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sulphur Dioxide

Titanium100 - 20-5-2015 at 16:48

Hi all,
Could someone help me to understand how much sulphur could be burned in a volume of 200 litres of air?
I am wanting to try making sulphuric acid from basic starting materials, and am just thinking about the volumes and masses involved.
Also if all the sulphur dioxide converted to trioxide and dissolved in water, how much sulphuric acid would be produced?

Thanks for any help

Tim

blogfast25 - 20-5-2015 at 17:07

Quote: Originally posted by Titanium100

Hi all,
Could someone help me to understand how much sulphur could be burned in a volume of 200 litres of air?
I am wanting to try making sulphuric acid from basic starting materials, and am just thinking about the volumes and masses involved.
Also if all the sulphur dioxide converted to trioxide and dissolved in water, how much sulphuric acid would be produced?

Thanks for any help

Tim

Wow. You want to prepare sulphuric acid on quite a scale but can't even do the basic stoichiometry?

Just when I thought I'd heard it all...

Titanium100 - 20-5-2015 at 17:20

Thanks for the helpful reply Blogfast!

This is a thought experiment only at this stage.
But if you don't want to offer any assistance, why do you bother replying at all? Ask yourself what the real reason for your reply is??

Amos - 20-5-2015 at 17:40

Well, it seems you know how to find the equations for the conversion of elemental sulfur to sulfur dioxide, and you can probably work out the further oxidation of sulfur dioxide to sulfur trioxide. From there you can find out what ratio of oxygen to sulfur by mass will be needed for the amount of reactants to be stoichiometric.

The density of air varies widely depending on temperature; you'll need to find a value for the temperature you're imagining, as this will allow you to calculate the mass of 200 liters of air.

The earth's atmosphere is about 21% oxygen by mass, but according to this paper, the limiting oxygen concentration of sulfur dust is about 9% oxygen. So only 12% of the mass of atmospheric air will be usable oxygen in this reaction.

So all you need to do is determine the mass of oxygen that will be consumed given you insert an ideal amount of sulfur into the system, and use the mass ratio of the reactants in the equations you derived earlier to find the mass of sulfur you would need. Assuming everything goes as planned and you allow the reaction adequate time to complete.

As far as finding the mass of sulfuric acid produced, just calculate how much sulfur trioxide would be produced from the burning of the amount of sulfur you use, and then use the reaction equation for the production of sulfuric acid to find how much would be produced at 100% conversion.

You may want to use this calculator for your math and for balancing equations, if that gives you trouble.

[Edited on 5-21-2015 by Amos]

blogfast25 - 20-5-2015 at 18:20

Quote: Originally posted by Titanium100

Ask yourself what the real reason for your reply is??

Because you're probably an idiot?

Do you even have the foggiest idea how to oxidise SO2 to SO3? That you can't even 'dissolve SO3 in water'? Or are you just the latest troll?

[Edited on 21-5-2015 by blogfast25]

j_sum1 - 20-5-2015 at 19:03

Blogfast is right -- if a little terse in this particular instance.

The route that looks good on paper:
S --> SO2 --> SO3 --> H2SO4
This is fraught with difficulty and not an insignificant amount of danger.
If you are after H2SO4, there are numerous other feasible routes. None are without their challenges. Many have been discussed here extensively.

The simplest method for most people is simply to buy some. You might need to purify, but that can be done too.
If you can't buy, you might want to look at electrolysis of CuSO4. If that does not grab you then there is the lead chamber process. That method even has its own sticky. Happy reading.

Let me reiterate, there is a lot on these boards on sulfuric acid. There is no short-cut to reading. You will find answers to your questions and also answers to questions that you have not got the experience to even consider at this stage.

Titanium100 - 20-5-2015 at 19:24

Dear Blogfast,
You seem very quick to stoop to abuse? Why is that?

I was under the impression that Nitrogen Oxide will catalyse the reaction of sulphur dioxide to trioxde. Are you also saying that sulphur trioxide will NOT dissolve in water ?? Interesting.

Amos - 20-5-2015 at 19:29

Quote: Originally posted by Titanium100

Dear Blogfast,
You seem very quick to stoop to abuse? Why is that?

I was under the impression that Nitrogen Oxide will catalyse the reaction of sulphur dioxide to trioxde. Are you also saying that sulphur trioxide will NOT dissolve in water ?? Interesting.

You're coming off a little uppity for not having the established background or demonstrable knowledge of the person you're arguing with. For example, what do you mean by nitrogen oxide? There are 4 principal oxides of nitrogen that are discussed with some regularity.

Yes, sulfur trioxide can dissolve and react with water. But that's not typically what happens when you toss the two together. Instead, you get a container full of water and a huge cloud of sulfuric acid aerosol that exits the reaction vessel with sometimes near explosive force. Is that something you might have wanted to know before you went about trying to synthesize sulfuric acid? It is if you like having skin.

The industrial process for producing sulfuric acid involves dissolving the sulfur trioxide in concentrated sulfuric acid to produce oleum, which can be more safely diluted with water, causing the dissolved SO3 to react and form sulfuric acid.

[Edited on 5-21-2015 by Amos]

Bot0nist - 20-5-2015 at 19:32

Yes, very exothermic. Streams of produced SO 3 are typically bubbled through ice cold and concentrated sulfuric acid.

as Amos said.

[Edited on 21-5-2015 by Bot0nist]

j_sum1 - 20-5-2015 at 21:10

Quote: Originally posted by Amos

Yes, sulfur trioxide can dissolve and react with water. But that's not typically what happens when you toss the two together. Instead, you get a container full of water and a huge cloud of sulfuric acid aerosol that exits the reaction vessel with sometimes near explosive force. Is that something you might have wanted to know before you went about trying to synthesize sulfuric acid? It is if you like having skin.

There you go, Ti100. In explicit and unambiguous terms. SO3 is naaaaasty. Experience and knowledge is needed before handling it. (More experience and knowledge than I have!) Anything less is foolhardy. Or, as bloggy eloquently put it, you are probably an idiot to consider it.

Now, blogfast is one whose experience and knowledge I would trust. He is not abusing you. He is unequivocally stating that what you propose is bloody dangerous and you need a whole lot more understanding before you attempt such a route.

I'll say it again, do some reading. There is plenty here on the topic. (Search on me if you like -- within the last 24 hours I have spelled out in detail my method for producing H2SO4.)

Oscilllator - 20-5-2015 at 21:11

Titanium100 it sounds like you could do with a bit of context on how hard what you are proposing is. I think you should read through the entirety of This thread before posting further on the topic, as it shows in detail one member(Axehandle) efforts to achieve just what you are proposing.

woelen - 21-5-2015 at 02:11

A practical way to make sulphuric acid via intermediate SO2 from accessible compounds is as follows:

Get some sodium metabisulfite or potassium metabisulfite. These are easy to obtain in most countries. They are used for wine making (cleaning vessels and bottles) or as sulfite additive.

When you add 10% HCl to the solid, then SO2 is formed. Slight heating is needed to drive off the gas, as it is soluble quite well in water. If you use 20% HCl then the process runs even smoother. Do not use stronger than 20% HCl. Making SO2 from metabisulfite really is a nice and comfortable process. It is MUCH easier than burning of sulphur in air. You get the gas in pure form.

Lead the gas through ~10% H2O2. Best is to assure small bubbles, so that all gas is absorbed. The H2O2 quickly oxidizes the SO2 to H2SO4. Keep on bubbling SO2 through the H2O2 until you get a clear smell of SO2. Then heat the liquid to boiling in order to drive off all excess SO2. Boiling for a few minutes makes it practically free of SO2 (at least that much that there is no smell left if it, I tried this myself).

In the above way you can get H2SO4 at a concentration of 20% or so without too much hassle and using compounds which are easy to obtain in most parts of the world.

Amos - 21-5-2015 at 12:00

In case anyone still cares, the OP wasn't really suggesting a very large scale for the procedure. I think the amount of air they were looking to make use of only corresponds to about 19-20 grams of sulfur.

blogfast25 - 21-5-2015 at 12:21

Quote: Originally posted by Amos

In case anyone still cares, the OP wasn't really suggesting a very large scale for the procedure. I think the amount of air they were looking to make use of only corresponds to about 19-20 grams of sulfur.

Off the cuff: 200 L of air at STP is approx. 8.3 mol of air. Mol fraction of O2 approx. 0.2, so 1.7 mol O2, corresponding to about 53 g of sulphur (about 160 g of sulphuric acid).

Still pretty stupid if you don't know what you're doing, if you ask me.

GrayGhost - 21-5-2015 at 14:48

Quote: Originally posted by Titanium100

Sulphur trioxide is not disolve in water because the reaction is violent, disolve in sulphuric acid, this pass to fumante, then is add water, down concentration to 98%.

El trioxido de azufre no se disuelve en agua porque la reaccion es violenta, se disuelve en acido sulfurico, éste pasa a fumante, entonces se agrega agua, se baja la concentración a 98%.

[Edited on 21-5-2015 by GrayGhost]

Hawkguy - 21-5-2015 at 18:30

Hey I've had a the same project as you ("Tim") for awhile. I've found that it just isn't effective to measure volumes of air but instead do it differently. You want to burn Sulfur in a sealed container. The container should be metal. You want to have a tube feeding air into the container, to support the burn, and another tube leading out. The tube leading out of the container should be of a fairly unreactive material, I just used glass. It should lead into a vessel of very cold water, to dissolve as much Sulfur Oxides as possible. Air is also bubbled through the water. If you have it, it helps to mix some potassium nitrate with the sulfur. The water will lose its smell as it is oxidized from sulfurous to sulfuric acid. I hate this procedure because it is so bad and inefficient and difficult. Good luck. After you get the sulfuric acid solution, filter it oout, boil it down and distill it. You will need a high temperature heating mantle for this. Not as easy as it sounds.

annaandherdad - 22-5-2015 at 01:58

Quote: Originally posted by Titanium100

Dear Tim,

In my opinion your question is intelligent and worthy of a respectful answer, although you are certainly a beginner. And I note that you posted in Beginnings. SM is a moderated forum, but that doesn't compel people to be polite. The material on SM ranges from elementary to extremely sophisticated, as you will see if you just peruse some subjects. It's easiest to find them by searching on google; mention sciencemadness and the topic you're interested in (making sulfuric acid and sulfur trioxide have been covered in depth in this forum). The forum has expectations of the research that someone will do before posting a question, which you will find out about if you read the rules. If you don't do your research you are likely to get pounced on. Some people on SM think that if they know a lot it gives them an excuse to be rude to newbies, but there are others, like woelen, who is extremely knowledgeable and invariably polite. If you want to have some enjoyable reading, just search on woelen's posts, or visit his web sites.

If you want to calculate how much chemicals are used in a reaction, it's called stoichiometry; it sounded like you weren't familiar with it from your original post, although there are additional complications about what percentage of the oxygen in the air is usable when burning sulfur. This is taught in introductory chemistry courses, but I don't think lack of knowledge of that should prevent someone from posting here (in Beginnings). Everyone has to start somewhere. This is my opinion, some others may not agree, but in any case there is no excuse for rudeness.

Titanium100 - 23-5-2015 at 04:31

Thanks to those who offered constructive comments.
As I mentioned I am thinking about the process only, and yes I have read the threads on the lead chamber process, and the sulphuric threads many times. In fact I have been reading the posts on SM almost every day for about six years.
My desire to experiment with a chamber process for sulphuric, is that it is not available OTC here, and H2O2 is also difficult and expensive, otherwise I would follow the method suggested by Woelen. (BTW the reactions calculator that you linked is excellent Amos.)
I appreciate that some want to spell out the risks of experimenting with certain things, and I can assure you all that I have been exposed to some very hazardous situations in my 56 years, and have so far survived with everything intact, probably in part due to having a strong sense of survival and of thinking about what could possibly go wrong before doing anything and taking appropriate safeguards.
WRT the chamber process, I cant see the sense in mucking around for 50mls of product, and I don't believe that a 200 l container would be unmanageable. I have recently completed a birkland eyed reactor that is going well, and thought that I may be able to divert the NO2 from that as an input for the H2SO4
system. Once again thanks to those that are able to offer calm measured advice/assistance. I am a believer in the benefits of collaboration, and look forward to being able to give something back at some stage.
T

j_sum1 - 23-5-2015 at 05:01

Well, if that is where you are at, then I recommend working out the largest container that you can comfortably boil down (maybe in a large glazed ceramic pot or slow cooker) and electrolysing batches of CuSO4 of that size.
Say you are working with 4L batches. You can potentially dissolve nearly 1.3kg of copper sulfate pentahydrate in that. A strong plastic bucket would work as a cell. You can use a cylinder of lead roof flashing as an anode and a scrunched up ball of copper wire as a cathode. You can use a PC Power supply or other suitable power supply. You don't want the voltage too high because it gets inefficient if you electrolyse your water. You run it for a couple of hours to reach a steady state and then measure the current. From there you calculate the time it will take to fully electrolyse -- likely a few days. When the liquid is clear you run it for a couple of hours longer to make sure all of the Cu2+ is gone from the solution. Transfer to your boil-down vessel and reduce in volume. Once the volume is low enough you transfer to your lab glassware and keep boiling. Eventually thick white fumes are given off which represents a concentration around 85-90%. If you want higher concentration then you can keep boiling but you won't get better than 98%. At this stage you should be collecting your distillate which will be reasonable strength acid itself. 1.3kg of copper sulfate will give you about 275mL of concentrated acid per batch.

Yes it is tedious. But a lot of the process is set and forget, once you have your gear organised. It is a lot less mucking around than other methods I have looked at.

Amos - 23-5-2015 at 05:44

Quote: Originally posted by blogfast25

Quote: Originally posted by Amos

I took his post to mean that the sulfur dioxide needed to be further oxidized by the oxygen in the air: 2 S + 3 O2 = 2 SO3

And I'm not saying my source was perfect(see my first post in the thread) but it stated that a minimum of 9% atmospheric oxygen was needed to burn sulfur in powder form.

[Edited on 5-23-2015 by Amos]

Molecular Manipulations - 23-5-2015 at 07:14

OP could have been a little more specific. I and likely others thought he wanted to try the contact process (very hard, very dangerous), while he really wanted to attempt the much easier chamber process. This is why he received Bloggy's not-so-nice reply (having browsed here for six year, you should have seen it coming).
The chamber process was actually the first non-textbook experiment I did three or four years ago. You should learn chemistry basics first though.