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making sulfuric acid


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budullewraagh, making SO3 is not simple at all. The chemicals, needed for it, indeed are very cheap, but the process is incredibly difficult to perform at home. Have a look at sciencemadness, where they have a full thread devoted to making SO3. If you look at the huge problems in making that and see that every idea tends to fail or have pathetic yields at best, at the cost of lots of energy and with high risk, then I can only come to the conclusion that making SO3 in a home-lab is REALLY REALLY difficult. Beyond doubt, for 99+ % of all home-hobby chemists this simply is beyond reach. At least, for me it is.

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sorry about that ive just been kicked out of other forums cause of my age and am tired of getting taunted for it. and anyway i have an oxyacetelyne torch. ive been to the state science fair twice and

You will quickly consume all the oxygen in the flask and not much SO2 will be generated. I suggest you give this sciencemadness.org thread a read and try that instead: http://www.sciencemadness.org/ta

NaHSO4 definitely does not give off SO2. It gives off water at a few hundreds of degrees Centigrade:

 

2NaHSO4 --> Na2S2O7 + H2O

 

At much higher temperature (I think around 1000 C) it decomposes to give Na2SO4 and SO3.

 

The V2O5 catalyst method indeed can be used (and is used in industry) for oxidizing SO2 to SO3, but that it very difficult in a home setup. You need specially prepared V2O5 in a long and very hot tube (IIRC around 400C) and then the O2 from the air reacts with SO2 to form SO3. Not something one easily does at home.

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No, this does NOT work! Hydrogen simply bubbles through the acid. It does not form NaOH.

 

Just an exercise for you. Try to balance the equation you proposed and even on the basis of that, you can see that this cannot work. The hydrogen would go to the +1 oxidation state while nothing is reduced on the other hand. So, even without any knowledge of the reactants and their peculiarities one can conclude on the basis of some simple math already that this is not a possible reaction.

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It has been stated several times in this thread that the decomposition of sodium pyrosulfate (Na2S2O7) at high temperatures yields sulfur trioxide. This information was seemingly gained from the sciencemadness board. However, the members there concluded that is not what happens. Sulfuric acid is needed to drive the reaction forward. Hence one cannot attain sulfuric acid from sodium pyrosulfate because sulfuric acid is needed itself.

 

At http://www.sciencemadness.org/member_publications/SO3_and_oleum.pdf, the equation Na2S2O7 --> H2SO4 --> Na2SO4 + SO3 is given. "The sulfuric acid plays the role of a catalyst, it does not take part in the overall reaction, but without its presence no SO3 is formed."

 

I also found this from an encylopedia at the library: "NaHSO4, has been employed in the manufacture of sulphur trioxide. When heated it loses water to form sodium pyrosulphate, Na2S2O7, which on treatment with sulphuric acid yields normal sodium sulphate and sulphur trioxide"

 

Just wanted to set that straight. You cannot make sulfuric acid from sodium bisulfate even when all safety precautions are taken and with proper equipment.

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I was just looking for a good source of HCl' date=' and while checking out the plumbing department I saw a whole bunch of drain cleaners with the ingredients as follows:

 

Active Ingredients: 98% Sulfuric Acid

Inactive Ingredients: Water.[/quote']

 

HAHAHA! That's awesome! Brand names?

 

Wal-Mart is cracking down on stuff though, they classify superglue as a dangerous chemical and will not sell it to minors, believe me I've tried. I doubt they'd sell drain cleaner to one, but I could be wrong.

 

But again, that's what parents and friends are for.

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[...']Just wanted to set that straight. You cannot make sulfuric acid from sodium bisulfate even when all safety precautions are taken and with proper equipment.

You're right, it is correct what you pointed out. H2SO4 is needed as a catalyst.

 

In fact, you CAN get SO3 from Na2S2O7, but you need also some H2SO4 as a start. The net reaction of course still is

 

Na2S2O7 ---> Na2SO4 + SO3.

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I was flipping through a very old (published in 1879) and found something relevant to this thread. In regards to NaHSO4 or as the book calls it "Hydrogen Sodium Sulphate", "like the corresponding potassium salt, it is decomposed by alcohol at once into sulphuric acid and the normal salt".

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Sounds interesting. Could you please elaborate more on the conditions under which this decomposition occurs. I have 99.9% methanol, free of water, and I have NaHSO4, which I can easily make free of water by heating it. I could give this a try, but if you could give more details on the reaction, then my try will not be a random shot.

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Sorry woelen, that quote is all the chem book says. It was under an extremely short section describing the properties of NaHSO4. The word "alcohol" refers to ethanol. In other passages, it will decipher between "absolute alcohol" and "alcohol", the latter probably refering to 95% ethanol.

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OK, I'll just add some NaHSO4 then to some 96% ethanol and see what happens. I expect to obtain a very acidic solution of H2SO4 in ethanol. In due time this will obtain an orange, later even a brown color (due to dehydration, followed by condensation reactions). This is a fairly nice test for H2SO4 in ethanol. if the liquid remains colorless for many days, then I severely doubt what happens. I'll give it a try and come back on this.

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How efficient is bubbling Sulfur Dioxide in 35-30% Hydrogen Peroxide?

 

If I put some H2O2 in a balloon and put the balloon on a tube connected to sulfur in a test tube, and heat the sulfur to emit SO2. The SO2 goes in the balloon then you take the balloon and shake it form like ten minutes than you pour the contents out.

 

Would that be less efficient or not as easy as then just bubbling SO2 in a beaker with H2O2?

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Bubbling SO2 is much easier. SO2 is very soluble in water and when you bubble this through water, it certainly will dissolve.

 

From a practical point of view, you burning sulphur method is not easy. You will not have real pressure behind the SO2 formed, so, it is not easy to bubble it through water.

 

You really cannot obtain H2SO4, or do you just want to do this experiment in order to educate and entertain yourself?

 

Making H2SO4 from SO2 can be done in another way. Take some dilute HCl (not more than 15%), add sodium bisulfite (or potassium bisulfite), which can be purchased at wine/beer making stores as disinfectant and preservant. Then heat the solution to (almost) boiling. The SO2 now bubbles out of this solution, it does not dissolve well in hot water. And, if your HCl is not of too high concentration, you hardly will get any HCl with the SO2. By bubbling the SO2 (with small amounts of HCl in it) through H2O2 you make H2SO4. Heating this in turn drives off any small amounts of HCl, decomposes excess H2O2 and concentrates the solution.

 

Be careful with leading SO2 through H2O2 in this case. Use only 10% H2O2 (you can boil off water lateron) and be careful for suckback. When production of SO2 is not fast enough or when heating is stopped, while the bubbling tube is still in the H2O2, then quickly liquid is sucked back and eventually may come in your HCl/bisulfite mix. This is a real risk, so be VERY careful.

You can use plastic/rubber tubing. The transparent tubes, sold in aquarium stores for oxygen pumps or filters is perfectly suitable. SO2 is not corrosive to those tubes (but it is to your lungs!, be careful).

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A note about heating H2O2 and H2SO4: this may be a bit dangerous and/or scary due to the potential formation of some quantity of peroxymonosulfuric acid- it probably depends on temperature and won't be formed in mass quantity unless for some reason your H2O2 doesn't decompose quickly at all and you end up with >50% H2O2 and >85% H2SO4 but still, be wary. Plus, even without the peroxymonosulfuric acid being formed, a hot solution of H2O2 and H2SO4 will act similar enough to pirahna bath to burn you and/or other things very easily.

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From a practical point of view' date=' you burning sulphur method is not easy. You will not have real pressure behind the SO2 formed, so, it is not easy to bubble it through water.

[/quote']

 

I'm not pressure behind SO2 and H2O2. I am putting H2O2 in a balloon and shaking it with SO2 so that it mixes better then just bubbling it and waste some of the SO2.

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It's just more hassle than is really necessary. SO2 is soluble in water so you shouldn't loose much if you take it slow.

 

Woelen, how do you prevent that suck back from occuring? The exact some problem happened to me several times when I was dissolving NO2 in water (from the decomposition of Mg(NO3)2). As soon as I stopped heating, my liquid would get sucked back into the Mg(NO3)2. I quickly learned to remove the tube from the beaker before I stopped heating. I also tried to clamp the tube after heating, but this didn't work too well (the suction was quite powerful).

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I agree with cesium with the SO2 bubbling. Burning sulphur and trying to get that SO2 in water+H2O2 is a lot of hassle. Besides that, you also get sulphur vapor with the SO2, which condenses at other places, so you probably also get a turbid solution. The bisulfite/HCl route is much cleaner and much easier to perform.

 

Preventing suckback can be done with nice lab equipment, but I don't have that. What I do is take an extra small bottle (I have a wide necked 50 ml erlenmeyer for this purpose). The erlenmeyer is stoppered with a rubber stopper, with two holes. The tube, going to the gas generator is connected to one hole and that tube ends just below the stopper (somewhere in the higher part of the erlenmeyer). The tube, going to the liquid into which you want to bubble gas is connected through the other hole and that tube ends at the bottom of the erlenmeyer. This method does not prevent suck back, but when it happens, the liquid is sucked into the erlenmeyer and not into your gas generator. You then have time to disconnect the stopper or take out the tube before the liquid sucks back into the gas generator. If you take care of cleaning that erlenmeyer very well before the experiment, then suck back is not an issue at all. In that case you also will not have contamination of your liquid.

 

With some setups, for extremely soluble gases, I use yet another strategy. E.g. I once made some HBr from KBr and H3PO4. My receiver flask was filled with water, and it was loosely stoppered with a rubber stopper, having a single hole. The gas generator also has a single hole and that is tightly stoppered, such that air cannot escape. A tube is connected through both holes and hence the system is loosely closed. The tube is NOT immersed in the receiving liquid, it ends a few cm above it. Then I start heating. Pressure wants to built up, but gas can escape from the receiver flask. As soon as the generation of soluble gas has continued for a long enough time (e.g. fumes of HBr are visible near the receiver flask), the stopper at the receiver flask also is pressed firmly into the neck of the flask and then I continue heating. The soluble gas quickly dissolves in the receiver's water and no pressure is built up. When no more HBr is formed, I simply stop heating, let cool down and disassemble the setup. No suck back at all, because the tube is not immersed in water.

Beware: this setup ONLY can be safely used with very soluble gases (HBr, HI, HCl, NH3). I would not use it for NO2 or SO2, because these gases probably do not dissolve quickly enough to avoid pressure buildup.

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although gas traps do work, it`s a little elaborate, I use a Funnel just under the liquid surface to bubble the gas through, it has the advantage of a greater surface area and if any liqud tries to suck up, the volume of the funnel is so great the liquid lev drops bellow contact area and breaks the seal :)

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With some setups' date=' for extremely soluble gases, I use yet another strategy. E.g. I once made some HBr from KBr and H3PO4. My receiver flask was filled with water, and it was [i']loosely [/i]stoppered with a rubber stopper, having a single hole. The gas generator also has a single hole and that is tightly stoppered, such that air cannot escape. A tube is connected through both holes and hence the system is loosely closed. The tube is NOT immersed in the receiving liquid, it ends a few cm above it. Then I start heating. Pressure wants to built up, but gas can escape from the receiver flask. As soon as the generation of soluble gas has continued for a long enough time (e.g. fumes of HBr are visible near the receiver flask), the stopper at the receiver flask also is pressed firmly into the neck of the flask and then I continue heating. The soluble gas quickly dissolves in the receiver's water and no pressure is built up. When no more HBr is formed, I simply stop heating, let cool down and disassemble the setup. No suck back at all, because the tube is not immersed in water.

Beware: this setup ONLY can be safely used with very soluble gases (HBr, HI, HCl, NH3). I would not use it for NO2 or SO2, because these gases probably do not dissolve quickly enough to avoid pressure buildup.

 

I have used this many times. Just place end of gas delivery tube some millimeters above water level. Adjust this distance when level of liquid increases. In fact no stopper is needed for the bottle where gas is dissolved, it just has to have high walls. This worked well when i made HCl, HNO3 and HBr.

 

Although it seems that there is not much gas lost you need fume hood (or very good ventilation at least) as otherwise you can be exposed to escaping gases.

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although gas traps do work, it`s a little elaborate, I use a Funnel just under the liquid surface to bubble the gas through, it has the advantage of a greater surface area and if any liqud tries to suck up, the volume of the funnel is so great the liquid lev drops bellow contact area and breaks the seal :)

Hey, that is a good one! Never though of that. With some careful adjustment, this can be made into a very safe construction! Only for the REALLY soluble gases I can expect this to be somewhat cumbersome, because of frequent unwanted breaking of the seal, but for many moderately soluble gases it sounds like a great idea.

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