making sulfuric acid

[YT2095]

you will need to have an outlet (second open hole in the stopper) incase some of the gas does not react for some reason like low concentration of H2O2 or similar.

 

indeed, the unreacted gas should be passed through a scrubber, a soln of NaOH will work quite nicely for this.

[hydraliskdragon]

Thanks for the help Frosch45 and YT2095,

 

Now I can finally, safely bubble SO2 in to H2O2!

 

PS. What concentration of H2SO4 would I be ending up with, saying that the H2O2 would be about 15%?

[frosch45]

Thanks for the help Frosch45 and YT2095,

 

Now I can finally, safely bubble SO2 in to H2O2!

H2O2 would be about 15%?

 

its only as safe as you make it

PS. What concentration of H2SO4 would I be ending up with, saying that the

H2O2 would be about 15%?

 

you want us to do all the work for you!?

depends a little bit on how well you do the reaction.

 

H2O2 + SO2 --> H2SO4

 

so if you started with 100% pure 15% H2O2 and then you have all of your so2 react, you should end up with a 15% solution of H2SO4 wich can be further purified by vacuum distillation

[offset442]

OK real simple, I not the type to flame a 13 old i was handlng fuming nitric acid when i was in pullups

 

ok 2 possible methods both require a plating power supply around 3 to 5 volts at 10 amps and some not so cheap and easy to get supplies....

 

first is the kestler ckline methiod: (well sort of) get a volume of mercury metal placed in the bottom of a pyrex container enough to at least cover the entire bottom of the container make the mercury the cathode (lower an insulated wire with a short striped end into he pool of metal) next you will need an anode preferably MMO coated Ti, but carbon/ graphite from a dry cell battery will do. then add saturated sodium bi-sulfite to the cell and turn on the power, the the sodium metal will plate into the mercury and the SO3 ion will be free to rip appart a nearby H2O making the acid.

(OK i know mercury is hard to get, I had to mine the 100 lbs that i have and retort it from raw sulphide ore(cinibar), took weeks of tinkering)

 

 

after some time the mercury will saturate with sodium, no problem, just decant the cell and rinse the acid off the top by over flowing the container, leave clear water in the container and stir until the all the sodium leaves the mercury making aqueous sodium hydroxide another nice chem to have!!! with this cycle you can work your way to most any inorganic acid directly or indirectly

 

the second in membrane electrolysis:

 

you will need a container that can be divided into to compartments by a membrane, with an electrode on each side of the membrane, ok so its a simple membrane cell you can google for ideas same voltage as above use a graphite or MMO Ti anode and the cathode can be any type of stainless steel (as long as you keep the power on when the cell has brine in it)

the only 2 possible membranes that i know of that a kid could get that will have any chance of sucess will be the following:

 

1. the seperater membrane found in the common DEEP CYCLE car style battery

2. DUPONT NAFION

3. TCF from a reverse osmosis membrane cartage

 

ok you mentioned fuel cell experiments, your fuel cell if its all plastic could make the acid just fine...........

 

OK everybody start making fun of my typos, i dont care....tata

[albgk]

I recently attempted to make H2sO4 here was my procedure:

I heated Sodium Bisulfate (NaHSO4) in a test tube- this should have made gaseous SO3

the test tube had a tube on the end that i used to bubble the SO3 through water in a collection beaker with a peice of Lithmus paper in it (so i could tell if it was working)....

Then came my problem... I was heating it and it was letting off water vapor (another bi product of heating NaHSO4) then i smelt sulfur and suddently the water (or possibly sulfuric acid) shot back up the tube and blew up my test tube. I was cleaning up this mess (thinking it was a complete failure) when i noticed my Lithmus paper had turned red. I think this could have been sulfuric acid, but i couldnt do any tests with it because i was busy picking up broken glass, and wiping up mystery liquid. The verdict- It is hard (and dangerous) for a home chemist with basic supplies (chemicals, testubes, burners, ect.) to make sulfuric acid.

 

Now that i have told my story, if any of you chemistry wizzes notice a flaw in my procedure please let me know because i would like to perfect an easy way to make sulfuric acid

 

thanks

 

albgk

[UC]

I recently attempted to make H2sO4 here was my procedure:

I heated Sodium Bisulfate (NaHSO4) in a test tube- this should have made gaseous SO3

the test tube had a tube on the end that i used to bubble the SO3 through water in a collection beaker with a peice of Lithmus paper in it (so i could tell if it was working)....

Then came my problem... I was heating it and it was letting off water vapor (another bi product of heating NaHSO4) then i smelt sulfur and suddently the water (or possibly sulfuric acid) shot back up the tube and blew up my test tube. I was cleaning up this mess (thinking it was a complete failure) when i noticed my Lithmus paper had turned red. I think this could have been sulfuric acid, but i couldnt do any tests with it because i was busy picking up broken glass, and wiping up mystery liquid. The verdict- It is hard (and dangerous) for a home chemist with basic supplies (chemicals, testubes, burners, ect.) to make sulfuric acid.

 

Now that i have told my story, if any of you chemistry wizzes notice a flaw in my procedure please let me know because i would like to perfect an easy way to make sulfuric acid

 

thanks

 

albgk

 

You've just experienced the magic of suckback, a common phenomenon. SO3 is extremely soluble in water, so there is no reason to submerge the exit tube. Just place it slightly above the surface of cold water. This tends to produce mists though, which is why SO3 is usually condensed neat or dissolved into concentrated sulfuric acid, then diluted with water.

 

Were you using a rubber stopper? I sure hope not, because SO3 attacks everything and anything organic. By smell of sulfur, do you mean hydrogen sulfide or sulfur dioxide? I would suspect that either is coming from the stopper being destroyed if you used one. You should be using all glass (preferably fused quartz/vycor) apparatus. Concentrated sulfuric acid is an acceptable joint grease for this operation.

[albgk]

I do believe that i was using a rubber stopper and ruber tubing... at the time of the suckback both rubber items seemed to be fine. While i was cleaning up the mess i did wash the stopper and tubing and they still seemed fine. Im about to go check them out again (hopefully they are still there).

 

Im also going to try something else (please tell me if you think this will work)

 

Im going to put H2O2 in a erlenmyer (bad spelling) flask then i will take a bent metal spoon with sulfur in it, light the sulfur on fire, and then lower the flaming sulfur (producing SO2) into the flask of H2O2. Ill give it a shake (carefully ) and i hope that will produce H2SO4.

 

I think this will work (not so efficiently as bubbling SO3 with water)

 

If i have a problem in this procedure please tell me (im going to wait on the ok from you guys this time )

 

 

Thanks

 

albgk

[UC]

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/talk/viewthread.php?tid=2824

[albgk]

thanks UC

ill try that, but i have another idea for a process, ill tell everyone if it works

it seems to be safe but it will take me awhile but ill try.

 

wish me luck

 

thanks albgk

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Ok heres what I have from my procedure

It was a partial success.

My procedure was like my 2nd to last post (with the flask and the buring sulfur) except that i put an angled glass tube in it to pull in more oxygen (i dont know if it worked). Also i attempted to boil the H2O2 to remove the H2O in the solution (i only had 3%) but that did not work well.

Now to the good part... i lit the sulfur and lowered it in and hey the blue flame was still there... once i felt i had a sufficient ammount of SO2 in the flask i removed the spoon and corked the flask... then i gave the flask a shake, for a very long time, and I still had some SO2 left in there. I then took the liquid and put it on some lithmus paper and it turned red, but when i put i put it in sugar... the sugar turned yellow and did nothing else =(.

The verdict-

i believe that i created alot sulfurous acid and a trace ammounts of sulfuric acid (i hope)

 

I now need suggestions on how to perfect this procedure

 

thanks albgk

[HeXx]

Hi,

About the so3, my chem teacher told me that it was possible to create it by burning S2 which leads to so2 and then burn the so2 it self in the presence of pure O2 which would then create so3.

 

Is that right?

[frosch45]

Its pretty hard to burn pure sulfur just in air. You need higher temperatures and a high concentration of oxygen.

 

At high temperatures and with a Vandium (V) Oxide catalyst your SO2 will be converted to SO3, but that reaction is reversable.

[John Cuthber]

Its pretty hard to burn pure sulfur just in air. You need higher temperatures and a high concentration of oxygen.

 

 

Bollocks.

Sulphur is relatively easy to burn.

Estimates vary but the ignition temperature is about 200 to 250 C

[frosch45]

http://www.angelo.edu/faculty/kboudrea/demos/burning_sulfur/burning_sulfur.htm

 

Don't read the stuff at the top of the page, just look at the picts/video

 

I don't know, if you look at the pictures/video on that page, I wouldn't call it "easy" to burn sulfur in just regular atmospheric conditions, but "easy" is a relative term. It is certainly do-able, but I was trying to emphasize the difference when you use pure oxygen.

[John Cuthber]

In the bad old days before they had maches they used tinder boxes. You struck a flint against a piece of steel then let the sparks fall onto tinder. Blowing on the glowing tinder could just about get a flame out of it. You then used this to light slivers of wood coated with sulphur because it's easy to light. Modern matches include sulphur for the same reason.

 

That's what I mean by "easy" in this context.

The fact that it burns better in pure oxygen isn't under debate here; most things do.

[frosch45]

Alright then, sulfur is "easy" to burn in pure air. I won't argue with you. When I wrote the first post, I was thinking relatively to what I considered to be "easy."

[Dr.]

The manufacture of sulfuric acid presents us with an interesting lesson in industrial economics. We have seen that the roasting of sulfide ores produces sulfur dioxide as a waste product. For example:

2 PbS(s) + 3 O2(g) -----> 2 PbO(s) + 2 SO2(g)

 

 

From the beginning of metal smelting to the mid 18th Century, sulfur dioxide was simply sent up a chimney into the atmosphere. Over long periods of time, the sulfur dioxide slowly reacts with oxygen and water in the atmosphere producing sulfuric acid:

2 SO2(g) + O2(g) -----> 2 SO3(g)

SO3(g) + H2O(l) -----> H2SO4(l)

 

 

This is one important source of acid rain. Consequently, a smelter was not the ideal place to build your dream home. But the world was big in those days, the wealthy simply didn't live near a smelter, and the environmental lobby was nonexistent.

 

The discovery that indigo could be used to dye wool changed the situation dramatically. Now there was a demand for sulfuric acid but no way to produce it cheaply in the quantities demanded by the textile industry. In 1746 John Roebuck developed the lead chamber process for the manufacture of sulfuric acid. Prior to this time, sulfuric acid had been produced in glass bottles several pounds at a time. But the lead chamber process could produce sulfuric acid by the ton.

 

In the lead chamber process, sulfur and potassium nitrate are ignited in a room lined with lead foil. Potassium nitrate, or saltpeter is an oxidizing agent which we have seen when we discussed explosives. The saltpeter oxidizes the sulfur to sulfur trioxide according to the reaction:

 

6 KNO3(s) + 7 S(s) -----> 3 K2S + 6 NO(g) + 4 SO3(g)

 

The floor of the room was covered with water. When the sulfur trioxide reacted with the water, sulfuric acid was produced:

 

SO3(g) + H2O(l) -----> H2SO4(aq)

 

Notice that this process depends on cheap supplies of saltpeter and produced yet another air pollutant, nitrogen monoxide. Thus we have to pay for a nitrogen source, saltpeter, and all of the nitrogen winds up going up the smokestack. Saltpeter could be replaced with its less expensive cousin, sodium nitrate ("Chile saltpeter") but nevetheless you wind up paying for nitrogen which winds up as a waste product. Saltpeter was a significant expense.

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Joseph Gay-Lussac invented a process for recovering the nitrogen in nitrogen monoxide and recycling it to replace the saltpeter as a source of nitrogen to generate sulfuric acid also:

 

 

4 NO(g) + O2(g) + 2 H2O(l) -----> 4 HNO2(l)

4 HNO2(l) + 2 SO2(g) -----> 2 H2SO4(aq) + 4 NO(g)

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Also, would this not be possible?

 

CaSO4(s) + 2 H2O(l) -------> Ca(OH)2(aq) + H2SO4(aq)

 

 

or

 

CaSO4(s) + H2O(l) + CO2(g) ------> CaCO3(s) + H2SO4(aq)

 

 

???

 

 

 

 

This may also be of interest:

 

You can replicate the most primitive production of sulfuric acid from sulfur and saltpeter. The reaction will take place in a test tube with very little air, so we must count on the saltpeter to supply all of the oxygen. Begin by balancing the reaction:

KNO3(s) + S(s) -----> K2S + N2(g) + SO3(g)

or,

H2O + KNO3(s) + S(s) -----> K2S + N2(g) + H2SO4(l)

 

Then use stoichiometry to calculate the number of grams of saltpeter needed to react with 1.0 g of sulfur. You will bring these calculations with you to the lab when you are ready to make sulfuric acid. You will also need a 2 L soda bottle to take the place of the original lead chamber.

 

Begin by weighing 1.0 g of sulfur and your calculated weight of saltpeter and place this mixture into a clean, dry, test tube. Close the test tube with a rubber stopper fitted with a piece of glass tubing. Rinse your 2 L bottle with water and drain it, leaving the walls of the bottle wet. Put the glass tubing into the bottle and hold the test tube over a Bunsen burner. The sulfur and saltpeter will react producing a mixture of nitrogen monoxide (brown gas) and sulfur dioxide (clear gas). These react further in the presence of water to produce nitrogen and sulfur trioxide, which dissolves in the water to produce sulfuric acid.

 

What you will see initially is a violent reaction as the sulfur and saltpeter are heated. Billowing clouds of gas will be produced and the material may even catch fire. This is nothing to worry about. The gas that is produced will be brown in color and will go out through the glass tubing and fill the 2 L bottle. The bottle will seem to be filled with brown, hazy mist: a mixture of nitrogen oxides and sulfur dioxide. The gas will turn clear and trasparent over the course of several minutes as the nitrogen oxides oxidize the sulfur dioxide to sulfur trioxide. Sulfur trioxide dissoves in water to produce sulfuric acid.

Edited June 24, 2009 by Dr.
Consecutive posts merged.

[UC]

Also, would this not be possible?

 

CaSO4(s) + 2 H2O(l) -------> Ca(OH)2(aq) + H2SO4(aq)

or

CaSO4(s) + H2O(l) + CO2(g) ------> CaCO3(s) + H2SO4(aq)

???

 

This little bit leads me to believe that you copied all the rest of that text from somewhere, because this shows next to no knowledge of chemistry.

 

Ah yes. At least some of it is from here: http://www.cavemanchemistry.com/oldcave/projects/acid/

 

post reported.

 

Preserved for posterity:

The manufacture of sulfuric acid presents us with an interesting lesson in industrial economics. We have seen that the roasting of sulfide ores produces sulfur dioxide as a waste product. For example:

2 PbS(s) + 3 O2(g) -----> 2 PbO(s) + 2 SO2(g)

 

 

From the beginning of metal smelting to the mid 18th Century, sulfur dioxide was simply sent up a chimney into the atmosphere. Over long periods of time, the sulfur dioxide slowly reacts with oxygen and water in the atmosphere producing sulfuric acid:

2 SO2(g) + O2(g) -----> 2 SO3(g)

SO3(g) + H2O(l) -----> H2SO4(l)

 

 

This is one important source of acid rain. Consequently, a smelter was not the ideal place to build your dream home. But the world was big in those days, the wealthy simply didn't live near a smelter, and the environmental lobby was nonexistent.

 

The discovery that indigo could be used to dye wool changed the situation dramatically. Now there was a demand for sulfuric acid but no way to produce it cheaply in the quantities demanded by the textile industry. In 1746 John Roebuck developed the lead chamber process for the manufacture of sulfuric acid. Prior to this time, sulfuric acid had been produced in glass bottles several pounds at a time. But the lead chamber process could produce sulfuric acid by the ton.

 

In the lead chamber process, sulfur and potassium nitrate are ignited in a room lined with lead foil. Potassium nitrate, or saltpeter is an oxidizing agent which we have seen when we discussed explosives. The saltpeter oxidizes the sulfur to sulfur trioxide according to the reaction:

 

6 KNO3(s) + 7 S(s) -----> 3 K2S + 6 NO(g) + 4 SO3(g)

 

The floor of the room was covered with water. When the sulfur trioxide reacted with the water, sulfuric acid was produced:

 

SO3(g) + H2O(l) -----> H2SO4(aq)

 

Notice that this process depends on cheap supplies of saltpeter and produced yet another air pollutant, nitrogen monoxide. Thus we have to pay for a nitrogen source, saltpeter, and all of the nitrogen winds up going up the smokestack. Saltpeter could be replaced with its less expensive cousin, sodium nitrate ("Chile saltpeter") but nevetheless you wind up paying for nitrogen which winds up as a waste product. Saltpeter was a significant expense.

---

Merged post follows:

Consecutive posts merged

Joseph Gay-Lussac invented a process for recovering the nitrogen in nitrogen monoxide and recycling it to replace the saltpeter as a source of nitrogen to generate sulfuric acid also:

 

 

4 NO(g) + O2(g) + 2 H2O(l) -----> 4 HNO2(l)

4 HNO2(l) + 2 SO2(g) -----> 2 H2SO4(aq) + 4 NO(g)

---

Merged post follows:

Consecutive posts merged

Also, would this not be possible?

 

CaSO4(s) + 2 H2O(l) -------> Ca(OH)2(aq) + H2SO4(aq)

 

 

or

 

CaSO4(s) + H2O(l) + CO2(g) ------> CaCO3(s) + H2SO4(aq)

 

 

???

 

 

 

 

This may also be of interest:

 

You can replicate the most primitive production of sulfuric acid from sulfur and saltpeter. The reaction will take place in a test tube with very little air, so we must count on the saltpeter to supply all of the oxygen. Begin by balancing the reaction:

KNO3(s) + S(s) -----> K2S + N2(g) + SO3(g)

or,

H2O + KNO3(s) + S(s) -----> K2S + N2(g) + H2SO4(l)

 

Then use stoichiometry to calculate the number of grams of saltpeter needed to react with 1.0 g of sulfur. You will bring these calculations with you to the lab when you are ready to make sulfuric acid. You will also need a 2 L soda bottle to take the place of the original lead chamber.

 

Begin by weighing 1.0 g of sulfur and your calculated weight of saltpeter and place this mixture into a clean, dry, test tube. Close the test tube with a rubber stopper fitted with a piece of glass tubing. Rinse your 2 L bottle with water and drain it, leaving the walls of the bottle wet. Put the glass tubing into the bottle and hold the test tube over a Bunsen burner. The sulfur and saltpeter will react producing a mixture of nitrogen monoxide (brown gas) and sulfur dioxide (clear gas). These react further in the presence of water to produce nitrogen and sulfur trioxide, which dissolves in the water to produce sulfuric acid.

 

What you will see initially is a violent reaction as the sulfur and saltpeter are heated. Billowing clouds of gas will be produced and the material may even catch fire. This is nothing to worry about. The gas that is produced will be brown in color and will go out through the glass tubing and fill the 2 L bottle. The bottle will seem to be filled with brown, hazy mist: a mixture of nitrogen oxides and sulfur dioxide. The gas will turn clear and trasparent over the course of several minutes as the nitrogen oxides oxidize the sulfur dioxide to sulfur trioxide. Sulfur trioxide dissoves in water to produce sulfuric acid.

Edited June 24, 2009 by UC

[Dr.]

Indeed I did, I was simply fishing through the internet attempting to help the kid with his project, I have a PhD. in History, I made no claims to be a chemist.

 

oh yes, and of course, a copyright must be applied for in the united states to be considered legal, but I'm sure you already knew that.

 

can I not post information from another website on this forum? or does it require citation? I am unfamiliar with Internet blogs and forums, as I spend most of my time doing productive things, such as teaching about the sengoku period of Japan, or the siege of Alamut, or simply reading books.

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and it is not plagiarism, the internet is protected by a federal act called the "fair use" clause.

 

also, as clearly defined in US law, expression is protected, but ideas and facts are not. The information contained within that article is chemistry, is chemistry not fact? and at least the historical aspects of it are fact. So, you should not accuse of plagiarism until you understand intent, which mine was clearly to aid this young man in his research on the given topic.

Edited June 24, 2009 by Dr.
Consecutive posts merged.

[Phi for All]

We like to use citations here. Next time, please.

[mooeypoo]

Post #116 should have a citation along with it, as it is taken from this site:

http://www.cavemanchemistry.com/oldcave/projects/acid/

 

Even "fair use" clause insists you *CITE* your sources.

 

Please avoid plagiarism, people.

[Tugrul]

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 am 13 and in senior chemistry. i cant do crap in my home lab because my parents wont buy anything online for me and you cant get anyhting in stores thats good. nothing better than maybe occassionly KOH. i have to make everyhting myself, now i need help making sulfuric acid for a fuel cell catalyst for sci fair and for a few other experiments where the chemicals need a reagent. i have about 2kg of sodium bisulfate and an acetylene torch. can you plz tell me how to do this, ive tried to work this out on paper but i keep gettinfg fried answers. can you help me plz witha straightforward answer.

 

I feel for you, am 16 and i cant get any equiptment and the school does not let me conduct any experiments becouse of "health and safety". I can not buy any chemical without getting dirty looks and I am sick of it. Oh well nearly a year to go then its AS chemistry, it might get more interesting.

[ignilc]

The easiest way to make H2SO4 is to just heat up CuSO4 causing it to decompose into CuO and SO3, bubble the SO3 through water and voila. Very efficient and cheap, and way faster than any electrochemical method. And the side product, CuO is very useful as well for the home scientist.

[lucky45]

Sufulic acid is so readibly availible Dont; try to make it or you might end up with a severe case of Bronkites or resperitory problems. (con-suferic acid has no oder or fumes, but itis a very strong acid which will react with almost any metals.

[frosch45]

The easiest way to make H2SO4 is to just heat up CuSO4 causing it to decompose into CuO and SO3, bubble the SO3 through water and voila. Very efficient and cheap, and way faster than any electrochemical method. And the side product, CuO is very useful as well for the home scientist.

 

At 650 °C, copper(II) sulfate decomposes into copper(II) oxide (CuO) and sulfur trioxide (SO3).

 

At 650 °C

 

A home chemist's hot plate can definitely not reach this, even the Corning PC-351 lab hotplate I have can only hit about 540. The only way you could do it would be with a small crucible over a very hot flame.

 

At those temperatures it is hardly feasible to capture the SO3. You couldn't feed it directly into water either because the water would instantly boil because the gas is so hot. You certainly couldn't use plastic tubing, or that cheap stuff you can melt to bend and shape.

 

And SO3 is toxic enough at regular temperatures, but at 650 °C!? YOU WOULD DIE IF YOU INHALED IT, and high temperature gasses (and therefore high pressure gasses) tend to make glassware break or explode, releasing themselves into your workspace.

 

Definitely not a good idea.

[ignilc]

I use this method to make me CuO. Actually anhydrous cuso4 will start to decompose at about 300C. The 650 figure is the temperature at which it decomps vigorously. The so3 fumes are indeed very choking and irritating. I make it in my room without a fumehood, i just open the window . wouldn't it be possible to condense the so3 (boiling point 45C) and then add water?