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12AX7
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Salts leeched from ground?
This might've been discussed already, but I have absolutely no idea how I'd find it...
I have a few gallons of a solution I leeched (indirectly) from the ground. I say indirectly because I was shovelling up material which flaked off our
old (circa 1890s) limestone foundation, which was caked with a white crystalline (fibrous) efflorescence, due to groundwater seeping through and
drying inside the basement. I soaked all the material and decanted off the water. That was a few years ago, it's been drying slowly since.
It's a dark reddish color, related to biological content no doubt.
Anyway, a few days ago I took a half gallon sample and froze it; I let it go too long and it froze solid, with long white crystals embedded in it. So
on thawing, I was left with some of this material. It does not appear to burn, and only melts momentarily below 800°F, presumably hydration.
I'm going to do some more tests yet. Any thoughts?
Tim
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Mr. Wizard
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My first guess would be Calcium Carbonate. Try treating with Hydrochloric Acid and if it foams up it's a Carbonate. It could also be Potassium
or Calcium Nitrate. Where did the water come from that caused the efflorescence on the limestone walls?
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12AX7
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Well, calcium carbonate is insoluble, and I doubt there's enough CO2 for bicarbonate to be present (though I haven't boiled it).
Just rainwater draining through the local Wisconsin dirt a few feet eventually soaking slowly through the walls.
I'll see if it releases any gasses with acid.
Tim
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Darkblade48
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My guess would be a nitrate of some sort too, especially if there used to be heavy fertilizer addition to the soil
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Esplosivo
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It's a dark reddish color, related to biological content no doubt.
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It's a strong indication of bacterial life, most probably of the nitrifying type, especially if your old house was once close to a barn or
something similar where animals were present. The 400 deg C melting point might show the presence of low boiling point nitrates, circa 350 deg C - the
increase in temp. could be due to impurities. If you are in to experimentation mixing some of the red mixture with a diluteish
sugar/urea/ammonia/small quantity of fertilizer mixturewould cause bacterial growth (the quantity of red matter would increase) if the bacteria are
'strong' enough.
Theory guides, experiment decides.
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12AX7
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I don't know if there would be any bacteria left, seems to me it'd be long gone given the strength of the salt. Probably 10 pounds of salts
in the two gallons or so of liquid. Nonetheless, I'll have to try that 
I took a sample of the frozen-out crystals and heated it in flame to redness. It melted and dehydrated at around 300C (not much hydration, in
contrast to epsom salts which fully liquifies on heating and gypsum which is neither soluble enough nor hydrated enough) then did absolutely nothing.
If it were an alkali salt, it would've melted by 1200-1600°F (though I didn't take it quite that high, some edges were glowing yellow in
the same way lime incandesces brightly).
HCl addition to the parent solution causes a little release of bubbles, probably CO2 (what else could it be?  ). After this, an addition of sulfate ions (alum) did not precipitate anything. Calcium hypochlorite of course
bleached the color to clear (with no sign of precipitate (CaSO4 etc.), oddly).
Tim
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Mr. Wizard
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"Well, calcium carbonate is insoluble"
It isn't very soluble, but it does dissolve: the hard water at my house builds up a solid ring around the drip edge of the tap over 3-4 months,
requiring a vinegar wash to get water through the aerator.
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12AX7
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3.3 x 10^-9 solubility constant, not much, but it goes up a lot with CO2 (bicarbonate). That's how caves form, CO2 being released on reaching
the opening.
Tim
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12AX7
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Alright, the solution is concentrated. I noticed crystals growing around the water edge so I picked off a large one and hung it on a string, just
below the surface. Not bad growth for only a few days.
The colorless crystal measures 3/8 x 7/16 x 1 3/8", rectangular prismatic with minor lengthwise striations. Two opposite corners are truncated
along the length, one more than the other. The smaller, more defined face is under 1/16" wide, while the wider, more striated face is 3/16"
wide. Both corner faces appear to be 45° to the major faces. The two remaining corners are very square; nearly sharp. The ends are truncated
(though not well defined in this sample), the wider faces draw together to a wedge point on one end, while the narrower faces draw to a similar point
on the opposite end. Thus the two ends are wedge-shaped, but at 90° to each other.
Taste: bitter, ever so slightly burning (basic?). Basicity would be supported by the greasy feeling, though that may be due to it still being moist
(I removed it from solution an hour ago).
As already mentioned, it appears to be mildly hydrated. Some smaller crystals I left in a warm place have turned white and become weaker. I did not
measure weight loss (doh). I do not know if the dehydrated material is hygroscopic; so far it appears not.
I will sift the solution (for crystals stuck in the possibly bacterial sludge) and measure solubility in water and later, water of hydration.
Edit: I added 30g of crystals to 25g of H2O. After an hour of lazy stirring, it nearly all dissolved. This is room temperature, maybe 72F = 22C. So
some 120g/100ml.
FWIW, It forms white precipitate with lead acetate. Which narrows it down to only everything but nitrate...
Tim
[Edited on 6-29-2005 by 12AX7]
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MadHatter
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Nitrates
That's what it sounds like given the solubility you mentioned. KNO3 has a cool and
bitter taste to it. Of course there's impurities in it because you said it burned.
From opening of NCIS New Orleans - It goes a BOOM ! BOOM ! BOOM ! MUHAHAHAHAHAHAHA !
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12AX7
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If it were a nitrate, it could be NaNO3, given the solubility. But that doesn't explain the lead precipitate or lack of melting (besides
hydration).
Tim
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Pyridinium
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crystals
I'd be curious to see if you got a precipitate from barium nitrate or barium chloride. You might have a sulfate there.
Either that, or a phosphate.
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12AX7
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Don't have any barium on hand, can try calcium though.
Tim
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12AX7
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:Blink:
Oh yeah, I'm recrystallizing the stuff again. I'm getting nice size crystals (1/16 to 1/2" across), clear to white. The remaining solution is
starting to get darker again so the crop is now offwhite, worth recrystallizing again for absolute purity.
I took a 5g sample, heated it until it melted, then all boiling stopped and I weighed the sample at a styrofoam-like 0.7g, representing a lot of moles
of water lost. Judging by the smell I did not lose much (be it CO2, SO3, etc.) to overheating. Anhydrous sample dissolves completely in water, and
has since mostly crystallized (will weigh yield).
To recap;
Ppt with Pb, Ba (suggesting SO4, PO4, CO3, etc.), OH (have not performed gravimetry on OH)
OH is not amphoteric (strongly suggesting alkaline earth)
Salt is extremely soluble (>100g/100ml)
Contains large amount of hydration; hydrate is somewhat efflorescent
Does not appear to dissolve in conc. H2SO4
Tim
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guy
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It is most probably a phosphate (H2PO4-, or HPO4 2-). Those precipitate lead acetate.
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UnintentionalChaos
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Given everything but the melting points you listed early on, it looks very much like sodium sulfate decahydrate. I have an unnecessarly large amount
of experience working with it. Crystal structure as well as the striations looks very much the same as it easily forms massive crystals.
I don't know about the taste, but the really greasy feeling is familiar.
Precipitate with lead nitrate would be insoluble, white, lead sulfate.
"Crystals left in a warm place have turned white"
Na2SO4 autodehydrates in dry air, especially if kept warm.
"took a 5g sample, heated it until it melted, then all boiling stopped and I weighed the sample at a styrofoam-like 0.7g, representing a lot of moles
of water lost."
Is this a sample of solution? the decahydrate will fully dehydrate upon heating to less than half its original mass, but I can't possibly explain that
loss of water if it were a crystal sample.
It would not dissolve in conc. H2SO4 because it is already a sulfate. I imagine the hydrated crystals would fall apart as they get dehydrated by it
though.
I have the ultimate test for the substance, however (to see if its Na2SO4). sodium sulfate has the weirdest solubility curve you will ever see. From
0-32.4C it goes from about 5g/100g H2O all the way to almost 50g/100g H2O. After that, the solubility drops and at 100C, its roughly 43g/100g H2O. It
will be very easy to exploit the first part of the solubility curve for testing. Simply take a large sample of the solution (room temperature would
have solubility of 20g/ 100g H2O) and boil it to half volume, then chill and hold at roughly 5C overnight (works fine in a cold fridge too). If you
now have crystal slush in the beaker/jar its definetly sodium sulfate.
[Edited on 1-20-07 by UnintentionalChaos]
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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indigofuzzy
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Have you tried a flame test to see what color it emits? If it is Na2SO4 then it should turn a flame intensely yellow.
Depending on what color(s) you get, this may give you a hint at what the metal ion in the salt is.
*edited: the forum did not like unicode subscripted numbers. (aargh!)
[Edited on 1.21.2007 by indigofuzzy]
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unionised
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My best guess is Epsom salts ie hydrated MgSO4.
They solubility is about right, the melting and loss of water is right the crystal shape is about right and it would give a ppt of PbSO4.
Also I have got MgSO4 from the efflorescence in my cellar.
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12AX7
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Sodium sulfate -- I have some crystallizing (and drying, and efflorescing) as we speak. It is most definetly not sodium sulfate, as it does not
dehydrate nearly as quickly. The last crop of "ground salt" currently has some white spots on it (nothing covering a crystal), wheras the sodium
sulfate (a somewhat more recent crop) has just a few bits of hydrate left to fall apart! The hydration is about right, but I don't get the strong
flame color or melting (at orange heat) and sodium sulfate doesn't precipitate in base, nor is the hydrate nearly this soluble.
The cation seems like magnesium or some other alkaline earth, but magnesium sulfate is only half as soluble, if I got the solubility right. It's also
only half hydrated by weight. When I recrystallized some epsom salts, I got a very long crystal form, but I may not have this at enough purity. The
dehydrated and dissolved sample did form some unusually long crystals as it dried. The counter experiment would be to take some of the impurity (once
I refine the liquor enough that it's a majority component) and add it to a solution of epsom salts.
Tim
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unionised
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A solution of calcium chloride or nitrate would give fairly good confirmation of sulphate (Barium might be better but it's not as easy to get).
There's a mixed NH4 Mg sulphate; ammonium sulphate would disapear on heating so it would look like a very hydrated material but I'm sure you would
have noticed the ammonia when you added base to it (and you'ld probably have spotted ammonium sulphate fumes too).
"took a 5g sample, heated it until it melted, then all boiling stopped and I weighed the sample at a styrofoam-like 0.7g, representing a lot of moles
of water lost."
That's a remarkable result- very few things hold that much water. Are you sure there were not other losses?
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YT2095
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I`ve had a look through my CRC type handbook, and there`s not much to choose from in way of soring it the Cation part of Mg or Ca.
it might be worth makeing a strong hydroxide soln such as KOH and obtaining a ppt.
the only distinctive part that may be worthwhile is that if you make a Phosphate of it, Mg will be a bluish color.
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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unionised
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I`ve had a look through my CRC type handbook, and there`s not much to choose from in way of soring it the Cation part of Mg or Ca.
Try adding SO4--
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YT2095
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it should have been "ETC..." actually, I was sidetracked with the 2 possibilities most frequently mentioned thus far.
I`de still work from the hydroxide PPT though, and then work with the filtrate, it does at least give you a KNOWN value to work from.
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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unionised
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Come to think of it you should be able to titrate the stuff against NaOH* and get a rough value for the equivalent weight, filter, then assay the
solution for the anion. The mass of the hydroxide might give more data too.
* titration by noting the absence of further precipitation- the least popular, slowest, and least accurate titration technique.
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YT2095
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sometimes it`s immensely helpful to split a one big problem into 2 smaller ones with Known values involved.
in this case: x-OH and Na-x.
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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