Sciencemadness Discussion Board - Catalytic oxidation of sulfurous acid

Catalytic oxidation of sulfurous acid

weiming1998 - 16-1-2012 at 02:52

I have tried to make sulfuric acid by bubbling sulfur dioxide through water. The process made sulfurous acid, by the reaction: H2O+SO2<==>H2SO3. The sulfurous acid is slowly oxidized by air into sulfuric acid with the equation: 2H2SO3+O2====>2H2SO4. But the oxidation is very slow, even
with a large surface area, and it is often prone to decomposition into SO2 and H2O. Is there a catalyst for the oxidation of H2SO3, and if there is, will catalysts for the oxidation of anhydrous SO2 work on sulfurous acid(Fe2O3, NO2, V2O5, etc) Also, how can the equilibrium be stopped from shifting to the left?

Note: I know of the H2O2 method, but the H2O2 available at where I live is very dilute (3%-6%) and very expensive ($7 dollars for 100 ml of 6% H2O2!) I also know the electrolysis method, but it is a waste of batteries for me, as I don't have an AC/DC converter.

Pulverulescent - 16-1-2012 at 03:27

The oxidation can be speeded up, somewhat, by bubbling dry O2 in (at low temperature to minimise SO2-loss)!
Ozone would be the ideal, but affordable, efficient ozone generators are pretty thin on the ground . . .

P

weiming1998 - 16-1-2012 at 03:32

Yes, I do have an ozone generator that bubbles small amounts of ozone (400mg generated/hour) and air. I turned it on for 10 minutes. The yield increased, but only very slightly. I guess I should try and bubble ozone for a few hours instead. Is there any other catalyst methods?

LanthanumK - 16-1-2012 at 06:31

Try addition of hydrogen peroxide. This reaction most likely occurs: H2SO3 + H2O2 -> H2SO4 + H2O

Pulverulescent - 16-1-2012 at 06:59

Quote:

Note: I know of the H2O2 method, but the H2O2 available at where I live is very dilute (3%-6%)

You can get 35%H2O2 here!
Bit pricey, though, and you'll need to boil the extra water off afterwards to get to 98% . . .

P

AJKOER - 16-1-2012 at 08:22

Here are some possible routes to oxidizing SO2 to make H2SO4:

1. 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

Adding more water forms Sulfuric acid and Hydrogen chloride:

HSO3Cl + H2O --> H2SO4 + HCl

2. Prepare HOCl/Cl2O. This route takes time, but is cheap. Add a weak acid (vinegar) to the highest available Bleach available (but the cheap 6% would also work, but requires more distillation to concentrate). This forms dilute HOCl which upon distilling off half of the solution will also double the HOCl concentration. This is because most of the Cl2O and HOCl as vapor comes over first (see Watt's). Concentrated HOCl (30%) is unstable but will oxidize S in water (or SO2) to H2SO4! Variations of this method employ Dichlorine mono-oxide. Cl2O can be prepared with some effort (and lab equipment) by directly passing Cl2 over heated moist Na2CO3, or heated NaOH and collecting the gas in a small quantity of H2O or CCl4. The reaction between Cl2O (or dry Ca(ClO)2 as it forms Cl2O with CO2 from the air) and dry Sulfur is explosive, but treating a thick Sulfur water suspension with Cl2O (in essence making conc HOCl) cautiously should work.

Caution: Explosion hazard avoid high concentrations of Cl2O in CCl4, contact with organic compounds, heat, shock and/or light.

Pulverulescent - 16-1-2012 at 11:17

Dilute nitric acid is an excellent oxidiser but some (approximate) stoichiometry would be needed . . .
I should say I haven't tried it myself, so caution is advised!
Heating the solution in a fumehood (or downwind) would drive off any unreacted HNO3!

P

AJKOER - 16-1-2012 at 14:50

A well know reaction:

HIO3 + 3 H2SO3 --> HI + 3 H2SO4

upon heating, the HI should come across as vapor.

Iodic acid, HIO3, can be prepared by the slow reaction of Iodine with water:

I2 + H2O <----> HI + HIO

and very rapidly:

3 HIO ----> HIO3 + 2 HI

Bubbling air into the above solution:

4 HI + O2 --> 2 H2O + 2 I2

and the reaction cycles creating HIO3.

plante1999 - 16-1-2012 at 14:55

Ad NO2 in the H2SO3 , this , with O2 will make H2SO4, this processe is called the lead chamber process.

weiming1998 - 16-1-2012 at 15:39

I am not allowed to buy things off the internet, and it would be very difficult for me to combine SO2 and Cl together. I don not have a distillation setup, so I can't distill the hypochlorous acid. Also, if I have nitric acid, why should I worry about getting sulfuric acid? Finally, the I2+H2O method seems good, as I will be trying to generate some iodine from 600 grams of dried seaweed. But I don't think there would be more than a few grams, and there is no supplier that will dare sell iodine. Also, I have no access to NO2. I tried to bubble chlorine through sulfurous acid before, and the yields are pretty high, but when I concentrate it, even if I boil it from 500mls to 25 mls, it still didn't get any purer! I would guess the ozone method is the most available to me right now. By the way, would Fe2O3 be a good catalyst?

[Edited on 16-1-2012 by weiming1998]

plante1999 - 16-1-2012 at 15:49

Wath apparatus does you have for generating O3? Fe2O3 will disolve in H2SO4 so no it won't be a good catalist. If I was gessing , I would say cobalt sulfate + O3 should oxidise realy fast H2SO3 to H2SO4.

[Edited on 16-1-2012 by plante1999]

[Edited on 16-1-2012 by plante1999]

weiming1998 - 16-1-2012 at 15:53

Somehow, I have an ozone generator that my parents took from China. It produces mainly air with a little bit of ozone mixed in it. Anyway, I don't think Fe2O3 will be dissolved very fast in very dilute H2SO4.

Pulverulescent - 16-1-2012 at 16:45

Quote:

I would say cobalt sulfate + O3 should oxidise realy fast H2SO3 to H2SO4.

O3 is just about the most effective oxidiser there is ─ why on earth would anyone want Co compounds in it, FFS? (

)

P

BromicAcid - 16-1-2012 at 17:05

http://www.sciencemadness.org/talk/viewthread.php?tid=6911&a...

There are two posts in this thread, one by me and one by Polverone regarding metal salt catalysis of SO2 solutions to give sulfuric acid. Check out the links for more info.

Sedit - 16-1-2012 at 17:47

Mn sulfate seems like it might be of some use here.

weiming1998 - 16-1-2012 at 18:06

Quote: Originally posted by BromicAcid

http://www.sciencemadness.org/talk/viewthread.php?tid=6911&a...

There are two posts in this thread, one by me and one by Polverone regarding metal salt catalysis of SO2 solutions to give sulfuric acid. Check out the links for more info.

Thank you, now I will try HClO and sulfur, then manganese sulfate.

weiming1998 - 16-1-2012 at 21:45

I just tried sulfur+hypochlorous acid. Here's what happened:
50mls of white vinegar is poured into a flask.
About 5-10 grams of calcium hypochlorite is added. The mixture is left to react for 30 seconds. The liquid is then separated from the solid, unreacted calcium hypochlorite and poured into a bottle. Sulfur is added in considerable amounts. The mixture has gotten warm and bubbling has started. The sulfur is now moving around in the solution. Floating sulfur sinks down for a while, then floats back up. The bubbles are chlorine gas (I think) because of what it smells like. Now, the mixture is acidic enough to react slightly with an iron wire, though it could just be the HCl that is formed corroding the iron.

[Edited on 17-1-2012 by weiming1998]

AJKOER - 17-1-2012 at 05:06

You may be short on the acetic acid as if you are using Bleaching powder, it contains some Ca(OH)2 (and CaCl2 hydrate) as well.

Also, you can start the oxidation chain via HOCl (created per the Acetic acid + Ca(ClO)2 ) not from the very beginning at Sulfur, but at H2SO3 (if you have so):

S + 2 HOCl --> SO2 + 2 HCl

SO2 + H2O + HOCl --> H2SO4 + HCl

Note, the reaction releases HCl as well so if insufficient Acetic acid, some Cl2 may be formed.

You can also employ (with precautions) H2S in water in place of H2SO3. Hydrogen sulfide can prepared per the reaction of heating Al + S (thermite in nature so do not attempt to heat the whole mixture; use very small quantities as the reaction is very exothermic capable of melting thru steel at 1,100 C) and then adding water to the Aluminum Sulphide (which is very sensitive to moisture forming H2S on contact). I would design the experiment so that the HOCl is available in a closed reaction chamber with the Al2S3 and water (that produces H2S and Al(OH)3 ). Note, Aluminum hydroxide does not readily react with HOCl. Caution, H2S is very toxic and will numb the senses (allowing a fatal dose via inhalation or skin exposure) hence the closed vessel approach. Nevertheless, take precautions and certainly do outdoors if considering this approach.

AJKOER - 17-1-2012 at 09:34

OK, I just noticed a major problem. You should have created the HOCl by reacting Ca(OCl)2 with H2CO3 and not Acetic acid (as you did not distill to get pure HOCl). As a result, you have some Calcium acetate (a soluble ionic salt) in solution. So free H2SO3 or H2SO4 could react with it to form the insoluble CaSO4.

Try again, use pure carbonated water, and filter out the CaCO3.

Also, the process with dilute HOCl should be slow (a gradual rise in the solutions pH with time).

Do not use Fe to test solution strength as even dilute HOCl attacks iron.

Endimion17 - 17-1-2012 at 13:34

Maybe you should check your premise first. There's no such thing as sulphurous acid being made by blowing sulphur(IV) oxide in water. The H2SO3 molecule doesn't exist in water.

Burning sulphur and stuffing the produced fumes in the water produces a solution of SO2 and sulphuric acid (from traces of SO3 made by the reaction of hot SO2 with atmospheric oxygen).
Bubbling sulphur(IV) oxide through water doesn't produce sulphurous acid.

Pulverulescent - 17-1-2012 at 13:50

Quote:

Maybe you should check your premise first. There's no such thing as sulphurous acid being made by blowing sulphur(IV) oxide in water.

LoL! 'Been reading the 'ever trustworthy' wiki again, have we???

AndersHoveland - 17-1-2012 at 15:23

It was my understanding that SO2 in water behaves analogously to CO2 in water, which is to say that the reason for the high solubilities is the equilibrium with their respective acids, H2SO3 and H2CO3.

[Edited on 17-1-2012 by AndersHoveland]

ScienceSquirrel - 17-1-2012 at 15:29

A solution of sulphur dioxide in water contains sulphur dioxide, bisulphite anions and solvated protons.
It is effectively a solution of sulphurous acid in so far as it forms salts with bases, acts as an acid towards indicators, it is a strong reducing agent, etc.
Some of chemistry is about sticking labels on things, some labels are better than others but very few of them really describe the underlying state of affairs.

AndersHoveland - 17-1-2012 at 15:37

Quote: Originally posted by ScienceSquirrel

Some of chemistry is about sticking labels on things, some labels are better than others but very few of them really describe the underlying state of affairs.

Yes, but chemists almost never refer to aqueous SH2 solutions as "hydrosulfuric acid". It is always referred to as "hydrogen sulfide". Similarly, when zinc chloride is dissolved in ether, it should technically be referred to as "zinc dichloride adduct", but chemists just do not do this.

weiming1998 - 17-1-2012 at 15:58

Quote: Originally posted by AJKOER

OK, I just noticed a major problem. You should have created the HOCl by reacting Ca(OCl)2 with H2CO3 and not Acetic acid (as you did not distill to get pure HOCl). As a result, you have some Calcium acetate (a soluble ionic salt) in solution. So free H2SO3 or H2SO4 could react with it to form the insoluble CaSO4.

Try again, use pure carbonated water, and filter out the CaCO3.

Also, the process with dilute HOCl should be slow (a gradual rise in the solutions pH with time).

Do not use Fe to test solution strength as even dilute HOCl attacks iron.

Ok, I will try this with carbonated water instead.

weiming1998 - 17-1-2012 at 16:06

Quote: Originally posted by Endimion17

Maybe you should check your premise first. There's no such thing as sulphurous acid being made by blowing sulphur(IV) oxide in water. The H2SO3 molecule doesn't exist in water.

Burning sulphur and stuffing the produced fumes in the water produces a solution of SO2 and sulphuric acid (from traces of SO3 made by the reaction of hot SO2 with atmospheric oxygen).
Bubbling sulphur(IV) oxide through water doesn't produce sulphurous acid.

Technically it is not H2SO3, but it is more convenient to call it that instead of "an aqueous solution of SO2" or "SO2(aq)"

AJKOER - 17-1-2012 at 21:56

Interesting point on H2SO3.

But, as a matter of chemistry, HOCl reputedly will oxidize S, SO2(aq) or H2S(aq) to H2SO4 (See Watt's Dictionary of Chemistry, Vol 2, page 16).

The solution, certainly at first, is most likely dilute H2SO4.

Note, any free Cl2 will react with any CaCO3 suspension (from H2CO3 + Ca(OCl)2 ) that was not completely removed upon filtering as follows:

CaCO3 (Suspended in water) + H2O + 2 Cl2 --> CaCl2 (aq) + CO2 + 2 HOCl (aq)

This reaction is also referenced in Watt's (page 12).

weiming1998 - 17-1-2012 at 23:28

Quote: Originally posted by AJKOER

Interesting point on H2SO3.

But, as a matter of chemistry, HOCl reputedly will oxidize S, SO2(aq) or H2S(aq) to H2SO4 (See Watt's Dictionary of Chemistry, Vol 2, page 16).

The solution, certainly at first, is most likely dilute H2SO4.

Note, any free Cl2 will react with any CaCO3 suspension (from H2CO3 + Ca(OCl)2 ) that was not completely removed upon filtering as follows:

CaCO3 (Suspended in water) + H2O + 2 Cl2 --> CaCl2 (aq) + CO2 + 2 HOCl (aq)

This reaction is also referenced in Watt's (page 12).

That way of making HClO is very interesting, I will try it some time.

weiming1998 - 17-1-2012 at 23:38

So I repeated my experiment yesterday with unflavoured soda water today. This time, 100mls of soda water is poured into a flask. Again calcium hypochlorite is added. This time
chlorine evolution is very low/non existent. A very fine mist of CaCO3 is suspended in the solution. I filtered it but only managed to filter off some of the mist. The solution is added to the leftover sulfur from yesterday. I waited for about 5-6 hours before pouring the solution out, now which appears to be clear, probably because calcium carbonate is insoluble in water, but calcium sulfate is slightly soluble. The solution is heated to about 25mls. This time, when added to dry sodium carbonate, the sodium carbonate turned green, but there was minimal fizzing! What's more strange is that when sodium carbonate solution is added to the acid, instead of fizzing, a white precipate forms! Is the precipate sodium sulfate? or something else entirely? The precipate formed on contact with the acid solution.

[Edited on 18-1-2012 by weiming1998]

AJKOER - 19-1-2012 at 06:15

OK, starting with Bleaching powder, Ca(OCl)2/CaCl2.nH2O/Ca(OH)2, and added H2CO3. Products are:

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

Ca(OH)2 + H2CO3 --> CaCO3 (s) + 2 H2O

Also, assuming a limited partial hydrolysis of the CaCl2 hydrate:

CaCl2 + 2 H2O <----> Ca(OH)2 (s) + 2 HCl

and the created Ca(OH)2 reacting with CO2 to form more CaCO3. Also, a corresponding small amount of Chlorine from any CaCl2 hydrolysis:

HCl + HOCl <---> H2O + Cl2

I would try and separate out the CaCO3 suspension (by adding some H2O and shaking and waiting till it settles), and not filter, as the organic filter paper may be bleached and consume some of the HOCl. As there is still some CaCl2 free in solution with any formed H2SO4:

CaCl2 + H2SO4 --> CaSO4 (s) + 2 HCl

so this is a visible check on the ability of the solution to produce a sulfate (including H2SO4), but as long as there is free Ca ion (from the CaCl2), only or mostly CaSO4 will be created. Hence, the need to remove the CaCl2 (free Calcium ion), which could be done via distillation.

Also, again some Chlorine formation:

HCl + HOCl <---> H2O + Cl2

[Edited on 19-1-2012 by AJKOER]

weiming1998 - 19-1-2012 at 16:57

I just tried bubbling ozone/air through sulfurous acid, then boiled it down from 500mls to 70mls. It increased my yield, but I don't know the concentration. It was already slightly fuming, but maybe that's water vapour. It didn't work when I tried to ignite potassium chlorate/sugar. By this stage, it can already attack iron quite vigorously, much like aluminum in HCl, when heated.

Also the heat coming from the beaker was much hotter than normal water vapour. It was already very hard for me to put my hand near the opening of the beaker, as the full heat blasted with enough force to scorch skin. Since I don't have a thermometer, I can't tell the exact heat. I might go buy a thermometer.

[Edited on 20-1-2012 by weiming1998]

Sedit - 19-1-2012 at 19:14

I don't understand why we are using oxidizing agents here, wouldn't electro-oxidation be cleaner and cheaper?

weiming1998 - 19-1-2012 at 22:15

Quote: Originally posted by Sedit

I don't understand why we are using oxidizing agents here, wouldn't electro-oxidation be cleaner and cheaper?

As I said before, I don't have a power switch that can do electrolysis without short-circuiting the whole house. Lantern batteries can only do a day or so of electrolysis, and is only 6 volts.

Sedit - 20-1-2012 at 00:26

Dude any adapter that plugs into the wall could be used for such a simple synthesis. You can by these adapters for 12 volts and over 1Amp with ease and they are dirt cheep. Yes I would recommend a resistor in the line to prevent it from over heating but for the most part there should be almost no issues what so ever.

weiming1998 - 20-1-2012 at 00:45

How much do they cost and where do I buy those adaptors? They are certainly very useful for electrolysis.

Sedit - 20-1-2012 at 00:57

They are on your cell phones your old nintendos (over 1 amp) and almost anything that plugs into the wall. They are nothing special they are just a transformer and rectified.

weiming1998 - 20-1-2012 at 01:24

So a cell phone charger?

Sedit - 20-1-2012 at 17:31

Any wall adapter, some will be more powerful then others. They are the worst power source because there low voltage and low amps but its better then nothing when your in a pinch.

weiming1998 - 20-1-2012 at 18:06

Ok, thanks. Yesterday I combined 2 batches of dilute sulfuric acid. I boiled it down to 25mls, still not fuming! But all the waste in there has precipated (some has even carbonized.) I guess I would have a better yield if I electrolyzed copper sulfate.

Sedit - 20-1-2012 at 21:04

I'm going to have a go soon at making a SO2 solution and using Lead electrodes to convert it to H2SO4 to see if there is Viability in what I am pretty sure would be the cheapest and simplest method of performing this oxidation.

weiming1998 - 20-1-2012 at 21:20

Use CuSO4 solution instead. Lead cathode+copper anode? Becuase I tried making lead sulfate with copper sulfate and lead, and it didn't work. The lead barely got a copper plating. Also, PbSO4 is insoluble and heavy, easy to remove by filter/pouring off.

AndersHoveland - 21-1-2012 at 13:08

Solutions of sodium sulfite in water can be spontaneously oxidized by air to sulfate, but I do not know if sulfur dioxide solutions can similarly be oxidized to dilute sulfuric acid.

But I found this:

Quote:

The air-oxidation of bisulfite can not be completely suppressed, and this appears to be the only factor pre- venting complete utilization of the bisulfite in sulfoacid.

Quote:

This is attributed to oxidation of bisulfite by air, but it is not clear why such incidental oxidation did not similarly affect the results for formaldehyde.

http://pubs.acs.org/doi/abs/10.1021/ac50092a021

Quote:

Sodium bisulfite solutions undergo oxidation on standing in air.

http://worldaccount.basf.com/wa/NAFTA/Catalog/ChemicalsNAFTA...

If bisulfite can be oxidized by air, then likely sulfurous acid can also be oxidized to sulfuric. But the concentration of sulfuric acid obtainable would in practice be very low, because lower pH would shift the equilibrium away from bisulfite.

SO2 + H2O <==> H2SO3 <==> HSO3[-] + H[+]aq

weiming1998 - 21-1-2012 at 14:49

Yes, the concentration is extremely low. 1000mls of dilute acid can be condensed down into 25mls of somewhat concentrated acid.

Sedit - 21-1-2012 at 21:06

I think one electrode should be made of a porous conductive material so that the SO2 bubbles out into the solution through the electrode so that oxidation occurs as its being feed into the system.

Placing materials like Mn Sulfate and what not would generate per sulfates causing the oxidation to take place much more efficiently.

weiming1998 - 21-1-2012 at 21:26

Quote: Originally posted by Sedit

I think one electrode should be made of a porous conductive material so that the SO2 bubbles out into the solution through the electrode so that oxidation occurs as its being feed into the system.

Placing materials like Mn Sulfate and what not would generate per sulfates causing the oxidation to take place much more efficiently.

I am probably now going to try electrolytic oxidation soon, but I am going to try a mixture of MnSO4 and FeSO4 for catalyst. Also, How much should I put in if I want it to act like a catalyst?

AJKOER - 24-1-2012 at 13:35

Interestingly, Chlorine can be used to oxidize dilute Sulfurous acid:

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

Note, water is consumed and the solution becomes more acidic.

Chlorine can be generated by adding FeSO4 to Bleach (NaOCl/NaCl), but please comply with local laws.

Polverone - 24-1-2012 at 14:08

Quote: Originally posted by AJKOER

Chlorine can be generated by adding FeSO4 to Bleach (NaOCl/NaCl), but please comply with local laws.

The incompatibility list of a material safety data sheet should never be mistaken for a laboratory preparation. Please do not share preparative advice that you cannot cite from a reputable source (e.g. a journal paper or scientific book, not Wikipedia or Loompanics) unless you have verified it yourself.

AJKOER - 25-1-2012 at 06:26

Quote: Originally posted by Polverone

Quote: Originally posted by AJKOER

Chlorine can be generated by adding FeSO4 to Bleach (NaOCl/NaCl), but please comply with local laws.

OK, a few points.

First, the source of the method of employing FeSO4 is actually an old Sciencemadness thread (Topic: "Chlorine" LINK:
http://www.sciencemadness.org/talk/viewthread.php?tid=1305&a...) reputedly based on repeated direct observations. I apology for not citing it, but given the poor science per some of our participants, I did not place it (in concurrence with your point) in the league of a journal paper or scientific book. Our threads are, however, at times good in reporting observations and free of reporting bias if the mechanism of the reaction(s) is(are) uncertain.

To my knowledge, there is no mention in the literature and I do agree with you that the MSDS should not be the source of choice.

Interestingly in that thread, several of participants debated the precise chemistry. I, myself, have also attempted, via a several stage reaction, in my recent thread on using FeSO4 to prepare dilute H2SO4 (pardon my arrogance in assuming everyone had read it) to explore the chemistry. If you read one of my postulated explanations, however, based on a hypothetical reaction with an unstable/unknown Iron hypochlorite, you may surmise possibly why it is not in the main stream literature.

Also, I should mention, there is a Cl2 preparation via HCl/H2O2 and Al foil also in that same Sciencemadness thread. I have cited in a prior thread a YouTube video on this reaction (also not reported/recommended in the literature) and dared to explain the chemistry of why (pardon me again).

[Edited on 25-1-2012 by AJKOER]

weiming1998 - 25-1-2012 at 06:56

Why FeSO4, which is 10 dollars for 500 grams. Why not use NaHSO4, which is more effective/cheap when generating the gas? Unless you want the by-product of the generation?

Also, the decomposition is really bothering. Unless you can drive the HCl/Cl2 out of the solution, however, without driving the SO2 out as well. Maybe freezing in a cold freezer?

[Edited on 25-1-2012 by weiming1998]

AJKOER - 25-1-2012 at 09:04

OK, per this link FeSO4, a common fertilizer, costs $60-80 / Metric ton (please comply with local laws). It is reputedly 98% pure!

http://www.alibaba.com/showroom/fertilizer-feso4.html

I should mention that in my previously cited method for the oxidation of H2SO3 via HIO3, I suggested the Iodine and water method:

I2 + H2O <---> HI + HIO

and rapidly: 3 HIO --> 2 HI + HIO3

where the Iodine is to be boiled in water until the solution becomes colorless. I also recommended adding O2 (via air or even H2O2) to convert the HI:

2 HI + H2O2 --> I2 + 2 H2O

and the reaction cycles to the beginning. For the record, the most recommended method is via adding Cl2 (see Wikipedia: http://en.wikipedia.org/wiki/Iodic_acid), but this results in HCl and HIO3.

But, per my route, it is not necessary to separate out the dissolved HIO3 since eventually all the Iodine is converted into HIO3 (aq).

Now, just treat the HIO3 solution with SO2 (or H2SO3):

5 H2SO3 + 2 HIO3 --> 5 H2SO4 + I2 + H2O

The reference cited below notes that the crude Iodine separates from the solution. Note, if we have excess H2SO3:

3 H2SO3 + HIO3 --> HI + 3H2SO4

so treating with O2 (or a little H2O2) also would remove any HI still remaining in solution (evident if solution gets darker after oxidation).

As an excellent reference: "Inorganic chemistry" by Egon Wiberg, Nils Wiberg, Arnold Frederick Holleman, pages 415-417. This good source discusses the extraction of Iodine, its solubility and applicable solvents.

http://books.google.com/books?id=Mtth5g59dEIC&pg=PA415&a...

weiming1998 - 25-1-2012 at 15:47

Iodine is a known drug precursor, so it and it's salts are more hard to find than nitrates. The only thing I found was a 50ml iodine tincture bottle that has 25mg/ml iodine and 25mg/ml potassium iodide. That would be only useful in microchemistry. I have tried extraction from 600 grams of DRIED seaweed and boiling down the salts formed. The salts are mainly carbonates, which fizzed when I added sulfuric acid to precipate the iodine. There is only about 10 grams of the salt.

[Edited on 25-1-2012 by weiming1998]

AJKOER - 25-1-2012 at 18:31

On Amazon KI pills, about 1 cent for each 10 mg. So, 1 dollar gets you 1 gram. Here is the link:

http://www.amazon.com/gp/offer-listing/B004SF6AEI/ref=sr_1_2...

Shipping is extra, but if you order like $25 you may get free shipping!

Or, 6 cents per ounce and $5.99 shipping.

http://www.amazon.com/gp/offer-listing/B000CFGNTU/ref=sr_1_1...

What do you think?

Also, you may be able to buy HIO3 directly at a good price and save some work in making it (min order 1 kilogram). See LINK:

http://www.alibaba.com/product-gs/490813026/Iodic_acid.html

Now, per our synthesis equation:

5 H2SO3 + 2 HIO3 --> 5 H2SO4 + I2 + H2O

each mole of HIO3 yields 2.5 moles of H2SO4 and half a mole in Iodine. Theoretically, if the HIO3 is cheap enough, one could get H2SO4 for nothing as you resell the Iodine for a profit after the cost of preparing the H2SO3. If the HIO3 is too expensive, one could get H2SO4 for nothing after you sell Iodic acid at a competitive price & profit including the cost of H2SO3 and I2 purchased. Just a thought for those would be cottage industry chemist.

[Edited on 26-1-2012 by AJKOER]

weiming1998 - 25-1-2012 at 20:10

If I have access to buying things online, I can buy H2SO4 instead. And not very expensive at that. Cheaper than buying KI or HIO3.

symboom - 1-2-2012 at 18:07

what about bubbling so2 first making sulfurous acid then chlorine into water. Cl2 + H2O + H2SO3 > HCl + H2SO4

AndersHoveland - 1-2-2012 at 18:26

Quote: Originally posted by symboom

what about bubbling so2 first making sulfurous acid then chlorine into water. Cl2 + H2O + H2SO3 > HCl + H2SO4

Yes, that will oxidize the SO2 to sulfuric acid.
But the question is what is the maximum concentration of sulfuric acid obtainable?

For example, will bubbling chlorine and sulfur dioxide into 50% concentrated sulfuric acid solution absorb more of the water, further concentrating it?

Also, there might be a maximum concentration achievable, because there is reason to believe that 99 percent concentrated sulfuric acid oxidizes HCl to chlorine. And similarly, pure anhydrous sulfuric acid (free from water) might not be obtainable from the hydrolysis of sulfuryl chloride with water.

AJKOER - 1-2-2012 at 21:08

Quote: Originally posted by symboom

what about bubbling so2 first making sulfurous acid then chlorine into water. Cl2 + H2O + H2SO3 > HCl + H2SO4

I share the concerns on how far the reaction can be pushed to the right.

However, treating a hot solution with air (O2) could remove the HCl (see Lead Chambers thread somewhat detailed discussion on this reaction). Also, it is know that:

4 HCl + O2 --Heat--> 2 H2O + 2 Cl2

and could further move the reaction to the right. So overall:

4 Cl2 + 4 H2O + 4 H2SO3 + O2 --> 2 H2O + 2 Cl2 + 4 H2SO4

or, upon netting:

2 Cl2 + 2 H2O + 4 H2SO3 + O2 --> 4 H2SO4

or, an interesting equivalent:

2 Cl2O + 2 H2O + 4 H2SO3 ---> 4 H2SO4

which equates to a known reaction (see Watt's):

4 HOCl + 4 H2SO3 ---> 4 H2SO4

namely, the oxidation of Sulfurous acid by Hypochlorous acid.

weiming1998 - 2-2-2012 at 03:18

Quote: Originally posted by AJKOER

Quote: Originally posted by symboom

what about bubbling so2 first making sulfurous acid then chlorine into water. Cl2 + H2O + H2SO3 > HCl + H2SO4

Air in a hot solution? The problem is, when you heat a solution of H2SO4 and HCl, the equilibrium is shifted to the left: Cl2+SO2+2H2O<======>2HCl+H2SO4. Both sulfur dioxide and chlorine will bubble out of the solution before you manage to oxidize the HCl (heat reduces solubility of these gases). So you would end up with a (very) dilute solution of H2SO4.

Also, more HOCl chemistry?!?!

AJKOER - 3-2-2012 at 06:32

There is some validity in your comment as, for example, the quoted standard preparation of HIO3 via Iodine in water treated with Cl2:

I2 + H2O <===> HIO + HI

3 HIO --> 2 HI + HIO3

Cl2 + 2 HI --> I2 + 2 HCl

but, per one of my sources, the overall reaction does not completely move to the right without the presence of Ag2O to remove the HCl!

However, in the current case, to quote from the allured to Lead Chamber thread:

Quote: Originally posted by AJKOER

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),
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

symboom - 3-2-2012 at 13:50

my thought is the more chlorine is added to the sulfurous acid of course cooling it with an ice bath so SO2 is absorbed more in the water forming Sulfurous acid at the same time bubbling clorine gas
forming HCL and Sulfuric acid.

now i got a question there must be a point where the concetration of suluric acid gets strong enough to release

just like this reaction
2 CuCl2 + SO2 + 2 H2O → 2 CuCl + 2 HCl + H2SO4

just with out the copper
Cl2 + SO2 + 2 H2O → 2 HCl + H2SO4 so you dont have to filter out the copper 1 chloride

as the concentration of sulfuric acid increases as does the hcl
the sulfuric acid takes the water from the hcl and the hcl just comes off as a gas as the concentrated sulfuric acid forms

so i think this happens
first Cl2 + SO2 + 2 H2O → 2 HCl(liquid) + H2SO4(liquid)
as the hcl dissolves in the water to and the acid.

the concentration of both acids increases much as the sulfuric acid takes the water from the hcl as the hcl has nothing to dissolve in it comes of as a gas.
Cl2 + SO2 + 2 H2O → 2 HCl(gas) + H2SO4(liquid)

weiming1998 - 3-2-2012 at 16:23

Quote: Originally posted by symboom

my thought is the more chlorine is added to the sulfurous acid of course cooling it with an ice bath so SO2 is absorbed more in the water forming Sulfurous acid at the same time bubbling clorine gas
forming HCL and Sulfuric acid.

now i got a question there must be a point where the concetration of suluric acid gets strong enough to release

just like this reaction
2 CuCl2 + SO2 + 2 H2O → 2 CuCl + 2 HCl + H2SO4

just with out the copper
Cl2 + SO2 + 2 H2O → 2 HCl + H2SO4 so you dont have to filter out the copper 1 chloride

as the concentration of sulfuric acid increases as does the hcl
the sulfuric acid takes the water from the hcl and the hcl just comes off as a gas as the concentrated sulfuric acid forms

so i think this happens
first Cl2 + SO2 + 2 H2O → 2 HCl(liquid) + H2SO4(liquid)
as the hcl dissolves in the water to and the acid.

the concentration of both acids increases much as the sulfuric acid takes the water from the hcl as the hcl has nothing to dissolve in it comes of as a gas.
Cl2 + SO2 + 2 H2O → 2 HCl(gas) + H2SO4(liquid)

So a chloride salt can be used instead of chlorine? But the hydrochloric acid in there wouldn't come off as a gas, it would come off as chlorine. the hydrochloric acid will also decompose the sulfuric acid formed to SO2 and H2O. If there is no way of removing the hydrochloric acid, then this method won't work.

aliced25 - 17-4-2012 at 08:30

weiming - there is a paper online about the freeze-thaw cycles with sulfurous and nitrous acids, to give sulfuric & nitric acids respectively after several cycles.

I've done it and the pH of the solution changed dramatically over several cycles, check it out and see what you find. You'll also find out what clathrates are (I did

), when the funky ice melts it is time to refreeze it.

The other alternative is to find an old PC power supply, then use that to build a lab power supply (I believe there is a silicon chip magazine issue online that costs about $7AU detailing precisely that). Then use that to oxidize the SO2. I may have to go that road myself in the near future, getting H2SO4 in Oz is a biatch.