Sciencemadness Discussion Board - Sulphuric acid

Sulphuric acid

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blogfast25 - 21-5-2014 at 04:46

98% nitric acid will burn within 5 seconds, and will leave yellow stains no matter how fast you wash it off. It catches all OTC gloves on fire within 10 seconds, so wearing no gloves is probably better than any at all.

You have evidence for this 'catches all OTC gloves on fire within 10 seconds'? Sounds exaggerated to me, it's not THAT strong an oxidiser. A reputable link would be nice...

Even 35 % nitric will leave yellow stains on skin fairly quickly but they're only stains and fade to nothing over a week or so. It's very superficial attack of the epidermis only.

Zyklon-A - 21-5-2014 at 08:28

Hmm, I probably spoke too soon. No, I don't have evidence. The only gloves I can find locally are Latex or Nitrile - both of which catch on fire spontaneously upon contact with fuming nitric acid.
https://www.youtube.com/watch?v=aBVdGGml6bU

[Edited on 21-5-2014 by Zyklonb]

macckone - 21-5-2014 at 08:36

Quote: Originally posted by blogfast25

Quote: Originally posted by Zyklonb

The three common OTC gloves are nitrile, latex and some
kind of clear plastic for food service that isn't identified.
As opposed to mail order that are more resistant and
designed for chemical use (butyl and pfa for example).
Those three OTC gloves do react 'vigorously' with RFNA and
are too thin to provide protection. The thicker dishwashing
gloves also react (latex). They do 'smoke' almost immediately
but I wouldn't call it 'catching fire' although they will combust
and don't provide any protection given RFNA.

Hardware stores sometimes carry other varieties of gloves as
do cleaning supply places but they aren't as readily available.
I guess it depends on how you define OTC and 'catch fire'.
Someone should probably do an experiment with video.

macckone - 21-5-2014 at 08:47

Speaking of sulfuric acid and combustion, has anyone tried
burning MSM (Methylsulfonylmethane aka dimethyl sulfoxide).
It is much more readily available than sulfur or sulfuric acid
drain cleaner or even battery acid in a box. It is available
at practically every health food store. In theory it could
be used like sulfur (although more expensive) but it isn't
as flammable and may not burn cleanly.

I haven't found anything in literature other than on the MSDS
that it yields sulfur dioxide and carbon monoxide when burned.
Any literature references or practical experience would be helpful.

blogfast25 - 21-5-2014 at 09:40

Zb:

Well, I didn't expect that, TBH.

Zyklon-A - 21-5-2014 at 10:32

When I first burned my hand with fuming nitric acid, my mom asked me why I wasn't wearing gloves - rather than tell her, I showed her. All it took was about 5 drops of RFNA on a Nitrile glove. (They caught fire and cracked the glass dish from the heat). She was quite surprised to say the least.
Ok I just got some of those annoyingly large clear plastic food grade gloves. When I get home I'll test them against 98% sulfuric acid and RFNA. I doubt they'll catch fire, but they will almost certainly get holes burned through them. (With FNA, not con. sulfuric acid).

aga - 21-5-2014 at 12:22

The drain cleaner + peroxide worked.

Now have some sulphuric acid at last, whch is clear and fairly pure.
Certainly purer than the garbage iwas playing with before.

My hotplate can't reach 300 C +, so i stopped after about 30 minutes of what i believe was 'fuming'.

Titration with NaOH solution shows it to be about 12 Molar, which his good enough for me, for now.

Zyklon-A - 21-5-2014 at 14:18

Quote: Originally posted by Zyklonb

When I get home I'll test them against 98% sulfuric acid and RFNA. I doubt they'll catch fire, but they will almost certainly get holes burned through them. (With FNA, not con. sulfuric acid).

Incredible news! I just tested the "annoyingly large clear plastic food grade gloves". I made a flat surface by placing the gloves over the mouth of a beaker. Then I added several drops of ~97% nitric acid. Nothing happened! 15 minutes later the gloves were not even etched. Then I added several drops of 98% sulfuric acid. Still nothing. I'm very amazed that the crappy food grade gloves managed while the "ultra chemical resistant" gloves all failed.
Anyway, I hope this answered your question aga!

[EDIT] Here's the same brand of gloves I used: http://www.anything4restaurants.com/products/textra-textured...
They are called Textra ^® Cast poly gloves. By FOODHANDELER ^™.

[Edited on 21-5-2014 by Zyklonb]

blogfast25 - 22-5-2014 at 04:55

Quote: Originally posted by aga

My hotplate can't reach 300 C +, so i stopped after about 30 minutes of what i believe was 'fuming'.

Titration with NaOH solution shows it to be about 12 Molar, which his good enough for me, for now.

Concentrated H2SO4 is about 18 M. Are you sure what you saw was real fuming and not just water vapour (steam) coming off? At 70 % I didn't see anything that qualifies as fuming.

Do a density measurement? Weigh a small measuring cylinder (w0

. Then weight it full with water (w1

. Then weight it full with acid (w2

.

dacid = (w2 - w0

/(w1 - w0

(density relative to water)

[Edited on 22-5-2014 by blogfast25]

blogfast25 - 22-5-2014 at 05:02

Zb:

It's perhaps less surprising than you think. These gloves are almost certainly made of LDPE and that is very chemically resistant. HDPE is routinely used for permanent storage of 70 % nitric acid and 98 % H2SO4. LDPE and HDPE are chemically almost identical, but HDPE is slightly more crystalline and therefore mechanically tougher.

Zyklon-A - 22-5-2014 at 06:59

So I guess my question is, why don't chemical resistant gloves imploy such plastics? Or at least mention in papers (like the one I linked) that those gloves do exists and are suitable for handleing fuming nitric acid.

blogfast25 - 22-5-2014 at 08:43

Quote: Originally posted by Zyklonb

So I guess my question is, why don't chemical resistant gloves imploy such plastics?

One possible explanation is that both latex (usually NR; thus poly isoprene, or SBR; thus poly butadiene styrene) and nitrile (more scientifically acrylonitrile butadiene rubber, aka NBR) contain a lot of double bonds, which obviously makes them prone to attack by oxidisers.

These rubbers would normally not be recommended as 'highly chemically resistant' although nitrile does have the advantage of being quite resistant to paraffinic and aromatic solvents due to its polarity (hence the use of NBR in fuel lines, for example).

LDPE, HDPE, PP and EPR are fully saturated and well known to resist most chemical attacks well.

http://en.wikipedia.org/wiki/Nitrile_rubber

http://en.wikipedia.org/wiki/Synthetic_rubber

[Edited on 22-5-2014 by blogfast25]

blogfast25 - 23-5-2014 at 07:26

According to one source, adipic acid (hexanedioc acid) can be prepared by oxidative cleavage of cyclohexene with nitric acid, an example of oxidation of a double bond with nitric acid.

Since as oxidations are usually strongly exothermic, reaction of a material plentiful with double bonds with fuming nitric may be sufficiently exothermic to actually ignite the material (and combustion in air to take over).

macckone - 23-5-2014 at 08:19

Interesting that the food grade gloves survived the nitric.
I guess different brands are going to give different results
based on what they are made of.

Still curious about combustion of MSM though.

aga - 23-5-2014 at 14:02

Quote: Originally posted by macckone

Still curious about combustion of MSM though.

Gay, Straight or just MSM Curious, experimentation is the way forward !

macckone - 23-5-2014 at 14:47

Quote: Originally posted by aga

Quote: Originally posted by macckone

Still curious about combustion of MSM though.

Gay, Straight or just MSM Curious, experimentation is the way forward !

I don't currently have a hypothesis because of the lack of literature.
It isn't scientific experimentation without a hypothesis.

blogfast25 - 24-5-2014 at 04:07

Quote: Originally posted by macckone

Still curious about combustion of MSM though.

Garden variety sulphur is even more available and will be cheaper per g of SO2 generated.

Fulmen - 12-11-2014 at 05:36

This seems to be the most current thread on concentrating H2SO4, so I'll try here rather than starting a new.

I just tried boiling in some SA (needed to make nitric acid from SA and calcium nitrate), and it worked much better than expected. My hotplate (Ikamag RCT) is limited to appr 300°C, so I just let it run until it stopped boiling. According to this: http://www.generalchemical.com/assets/pdf/Sulfuric_Acid_Wate... you should get huge losses long before you get anywhere near 90%, yet I got a density > 1.8 with only mildly unpleasant vapors coming off.

And then it struck me. I started with 250ml of spent battery acid in a 300ml EM-beaker, by the end I was down to appr 50ml. Considering the high temperature and almost empty beaker, wouldn't it act like a crude rectifying column? The Liquid/vapor equilibrium is a bit odd compared to simple mixtures like water/ethanol, am I wrong in assuming that it makes SA very easy to distill to high concentrations?

[Edited on 12-11-14 by Fulmen]

bbartlog - 12-11-2014 at 06:40

MSM seems unlikely to me to be a good source material for anyone who wants well-defined products. Too much carbon and hydrogen in the mix. I guess if you were careful to provide excess oxygen and high temperatures you could at least avoid contaminating your sulfuric products with unburned carbon products aka soot, but you would still end up with added water as one of the products of combustion.

greenlight - 12-11-2014 at 08:21

Fulmen, Sulphuric acid that is intended for use in batteries is very dilute to start so that may explain why you were only left with 50ml.
Using an Erlenmeyer flask (if that's what you mean by EM flask) I think there would be some condensation on the sides of the flask.
You say you used a 300 degree max hotplate and just turned it on till boiling temp. The boiling point for concentrated Sulphuric acid is 337 degrees Celcius so if you had the hotplate turned up all the way and you shot near that temperature maybe some of the acid decomposed too.
When I used to concentrate 50% Sulphuric acid with heat I used to do it in a glass saucepan or glass heat resistant dish with a thermometer in the dilute acid. I heated the acid to just past 100 degrees Celcius usually to about 110 to 115 degrees Celcius, just above the BP of water. This way there were no nasty fumes from the acid boiling when concentrating and I didn't need to bring the acid to a boil at all.
You can tell when it is concentrated because slight white fumes start to come off and the colour changes to much darker. If you hold a piece of flat glass like the bottom of a large beaker or a watch glass over the heating vessel for like 10 seconds you will notice the water from the acid mix boiling off and condensing on the glass making it foggy. This is another way to check if it is finished because once it stops, no more water is coming off and it is concentrated.

[Edited on 12-11-2014 by greenlight]

Antiswat - 12-11-2014 at 08:29

i recall 50% H2SO4 is called ''tower acid''
some time back i found people on ebay selling 50% H2SO4 for swimming pools use, it was pretty large amounts, and surely they are still trading it

i guess for concentrating it down, if you are lucky you could perhaps get some high temperature resistant glass tube and fill with some sort of a substrate with V2O5 on it, and run SO2 / air through it
yields will be increased with dry air
then run through H2SO4.. the SO3 formed should remove water from the H2SO4 and may form oleum if let arise in concentration past 98%

one guy on SM talked alot about a machine like this he had he used alot, but he never got to post pictures about it, sadly.. he said it was full glass, and SO3 through H2SO4 is ofcourse a preferred method for concentrated sulfuric acid, considering SO3 everywhere as the alternative

Fulmen - 12-11-2014 at 08:51

The reduction in volume was to be expected, and I also assume there has been some decomposition as it turned yellow at the end. The interesting bit was how concentrated the result seemed to be with very little acid vapors. Judging the fumes is hard as I do this outside at windy 40-50F temps, even water produces a white fog. I'm going to run a few more batches and see if I can't do some temperature measurements of both the liquid and vapor coming off and perhaps a simple titration.

Fulmen - 12-11-2014 at 12:09

Crud, my only suitable thermometer has died, so no luck there. But a sample of the concentrated acid started freezing at appr 9°C (40F), which should put it in the 80-90% range. Not bad really, considering the setup. This warrants further experimentation.

vmelkon - 12-11-2014 at 17:09

Quote: Originally posted by Fulmen

Crud, my only suitable thermometer has died, so no luck there. But a sample of the concentrated acid started freezing at appr 9°C (40F), which should put it in the 80-90% range. Not bad really, considering the setup. This warrants further experimentation.

I suggest measuring the density. You can buy a cheap digital scale for 10$ on ebay. Mine goes from 0.01 g to 500.00 g but precision is +/- 0.03 g.
There is another one that does 0.001 g to 50.000 g but.... how good is the precision on these things.

Also, get yourself a good graduated cylinder. Preferable a thin one. Or get a burret (I have a bunch with accuracy to 0.02 mL or something like that).

Fulmen - 13-11-2014 at 03:46

I have scales, the problem is to measure volumes accurately. I think a weight-based titration will be the best route, I'll see what I can whip up.

Zyklon-A - 13-11-2014 at 08:55

You could use a syringe or something like that. Or you could tintrate it against a know amount of barium nitrate solution.

Fulmen - 13-11-2014 at 09:35

What sort of accuracy can one expect from a disposable plastic syringe? I tested the first batch this way with a 1ml syringe, it gave me a density of 1.8. But once you reach 80-90% you really need a resolution of 0,01g/ml to determine concentration, and I just don't believe this is achievable with make-shift equipment.

Right now I'm pondering if it's possible to use the weird freezing points to both concentrate and determine concentration. A combination of boiling and fractionate freezing could possibly give a useful result without having to breathe in too much noxious gas.

blogfast25 - 13-11-2014 at 11:10

Quote: Originally posted by Fulmen

What sort of accuracy can one expect from a disposable plastic syringe? I tested the first batch this way with a 1ml syringe, it gave me a density of 1.8. But once you reach 80-90% you really need a resolution of 0,01g/ml to determine concentration, and I just don't believe this is achievable with make-shift equipment.

A glass measuring cylinder of 10 ml, used relative to deionised water, should give fairly good accuracy. You can't beat a decent pycnometer with mg scales of course, for 4 or more significant digit densities.

And there's always good old acidometry, easier and faster than you might think if you're well organised.

Fulmen - 13-11-2014 at 13:05

The acid is undergoing some freezing experiments now, so it'll have to wait until morning. I'm just so used to quality precision equipment I have a real time trusting these improvised methods. A titration seems like the best option, I'll have to dig around to find a suitable indicator.

After a bit of reheating the acid now freezes partially (perhaps 3/4) at -20°C. This should indicate appr 90%, decanting the liquid should give me 92% according to this: http://www.generalchemical.com/assets/pdf/Sulfur_Trioxide_Wa...

This seems to be the limit with the lab hotplate, I'll try again with another hotplate and a 500ml flat-bottom florence flask. Hopefully the long stem will act like a crude column and add a few % of separation without too much loss. Would be nice to reach 05%, but I'm not complaining abut the results so far.

Fulmen - 14-11-2014 at 05:36

This is weird. After freezing at -20°C I got a waxy solid precipitation at the bottom that would not melt even at room temperature. So I did a couple of titrations using 1 & 10% NaOH (w/w) and litmus indicator, the NaOH was technical grade (drain opener, not sure what purity one can expect).

According to this the liquid phase was 95% while the solid was 87%. I'll have to recheck my results, but it could indicate that fractional freezing could work. Only problem is the fact that sulfuric acid should be liquid at room temperature, regardless of concentration. Does anybody have a better theory?

I also found some K2CO3 puriss (min 98%) and haematoxylin indicator (pH 5-7). Would titrating the acid with the carbonate work reliable? I can't seem to find any examples of this.

bbartlog - 14-11-2014 at 07:38

Sulfuric acid has a monohydrate (corresponding to 84.5% concentration) which apparently has a fairly high melting point - above freezing, anyway. Phase diagram: http://image.bayimg.com/aakdgaabi.jpg

I'm guessing that's what you have, though 'room temperature' still seems like it should be high enough to melt this.

Fulmen - 14-11-2014 at 08:57

Actually this is something else completely. Heating to 80C partially dissolved the solid, but it solidified again upon cooling. I have no idea what this is, perhaps a salt of some sorts? It's soluble in water but not cold concentrated acid.

Metacelsus - 14-11-2014 at 09:12

Sounds like a bisulfate of some sort.

Fulmen - 14-11-2014 at 09:45

Maybe. Not sure where the cation would come from, but it's pretty obvious it's some unwanted contaminate. Don't think I'll be wasting more time on it, it's just a matter of freezing it out and decanting.

I'd like to find a better method for determining strength, I don't trust the purity of the sodium hydroxide. I'll have to try titrating K2CO3 with diluted acid to see if that gives a different result.

blogfast25 - 15-11-2014 at 06:55

Quote: Originally posted by Fulmen

I also found some K2CO3 puriss (min 98%) and haematoxylin indicator (pH 5-7). Would titrating the acid with the carbonate work reliable? I can't seem to find any examples of this.

For titrating the salt of strong base and a weak acid (carbonic acid) with a strong acid you need an indicator like Methyl Orange because at end point the pH will be slightly acidic, due to carbonic acid.

Recrystallise the K2CO3 at least once, then dry it at 200 C or so for 2 h. Prepare a solution 0.1 N (i.e. 0.05 M) K2CO3 using accurate scales and a calibrated volumetric flask. Note the titer of the solution, calculated from the precise weight of K2CO3 used. This is your Standard Solution.

Prepare a solution of the sulphuric acid with an accurately weighed sample of the acid. Using your expected value of H2SO4 concentration in your acid, aim for a solution strength of 0.1 N (0.05 M). Use a closed weighing boat (H2SO4 is hygroscopic), accurate scales (1 mg, 0.1 mg is better of course) and a calibrated volumetric flask.

Titrate, repeatedly, 20.0 ml (pipette!) of Standard Solution with the H2SO4 solution and Methyl Orange, from yellow to orange (pH = about 3.8).

A relative measuring error of 1 - 2 % is within reach if executed properly.

You can use low grade NaOH but NOT without standardising the titrant solution, for instance against pure potassium hydrogenphthalate (KHP).

[Edited on 15-11-2014 by blogfast25]

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