-2012 at 21:57

I really think I should stop talking to you now.

Also, sorry for sounding a bit "rude" or begging for help. I didn't mean to word it like that. Finally, no matter what you or anybody else thinks, I am not a k3wl that wants to blow things up with home-made explosives.

[Edited on 23-1-2012 by weiming1998]

[Edited on 23-1-2012 by weiming1998]

AJKOER - 22-1-2012 at 22:43

For method 3, wouldn't the H2SO4 formed immediately react with the Fe(OH)3?

First, here is another reference: "Hydrometallurgy in extraction processes", Volume 1, by C. K. Gupta, T. K. Mukherjee, pages 59 to 60:

"In some operations, the ferrous sulfate solution is left in ponds exposed to the air for several months so that oxidation and hydrolysis take place, thus regenerating the acid. The overall reaction is:

4 FeSO4 + O2 + 10H2O --> 4 Fe(OH)3 + 4 H2SO4 (22)

The oxidation of the ferrous sulfate is catalyzed by the autotrophic microorganism Thiobacullis ferrooxidans."

LINK:
http://books.google.com/books?id=F7p7W1rykpwC&pg=PA60&lpg=PA60&dq=4+FeSO4+%2B+O2+%2B+10+H2O+4+Fe(OH)3+%2B+4+H2SO4&source=bl&ots=fi SSl16x5f&sig=VsoTmxn7wDiLNncwWFwZPJB0SYU&hl=en&sa=X&ei=O_ocT-2DDefj0QHM1JixCw&sqi=2&ved=0CGkQ6AEwCQ#v=onepage&q=4%20FeSO4% 20%2B%20O2%20%2B%2010%20H2O%204%20Fe(OH)3%20%2B%204%20H2SO4&f=false

Note, the equation is nearly the same as the one I quoted. As I suspected, there is a time element and to my surprise, even a bacterial catalysis.

To answer your good question, I suspect that the jelly like Fe(OH)3 is physically separated from the H2SO4, although there could be, per your point, some Iron sulfate at the boundary itself, depending on the acid strength of the solution.

weiming1998 - 22-1-2012 at 23:02

Now I get why a freshly-made solution of FeSO4 is green but turns yellow over time. It's the gradual formation of Fe(OH)3!

Also, this reaction FeSO4 + 2 NaCl ---> Na2SO4 + FeCl2 works because FeCl2 is much more soluble in cold water than Na2SO4. So, in theory, this reaction (FeSO4 + 2 HCl ---> FeCl2 + H2SO4) wouldn't occur, because:
1, H2SO4 is a stronger acid
2, FeCl2 cannot precipate without his happening: FeCl2+H2SO4===>2HCl+FeSO4, driving the reaction back where it started.
3, The volatility of HCl is significantly higher than H2SO4
If this reaction really occurs, then it would be in an equilibrium, and a certain condition is required to prevent FeCl2 from reacting with the H2SO4.

AJKOER - 23-1-2012 at 00:00

Quote: Originally posted by White Yeti

Method 2. By reacting with HCl (or other strong acid):

FeSO4 + 2 HCl ---> FeCl2 + H2SO4

OK, here is a YouTube on FeSO4 + 2 HCl:

http://www.youtube.com/watch?v=MDm37HPc1X8

Yes, I agree with you on the FeSO4 + 2 NaCl reaction, and the comments at:

http://www.mychemistrytutor.com/questions/3197

However, this spectator reaction is presented in the context of:

FeSO4 + 2 NaOCl ---> Fe(OCl)2 + Na2SO4

to possibly explain:

FeCl2 + Fe(OCl)2 ---> 2 Cl2 (g) + 2 FeO

which actual makes sense in water as:

2 Cl2 + 2 H2O <---> 2 HCl + 2 HOCl

and upon adding in the FeO and grouping terms:

2 HCl + FeO ---> FeCl2 + H2O

2 HOCl + FeO ---> Fe(OCl)2 + H2O

(this last equation being the classic preparation of a hypochlorite by direct reaction with HOCl although even the temporary existence of Iron hypochlorite is debatable, but would be congruent with the known observed stability issues associated with Aluminum hypochlorite and also Zinc hypochlorite)

It would even appear that the following speculated equation is reversible:

FeCl2 + Fe(OCl)2 + 2H2O <---> 2 Cl2 + 2 FeO + 2 H2O

in a closed system.

Note, the evolution of Chlorine is definitely observed when FeSO4 is added to Bleach, and I am just speculating on possible paths. The fact that the literature appears to be uncomfortable with speculating on reaction products and paths here (and similarly in the case of the reaction of Fe in dilute HOCl forming Cl2 and FeCl3) leaves it to the practicing chemist to take his best guess (please correct me if I have mis-spoke).

weiming1998 - 23-1-2012 at 01:02

By watching the video you've posted, I think I know what's happening. What's happening is that the Fe+,SO4-2, Cl- and H+ are all floating in the solution. Because they are all free ions, some H2SO4 or FeCl will be formed. But when you boil down the solution, the equilibrium shifts back to HCl and FeSO4 because HCl is more volatile. It won't even work if you try to dehydrate it at room temperature, because the bonds connecting Fe to SO4 is stronger than the bonds connecting it to Cl-, and the ion rejoins with the SO4-2. However, if you place some CaCO3 in the solution, some CaSO4 precipate will form, evidence that sulfuric acid exists in the solution.

The reason why FeSO4 is oxidized and hydrolyzed to Fe(OH)3 and H2SO4 is also because of this. As FeSO4 is a very weak base and also insoluble in water, It will precipate, driving the reaction forward. This can only form small amounts of H2SO4, as small amounts doesn't react with the Fe(OH)3/reacts so slowly that the oxidation and hydrolysis of FeSO4 is faster. However, once the concentration of H2SO4 gets to a certain point, the speed of acid-base neutralization and the hydrolysis of FeSO4 are the same, enabling no more H2So4 to form. If too much H2SO4 is in the solution, it will revert back to FeSO4 and H2O until it reaches this point, where both reaction stops.

blogfast25 - 23-1-2012 at 07:39

Is it academic interest? Or what? I hate to use the T word, but your posts kinds of have that aroma.

You don't say...

'AJKOER's world of chemistry' revolves aroung hypochlorite and the mysterious 'NH4OH'.

AJKOER - 23-1-2012 at 15:29

blogfast25:

FYI, some of my recent threads and responses had included with specific reference to the halogen family only, the following compounds:

Cl2, SOCl2, HCl + H2O2, HOCl, Cl2O, ClO2, NaCl, NaClO, CaCl2, Ca(OCl)2, Al(ClO)3, NaClO2, NaClO3, Mg(ClO)2, Mg(ClO3)2, Zn(ClO)2, KCl, KClO3, Zn(ClO3)2, FeCl2, FeCl3, Fe(OCl)2, HClO3 and even HClO4 (a world patent describing preparation of Perchloric acid via electrolysis starting with chloride fee HOCl).

With respect to Iodine:

I2, HI, CaI2, KI, HIO, HIO3, Ca(IO3)2 and KIO3.

so one would think you would have second thoughts on your comment "AJKOER's world of chemistry' revolves aroung hypochlorite and the mysterious 'NH4OH' " (surprise me!).

With respect to your academic/theoretical issue with this thread, with all the equations and postulates I have put forth, one would think you would offer at least one counter point or reference an error (surprise me!).

[Edited on 24-1-2012 by AJKOER]

AJKOER - 23-1-2012 at 20:08

The reason why FeSO4 is oxidized and hydrolyzed to Fe(OH)3 and H2SO4 is also because of this. As FeSO4 is a very weak base and also insoluble in water, It will precipate, driving the reaction forward. This can only form small amounts of H2SO4, as small amounts doesn't react with the Fe(OH)3/reacts so slowly that the oxidation and hydrolysis of FeSO4 is faster. However, once the concentration of H2SO4 gets to a certain point, the speed of acid-base neutralization and the hydrolysis of FeSO4 are the same, enabling no more H2So4 to form. If too much H2SO4 is in the solution, it will revert back to FeSO4 and H2O until it reaches this point, where both reaction stops.

Actually, perhaps not as those little microbes essentially may be driving the process and actually creating a fairly acidic product. To quote:

"Thiobacillus ferrooxidans is the most common type of bacteria in mine waste piles. This organism is acidophilic (acid loving), and increases the rate of pyrite oxidation in mine tailings piles and coal deposits. It oxidies iron and inorganic sulfur compounds. The oxidation process can be harmful, as it produces sulfuric acid, which is a major pollutant."

Also:

"This genus is thermophilic, preferring temperatures of 45-50 degrees Celsius. In addition, this is an acidophilic genus, preferring a pH of 1.5 to 2.5. A few species, however, only grow in a neutral pH."

And:

"Thiobacillus are strictly aerobic bacteria. All species are respiratory organisms.
Thiobacillus are obligate autotrophic organisms, meaning they require inorganic molecules as an electron donor and inorganic carbon (such as carbon dioxide) as a source. They obtain nutrients by oxidizing iron and sulfur with O2."

The way I interpret the last sentence is that the bacteria are using the oxygen, by assimilation and employing it as a tool, to create nutrients and H2SO4. Otherwise, O2 would not be reacting with the FeSO4.

Source: MicrobeWiki, the student-edited microbiology resource.

LINK:
http://microbewiki.kenyon.edu/index.php/Thiobacillus

weiming1998 - 24-1-2012 at 03:17

Hang on, this article says that this bacteria oxidizes Fe2+ to Fe3+ using oxygen and H+ ions! H+ ions? It would be consuming the H2SO4 formed by the oxidation of FeSO4. The formula would be this: 4FeSO4+2H2SO4+O2===>2Fe2(SO4)3+2H2O! It consumes the H2SO4 formed, not create it! So after a long time, you would end up with a solution of mainly Fe2(SO4)3, with traces of H2SO4 and Fe(OH)3. The solution would likely to end up less acidic than before.

Neil - 24-1-2012 at 04:59

There are bacteria that consume sulfide and produce sulfate, the blue skin on stagnant water which you will find in healthy bogs is the residue of such bacteria. Incredibly low pH levels can be utilized by these little guys. The masses of iron oxide/hydroxide they leave behind is known as 'bog iron'.

By some mechanism, ether very low pH or other, I have seen a lot of evidence that these colonies are able to literally dissolve rocks and sand.

got fools gold? got to much time on your hands? got a nearby bog?

weiming1998 - 24-1-2012 at 05:37

There are bacteria that consume sulfide and produce sulfate, the blue skin on stagnant water which you will find in healthy bogs is the residue of such bacteria. Incredibly low pH levels can be utilized by these little guys. The masses of iron oxide/hydroxide they leave behind is known as 'bog iron'.

By some mechanism, ether very low pH or other, I have seen a lot of evidence that these colonies are able to literally dissolve rocks and sand.

got fools gold? got to much time on your hands? got a nearby bog?

If my formulas are correct, only FeS2 would be oxidized by the bacteria by:
2FeS2+2H2O+7O2==bacteria===> 2FeSO4+ 2H2SO4. FeSO4 would further be oxidized/hydrolyzed into Fe(OH)3 and more H2SO4. But this time, the concentration of H2SO4 would increase dramatically because the bacteria produces H2SO4, which can, in theory, increase indefinitely until it gets too acidic and bacteria dies. But FeS, which is more available to me and easily made by iron powder+sulfur won't work. It would be: FeS+2O2===>FeSO4 . Anyone got an idea of the artificial synthesis of fool's gold?

AJKOER - 24-1-2012 at 06:00

Actually, I think both Neil and Weiming1998 are correct. I have seen the reaction written to "completion". Here is an example, based on a quick search:

12FeSO4 +3 O2 +6 H2O = 4Fe2 (SO4) 3 +4 Fe (OH) 3

LINK (hit translate unless you do speak Chinese):
http://wenwen.soso.com/z/q168566732.htm

But, the partial reaction (to H2SO4) does occur as I have seen the reaction (see above citation in "Hydrometallurgy in extraction processes" by C. K. Gupta, T. K. Mukherjee, page 60), and also H2SO4 production is clearly cited in the article as an environmental hazard.

Remember the reaction time is in months!

[Edited on 24-1-2012 by AJKOER]

Neil - 24-1-2012 at 06:08

From wiki

"Pyrite has been used since classical times to manufacture copperas, or iron sulfate. Iron pyrite was heaped up and allowed to weather as described above (an early form of heap leaching). The acidic runoff from the heap was then boiled with iron to produce iron sulfate. In the 15th century, such leaching began to replace the burning of sulfur as a source of Sulfuric Acid. By the 19th century, it had become the dominant method."

For our next trick we will knock the edges off a square and attach it to a cart.

Maybe I'm daft but I thought that heating Fe + 2S gave FeS2

The bacteria that use sulphate as an energy source leave behind the iron hydroxides, they live on the wastes of bacteria that feed actual sulfides.

Edit: http://technology.infomine.com/enviromine/ard/Microorganisms...

[Edited on 24-1-2012 by Neil]

Neil - 24-1-2012 at 06:16

Also, if I tried to talk about those pool workers about how I am doing home chemistry and need sulfuric acid from the low- fume pool acids, then they will either:
.Sell nothing to me
.Sell things to me and act normal, then ring the police.
.Sell nothing to me AND kick me out of the shop.
.Sell nothing to me, then ring the police.
. Actually having a discussion, and selling me SOME of the things.
. Lecture about how these things are dangerous and telling me to go play/get a girlfriend (like you people have done)

"Hello kind sirs, I like working with copper and making copper inventions/jewellery/art and I really need some sulphuric acid for my pickling bath."

Also check around for jewellery supply shops and welding shops. Look for pickling solution.

weiming1998 - 24-1-2012 at 06:26

From wiki

"Pyrite has been used since classical times to manufacture copperas, or iron sulfate. Iron pyrite was heaped up and allowed to weather as described above (an early form of heap leaching). The acidic runoff from the heap was then boiled with iron to produce iron sulfate. In the 15th century, such leaching began to replace the burning of sulfur as a source of Sulfuric Acid. By the 19th century, it had become the dominant method."

For our next trick we will knock the edges off a square and attach it to a cart.

Maybe I'm daft but I thought that heating Fe + 2S gave FeS2

The bacteria that use sulphate as an energy source leave behind the iron hydroxides, they live on the wastes of bacteria that feed actual sulfides.

FeS2 is formed by Fe/oxides+S under pressure. So it is very common in the earth's crust, but not very easily synthesized. The amount of H2SO4 formed would be consumed in two side reactions: The reaction: 4FeSO4+ 2H2SO4+ O2?===> 2Fe2(SO4)3+2H2O will consume some of the reactants, unless Fe2(SO4)3 is already formed by the bacteria by 4FeS2+15O2+2H2O===>2Fe2(SO4)2+ 2H2SO4 (so hard to balance these equations!) Then the bacteria consumes the iron sulfate by Fe2(SO4)3+6H2O====> Fe(OH)3+3 H2SO4. The other possible reaction would have numerous side-reactions destroying the H2SO4.

weiming1998 - 24-1-2012 at 06:28

I suck at lying though. I even got nervous last time buying the pool chemicals, and all I had to say was "yeah!" and listing some of the uses for the chemical.

AJKOER - 24-1-2012 at 06:29

OK, so heap leaching as a mode to H2SO4, given its historic significance, is a valid method of Sulfuric acid production.

Perhaps, however, not the 1st to be recommended and yes, we are indeed turning back the clock!

weiming1998 - 24-1-2012 at 06:31

Actually, I think both Neil and Weiming1998 are correct. I have seen the reaction written to "completion". Here is an example, based on a quick search:

12FeSO4 +3 O2 +6 H2O = 4Fe2 (SO4) 3 +4 Fe (OH) 3

LINK (hit translate unless you do speak Chinese):
http://wenwen.soso.com/z/q168566732.htm

But, the partial reaction (to H2SO4) does occur as I have seen the reaction (see above citation in "Hydrometallurgy in extraction processes" by C. K. Gupta, T. K. Mukherjee, page 60), and also H2SO4 production is clearly cited in the article as an environmental hazard.

Remember the reaction time is in months!

[Edited on 24-1-2012 by AJKOER]

So the Fe2(SO4)3 and H2SO4 generating reaction might be a side reaction! But the large amounts of Fe(OH)3 would destroy it easily since that is the main reaction. I would end up with a pool of Fe2(SO4)3, Fe(OH)3 and dead bacteria.

Neil - 24-1-2012 at 06:35

Why lie? As a home chemist you are inevitably going to want to make things that real labs can order. Making a still or copper retort is fun and you actually will want a pickling bath.

Besides when you deside to take a break from chemistry to focus more of your attention on girls, being able to make jewellery is helpfull. Learning to work a metal means learning the metals chemistry.

weiming1998 - 24-1-2012 at 07:22

Why lie? As a home chemist you are inevitably going to want to make things that real labs can order. Making a still or copper retort is fun and you actually will want a pickling bath.

Besides when you deside to take a break from chemistry to focus more of your attention on girls, being able to make jewellery is helpfull. Learning to work a metal means learning the metals chemistry.

A retort on any kind, especially a thermal shock-resistant copper retort, will be very useful to me to collect heated gases. But don't you need a very hot furnace to shape the molten copper metal? Focusing my attention on girls is doomed to fail, as I do not know all the complex dating strategies/ I am not interested in hobbies a normal male would be interested in, like sports. Also, I like to talk about chemistry and other technicalities. So no.

And a sulfuric acid pickling bath might be useful from isolating noble/semi noble metals from a mixture of metals.

Neil - 24-1-2012 at 07:30

We are getting a bit off topic

AJKOER - 24-1-2012 at 08:23

With reference to Neil point on making things that real labs could order and to refocus on the thread, employing FeSO4, may I suggest the following path to pure H2SO4:

1. React FeSO4 with Bleach (NaClO) to produce Cl2 (conform with local laws).

2. Create Chlorine water (use freshly boiled distilled H2O to remove CO2).

3. Add SO2 to the Chlorine water. Reaction:

H2SO3 + Cl2 + H2O --> H2SO4 + HCl

or prepare SO2Cl2, a colorless fuming liquid. Sulfuryl chloride can be synthesized by combing SO2 and Cl2 in sunlight, or in the presence of camphor or activated charcoal as a catalyst:

SO2 + Cl2 --Catalyst--> SO2Cl2

adding water forms Chlorosulfuric acid and Hydrogen chloride:

SO2Cl2 + H2O --> HSO3Cl + HCl

and adding more water forms Sulfuric acid and Hydrogen chloride:

HSO3Cl + H2O --> H2SO4 + HCl

Actually, this is oxidation via HOCl as:

Cl2 + H2O <==> HOCl + HCl

and if we add a suspension of Silver acetate (dissolve Ag in vinegar and dilute H2O2, evaporate), then only HOCl remains in solution as the Silver chloride and unreacted Silver acetate are insoluble.

Also, even more potent than HOCl is Cl2O.

weiming1998 - 24-1-2012 at 16:42

The problem is, using the SO2+Cl2 method can only produce a weak sulfuric acid/hydrochloric acid solution. When you try to boil it down, it decomposes (H2SO3 + Cl2 + H2O <--> H2SO4 + HCl) This reaction is at an equilibrium so when the concentration of the acids gets up, it decomposes. That's what happened last time I tried it.

The silver acetate is a quite an interesting method to get HClO, but my mum would definitely not like to see her precious silver jewelry
dissolved in vinegar/H2O2!

[Edited on 25-1-2012 by weiming1998]

AJKOER - 24-1-2012 at 17:03

weiming1998:

You are right, an oversight on my part:

H2SO3 + Cl2 + H2O <--> H2SO4 + HCl

leads to a dilute acid solution, but the Sulfuryl chloride method, also cited, is capable of producing more concentrated products.

weiming1998 - 24-1-2012 at 17:23

How big of a container would I need to contain enough SO2/Cl2 to make even 10mls of sulfuryl chloride?

AJKOER - 24-1-2012 at 19:58

Density SO2Cl2: 1.6674

My calculations suggest that using a 2 liter bottle, you could produce 3.6 ml of SO2Cl2.

However, SO2 boils at -10.0 C, and a solution of 23% salt and water can achieve much colder temperature (-51 C), so one may be able to liquefy it. You may even consider combing the liquid SO2 with Chlorine hydrate (see discussion below).

Chlorine apparently forms a solid hydrate precipitate (Cl2 7.3H2O). To quote: "If chlorine gas is passed into a dilute solution of CaCl2 at about 0 °C, greenish, feathery crystals appear that can be removed from the solution, dried, and kept indefinitely at room temperature. " LINK: http://mysite.du.edu/~jcalvert/econ/hydrates.htm

I also recall reading in a Sciencemadness thread, that CaCl2/ Chlorine water is the preferred method to actually obtain Chlorine hydrate crystals.

Another older source:"Chlorination of water" by Joseph Race, page 103:

"Although chlorine water appears to be of little value because of its instability there appears to be no reason why chlorine hydrate should not be successfully employed. The hydrate was first prepared by Faraday [9] by passing chlorine into water surrounded by a freezing mixture. A thick yellow magma resulted from which the crystals of chlorine hydrate were separated by pressing between filter paper at 0° C. .......
Chlorine hydrate separates into chlorine gas and chlorine water at 9.60 C. in open vessels and at 28.70 C. in closed vessels. "

Another source: "Treatise on general and industrial inorganic chemistry", Volume 1, by Ettore Molinari, page 122:

"The system chlorine-water exhibits the following interesting cases: saturation of water at 0° with gaseous chlorine leads to the separation of greenish crystals of chlorine hydrate (Cl2 + 8H20), these being stable only up to the temperature 9.6°, above which they are decomposed into gaseous chlorine and water saturated with chlorine; this hydrate is, however, stable at higher temperatures if the pressure is raised and, on the other hand, decomposes below 9° if the pressure diminishes. As a rule there are three phases for chlorine hydrate, but at a certain definite point, namely, — 0.26°, four phases are possible, i.e., ice, chlorine hydrate, aqueous chlorine solution, and gaseous chlorine. The equilibrium is, however, easily altered, and at the least rise of temperature the ice disappears, whilst lowering of the temperature causes transformation of the aqueous chlorine solution into ice and chlorine hydrate and the consequent suppression of one of the four phases; this equilibrium depends also on the pressure (24.4 cm.), and if this increases, the gaseous chlorine disappears, whilst if it diminishes, another phase (the ice or the hydrate) disappears."

[Edited on 25-1-2012 by AJKOER]

weiming1998 - 24-1-2012 at 20:36

It would be difficult to keep the solution cold, considering how hot is is now in Perth
(42 degrees celsius at noon!) But with a refirigerator, some planning and some time, I might be able to do it. Also, would a solution of SO2, put in a freezer (under -18 degrees celsius) be able to liquefy the SO2?

blogfast25 - 25-1-2012 at 06:39

With reference to Neil point on making things that real labs could order and to refocus on the thread, employing FeSO4, may I suggest the following path to pure H2SO4:

1. React FeSO4 with Bleach (NaClO) to produce Cl2 (conform with local laws).

2. Create Chlorine water (use freshly boiled distilled H2O to remove CO2).

3. Add SO2 to the Chlorine water. Reaction:

H2SO3 + Cl2 + H2O --> H2SO4 + HCl

or prepare SO2Cl2, a colorless fuming liquid. Sulfuryl chloride can be synthesized by combing SO2 and Cl2 in sunlight, or in the presence of camphor or activated charcoal as a catalyst:

SO2 + Cl2 --Catalyst--> SO2Cl2

adding water forms Chlorosulfuric acid and Hydrogen chloride:

SO2Cl2 + H2O --> HSO3Cl + HCl

and adding more water forms Sulfuric acid and Hydrogen chloride:

HSO3Cl + H2O --> H2SO4 + HCl

Actually, this is oxidation via HOCl as:

Cl2 + H2O <==> HOCl + HCl

and if we add a suspension of Silver acetate (dissolve Ag in vinegar and dilute H2O2, evaporate), then only HOCl remains in solution as the Silver chloride and unreacted Silver acetate are insoluble.

Also, even more potent than HOCl is Cl2O.

This just HAS to get a prize, IMHO…

What about :

‘Most far-fetched, most untested, most wasteful, least needed synth EVER!’, worthy possibly of an ‘Ig-noble Prize!’

But I appreciate your attention to detail: ‘conform with local laws’ was priceless; do you normally disregards local laws?

You’re not a chemist, you’re a pseudo-scientific contortionist! But as long as it’s got bleach in it…

weiming1998 - 25-1-2012 at 07:01

Quote: Originally posted by blogfast25

The method is far from useless, for me, if I can drive the Cl2/HCl out of the solution without destroying the SO2/H2SO4. But I don't understand the reason for making Cl2 from FeSO4. Why not from NaHSO4, which is cheaper?

ScienceSquirrel - 25-1-2012 at 07:36

This is not going to make significant quantities of sulphuric acid, pure or otherwise.

AJKOER - 25-1-2012 at 09:52

Actually, I think you guys have done it this time.

Science Squirrel and Weiming1998, please do yourself a big favor and take anything blogfast25 says lightly.

In fact, SO2Cl2 is not my favorite either, but in a prior hypochlorite method to H2SO4, it received a favorable comment by Woelen as having the potential to produce something other than dilute H2SO4. Excuse me for including it.

Please look it up yourselves and just ignore blogfast25 (no adjectives necessary).

Now, if someone has an issue with the synthesis of chlorosulfuric and sulfuric acid, please address your comments to my source: "Concise Encyclopedia Chemistry" by Walter de Gruyter, 1994, page 1058.

If you all don't have enough mad on your faces, I know of a good source for Black Sea mud (really, it is a big seller here!)

[Edited on 25-1-2012 by AJKOER]

White Yeti - 25-1-2012 at 10:01

Ahem....
What blogfast25 is saying is very reasonable and there is no reason to brush him aside.

It's true that in all your posts hypochlorite seems to be the fix for everything. Give it up, what do you need sulfuric acid for anyway?

If you really want to make sulfuric acid, set up a big heated pressure vessel loaded with V2O5 garden sulfur, and some air.

Fire it up and see what happens, your yields will be much better than through any other process (except maybe the lead chamber process).

There's no need to re-invent the wheel. For most purposes, you can swap sulfuric acid for another acid.

rstar - 25-1-2012 at 10:02

How about electrolysis of FeSO4 solution ?

A similar NurdRage video <a href="http://www.youtube.com/watch?v=5dUSF9Gl0xE"> here </a>

White Yeti - 25-1-2012 at 10:04

Quote: Originally posted by rstar

How about electrolysis of FeSO4 solution ?

A similar NurdRage video <a href="http://www.youtube.com/watch?v=5dUSF9Gl0xE"> here </a>

Won't work, iron re-dissolves in sulphuric acid.

Neil - 25-1-2012 at 11:19

Electrolysis is used to regenerate sulfuric acid in titanium refining from the otherwise wasted Iron sulfate.

They use membranes, pH control and fancy anode cathode materials to plate iron out of the solutions.

Neil - 25-1-2012 at 11:35

Actually, I think you guys have done it this time.

Science Squirrel and Weiming1998, please do yourself a big favor and take anything blogfast25 says lightly.

In fact, SO2Cl2 is not my favorite either, but in a prior hypochlorite method to H2SO4, it received a favorable comment by Woelen as having the potential to produce something other than dilute H2SO4. Excuse me for including it.

Please look it up yourselves and just ignore blogfast25 (no adjectives necessary).

Now, if someone has an issue with the synthesis of chlorosulfuric and sulfuric acid, please address your comments to my source: "Concise Encyclopedia Chemistry" by Walter de Gruyter, 1994, page 1058.

If you all don't have enough mad on your faces, I know of a good source for Black Sea mud (really, it is a big seller here!)

[Edited on 25-1-2012 by AJKOER]

What Blogfast is errr 'alluding' to is that your posts are a bit more mental masturbation then useful chemistry. They are interesting, or sometimes loony, but not useful. While that may have been a bit blunt of him, his point about the applicability of the synthesis is valid.

The cost of getting NaClO from bleach makes it fiscally impractical for use in any synthesis, even straightforward ones like haloform reactions.

While the interactions of bleach are interesting they are, at the costs and grades available OTC, useless as anything other then a novelty.

I do find the idea of making sulphuric from fools gold interesting, but I find a perverse joy in making things out of illogical baser ingredients.

blogfast25 - 25-1-2012 at 11:56

AJKOER:

Really you should start with some simple experiments and a basic text book. Get some joyful mileage out of these. Instead you concoct the most bizarre attempts at ‘synthesis’ (this one not being the first by a long shot and nor will it be the last in all sad likelihood), thrown together from quoted sources you clearly either not understand or pull out of context.

As regards mudslinging, for the most part I ignore your scribbles because rebutting them could cost hours of my time I’d never get back… But you are seriously getting on people’s nerves here, it really isn’t just me alone.

Now seriously, go and BUY some concentrated H2SO4 drain unblocker, some base chemicals and do some basic stuff. Well reported even the simplest of things are a joy to read. Not so your pseudo-chemical houses of cards in the sky…

Thanks to you I may never be able to look at a bottle of bleach in the same way…

[Edited on 25-1-2012 by blogfast25]

neptunium - 25-1-2012 at 12:00

so i guess that closes the topic??? moderator ?

AJKOER - 25-1-2012 at 12:03

A little off topic, but my favorite oxidizers for the home chemists is H2O2 + Ferrous salt, or Sodium percarbonate (in the presence of a catalyst like Na2CO3, Zn dust,...), which appears to enter the realm of Fenton chemistry with the production of the free hydroxyl radical.

Next is Cl2O many times more reactive than HOCl (example, HOCl oxidizes Sulfur safely, while Cl2O and dry S is reportedly explosive). The home chemist can make it by dehydrating concentrated HOCl (from repeated distillations of half a solution of the starting HOCl) with Ca(NO3)2. Also with more effort, by passing Cl2 over heated moist NaHCO3 and collecting the dried Cl2O gas in CCl4 (cooled for safety and avoid strong light), or better, not stored but used immediately given its explosive nature with exposure to heat, organic substances, light and shock.

Next is HOCl due to it easy preparation from inexpensive household ingredients.

Next is Cl2, then H2O2, Ferrates and finally HIO3.

Note, the list is short as many substances are not normally available at home or, subject to security questions, before being able to purchase.

[Edited on 25-1-2012 by AJKOER]

AJKOER - 25-1-2012 at 12:32

Quote: Originally posted by blogfast25

AJKOER:

Now seriously, go and BUY some concentrated H2SO4 drain unblocker, some base chemicals and do some basic stuff. Well reported even the simplest of things are a joy to read. Not so your pseudo-chemical houses of cards in the sky…
[Edited on 25-1-2012 by blogfast25]

blogfast25: you are right, many of us can still go and buy H2SO4, for now. I have noticed, but I could be wrong, that my bathroom bowl cleaner has gone form 5% HCl to none. Perhaps part of the green revolution or not.

I viewed this thread more of a puzzle to be solved (I like puzzles). That is, given FeSO4 and some standard available chemicals, make some dilute H2SO4 (the more concentrated and purer the better).

May the best garage chemist win!

[Edited on 25-1-2012 by AJKOER]

Neil - 25-1-2012 at 13:04

I think Garage Chemist wins that award (and then some) with his manufacture of Oleum.

Start playing with HClOx in real life and not on paper and see if they are still your favourite

Quote: Originally posted by blogfast25

Well reported even the simplest of things are a joy to read.

[snip]

Thanks to you I may never be able to look at a bottle of bleach in the same way…

[Edited on 25-1-2012 by blogfast25]

Amen.
and
Amen.

[Edited on 25-1-2012 by Neil]

White Yeti - 25-1-2012 at 13:06

Puzzle solved:
http://www.sciencemadness.org/talk/viewthread.php?tid=2824

weiming1998 - 25-1-2012 at 15:37

Quote: Originally posted by White Yeti

Puzzle solved:
http://www.sciencemadness.org/talk/viewthread.php?tid=2824

Or not really.

That requires nitrate salts, remember, and people like me can't find them. Read this http://www.sciencemadness.org/talk/viewthread.php?tid=18588

But I got to admit, why use such an inefficient process as to use bleach on FeSO4? The NaOH in bleach should inhibit chlorine production. Why not use NaHSO4 solution and Ca(ClO)2 to generate chlorine gas? Why waste the time using a giant pressurized vessel (which most people don't have) and advanced catalysts, just to make a few tens of mls of SO2Cl2? If you had the condition to concentrate galleons of SO2 and Cl2, I'm sure you can find H2SO4. In fact, if you had that, I'm pretty sure you can find every chemical that is not a drug precursor.

Now, an idea to drive off HCl/Cl2 in a SO2/Cl2 aqueous solution without driving the SO2 off/decomposing the H2SO4?

Neil - 25-1-2012 at 18:16

http://www.sciencemadness.org/talk/viewthread.php?tid=2824&a...

The Lead chamber thread covers ideas for nitrate free runs and has ideas on sulfuric from Iron sulfates, as linked to above.

AJKOER - 25-1-2012 at 22:06

Actually the Lead Chamber thread on pages 8 & 9 says some very good things about the H2O + H2SO3 + Cl2 ==> H2SO4 + 2HCl route (on the middle of page 8 starting with un0me2), which is only proceduraly different from adding water to SO2Cl2. That is, (H2O + SO2) + Cl2 vs. H2O + (SO2 + Cl2). If adding Cl2 to H2SO3 is better, so be it. For the record, I did mention both of these routes, with SO2Cl2 mentioned last, but for those with short term memory, here is an extract:

Here is a convenient link to the Lead Chamber thread:

http://www.sciencemadness.org/talk/viewthread.php?tid=2824&a...

In particular on page 9 of Lead Chamber thread, and off topic for that thread, but on key for this one:

Quote: Originally posted by Formatik

References in Gmelin verify the reaction goes as thought: when SO2 and Cl2 are led into water, this exotherms a bit and accumulates the H2SO4 as the HCl concentration decreases. Neumann described the reaction is going rapidly and almost completely (95-100% theoretical amounts were converted), the sulfuric and hydrochloric acids result immediately as fine droplets/fog, these are difficult to absorb and also pass over, as gases and water initially interact.

The patent mentioned of Stolle, leads same parts SO2 and Cl2 into water, eventually raising the temperature to 250 deg., yielding 90% H2SO4 and conc., free from Cl2 and SO2, aqueous HCl. Neumann's process is much more descriptive.

Neumann also described despite having used a Cl2-excess, a significant amount of SO2 got solubilized in H2SO4, since SO2 solubility increases with H2SO4 concentration. Though experiments also showed conc. H2SO4 which had Cl2 or SO2 solubilized in it, after blowing in air for 15 minutes, were almost completely removed.

Quote: Originally posted by S.C. Wack

What does this have to do with the lead chamber process?

It seems this thread is the designated stickied sulfuric acid thread. I would retitle it as the sulfuric acid preparation thread, or remove the non-Chamber discussions and sticky those with said title instead. Good eye on that reference, I also found it through Gmelin.

Quote: Originally posted by 497

Do you think bubbling Cl2 + SO2 into water could get H2SO4 concentrated past azeotropic? I don't remember if there are any side reactions that could occur in a mixture of HCl, conc. H2SO4, Cl2, SO2 and H2O... I can't think of any off the top of my head, besides formation of sulfuryl chloride, but that should not occur without a catalyst right?

I doubt it's of concern. Neumann described that after the reaction heat slows down, that the gases come out ununited. This heat is especially large when water is first consumed in the reaction. Their later experiments used additional heat (60-92 deg), to make the reaction go much faster.

Concerning the concentration of H2SO4 obtained by combination of SO2 and Cl2 with H2O, Neumann says it is that of the Chamber acid or Glover acid (66-88%). That's the raw figure then, it can be concentrated further by regular means. For practical purposes, instead of H2O, conc. HCl was recommended. Then when a specific gravity of 1.6 is reached, the hydrochloric acid content has been nearly completely removed.

[Edited on 20-8-2010 by Formatik][/rquote

[Edited on 26-1-2012 by AJKOER]

AJKOER - 26-1-2012 at 08:00

One manner a garage chemist could implement the

H2O + H2SO3 + Cl2 ==> H2SO4 + 2 HCl

reaction with limited equipment in a safe manner is as follows:

1. Prepare Chlorine hydrate, Cl2 +7.3 H2O (note, for calculation purposes, the actual value is not 7.3, but 7.27 plus or minus .17). The recommended process is bubbling Cl2 into a cold (below 0 C) solution of CaCl2. My variation is to prepare HOCl (weak acid like H2CO3 to Ca(OCl)2/CaCl2 solution and filter). Then add an appropriately small amount of H2O2. Reactions:

Ca(OCl)2 + H2CO3 --> CaCO3 (s) + 2 HOCl

2 HOCl + H2O2 ---> HOCl + HCl + H2O + O2

HOCl + HCl <---> Cl2 + H2O

Now, freeze the Chlorine water and salt solution, and harvest the long green crystals (I done this). Let dry in the cold (stable at under 9 C in an open vessel). In a closed vessel, also stable at/near room temperature.

2. Fill a glass vessel with some SO2 (liquid or gas or use H2SO3) and add the Chlorine hydrate and seal. Place in sunlight and observe the reaction from a distance (ice bath may be necessary to control temperature).

3. Upon completion, add/remove H2O to obtain desired H2SO4 strength. Heat/bubble air into the solution to remove unwanted HCl.

[Edited on 26-1-2012 by AJKOER]

[Edited on 27-1-2012 by AJKOER]

White Yeti - 26-1-2012 at 09:46

Ca(OCl)2 + H2CO3 --> CaCO3 (s) + HOCl

2 HOCl + H2O2 ---> HOCl + HCl

reactions not balanced are

AJKOER - 26-1-2012 at 10:45

Thanks, I was in a hurry as I actual changed from NaClO to Ca(OCl)2 (equations now edited).

As a self criticism, I wonder if the Chlorine hydrate would be completely free of CaCl2 (as not, it would increase HCl and reduce the H2SO4 yield). One could dilute the solution, however, before cooling and retrieving the Cl2+7.3H2O, and thereby reduce the potential contamination. Given the reaction:

Cl2 + H2O <----> H(+) + Cl(-) + HOCl

as the solution starts off alkaline, dilution would lower the pH and move the reaction to the left (more Chlorine water).

Or, one could also re-dissolve the Chlorine hydrate in distilled water and attempt re-crystallization.

[Edited on 26-1-2012 by AJKOER]

White Yeti - 26-1-2012 at 11:45

...would increase HCl and reduce the H2SO4 yield.

if any at all

blogfast25 - 26-1-2012 at 12:45

3. Upon completion, add/remove H2O to obtain desired H2SO4 strength. Bubble air into the solution to remove unwanted HCl.

Really? You're gonna separate H2SO4/HCl by 'bubbling air through it'? You're quite the Chief Engineer!

AJKOER - 26-1-2012 at 13:33

blogfast25:

Please re-read my extract from the Lead Chamber thread above. Per Formatik:

"The patent mentioned of Stolle, leads same parts SO2 and Cl2 into water, eventually raising the temperature to 250 deg., yielding 90% H2SO4 and conc., free from Cl2 and SO2, aqueous HCl. Neumann's process is much more descriptive.

Neumann also described despite having used a Cl2-excess, a significant amount of SO2 got solubilized in H2SO4, since SO2 solubility increases with H2SO4 concentration. Though experiments also showed conc. H2SO4 which had Cl2 or SO2 solubilized in it, after blowing in air for 15 minutes, were almost completely removed."

So the combination of heat (unclear to me as to whether the source is only the reaction or external also) and air removes Cl2, SO2 and aqueous HCl yielding 90% H2SO4.

However, you may have a point and I have edited the process to include treating with heat and air.

[Edited on 27-1-2012 by AJKOER]

weiming1998 - 26-1-2012 at 17:23

Bubbling air/ozone through water is not going to remove HCl. What about silver sulfate in the solution? It is slightly soluble in water and will actually increase H2SO4 yields. But where do you get something like silver sulfate?

entropy51 - 26-1-2012 at 18:25

blogfast25:

Please re-read my extract from the Lead Chamber thread above.

As best I can discern, no one who posted in that thread actually made any H2SO4. I find it amusing, but not surprising, that you cite such balderdash as a reference.

No one who actually works with chemicals could actually believe that there is any practical utility in these Rube Goldberg schemes.

If you perform a six step synthesis, and the yield of each step is 50%, then the overall yield is 1.5%.

If you're so sure of yourself, why not perform the experiment and post the data? Otherwise this is all just brain farts.

AJKOER - 26-1-2012 at 18:29

As the H2SO4 becomes more concentrated, being a powerful dehydrating agent, we may have essentially hydrogen chloride in H2SO4. The combination of heat and aeration may be able to dislodge it from the H2SO4.

I do agree with weak H2SO4, no such luck.
-----------------------------------------------------------------------------

To prepare Ag2SO4, assuming no access to H2SO4, I would try Silver acetate (AgC2H3O2) in Acetic acid (where it is soluble) and add Epsom's Salt (MgSO4).

Possible reaction:

2 AgC2H3O2 + MgSO4 --> Ag2SO4 (s) + Mg(C2H3O2)2

where I suspect/hope that Silver sulfate does not dissolve in weak Acetic acid. Fortunately, per one source, it does dissolve in dilute H2SO4, making it possibly effective for our purpose.

Ag2SO4 + HCl --> 2 AgCl (s) + H2SO4

Assuming AgCl is not soluble in H2SO4.

[Edited on 27-1-2012 by AJKOER]

AJKOER - 26-1-2012 at 18:46

Entropy51:

I agree with you.

Until someone on Sciencemadness can successfully perform a synthesis, even on a micro level, lets honestly state it is just theoretically possible.

[Edited on 27-1-2012 by AJKOER]

White Yeti - 26-1-2012 at 18:52

Is it time yet for a moderator to lock this thread? The OP is deaf to any criticism, most of which is justified. We are just wasting breath and heading nowhere but deeper into the abyss of untruth and armchair speculation; an abyss to which there is no bottom

weiming1998 - 26-1-2012 at 19:01

Yes, I have nothing to say, because everything is speculation in this thread.

Bot0nist - 26-1-2012 at 19:06

Quote: Originally posted by White Yeti

Is it time yet for a moderator to lock this thread? The OP is deaf to any criticism, most of which is justified. We are just wasting breath and heading nowhere but deeper into the abyss of untruth and armchair speculation; an abyss to which there is no bottom

Crying out for a moderator in a thread does not signal them. If you have enough disdain for a thread to make multiple posts in it, just to call for it's closure, why not simply report the thread to the moderators. Unlike what your doing here, this will both call the moderators attention to the thread you have a problem with, and stop you from inflating your post count with useless posts. (since they don't aid in closing a problem thread, and they only increase it's size and bumps it up)

[Edited on 27-1-2012 by Bot0nist]

AJKOER - 26-1-2012 at 22:04

SUMMARY OF THREAD MAKING DILUTE (OR BETTER) H2SO4 FROM FeSO4

1. Using just FeSO4, and heating yields SO2 and SO3, an agreed upon path to H2SO4.

2. Direct oxidation of aqueous FeSO4 with air via microbes is a slow process (measured in months), but tried and true having been replaced by modern methods.

3. The oxidation of H2SO3 with Cl2 is less favored by some and proclaimed by others. Following the precise procedure as laid out in a patent is recommended. Good off topic discussion in the Lead Chamber thread.

4. The direct reaction of SO2 and Cl2 in the presence of a catalyst had even less appreciation over questions of volume of H2SO4 that could possibly be created. Use of cooled SO2 and/or Chlorine hydrate, as possible solutions, drew no support.

Please comment if you have a significantly different opinion on any of these points.

497 - 26-1-2012 at 22:31

Very high yields of SO3 are said to be obtained by pyrolysis of basic iron sulfates at easily manageable temperatures. Why worry about chlorine compounds?

weiming1998 - 26-1-2012 at 23:11

Easily manageable? That's 700-800 degrees, according to posts here! (Though Wikipedia lists the decomp temperature of FeSO4 at 480 degrees, which might be achievable by a stove) Also, if you want to make H2SO4 with SO3, you must first have H2SO4. you also need a complete glass distillation set, as SO3 will attack pretty much everything else. It is also dangerous to handle, will crisp your skin instantly when spilt, and reacts violently with water (I have never seen the reaction of SO3 and water though).

AJKOER - 27-1-2012 at 06:10

An important point per Weiming that applies for all of these synthesis.

H2SO4 and nearly all preparatory routes are inherently dangerous. Should never be attempted by the inexperienced and certainly without ample safety precautions.

weiming1998 - 27-1-2012 at 06:37

Even conc.H2SO4 or NaOH is nothing compared to SO3, in terms of corrosiveness. Acids attack skin by denaturing protein and absorbing all the water in cells, which is what makes H2SO4 much more corrosive than HCl, because it's hygroscopic! SO3 is almost as attracted to water as P2O5 and will suck every molecule of water out of a compound. It will do the same to your skin. Also, I hear people comparing the violent reaction of it in water with cesium in water, though I suspect it to be exaggeration. It does form a fog of concentrated H2SO4 when in contact with water though, which, if you breathe it in, damage your lungs irreversibly.

Overall, some route of making H2SO4 can be "quite safe" ( to responsible and careful people) But SO3 is not one of them, unless you have complex equipment. I tried to make SO3 with NaHSO4 before on a stove. It didn't work. I'm glad it didn't, and I will never work with SO3 unless I have all the protective equipment I need.

497 - 27-1-2012 at 19:10

Try like 500-600* and the option of doing it in the microwave may simplify things a lot further. Also, just because one creates SO3 does not mean they must literally handle it. There is no reason you couldn't devise a simple system to absorb it as it is produced, while you are nowhere near it.

weiming1998 - 27-1-2012 at 19:26

Would a complete distillation equipment fit inside a microwave? Maybe if you have a tiny distillation setup that makes about 1ml of SO3. Also, if microwaves can achieve temperatures like that, then food coming from it would be turned to charcoal. The simple system thing works, if you already has concentrated H2SO4. SO3's direct reaction with water might be sufficient enough to crack the glass container.

497 - 27-1-2012 at 21:59

Okay you obviously don't have near the level of understanding required to do such a thing without being spoonfed every detail.

For anyone else out there interested:
It is easy to pass a small tube through the side of a microwave oven, it is easy to achieve temperatures up to 1000* using iron oxides, and/or graphite susceptors in a home microwave.

If you can't figure out how to get a small amount of concentrated sulfuric to absorb the SO3, there is no hope, come back in a couple years when you have finished puberty.

Edit
Since the hematite produced by decomposition of Fe sulfates absorbs microwaves, there may not even be a need to use any additional susceptor. Possibly just a few seed particles would be needed at the beginning. This means you could avoid all the bulk heat transfer constraints usually present. It also means a simple borosilicate flask could be used. A layer of microwave nonabsorbent insulation may be easily sintered on the inside of the flask to shield the flask from direct contact with hot oxide particles. A paste made with fine silica should do the trick. Such an internally insulated, internally heated flask could be valuable for a number of high temp reactions that are normally done in quartz tubes and such.

[Edited on 28-1-2012 by 497]

weiming1998 - 27-1-2012 at 23:02

Obviously you don't have the brains to make SO3 either without hurting yourself if you have to resort to direct insults.

Graphite might work to transfer the heat to the NaHSO4/FeSO4, but even then I doubt it would get hot enough for it to decompose. Making a miniature model of a complex distillation kit in a microwaveand sticking holes in it is not what I, or most people want for a few mls of concentrated H2SO4. If you need SO3 for experiments that must require it, fine, but just for H2SO4, there are much better ways that don't require SO3, high temperatures, poking holes in microwaves, miniature distillation kits, and idiots using ad hominem.

497 - 28-1-2012 at 08:10

Actually, travelling in foreign countries for the last few months makes any experimentation impossible.

Do you have any evidence at all to back up a single one of the assertions you've made in the last couple posts?? Negative speculation like that is entirely useless. If you find any evidence that pertains to my posts, I would be interested to see it.

Why do you keep talking about complex mini distillation kits? Last I checked, a 250ml rbf would fit in a microwave. The whole point of what I said earlier was that there would be no need for a condenser at all! On top of the rbf you could put a right angle inlet adapter such as the ones here: http://unitedglasstech.com/Adapters.htm
Then a couple feet of teflon tube could lead from the adapter, outside the microwave directly in to a flask of H2SO4. Drilling a hole in a microwave is as easy as drilling a hole in any other sheet metal. It is not dangerous if the hole is fairly small, and you should be far away while it is running anyway.

You could be making more like liters a day of SO3, not a few ml. Just because you happen to not care about oleum/SO3 does not mean anything. Based on the posts and number of views that the SO3 synthesis threads in prepuplication recieved, I'd say plenty of people would be interested.

blogfast25 - 28-1-2012 at 09:00

If you perform a six step synthesis, and the yield of each step is 50%, then the overall yield is 1.5%.

That’s a slanderous lie! I’ve personally checked all steps and they should all give yields between 114.5 and 121.1 %, resulting in an overall yield of about nearly a 1000 %!

{irony off}

weiming1998 - 28-1-2012 at 17:12

Quote: Originally posted by 497

Actually, travelling in foreign countries for the last few months makes any experimentation impossible.

Do you have any evidence at all to back up a single one of the assertions you've made in the last couple posts?? Negative speculation like that is entirely useless. If you find any evidence that pertains to my posts, I would be interested to see it.

Why do you keep talking about complex mini distillation kits? Last I checked, a 250ml rbf would fit in a microwave. The whole point of what I said earlier was that there would be no need for a condenser at all! On top of the rbf you could put a right angle inlet adapter such as the ones here: http://unitedglasstech.com/Adapters.htm
Then a couple feet of teflon tube could lead from the adapter, outside the microwave directly in to a flask of H2SO4. Drilling a hole in a microwave is as easy as drilling a hole in any other sheet metal. It is not dangerous if the hole is fairly small, and you should be far away while it is running anyway.

You could be making more like liters a day of SO3, not a few ml. Just because you happen to not care about oleum/SO3 does not mean anything. Based on the posts and number of views that the SO3 synthesis threads in prepuplication recieved, I'd say plenty of people would be interested.

That might work, sorry for insulting you.

497 - 30-1-2012 at 02:11

Thanks. Also it is worth mentioning that SO3 will dissolve in many organic solvents forming more or less stable complexes. DMF, pyridine, alkyl acetamides, dioxane, and DMSO are examples.

I'm really dying to try this out... but it is impossible for at least another couple months or more.

So if anyone out there has an expendable microwave and some FeSO4 I emplore them to try it out. Pretty much grind some FeSO4 with graphite, throw it in some sort of dish or flask, wrap some fiberglass around it and microwave it. Obviously outside, because if it works as I hope there will be quite a bit of thick white sulfuric mist produced. Dissolving the residue in water and weighing the insoluble fraction will tell you the yield.

[Edited on 30-1-2012 by 497]

AJKOER - 30-1-2012 at 08:19

Just asking, but wouldn't SO3 do a job on the microwave itself (or its circuitry)?

Neil - 30-1-2012 at 08:44

Just asking, but wouldn't SO3 do a job on the microwave itself (or its circuitry)?

Oh yes. It would do a pretty good job on it.

497 - 30-1-2012 at 11:08

Uhh, the SO3 will most definitely not make it far out of the crucible. Fortunately, I believe many microwaves have fans that suck air through the magnetron and then blow into the cavity (while turning the thin metal wave chopper in the process). So the positive pressure in the cavity should mostly protect the circuits. Also, as the resulting sulfuric mist is notoriously reluctant to condense on to anything, corrosion should be minimal. Even still, don't use your mothers $300 microwave for that experiment. Cheap/free functional microwaves are plentiful most places, and they are useful for so many other things.

Neil - 30-1-2012 at 11:19

And the potential for the vapour to ionize and give you a plasma sun made out of S and O?

Why would the vapours not want to leave the crucible?

497 - 30-1-2012 at 12:02

... because they would instantly react with humidity to form highly uncondensable sulfuric mist. I didn't mean they wouldn't leave, just that SO3 is unlikely to survive in the atmosphere long enough to go far. Especially since the open crucible experiment would just be a small scale test of concept. Or you could pipe it out of the microwave and not have to worry.

[Edited on 30-1-2012 by 497]

Neil - 30-1-2012 at 18:48

ahh you were being very literal, I conceded that yes you are likely right

But then you still have the problem of all that sulfuric acid/oxide fume blowing around and the chance of turning the microwave into a sulfur lamp.

497 - 31-1-2012 at 06:12

I guess that's why some would call it an experiment...
If you're that worried about the microwave, doing any chemistry in it is unadvisable.

Sulfur lamps operate under pressure, without oxygen present. But, yes some sort of plasma is still possible I suppose? Can anyone confirm? That would be pretty cool in itself.

watson.fawkes - 31-1-2012 at 07:32

Quote: Originally posted by 497

But, yes some sort of plasma is still possible I suppose? Can anyone confirm?

I can't confirm from direct experience, but thermal runaways are ordinary in microwave ovens with certain materials. The issue is that absorption changes by orders of magnitude when you get more free electrons. So the first drop of liquid can superheat, and if that evaporates and ionizes, the first bit of plasma absorbs even more heat. I don't know the absorption coefficients, but I have to imagine they've already been measured.

If you want to try for a plasma, you want as rapid heating as possible. Loosing heat by conduction minimizes the runaway effect. You want to dump many watts into the smallest volume. Two pieces of advice. First, promote standing waves in your oven. Easiest way to do this is to remove the stirrer. If you're really diligent, tune the cavity size to be integral multiples of the half-wavelength. Second, keep your sample at about the quarter-wavelength diameter, and locate it over one of the standing wave antinodes. You'll get a single point of high-concentration RF energy.

AJKOER - 2-2-2012 at 13:35

Before this thread is closed, it should be mentioned (albeit somewhat obvious), there is a potential path to dilute H2SO4 via direct hydrolysis of FeSO4 "under air". Source:

"Hydrolysis of Iron Ion in Chrysotile Nanotubules: A Template Effect on Crystal Growth"
by Sumio Ozeki, Hiroyuki Uchiyama, Motomi Katada
Langmuir, 1994, 10 (3), pp 923–928
DOI: 10.1021/la00015a051
Publication Date: March 1994

To quote: "When aqueous solutions of FeSO4 are hydrolysed under air at room temperature, various products,.." and other than FeOOH, Fe3O4, Fe2O3 listed, there could be some dilute H2SO4 along with aqueous FeSO4, I suspect.

Whether the amount of dilute H2SO4 created is significant may be debatable.

Note, this is a laboratory induced hydrolysis without the helping hands (figuratively speaking) of little microbes, previously discussed.

Link:
http://pubs.acs.org/doi/pdf/10.1021/la00015a051

Here is a full paper discussing Chrysotile-Nanotubes:

http://www.lebsc.it/wp-content/uploads/2011/02/2_Geoinspired...

[Edited on 2-2-2012 by AJKOER]

weiming1998 - 3-2-2012 at 20:57