Sciencemadness Discussion Board

NaOH dissolution fumes

aeacfm - 11-11-2010 at 06:57

every time i dissolute NaOH (but in high concentrations) some fumes generated which make me cant breath !!
what are these fumes

Jor - 11-11-2010 at 07:06

These are small droplets of water containing dissolved NaOH.
Boiling solutions (but also if the aqueous solution is close to the boiling point) release very small droplets, aerosols, in some salts may be dissolved, in this case NaOH. This is the irritating smell you notice. This is not a really a problem, but don't breath it too much.

This is exactly the reason why I never boil down solutions containing mercury or lead, because this contaminates the area with small amounts of the metal-salts.

woelen - 11-11-2010 at 13:21

Jor's answer is right, but there is more to this. Why do you get NaOH in the droplets? The solution is far from boiling, but still NaOH is in the fumes from the solution, while this is an ionic compound.

On the other hand, if I boil a mix of water and sulphuric acid with e.g. 30% H2SO4 then the boiling liquid does not contain a noticeable amount of H2SO4. Even when a dilute solution of a volatile compound like HBr is distilled then initially only water comes over.

The formation of aerosols with mercury or lead can easily be explained, because of the boiling of the liquid, small droplets are ejected into the air. This is just a mechanical phenomenon. But as long as no bubbles are produced, then no lead or mercury will be present in the mist from the solution.

An interesting experiment would be to take a little piece of paper, soaked with Na2S and keep this above a hot solution of Pb(NO3)2 or HgCl2. I expect that as long as the liquid is not boiling, the paper will not become dark.
On the other hand, if you take a piece of humid pH paper and keep that above a warm solution of NaOH do you think that the paper will display a color for high pH?

One thing which I can imagine is due to the choking fumes on dissolving of solid NaOH is that very small bubbles are produced in the solution (e.g. dissolved air which is driven out, or small amounts of air, trapped in the granules of NaOH) which slowly move to the surface and cause formation of an aerosol.

Sedit - 11-11-2010 at 13:39

"But as long as no bubbles are produced, then no lead or mercury will be present in the mist from the solution."

Sorry woelen this isn't true and you can repeat my experiment to prove me wrong if you like.

I was evaporating a solution of HgCl2 and in order to eliminate the possibility of partical release like we are talking about here I used a double boiler type setup with warm water in the lower section. This water isn't even hot enough to burn ones finger but still warm enough to slowly evaporate water. I placed a piece of glass over it just to check and low and behold when the condensation evaporated I was disturbed to see HgCl2 crystals had formed on the glass meaning no matter how hard you try the Hg slowly gets out.:o . I haven't messed with Hg salts since because im positive i can not avoid contamination.

DJF90 - 11-11-2010 at 15:44

HgCl2 sublimes, as such it was known in the old days as "corrosive sublimate"... probably not a "model compound" to test this theory with.

Sedit - 11-11-2010 at 16:12

But is the vapor pressure of HgCl2 close enough to that of waters that they should "sublime" together so to speak? The solution was not really saturated or anything so I would expect the water to come over before it started to sublime.

Iv been considering testing this even further on a hot summer day by placing a piece of glass over the solution and see if even the very slight heat from the sun is enough to carry it along with the water vapors.

MagicJigPipe - 11-11-2010 at 16:58

If this were true distillation of water would not work and would still contain significant amounts of dissolved solids. Of course a molecule might stray from the solution every now and then but surely not enough to be significant. I'm pretty sure I have held pH paper above hot NaOH soln. and I remember somewhat distinctly that it did not show increased pH.

If HgCl2 is subject to sublimation it will do so at any temperature (depending on pressure) but just to a lesser extent, right? In the same way that water evaporates even at 0*C just relatively slowly. In this sense won't many if not most ionic salts "sublimate" just at a rate that renders the sublimation insignificant? I keep asking questions like this about things that happen but are not considered significant and are therefore, for practical purposes, ignored but no one will answer me. Seriously, what's going on?

Sedit - 11-11-2010 at 17:20

I was under the impression that as a substance evaporated, especially in a closed system, it exerts a pressure with the force of the vapor pressure of that substance so that as two substances where in the same system the ones sublimination point so to speak could not be reached or atlest seriously reduced until the substance exerting the higher vapor pressure was eliminated in one way or another(evaporation ect...) It just seemed like simple logical conclusion to me so I never gave it much thought to be honest until I seen the HgCl2 come out of water regardless of the temperature used.

Surely the gurus around here could clear this up for us laymen once and for all....right?

MagicJigPipe - 11-11-2010 at 18:19

The key words here are "seriously reduced". The HgCl2 vapor is there, but in what concentration?

I think I remember Raoult's law(?) from Gen. Chem. 2. Now I will attempt to relearn it and apply it. Maybe it will tell us something about this problem (Sedit)?

According to an MSDS the vapor pressure of HgCl2 is:

1 @ 136.2C

But of course we can't use this because, in solution, it is split into Hg2+ and Cl- ions surrounding by water molecules.

But let's see. Sat. soln. would be 7.4 g per 100 mL. That's about a mole fraction of .005 which is just multiplied by less than 1 mmHg so this can obviously not be used to explain.

I'm thinking there must be some bubbling or splashing that you didn't notice. How long was the soln. heated?

Oh, well, see below for your answer. It is volatile, even at 100*C. Even with the water in the system this should still allow for the vapors to visibly condense over time.

[Edited on 11-12-2010 by MagicJigPipe]

Lambda-Eyde - 11-11-2010 at 18:21

Quote: Originally posted by MagicJigPipe  

If HgCl2 is subject to sublimation it will do so at any temperature (depending on pressure) but just to a lesser extent, right? In the same way that water evaporates even at 0*C just relatively slowly. In this sense won't many if not most ionic salts "sublimate" just at a rate that renders the sublimation insignificant? I keep asking questions like this about things that happen but are not considered significant and are therefore, for practical purposes, ignored but no one will answer me. Seriously, what's going on?


I'm not following the discussion at all, but I just want to add that HgCl<sub>2</sub> is a molecular compound, not ionic. Which may help in explaining its odd behavior with regards to sublimation.

entropy51 - 11-11-2010 at 18:21

The Merch Index says HgCl2 volatilizes at 300 C. Also that it is "slightly volatile at ordinary temperatures, appreciably so at 100 C".

Based on this, I would guess that if you heated it a bit, some might volatilize onto a cooled surface in close proximity. I have to agree with DJF90.

Sounds like an interesting quantitative experiment for someone with some HgCl2 and a good balance. Blogfast25??

Sedit - 11-11-2010 at 19:09

I got some HgCl2 that I wish to not touch with a ten foot pole and some HgSO4 as well. Either way the best scale I have goes to .1 grams which I dont feel is enough to perform an experiment of this nature.

Not to mention im a messy basterd.:D

Magpie - 11-11-2010 at 19:58

In one of my books there is a small graph of the vapor/liquid equilibrium composition vs temperature for the system NaOH-H2O. Along with it the text says "Note that a discernible concentration of sodium hydroxide in the vapor phase is not obtained until the liquid phase reaches 95 per cent NaOH and boils at some 700F (371C)." This is rather what I would expect for an inorganic ionic dissolved substance that has a mp of 318C and a bp of 1390C.

But what is meant by "discernible." This may depend on the substance and the sensitivity of the method used for measurement, ie, a titration, a conductivity meter, damp pH paper, the human nose, etc.

Many solids, eg, some metals, can be smelled at room temperature. All solids have some vapor pressure.

blogfast25 - 12-11-2010 at 07:28

I've often asked myself the same question when dissolving metals in acids: there too a characteristic smell can usually be observed, especially for metals that are easily oxidisable by H3O+ alone and lots of hydrogen evolves. A muff, metalicky odour, not very pleasant. With Al and Fe it's particularly similar, in my perception...

[Edited on 12-11-2010 by blogfast25]

DJF90 - 12-11-2010 at 08:08

I've noticed this with magnesium and zinc, although I began to wonder if it was metal particulates I could smell, or whether it was nascent hydrogen (pretty unlikely me thinks?)...

blogfast25 - 12-11-2010 at 08:33

Certainly hydrogen (H2) is odourless and nascent hyfrogen (H. radicals) is very shortlived. Metal particles? Unless extremely fine they can't tickle your olfactory system, you need some vapour pressure for that, as far as I know...

I had been thinking perhaps silane or other volatile metallo hydrides (from trace alloy elements e.g.) but many are highly inflammable/unstable in the presence of oxygen...

Tin metal with hydrochloric acid also smells peculiarly...

ScienceSquirrel - 12-11-2010 at 09:22

I have made iron salts by dissolving steel wool in sulphuric acid.
The gas that comes off is mainly hydrogen but it has a little hydrogen sulphide and maybe something more.
I doubt that a trace of arsine would inflame and it would give the gas a slightly garlicky odour.

blogfast25 - 12-11-2010 at 10:56

Maybe trace elements hydrides are causing it, it sounds plausible.

I've just taken receipt of 50 g of fairly pure tin (99.9 %), so I'll see how that whiffs with strong HCl... Where I wrote 'tin' previously I meant pewter of course (about 95 % Sn)...

[Edited on 12-11-2010 by blogfast25]

12AX7 - 12-11-2010 at 12:20

Theory:

Suppose the NaOH grains are porous. This seems reasonable, as NaOH is generally produced in prill form, and the water removed during crystallization will leave pores.

When prills are added to water, the outer surface is completely wetted and begins to dissolve. However, water is not able to penetrate the grain completely, and air bubbles become trapped inside.

As the prill dissolves, air bubbles are loosened. The bubbles are very small, corresponding to the prill's pore size.

When these bubbles reach the surface of the solution, they burst, releasing very fine particles of NaOH solution. This aerosol irritates the mucosa, leading to a burning sensation when inhaled.

The action of very fine bubbles bursting seems to be different from the bursting of large bubbles, which result in correspondingly large droplets, which are not able to form an aerosol.

If the fine bubbles are not produced by this mechanism, perhaps there is a chemical explanation. This seems unlikely, as the only gas that could be produced is ammonia, which has an odor. CO2, from hard water for instance, would be bound instead. Hydrogen or oxygen would require redox conditions, which are absent.

Tim

blogfast25 - 12-11-2010 at 13:01

I've often thought that very localised overheating (from solvation energy) could contribute to fine bubble formation...

One way to test Tim's theory would be to dissolve fused NaOH...

[Edited on 12-11-2010 by blogfast25]

woelen - 12-11-2010 at 15:13

With the HgCl2 I indeed must say I am wrong. This is not an ionic compound in aqueous solution and then it is sufficiently volatile to get out of aqueous solution even below the boiling point of water. This, however, is not the case with e.g. Hg2(NO3)2 or with Pb(NO3)2 and certainly not with NaOH (that's what the topic started with).

So, I indeed think that Tim's theory may be a good description of what happens.

An easy check is to take some NaOH, dissolve this in water (which gives a choking fume). Then wait for a few hours such that the liquid is completely clear and has cooled down. Then heat this liquid again but do not allow it to boil (e.g. let it heat up to 80 C or so) and carefully waft some fume from this liquid to your nose.

Maybe I'll try this little experiment tomorrow afternoon.

12AX7 - 13-11-2010 at 11:53

Quote: Originally posted by blogfast25  
I've often thought that very localised overheating (from solvation energy) could contribute to fine bubble formation...

One way to test Tim's theory would be to dissolve fused NaOH...

[Edited on 12-11-2010 by blogfast25]


The trouble with any other kind of gas formation, as I noted, is it's either reactive, or doesn't belong. If overheating caused localized boiling, it could release dissolved gasses, mostly CO2 and O2, of which only O2 could survive in this solution. But the very fine bubbles would also redissolve easily, since we're talking transient formation here.

On the other hand, as temperature rises, gas solubility drops. So it could be a contributing factor.

Easy test: use preboiled water and fused NaOH.

Keep in mind that NaOH contains nucleation sites. Merely reheating the solution may not induce bubble formation. Fused NaOH may not nucleate as well, either.

Tim

aeacfm - 13-11-2010 at 14:46

Quote: Originally posted by 12AX7  
Theory:

Suppose the NaOH grains are porous. This seems reasonable, as NaOH is generally produced in prill form, and the water removed during crystallization will leave pores.

When prills are added to water, the outer surface is completely wetted and begins to dissolve. However, water is not able to penetrate the grain completely, and air bubbles become trapped inside.

As the prill dissolves, air bubbles are loosened. The bubbles are very small, corresponding to the prill's pore size.

When these bubbles reach the surface of the solution, they burst, releasing very fine particles of NaOH solution. This aerosol irritates the mucosa, leading to a burning sensation when inhaled.

The action of very fine bubbles bursting seems to be different from the bursting of large bubbles, which result in correspondingly large droplets, which are not able to form an aerosol.

If the fine bubbles are not produced by this mechanism, perhaps there is a chemical explanation. This seems unlikely, as the only gas that could be produced is ammonia, which has an odor. CO2, from hard water for instance, would be bound instead. Hydrogen or oxygen would require redox conditions, which are absent.

Tim


good explanation
but you said i the chemical explanation that" the only gases could be produced is amonia" , ans i am asking what amonia here in NaoH solution or you mean other case????
other thing together with these fumes the temperature pf the solution becomes very high (with out heating or boiling i am working at room temperature ) may due to the heat of fromation or hydration but no external heat applied to the system
any way good shot sir



[Edited on 13-11-2010 by aeacfm]