Sciencemadness Discussion Board - Calcium Sulphate Destroys Glassware

Calcium Sulphate Destroys Glassware

Chemgineer - 27-9-2021 at 17:04

Recently i've been doing a dry distillation of Calcium Nitrate with Sodium Bisulfate. It is a really convenient procedure to produce Ntric Acid (without any foaming) and also the resulting Calcium Sulfate does dissolve and come out of the flask at the end.

The problem I have is that the Calcium Sulphate seems to be adhering to my flasks and literally pulling borosilicate glass off the surface, then the next time I use that flask it ends up broken or cracked.

I am using Indian Supertek 500ml round bottom flasks in a heating mantle, am I suffering from poor glassware or just punishing it too much?

[Edited on 28-9-2021 by Chemgineer]

MidLifeChemist - 27-9-2021 at 18:56

I have heard of others being similarly punished by their calcium sulfate deposits. You are not alone. Still, if nitric acid is expensive and flasks are cheap where you live, then it might be worth it

Chemgineer - 28-9-2021 at 04:45

I might just add some water next time and accept a lower concentration.

macckone - 28-9-2021 at 12:53

Anecdotal:
Try adding some 90+% sulfuric acid, it seems to protect the glass.
I originally tried it to help dissolve the mix and reduce the amount of water carried over from non-anhydrous salts.
It seems to reduce the water and improve the reaction but I haven't done a quantitative study.

Chemgineer - 28-9-2021 at 13:26

Something else I considered was adding a percentage of Potassium Nitrate, as that reaction with Sodium Bisulfate causes foaming, the foam might help.

Tsjerk - 28-9-2021 at 13:36

I think you heated the flask too much. I never distilled nitric acid from this mix, but when I did I never tried to pull the last bit.

Or it might have been the real anhydrous non-liquid combination between calcium sulfate and glass. I went with a percent of water and a second distillation from sulfuric acid anyway to get rid of the NO2. Although I don't know how a bit of water would have helped with CaSO4 as a drying agent in there

Chemgineer - 28-9-2021 at 13:54

I've done this with Potassium Nitrate and Sodium Bisulfate before several times using the same make of flask and no damage, I probably have kept the reaction going for longer with the Calcium Nitrate.. maybe it's because the calcium sulfate ends up drier than the mix of Sodium Sulfate and Potassium Sulfate and there is less heat sink for the glass which gets hotter. But to me it looks like the Calcium Sulfate has adhered to the glass surface and then retracted on cooling putting a layer with it.

[Edited on 28-9-2021 by Chemgineer]

Chemgineer - 28-9-2021 at 14:09

As in the photo there is a general dull area with an excavation of glass top right and then two cracks around the outside of the affected area.

Chemetix - 28-9-2021 at 15:10

What you are experiencing is ion exchange in the glass. Strong concentrated solutions that have different ionic composition to glass will cause some surface ions to migrate.

Osmotic potential is the force behind the exchange. Sodium ions will want to move into the aqueous solution and calcium will try to move into the glass. Calcium makes the glass brittle. It becomes a glass with different coefficient of thermal expansion to the rest of the flask and so heating and cooling will cause stress at the boundaries where ion exchange has most strongly occurred.

Chemgineer - 28-9-2021 at 15:30

Hey thanks, what you say makes allot of sense and I expect you are correct. I suppose the good thing is the heating is applied gradually by the mantle and failure only seems to only occur on cooling when the calcium sulfate is a solid mass and can't spill from the flask.

Amos - 28-9-2021 at 16:10

Does calcium sulfate not readily precipitate when you combine the two reagents? If it does, there's no reason to leave it in the flask while you distill; vacuum filter it first and distill the filtrate. If you lack vacuum filtration capabilities, I would add more water as you mentioned and still manually gravity filter. You won't distill the last bits of nitric acid out of the mixture until the dissolved salt content has nearly reached saturation, and this will be the most concentrated distillate of the entire process.

Texium - 29-9-2021 at 05:39

He said he’s doing it as a dry distillation, which would seem to be a horribly inefficient and potentially dangerous method, but oh well. So currently he’s not adding any water at all, just heating the solid mixture in a flask.

@Chemgineer: What sort of yields do you get with this dry distillation? If “foaming” is really such an insurmountable issue with the aqueous procedure, I would suggest investing in a larger boiling flask… Also, could you please stop Capitalizing Your Chemical Names? They are not proper nouns, and writing them that way just looks pretentious.

Metallophile - 29-9-2021 at 08:02

Quote: Originally posted by Texium

They are not proper nouns,

I have made this mistake before. I think it's actually kinda strange that elements are not proper nouns, even the ones named after people! But that's the rule, so I try to follow it.

Chemgineer - 29-9-2021 at 08:42

Quote: Originally posted by Texium

He said he’s doing it as a dry distillation, which would seem to be a horribly inefficient and potentially dangerous method, but oh well. So currently he’s not adding any water at all, just heating the solid mixture in a flask.

@Chemgineer: What sort of yields do you get with this dry distillation? If “foaming” is really such an insurmountable issue with the aqueous procedure, I would suggest investing in a larger boiling flask… Also, could you please stop Capitalizing Your Chemical Names? They are not proper nouns, and writing them that way just looks pretentious.

I use 160g sodium bisulfate and 100g of calcium nitrate and I typically get around 50ml of RFNA.

# Edited to remove my capital letters!

[Edited on 29-9-2021 by Chemgineer]

Fulmen - 29-9-2021 at 09:15

What's the temperature? With thoroughly dried chems I would guess you would need to melt the bisulphate (315°C). There are nitrate eutectics that melt at much lower temperatures, might be worth a shot.

KNO3 NaNO3 Ca(NO3)2 MP
24.............................76...........142
.................31............69...........142
26.............................74...........145
65............................ 35...........150
46.............24............30...........160
20.............................80...........162
.................19............81...........214
.................14............86...........214
40.............60...........................222

Texium - 29-9-2021 at 10:53

Quote: Originally posted by Chemgineer

I use 160g sodium bisulfate and 100g of calcium nitrate and I typically get around 50ml of RFNA.

That would be about a 98% yield based on calcium nitrate, if we assumed that what you collect is 100% nitric acid, but since RFNA contains additional dissolved NO2, that amount wouldn’t even be possible to obtain from that amount of reagents… Are you sure you’re using anhydrous calcium nitrate? If not, you likely aren’t getting true RFNA.

Chemgineer - 29-9-2021 at 11:56

Quote: Originally posted by Texium

Quote: Originally posted by Chemgineer

I use 160g sodium bisulfate and 100g of calcium nitrate and I typically get around 50ml of RFNA.

Not fully anhydrous no, this probably helps keep the temperatures lower, the product does fume though when opening a closed bottle.

Chemgineer - 29-9-2021 at 14:04

Ok i've just measured the specific gravity with a measuring cylinder and scales, the HNO3 from calcium nitrate seems to be 1.41 sg @ 18 deg C which would make it 68%..... so azeotropic.

My previous HNO3 from a similar dry process with potassium nitrate and sodium bisulfate comes in at 1.46 sg @18 deg C which would make it 81%.

rockyit98 - 29-9-2021 at 15:52

you doing it wrong. you need to add H2SO4 to concentrated Solution of Ca(NO3) slowly while mixing. CaSO4 will precipitate out. remove them and distill the dilute solution till azeotropic forms.

Amos - 30-9-2021 at 14:45

Quote: Originally posted by rockyit98

you doing it wrong. you need to add H2SO4 to concentrated Solution of Ca(NO3) slowly while mixing. CaSO4 will precipitate out. remove them and distill the dilute solution till azeotropic forms.

It should be very obvious that sulfuric acid isn't easy to come by for the OP, or they wouldn't be attempting to make nitric acid this way.

clearly_not_atara - 30-9-2021 at 16:13

If you want to do a dry distillation, you can probably protect your glassware by using a eutectic mixture of sodium and potassium bisulfate instead of just one or the other.

Eutectic (44% KHSO4 by moles; 50% by mass): 133 C [1]
NaHSO4: 315 C [2]
KHSO4: 197 C [2]

1:
https://www.researchgate.net/publication/230557353_Conductiv...
2: Wikipedia

Amos - 2-10-2021 at 10:09

Quote: Originally posted by clearly_not_atara

If you want to do a dry distillation, you can probably protect your glassware by using a eutectic mixture of sodium and potassium bisulfate instead of just one or the other.

Eutectic (44% KHSO4 by moles; 50% by mass): 133 C [1]
NaHSO4: 315 C [2]
KHSO4: 197 C [2]

If the reaction is successful though, bisulfates will still end up as solid sulfates

Fantasma4500 - 22-10-2021 at 04:19

i had some distillation mixture of H2SO4 and CaSO4 standing around
i dumped it into a plastic bottle to be situated next to my toilet, to help remove calcium deposits every now and then
about a year later i found the bottle again
and there was about a fist-sized chunk of crystals in there
i eventually found out it was CaSO4 crystals
https://gyazo.com/74e95b60263131dfad29986ace145977

now what i understand is that i hit a particular H2SO4 concentration in which CaSO4 is soluble in
and upon diluting this further, it precipitated out the CaSO4, very slowly
so there must be maybe 10-20% "sweetspot" with H2SO4/H2O concentration in which CaSO4 is soluble.

so that might work for ya.

i would say maybe 20% or even less is where the CaSO4 dissolve- its only a vague guess

Chemgineer - 12-11-2021 at 07:58

Quote: Originally posted by Chemetix

What you are experiencing is ion exchange in the glass. Strong concentrated solutions that have different ionic composition to glass will cause some surface ions to migrate.

Osmotic potential is the force behind the exchange. Sodium ions will want to move into the aqueous solution and calcium will try to move into the glass. Calcium makes the glass brittle. It becomes a glass with different coefficient of thermal expansion to the rest of the flask and so heating and cooling will cause stress at the boundaries where ion exchange has most strongly occurred.

I am wondering if the quality of the glass might have a bearing on this issue, the flasks which I have seen fail are all from the same manufacturer (supertek) and were quite cheap. I've noticed in other reactions that better quality flasks such as Quickfit or Glassco brand seem to be more none stick on the inside surface so possibly a smoother surface finish.

I'm just not quite ready to risk a better flask to failure on this reaction yet!

bluamine - 15-1-2022 at 18:27

In such a temperature degree you'll end up with some 'burnt' plaster which even doesn't get hydrated!!!
@clearly_not_atara: does this method make the glassware "passive"?
Wikipedia mentions that calcium sulfate is slightly soluble in warm hydrochloric acid if my memory is still good, it may help you @chemingeneer

chloric1 - 28-9-2024 at 18:24

Quote: Originally posted by Fantasma4500

i had some distillation mixture of H2SO4 and CaSO4 standing around
i dumped it into a plastic bottle to be situated next to my toilet, to help remove calcium deposits every now and then
about a year later i found the bottle again
and there was about a fist-sized chunk of crystals in there
i eventually found out it was CaSO4 crystals
https://gyazo.com/74e95b60263131dfad29986ace145977

now what i understand is that i hit a particular H2SO4 concentration in which CaSO4 is soluble in
and upon diluting this further, it precipitated out the CaSO4, very slowly
so there must be maybe 10-20% "sweetspot" with H2SO4/H2O concentration in which CaSO4 is soluble.

so that might work for ya.

i would say maybe 20% or even less is where the CaSO4 dissolve- its only a vague guess

.

I’ve gotten gypsum crystals many times larger than precipitated calcium sulfate from concentrated ammonium nitrate solution. Easily filtered with coffee filter! Calcium sulfate does have considerable solubility in ammonium salt solutions at room temperature. But a nice trick is that calcium sulfate solubility is diminished by higher temperatures as well as increasing ammonium salt concentration. So if you make ammonium nitrate or chloride by double displacement with calcium sulfate as the precipitated solid, know that boiling down ammonium salt filtrate with a couple gravity filtrations will reduce dissolved calcium sulfate to less than 1% concentration before you get a salt cake.

Random - 29-9-2024 at 18:54

There was this procedure, was it inside Wagner Chemical Technology book from 1800s

...

Something regarding Sulfuric acid or Nitric acid... or \\\both

...

It was a procedure something as if you put all this inside a how we could call it vessel and let it burn or what was it...

...

I cannot recall it from my head currently.

Also I have seen this reaction how CaSO4 along with SiO2 gives SO3 when heated

...

They change this reaction into another with no SO3

Look what Politics is.

...

But let me get again how I could say OnTopic.

I could write a lot regarding this topic. But if I only quickly read it, don't use GlassWare then.

You can use another \\\material

Edit:

Metal. Ceramic.

...

Edit2:

What about Silver Nitrate and HCl

...

[Edited on 30-9-2024 by Random]

RogueRose - 29-9-2024 at 19:12

I found a convenient source for "disposable" glassware is to use old metal halide (HPS/MH) light bulbs (250w - 1500w) which I've found to be made of borosilicate glass. The 1500w bulbs are 1.5-2.5L in size. I used a Dremel like tool with a diamond cutting disc to cut the glass near the bulb thread & remove the "guts" of the bulb.

These bulbs are used in industrial applications like large parking lots, warehouses or any other place that needs large areas lighted. I found tons of these (and the transformers) at various second hand building material stores (like Habitat for Humanity Re-Store & some Goodwill's). Many places are replacing these lights with LED so these stores often have lots of these lights & bulbs.

Random - 29-9-2024 at 19:38

Quote: Originally posted by RogueRose

I found a convenient source for "disposable" glassware is to use old metal halide (HPS/MH) light bulbs (250w - 1500w) which I've found to be made of borosilicate glass. The 1500w bulbs are 1.5-2.5L in size. I used a Dremel like tool with a diamond cutting disc to cut the glass near the bulb thread & remove the "guts" of the bulb.

These bulbs are used in industrial applications like large parking lots, warehouses or any other place that needs large areas lighted. I found tons of these (and the transformers) at various second hand building material stores (like Habitat for Humanity Re-Store & some Goodwill's). Many places are replacing these lights with LED so these stores often have lots of these lights & bulbs.

There are various ideas what could be used as Apparatus.

I have been doing Plumbing before I started attending College and everything concerning dealing with Apparatus became easier for me.

jackchem2001 - 1-10-2024 at 03:37

Very interesting, good to know. Could the cracking instead be due to a crust forming with a different coefficient of thermal expansion rather than direct embrittlement of the glass?

bnull - 1-10-2024 at 05:23

Quote:

Could the cracking instead be due to a crust forming with a different coefficient of thermal expansion rather than direct embrittlement of the glass?

No, it's really embrittlement. I found at least one paper describing the effects of substitution or removal of sodium from glass. If I'm not mistaken, it was N. J. Kreidl, B. F. Trumm, R. F. Scott, Electrolytical Replacement of Sodium by Ammonium in Glass (https://doi.org/10.1111/j.1151-2916.1941.tb14850.x). From the abstract:

Quote:

A review is given of the exchange of ions (1) in glasses containing water, (2) on the hydrated surfaces of ordinary glasses, and (3) in commercial glasses with and without the application of an electrical field, and it is shown that the ammonium ion offers unique features. It may be electrolytically introduced into soda glasses without impairing their structure as much as other ions.