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Author: Subject: Sulfuric Acid at Home
grndpndr
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[*] posted on 8-6-2010 at 23:45


Not a 100yr old book+ but a modern chemistry treatise w/references not as "i was led to believe" OR ANOTHER METHOD beyond simple heating in a beaker as is so often discussed here to concentrate H2so4 to 98% unequivacly
Dont blow me off on this one guys several modern sources state unequivacly the azetrope of a h2so4/water is 93.3% @338C at ambient pressure.Others vjust as insistent azetrope
is 98.3 at the same temp.

Beins this is an honest chemistry forum above the average lets have honest answers.How do you beat the stated azetrope with heaT and a beAker as has been so often reccomended possible?:(
If azetrope means what i assume above this temp as much h2SO4 is lost as H20.Others are commenting on this apparent innacuracy.Please dont shove me off to a treatise ill never find but use qoutes (RECENT) and explanations as I have mine simple sweet and I think, irrefutable given the circumstances.If the inflated info's BS let it die,this is a science forum.Totse/utube needless to say arent acceptable evidence.

To p0revent useless bickering ive cut to the chase it seems from what ive found theres a big discrepancy between modern information on the net (reliable sources) listing 98.3 and 93.3? aS THE AZETROPE AT IDENTICAL TEMPS,AMBIENT PRESSURE WASNT MENTIONED EXCEPT IN THE 93.3 FIGURE.
Now whaT?I Hope in a civil manner WHATEVER the outcome.Ive really no dog here ID JUST LIKE TO SEE THE TRUTH OR IS THERE A PROBLEM WITH METHODS OF MEASUREMENT/pressure?aLTHOUGH A 5% DIFFERENCe cant be ignored.

In retrospect there does seem to be more ofa consensis towards 98.3 but no pressures mentioned.No offense anyone just a interest in the truth.:)

New info added/some incorrect retracted.

[Edited on 9-6-2010 by grndpndr]
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BenZeen
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[*] posted on 9-6-2010 at 01:48


electrolysis of copper sulfate soln
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grndpndr
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[*] posted on 9-6-2010 at 02:12


I dont pretend to understand this but the handbook of chemistry also Design of adibatic Flash units, Ekberg.
Of course pressure has a difference maybe 1.5 % or thereabouts between 1bar and .05bar.Seems the main difference is again something I dont grasp is the composition measurements.The lower 93.3 azetrope seems measured like this,( mole %H2SO4), the higher 98.3,
(%w/w H2SO4), If anyone grasps this I would appreciate an explanation.If its a wrong turn also appreciate it.I was under the impression this was quite straighforward with some misunderstanding?
This was taken from theDesign of Aliphatic Flash Units table 1 Ekberg. HELP!Not that the small difference will have alot of difference in my backyard chem its an interestin point?

edited ;revisions content as better came to light.

[Edited on 9-6-2010 by grndpndr]

[Edited on 9-6-2010 by grndpndr]
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hissingnoise
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[*] posted on 9-6-2010 at 02:37


grndpndr, the figure 93.3% for conc. H2SO4 is obviously a typo - 98.3% is the correct value. . .

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grndpndr
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[*] posted on 9-6-2010 at 07:31


ok?THERE SURE IS ALOT OF TYPOS OUT THERE TO INCLUDE MODERN TEXTS.SAY APPROX 1/3 CONTAIN THE AZETROPE OF h2SO4 AS 93.3 @AMBIENT TEMP.tHE VERY SAME TEXT/SAME TABLE LISTED THE 2 AZETROPES AT 93.3 THE OTHER SIDE 98.3.tHE ONLY DIFFERENCE A MEASUREMENT DIFFERENCE AS i TRYED TO EXPLAIN ABOVE.iTS CONFUSING AT BEST ONE WOULD THINK A STD WOULD BE THAT HARD TO RECOGNIZE.:(

[Edited on 9-6-2010 by grndpndr]

[Edited on 9-6-2010 by grndpndr]
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rrkss
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[*] posted on 9-6-2010 at 07:38


I've done concentration of sulfuric acid using boiling. When titrated using a mohr pipet and Fisher 1.00 N NaOH solution, 1 mL of H2SO4 requires 37.0 mL of NaOH solution to make the phenolphalein turn pink. To me this gives me proof that my solution is at the 98% concentration and not 93%.

[Edited on 6-9-10 by rrkss]
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agorot
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[*] posted on 9-6-2010 at 08:06


its 98% max. really.
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bbartlog
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[*] posted on 9-6-2010 at 08:21


As you suggested, the 93.3% figure is based on mol/mol rather than weight/weight. So in one hundred moles of the azeotrope there will be 93.3 moles of H2SO4 and 6.7 moles of water. Such a mix would contain

93.3 x 98 ~ 9144 g H2SO4
6.7 x 18 ~ 114g H2O

for a total of 9256g, of which less than two percent by weight is water. On the other hand, my calculation suggests that this would end up being somewhat more than 98.3% by weight, so either I've made a mistake or else I really don't understand the mol% system; but I'm quite sure that the different percentages are related to the different measures, not to some mistake. Presumably there are circumstances where the mol/mol measure is more convenient for calculation.
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grndpndr
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[*] posted on 9-6-2010 at 10:58


Thanks all and particularly barrtlog for the expl i was looking for couldnt express well,(could it be a diff in atmospheric pressure)Also Q.S for the quick dirty Idea of a 1.84 auto hydrometer.Thanks all:)Auto supply hyd to cheap or expensive sciplus $9hyd and my own 100ml grad cyl.

Im curious the difference might not be bar's? Between 1bar and .o5bar the difference in is 1.5% concentration give or take.
im at 4500ft so..well see when the hydrometer gets here.

\One last thing do I have to achieve boiling to concentrate my h2so4 to azetropic.My inex. worn hotplate wont manage even the 290C BP?:( concerned I may need to go to a propane burner/asbestos.:(


[Edited on 9-6-2010 by grndpndr]
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[*] posted on 11-7-2010 at 19:59


Here's a method of making H2SO4 I just successfully completed. I'm not sure it's economical, but I'm new to chemistry so at least I'm learning something.

Basically I just add stoichiometric amounts of Copper and Sodium Bisulfate in a test tube and heat it and the following reaction takes place:
4NaHSO4 + heat ---> 2Na2S2O7 + 2H2O
2Na2S2O7 + heat ---> 2Na2SO4 + 2SO3
2SO3 + Cu ---> CuSO4 + SO2

Then I run the sulfur dioxide through hydrogen peroxide (3%) to make about a 2.9% H2SO4 solution and cook it down.

In the end I'm left with a nice chunk of copper (II) sulfate in the test tube I can grow crystals and concentrated sulfuric acid.

If anybody is interested here are the amounts I use as a guide:
NaHSO4 = 1 g
Cu = .132 g
H2O2 = 2.362 mL

PS. I strip copper electrical wire instead of using copper flashing since the wire has more surface area.
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Lambda-Eyde
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[*] posted on 11-7-2010 at 23:35


Err - why not bubble the sulfur trioxide directly into water to make sulfuric acid?
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[*] posted on 12-7-2010 at 01:33


Sulphur trioxide doesn't easily dissolve in water without forming an acid mist that gets everywhere you don't want it! So it's usual to form or buy sulphuric acid and add sulphur trioxide to it -which in practise works better. The product is a strong sulphuric acid or oleum, this is then extracted and the concentration adjusted to usable or saleable specifications.
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Formatik
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[*] posted on 13-7-2010 at 11:27


Quote: Originally posted by MttLsp  
Here's a method of making H2SO4 I just successfully completed. I'm not sure it's economical, but I'm new to chemistry so at least I'm learning something.

Basically I just add stoichiometric amounts of Copper and Sodium Bisulfate in a test tube and heat it and the following reaction takes place:
4NaHSO4 + heat ---> 2Na2S2O7 + 2H2O
2Na2S2O7 + heat ---> 2Na2SO4 + 2SO3
2SO3 + Cu ---> CuSO4 + SO2


The copper may have reacted with the bisulfate directly:

4 NaHSO4 + Cu = CuSO4 + 2 Na2SO4 + SO2 + 2 H2O

This is not an efficient use of the sulfate. Crystallization of some Na+ salt isn't hard, but you'll still have impure CuSO4.

If the pyrosulfate intermediate was needed then the reaction might need something like a quartz vessel to be useful, because when you are doing the reaction in larger amounts, decomposition of pyrosulfate happens only significantly with blow torch temperatures or an electric heating coil, etc. Don't know if the copper changes this.

Carbon can be used to reduce Na2SO4 to SO2 (alongside some Na2S), but this needs temperatures above 1000 C to give off SO2 at a significant rate. NaHSO4 might reduce similarily. Though I'm sure there are reductants that should work at lower temperatures. It should be tested carefully, e.g. at red glow aluminium decomposes Na2SO4 under detonation - Ch. Tissier, A. Tissier (C.r. 43 [1856] 1187; J. pr. Ch. 71 [1857] 76).

NaHSO4 should be converted to H2SO4 with aq. HCl. I've done it with Na2SO4 and HCl, but could only convert a bit less than half of its sulfate to sulfuric acid. I had to also distill the H2SO4.

[Edited on 14-7-2010 by Formatik]
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MttLsp
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[*] posted on 15-7-2010 at 18:32


Oh wow I didn't even think about how Na2SO4 is left over. And that HCl + NaHSO4 is a good idea, but the distillation seems dangerous unless you had a temperature controlled hotplate or mantle. I actually used table salt and hydrochloric acid to make the sodium hydrogen sulfate. But I have a question:

Is there a way to seperate the Na2SO4 and CuSO4 in the reaction i suggested. Perhaps by combining the SO3 from the NaHSO4 with Cupric Oxide in a seperate flask and then running the SO2 into the H2O2 as previously stated? Or would this still be inefficient?
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Formatik
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[*] posted on 16-7-2010 at 07:56


Quote: Originally posted by MttLsp  
Oh wow I didn't even think about how Na2SO4 is left over. And that HCl + NaHSO4 is a good idea, but the distillation seems dangerous unless you had a temperature controlled hotplate or mantle.


You would need a strong hot plate to distill even small amounts of H2SO4. I was actually heating less than 2mL under the bunsen burner, got impatient and starting heating with an additional second burner! Its high boiling point makes this problematic at regular atmospheric pressure.

Quote:
I actually used table salt and hydrochloric acid to make the sodium hydrogen sulfate.


Usually H2SO4 and table salt, or Chile saltpeter is used to make it. But it makes no sense to start out with H2SO4 only to end up later with H2SO4 with a lower yield.

Quote:
Is there a way to seperate the Na2SO4 and CuSO4 in the reaction i suggested.


Crystallization as mentioned. Again, it only leaves you with partly pure CuSO4. You could also find a solvent that one is soluble in, but the other is not (good luck).

Quote:
Perhaps by combining the SO3 from the NaHSO4


Again, try decomposing large amounts of NaHSO4 by itself, and you'll see its pyrosulfate doesn't decompose significantly until you start heating with temperatures that soften and ruin normal borosilicate flasks, beakers, etc (see thread on decomposing NaHSO4 in prepublications on why quartz is used).

Quote:
with Cupric Oxide in a seperate flask


That ought to leave you with CuSO4. SO3 + CuO = CuSO4.
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AndreiChim
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smile.gif posted on 15-6-2011 at 11:56


Hi. Could you give more details to the NaHSO4-ethanol/water method you described earlier? It sounds interesting. Thank you.



A chemical - friend or foe?
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albqbrian
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[*] posted on 15-6-2011 at 16:08
As for availability...


Luckily I see SA as a chemical that will be widely available. If for no other reason than its use in biodiesel production. Check out suppliers in that area for a variety of chems. Like:

www.dudadiesel.com
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