Sciencemadness Discussion Board

Salts leeched from ground?

12AX7 - 5-5-2005 at 06:52

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

Mr. Wizard - 5-5-2005 at 07:12

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?

12AX7 - 5-5-2005 at 07:59

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

Darkblade48 - 5-5-2005 at 10:08

My guess would be a nitrate of some sort too, especially if there used to be heavy fertilizer addition to the soil

Esplosivo - 5-5-2005 at 10:17

Quote:

It's a dark reddish color, related to biological content no doubt.


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.

12AX7 - 5-5-2005 at 11:50

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

Mr. Wizard - 5-5-2005 at 12:25

"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.

12AX7 - 5-5-2005 at 13:49

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

12AX7 - 28-6-2005 at 21:35

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]

Xtal.jpg - 29kB

Nitrates

MadHatter - 29-6-2005 at 03:50

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.

12AX7 - 29-6-2005 at 13:14

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

crystals

Pyridinium - 30-6-2005 at 11:24

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.

12AX7 - 30-6-2005 at 13:34

Don't have any barium on hand, can try calcium though.

Tim

12AX7 - 20-1-2007 at 12:26

: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

guy - 20-1-2007 at 12:35

It is most probably a phosphate (H2PO4-, or HPO4 2-). Those precipitate lead acetate.

UnintentionalChaos - 20-1-2007 at 13:22

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]

indigofuzzy - 21-1-2007 at 02:41

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]

unionised - 21-1-2007 at 04:14

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.

12AX7 - 21-1-2007 at 09:00

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

unionised - 21-1-2007 at 10:49

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?

YT2095 - 21-1-2007 at 10:49

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.

unionised - 21-1-2007 at 11:01

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--

YT2095 - 21-1-2007 at 11:12

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.

unionised - 21-1-2007 at 12:02

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.

YT2095 - 22-1-2007 at 02:52

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.

unionised - 22-1-2007 at 11:07

Yes, it certainly is, but you could also get the equivalent mass at the same time by measuring hom much NaOH you needed.

ordenblitz - 30-1-2007 at 11:31

Received a sample of the unknown substance today from 12AX7.

0.007 grams of the roughly ground crystal were mixed with 0.118 grams of powdered IR grade Kbr and gently ground in an agate mortar and pestle. This sample was then loaded into a McCarthy NAB die set and compressed until it remained only slightly opaque. This die set was then loaded into the beam path of a Nicolet 360 FTIR and analyzed.

The library search called it: MgSO4*7H2O however the comparison entry was tested in a nujol mull and not in Kbr tablet. So there were some slight differences in the spectra due to the different sampling method. I then prepared another sample die, this time with commercial magnesium sulfate. The second spectrum was compared to the first and they were so identical the second spectra looked like a shadow of the first.

So here Tim has Epsom salt of at least USP grade or better. Good guessing Unionized.

unionised - 30-1-2007 at 12:07

What do you mean "guessing"? :)
BTW, a few tens of ppm of lead or copper would put it outside the USP specification and I doubt that would show up on the IR so I don't think you can sell it as USP. On the other hand, I'm sure it makes a good drying agent if you bake it for a while.

12AX7 - 30-1-2007 at 13:27

Thanks!

Now a wider question...why am I getting Epsom Salts from the basement masonry, and why did I observe that 5g dehydrated to 0.7g? (I'll check that again with a larger measure, and do some gravimetrics yet.)

Tim

ordenblitz - 31-1-2007 at 09:32

Here is a photo of both spectra on the screen of the FTIR.
The red one is Tim's sample. The blue one is MgSO4*7H20, USP grade.
They don't often come out this close.

[Edited on 31-1-2007 by ordenblitz]

DSC02728.JPG - 122kB

unionised - 31-1-2007 at 11:14

"Now a wider question...why am I getting Epsom Salts from the basement masonry, and why did I observe that 5g dehydrated to 0.7g? "
You don't, by any chance live in Epsom do you? OK, i'm kidding but it isn't rare in groundwater- that's how it got a common name.
The stuff usually spits like crazy when I try to dry it

12AX7 - 31-1-2007 at 17:12

That is a pretty nice match, LOL!

If it were a combination salt, such as including sodium or potassium, would those also show up? (I would suppose such inclusion would "distort" it quite a bit away from the spectra of U.S.P. Epsomite though.)

I mean, the walls the stuff was/is seeping through are dolomite, and I can imagine something like sulfuric acid leeching it out preferrentially (hrm, is (Ca,Mg)CO3 + SO4(2-) <--> MgSO4 + ... preferred, I mean wouldn't CaSO4 be better?), but I don't really know where the sulfate comes from. I would be prone to expect chloride instead, due to what road salt is used around here (depending on time of year of course). We aren't in a particularly acidic rain region.

Tim

unionised - 1-2-2007 at 11:22

I should have mentioned that I also get MgSO4 in the efflorescence from the (un plastered) walls in my roof space (the roof has a few leaks).
To me that srongly sugests that it's from the bricks or cement. SO2 from a coal fired kiln might be a sulphur source here.
IR is not good at detecting low level impurities in materials. In some cases it's completely blind. Those IR spectra have no contribution from the N2 and O2 that the beam passed through.

[Edited on 1-2-2007 by unionised]

12AX7 - 10-2-2007 at 19:51

On a much larger sample (of somewhat less purity), I got a much more believable water loss. It wasn't heated as hot, though.

A pie tin was weighed as 20g (give or take 1.25g on all measurements). Loaded with the salt, 657.5g, therefore there were 637.5g of salt. This was melted over direct flame (low) until the consistency was putty-like (reaching a temperature of 214°F / 101.1°C... kitchen digital thermometer ;) ). It was then cooled and crushed. Total mechanical loss is estimated at 10+/-10g. The material was baked at 250°F for five hours, then 400°F for two hours. The tin was then weighed at 380g, or 360g dry salt, for a weight loss of 277.5g (257.5 to 277.5g including mechanical loss). Taking the best fit value of 267.5g, this corresponds to 42.0%. If the dry formula is MgSO4, molecular weight 120, this suggests 166.4 atw H2O, or about 9 hydrate. If the dehydration is indeed 51.3%, then this suggests I lost 33 grams mechanically, which I doubt.

I don't know where I got 5 > 0.7g before, I might've forgotten a tare weight or something.

So, MgSO4 not within experimental error, at least yet. Next I'll have to actually analyze it chemically.

Tim

unionised - 11-2-2007 at 07:17

Could there be MgSO4 solution ocluded in imperfections in the crystals?

12AX7 - 11-2-2007 at 11:14

I don't know how much, but I did find it curious that, on breaking a large crystal apart, it glistened with moisture...

Tim

unionised - 11-2-2007 at 11:21

Given the near perfect IR match, I'd say the water is roughly the equivalent of the difference between a 7 and a 9 hydrate.

not_important - 11-2-2007 at 11:25

Try treating some finely ground material with alcohol, even 70% rubbing alcohol, filter, evaportate, see if anything was dissolved.

Sulfates and neutral phosphates generally have low solubilities in alcohol, Mg and Ca halides and nitrates are reasonably soluble in alcohol.

12AX7 - 11-2-2007 at 20:23

I tried drying some ethanol (distilled, maybe 20%, if that) with the stuff I dehydrated. It seems to have turned to a microcrystalline mush.

I also didn't get very good results using sodium sulfate in the same way...

Tim

Unch - 11-2-2007 at 21:34

Quote:
Originally posted by 12AX7
Thanks!

Now a wider question...why am I getting Epsom Salts from the basement masonry, and why did I observe that 5g dehydrated to 0.7g? (I'll check that again with a larger measure, and do some gravimetrics yet.)

Tim


Faith's mysterys.
Damp?Humidity?



[Edited on 12-2-2007 by Unch]

not_important - 11-2-2007 at 22:28

Quote:
Originally posted by 12AX7
I tried drying some ethanol (distilled, maybe 20%, if that) with the stuff I dehydrated. It seems to have turned to a microcrystalline mush.

I also didn't get very good results using sodium sulfate in the same way...

Tim


I was looking more at attempting to see if there was Mg/Ca in excess of SO4, and any Cl or NO3. Combinations of those ions should dissolve in alcohol, filtering and evaporating the alcohol will show if that had happened.

The weight loss seen is suggestive of either NO3, or of ammonium salts, in addition to MgSO4. Either could be found in water from soil, especially with roof runoff.

12AX7 - 11-2-2007 at 22:59

Yeah, but would those crystallize? This is like second or third recrystallization, and I haven't seen crystals of anything else. In fact, less than 50 grams of junk (in a dark yellow to red color, corresponding to the yellow color of the mother liquor) remained after the first crystallization. I got some blocky crystals from that, probably salt.

Tim

not_important - 11-2-2007 at 23:57

I don't know if they would, or form a double salt. I was just suggesting a simple way of checking for minor components without anything fancy. Same way a quick test for ammonium would be to mix a bit of the mystery stuff with a bit of potassium or sodium hydroxide and a drop of water.

ordenblitz - 12-2-2007 at 19:01

After receiving the samples, I ground some in preparation of the KBr tablets. It turned into a wet mush from large amounts of occluded moisture in the crystal. I had to put the sample in an oven at 105c for a short time until dry.

If there were enough impurities to effect the dehydration weight, or any other characteristic different than the pure material, the IR would most certainly see it. Modern FTIR equipment is usually only fooled by PPM quantity impurity or less.

The absence of peaks from nitrogen and oxygen is because the little effects they have on this short beam path instrument are eliminated by taking a background before the actual test. What is visible however are the small peaks for CO2 and H2O vapor. I don't bother to purge the sample chamber and the software corrections for removal of same are only so effective.

unionised - 13-2-2007 at 13:03

The FTIR clearly ignores the great bulk of the sample you ran ie the KBr, do you still stand by this assertion "Modern FTIR equipment is usually only fooled by PPM quantity impurity or less."

As for "The absence of peaks from nitrogen and oxygen is because the little effects they have on this short beam path instrument are eliminated by taking a background before the actual test."
No it's because IR spectroscopy is blind to homonuclear diatomics like Cl2, N2 and O2 because they simply don't absorb. It's because there's no change in the dipole moment when the internuclear distance changes.
Think about it, if the thing spots the 330ppm or so of CO2 in the air why not the 210000ppm of O2?
It cannot be, as you say, the short pathlength and the compenstion by the background scan because those are just as true for the CO2 as for the O2.
Some things are just not very good absorbers so they don't show up well.

12AX7 - 13-2-2007 at 14:23

Hey unionized, congrats on the 1337 posts. :P

I don't think N or Cl or O is going to make a difference, in atomic terms. This isn't a Moseley plot. N2 and O2 aren't going to do much in the IR spectrum being nonpolar, reasonably high binding energy molecules. And that's fortunate so this doesn't have to be done in a vacuum. However, N as NO3- or NH4+ or others ought to show up somewhere, if present.

What I'm more concerned about, is the IR response more a matter of the crystal structure or the ions (properties of them, that is) inside it? If the ions, then you should be able to detect the spectra of impurities or compounds (Na, K, etc.) in addition to the same. If it's the structure, then surely a compound salt would show a different spectrum, necessarily being a different structure.

Whatever it is, it's pure. As I said, this is the only thing I'm crystallizing.

Say, can you subtract the stock MgSO4 spectrum from the measured plot and evaluate the resulting difference? It oughta work that way.. Looking at the screenshot, it looks like the only differences are the large peak, which is relatively shorter for the blue curve, and a small peak on the far right end of the spectrum. The overall shape is certainly close beyond a doubt though.

Tim

[Edited on 2-13-2007 by 12AX7]

ordenblitz - 13-2-2007 at 18:57

unionised:
My posts were clearly brief explanations to avoid getting into complicated FTIR theory. You may decide that instead of nitpicking each word of my argument to your assertion that IR is not good at detecting low level impurities and yourself go into the laborious explanation on all things FTIR which I decided not too. The reason KBr is used is because it is essentially invisible to IR in the useful frequency range as are nujol, NaCl and certain crystals. So what does this have to do with the accuracy of my machine? This Nicolet is good at seeing impurities, I use it each day to do just that! If you want Tim to send you some of his sample.. I'm sure he would and then you can do some actual lab work on this issue instead of armchair analysis of mine.

12AX7:
This machine does have a subtract feature. I think its the 7th button in the upper menu from the left. It will subtract one spectra from the other and you then can search the library for the result. It's usefulness depends on having a really good library of commercial spectra which are very expensive. I think I only have some 20,000 in my libraries. I didn’t spend a great deal of time in sample preparation as I was just looking to find out what it was. If I was going to do trace work I would have done a better job of it so as to be more accurate. The differences in the spectra you observed were due to auto-gain intensity variances since each KBr tablet is somewhat more or less opaque from variability in compression.

[Edited on 14-2-2007 by ordenblitz]

not_important - 13-2-2007 at 21:11

all of the alkali metal halides are mostly transparent for the IR spec range, which is why IR optics use them. All my experience with IR has been with organics, except for the unavoidable CO2 and H2O, so I don't really know how inorganics interact. (OK, I know diamond has a notch in the near IR, the how SiO2, Al2O3, Si, and Ge look).

From what I do know, MgCl2 would be difficult to detect in MgSO4, as it likely doesn't have strong bands in the normal wavelength range. On the other hand, nitrate should show if if not masked by a much larger amount of sulfate, same for carbonate.

Am I correct, or completely offbase on this?

unionised - 14-2-2007 at 11:49

Not important, you are right. A small peak can easilly be hidden by a big one. Some materials only have small peaks so they don't show up.
IR, therefore, cannot reliably detect impurities. For example it did not notice the roughly 95% of the sample ordenblitz looked at because (as he said) KBr doesn't absorb IR.
A perfect IR match does not prove that the compound is pure. NaCl is a very common material. It's a perfectly plausible impurity in the MgSO4. Does anyone think a few percent of it it would show up in the IR spectrum?


For the record nujol absorbs IR quite well, albeit in a couple of main bands. I've been using IR to identify things for roughly 20 years so I'm not just an armchair analyst. While it sometimes shows the impurities I don't bet on it for the reasons I gave earlier.
Ordenblitz if you want to do the subtraction 12AX7 sugested earlier you will need to convert the y axis from %t to absorbance. I'd be very grateful if you would post the spectrum in absorbance units please.

A Source for this Stuff?

BeanyBoy - 22-3-2007 at 12:49

Hey Tim,

Did you ever come up with a theory about the source of this precipitate, beyond it maybe just being in the ground water?

Did it seep into the basement predominately from one corner, or side, of the house? The house dates to the 1890s you said... do you know where the Facility was located in those days?

I'm amusing myself with the notion that a former owner of the home had to make regular and frequent use of Epsom Salts despite its side-effects...

-thinking of beans, always....

12AX7 - 22-3-2007 at 19:17

Heh!

No idea. AFAIK, this place has always had plumbing -- or almost always, anyway. The west wing is a few years newer and is where the (cast iron) sewage line exits the house. I don't know that I could find evidence of the original sewer line, if there was one.

Probably, it's a result of acid rain (not very acid, but enough) leaching through the dolomite foundation over decades. The salt was obtained from crud that had fallen from the decomposing wall, which must've been exposed to rainwater for many years (bad gutters).

I suppose a more definitive test on the source would entail removing a chunk of the wall's rock and determining the gypsum and free calcium and magnesium (as carbonates) content.

Tim