2NaOH from H2O, NaOH, and ? 5

[AlbertG]

Good day.

I have a question for better minds than mine; as I am not a chemicals person myself )

I am interested in the prospect of changing plain sulfur dioxide scrubwater, sulfurous acid (H2SO3), into a non-tox fluid which might even have some commercial value.

I have discovered the following in my stumblings about (goodness, I wish I'd paid closer attention to this stuff in high school).

OK. Here's what I think can be made to happen, based upon the noises around the web.

If I start with Na2O -- disodium oxide -- CAS #1313-59-3; the following can be made to happen (with much excitement at first):

Na2O + H2O ---> 2NaOH

Then, run the SO2 through the prepared scrub fluid (2NaOH):

H2SO3 + 2NaOH ---> 2H2O + Na2SO3

Here are some applications of sodium sulphite (Na2SO3):

Waste water treatment in metal plating and tanning;
Photographic industry as a fixer;
Bleaching products;
Preservative for certain foods;
Introduction agent of the SO 3 group into organic molecules (production of acrylic fibers for example).
Used as a decontaminating agent in fresh orange juice, and during sugar refining.

Solubility: 23 g/100 mL water (20° C)

Not considered a hazardous waste; see Auburn University's list at:

http://www.auburn.edu/administration/rms/pdf/nonhazlist.pdf

Now, here's the sticker: Na2O reacts rather badly with water, as I understand it. Plus, it seems to be as expensive as all getout.

Here's what I'd like to do instead.

Sodium hydroxide (lye - NaOH) ionizes in water as shown by the following equation:

NaOH + H2O <---> Na+ + OH- + H2O

Is there some way to yield forth the necessary 2NaOH from a mixture of water, lye (NaOH) and something common (good ol' Morton's sodium aluminosilicate waste material (table salt) would be handy ) without creating a substantial poison as a byproduct in the process?

Hope I haven't bored anyone with my question; as I'm sure the answer is probably as plain as can be.

Thanks for helping out!

 

[moltenmetal]

http://www.marsulex.com/proven_solutions/examples/petrocan_shell.htm

Too late!

PS: sodium hydroxide is a commodity chemical. You don't need to make it from sodium oxide, you simply produce it by electrolysis of salt water.

 

[AlbertG]

Woah!

Not too bad for just a guy with an idea . . .

That's the story of my life; late as usual (

Thanks for the tip about NaOH!

Now, just one last question for my info -- how do we get 2NaOH from our lye/waterbath? Do I need to add another something in the mix to make the transition from NaOH to 2NaOH?

Teriffic help. Thanks again!

 

[Zoobie]

2NaOH is the same as NaOH. The reaction is written this way to balance the equation. For the reaction that you wrote you will need 2 moles of NaOH and hence two moles of sodium. As you can see the sodium sulphite has 2 sodium atoms in its make-up. 2 sodiums on the left = 2 sodiums on the right. This applies to all of the other atoms in the reaction. Molecules can change but the atoms have to be the same on either side of the equation (unless you are exploding atomic bombs are have built a particle accelerator in your basement).

 

[AlbertG]

Thanks Zoobie.

I thought that this was the case (somewhere in the back of my brain).

But, here's something which I found that threw me:

H2SO3 + NaOH ---> H2O + NaHSO3

Am I misunderstanding the same type of thing as before, or is sodium bisulfite (NaHSO3) a different creature from sodium sulphite (Na2SO3)?

Or, is the above reaction/product simply someone's error???

Thanks, guys, for putting up with me here )

 

[Zoobie]

AlbertG,

Both reactions are possible and will occur. Unfortunately I have switched industries and my acid/base chemistry is quickly starting to get rusty. As I recall, sulphurous acid (H2SO3) will first be neutralized to sodium bisulphite. The bisulphite will then be neutralized to sulphite. Each reaction has its own Ka (I don't have a reference handy and I don't want to wade into the waters of Ka's since it has been a long time...I'll let someone else fill in those blanks). Using Ka values and the concentrations that you have, you can determine the concentrations of each species that you will have at equilibrium. From a practical point of view, whether or not you get sulphite or bisulphite, I believe, will depend on how much base you use.

 

[AlbertG]

Thanks again, Zoobie.

What you said makes sense as:

"Sodium bisulfite releases sulfur dioxide gas when added to water or products containing water."

If I understand this particular statement well enough in this context, it would seem to be the self-regulating part of the bisulphite/sulphite mechanism...

In other words, not enough lye, sodium bisulphite and SO2; adequate concentrations of lye, sodium sulphite and 2H2O (or water)?

Thanks for the head scratchings!

Love passing out those stars ;o)

 

[UmeshMathur]

AlbertG:

Sodium bisulfite is what is called an "acid salt", capable of being further reacted to the sulfite form. The final mixture ionic / molecular concentrations depend on the solution of a number of ionic equilibrium equations. The extent of these reactions is not possible to predict without using a process simulator with an "electrolyte data bank" that contains information on the ionic equilibria for all reacting species.

Aspen Plus is an excellent package for doing such calculations. However, it is expensive and requires a lot of training to learn to use, especially for electrolyte systems. This is definitely not an area for amateurs, and I would recommend that you seek expert advice on how to do this.

You can always conduct a series of lab experiments but this is also by no means a trivial task, as you need quite sophisticated laboratory analyses and quality controls.

 

[AlbertG]

Thanks for the cautions, UmeshMathur.

When (or if) I shall undertake an endeavor such as this, I'll bring in some competent help.

However, I have one last question -- is it simply possible to "guesstimate" the presence of the end sulfite form by simply testing the pH?

In other words, add lye content to the sulferous acid (or vice versa) solution until the pH is where one would expect for this concentration of Na2SO3?

Once again, thanks.

Stars for helping )

 

[IRstuff]

I think you might find that the demand for the product substantially less than the supply of available waste acid. Or, you might find that the expense of conversion outweighs the cost of the same product produced in bulk.

Additionally, most waste acid solutions contain other compounds that would be considered contaminants or are otherwise toxic and the cost of removal of those contaminants would substantially reduce the cost effectiveness of the process. We used sulfuric acid for removing photoresist, which contains a variety of organic compounds that would be difficult to remove from the waste acid.

Finally, while your product COULD be used for your listed, the question is whether it actually is used. I suggest that you go to your local stores and find these end products and see what's on their contents lists to get confirmation of their applicability.

TTFN

 

[AlbertG]

Thanks for the thoughts, IRstuff )

Actually, a spin on this methodology is being implemented already:

http://www.marsulex.com/proven_solutions/examples/petrocan_shell.htm

Truly an interesting concept: Take a substantial menace (sulfurous acid) and make a wine preservative out of it.

Don't think that I'd personally be drawn to a glass of vino ala smokestack (yuk).

But, there is a market for this approach.

SO2 -- it's not just for fertilizer anymore!

Do you have any spin on my latest q:

"...is it possible to "guesstimate" the presence of the end sulfite form by simply testing the pH?

In other words, add lye content to the sulfurous acid (or vice versa) solution until the pH is where one would expect for this concentration of Na2SO3?"

Thanks again.

 

[UmeshMathur]

I believe that such mixtures can contain a wide range of concentrations of the various ionic species, depending on starting solution strengths and of course how much alkali you have used. Also, note IRstuff's point that other impurities would likely create many complications - it is hard to imagine that you would ever obtain reagent grade sulfurous acid from an industrial scrubber.

Of course, no one can quarrel with your original point that you are simply trying to abate a nuisance and, in the process, attempting to generate byproducts that have some commercial value. A few well designed lab experiments may help answer the basic questions about usability / saleability of your end products, as well as overall economics. The kind of simulation work I referred earlier to would be justified only if you were going to do a commercial-scale design.

 

[25362]

Exhaust gases from combustion contain another acid gas: carbon dioxide. Because of the much higher concentration of
CO₂ than SO₂ in the flue gases, the concentration in scrubber water at, say, 30^oC, is expected to be about 3 times as much CO₂ as SO₂.

The needed amount of alkali is, thus, quite larger and so is its cost. The industrial trend points to the direction of using much cheaper limestone CaCO₃, quicklime CaO, or hydrated lime, Ca(OH)₂, being all sufficiently cheap to be used in throwaway processes.

There are other "regenerable" chemicals used for wet systems
such as sodium carbonate and sodium bicarbonate without the problem of calcium salt deposition in the scrubbers.

 

[AlbertG]

UmeshMathur: Thanks again for the input.

Indeed, stack filth is likely to have all sorts of unusual items in the waste stream. You nailed it as far as what might be the most useful endproduct; a relatively tame scrubwater mass.

25362: Interesting point about the CO2 ... and what about the Baking Soda approach!

H2SO3 + NaHCO3 ---> H20 + ?

Anything evil to know about here? Wouldn't we simply release SO2 from the H2SO3 fluidmass as the baking soda was added?

Inquiring minds want to know )

Stars all 'round . . .

 

[25362]

Soda ash or baking soda are both used in double alkali scrubbers and even in dry-wet treats. They can both be regenerated by limestone at least in part. They are more efficient reagents in catching SO2 than calcium-based chemicals.

The first reaction produces sodium sulfite or bisulfite and releases CO2. The regeneration is done by limestone and oxygen.

Barring secondary reactions, one could summarize them as follows:

Na₂CO₃ + SO₂ => Na₂SO₃ + CO₂ (1)​

Na₂SO₃ + CaCO₃ + [&frac12;]O₂ => CaSO₄ + Na₂CO₃ (2)​

 

[AlbertG]

Thanks again, 25362!

So, let me see if I understand this reasonably well.

If we work with the above line of reasoning with straight SO2 in solution (H2SO3), then we should have the following result:

Sodium Carbonate as the neutralizer --

Na2CO3 + H2SO3 ---> Na2SO3 + CO2 + H2O

Sodium Bicarbonate as the neutralizer --

NaHCO3 + H2SO3 ---> NaHSO3 + CO2 + H2O

(The boards felt a little wobbly under me for that last one, but it seemed logical )

Does this look reasonably close???

Stars and thanx . . .

 

[25362]

About right. The first reaction depends on the amount of soda ash. If insufficient one may obtain:

2 SO₂(g) + Na₂CO₃(aq) + H₂O => 2 NaHSO₃(aq) + CO₂(g)​

However, one cannot isolate the NaHSO₃ by boiling off water, since complex events occur. Some Na₂SO₃, some SO₂, and some Na₂S₂O₅ will form, but no NaHSO₃. Na₂S₂O₅ is sometimes called sodium metabisulfite.

Although chemical supply houses sell what they call sodium bisulfite (or sodium hydrogen sulfite) NaHSO₃, the material in the bottle is actually the metabisulfite.

In an acidifed solution, however, both are chemically equivalent because of the following equilibrium:

2 HSO₃^-(aq) <=> S₂O₅^2-(aq) + H₂O​

Returning to our first equation, if you continue adding soda ash, you get an acid-base neutralization. The resulting solution could be vaporized to dryness to give sodium sulfite either anhydrous or as the heptahydrate depending on the temperature. BTW, Any SO₃ present would convert to sodium sulfate.

A relatively concentrated aqueous solution of SO₂, "sulfurous acid" is unstable and releases gaseous SO₂. Thus, if you carry out trials on the lab bench, use a fume hood.

In sulfites the oxidation number of sulfur is +4, intermediate in its range from -2 to +6. Thus it can act either as reducing agent or oxidating agent.

 

[AlbertG]

Thanks for all the info!

I almost feel as though I've got a hold of a corner of this now...

It sounds like it would be better to scrub the SO2 through the prepared solution of soda ash or baking soda instead of just H2O -- that way we'd be able to have adequate amounts of neutralizer on hand for the process from the outset. SO2 releases would then be far less likely during the "neutralization" phase; and simply watching the pH or monitoring for relatively small SO2 "avalanching" from the scrub fluidmass would give adequate warning of solution exhaustion.

Am I fairly close to the bullseye here?

Stars and thanks again!

 

[25362]

When regenerating the solution by addition of limestone, you must add air to precipitate calcium sulfate (gypsum) with a solubility of ~1800 ppm at ~50^oC. Some limestone would also precipitate in the thickener.

The thickener's overflow goes to a tank where water and makeup bicarb or carb are added to account for the losses on the moist cake obtained from the thickener underflow upon filtration. The "regenerated" solution would have minute quantities of calcium (if the precipitaing conditions are carefully monitored) so that no solids will form in the scrubber.

Soda ash and baking soda are also used in "dry" treatments by injecting finely ground salts into the duct carrying the exhaust gas.

Could you explain: a. what kind of process are you dealing with, to get an idea on whether yours is a lean or rich waste gas; b. what is the size of the installation for you to decide on the economics of using a throwaway process rather than a regenerating one ?