Sciencemadness Discussion Board - Rock Molester's Club

Be prepared for lots of alkali fusions! (If you want to be serious about chemically analysing/characterising minerals)

https://en.wikipedia.org/wiki/Silicate_minerals

Quote:

Trust me, these bastards do not respond to either dilute or concentrated acids!

[Edited on 26-7-2015 by blogfast25]

I bet ammonium biflouride will put a hurt on them.

blogfast25 - 25-7-2015 at 18:56

Quote: Originally posted by hyfalcon

I bet ammonium biflouride will put a hurt on them.

Ammonium bifluoride would still would still have to be used in the molten form to get fast digestion. Fancy it?

Anhydrous HF is also used to digest silicates (SiF4 escapes). Fancy a bit of that, do you?

Seriously, digestion to fluorides isn't for hobbyists.

blogfast25 - 25-7-2015 at 18:59

its very well intentioned aga, but I have to agree with bf25. rocks are very difficult to dissolve. even piranha or AR take ages with heating. and Fe, Al, Si, O, Na are exceedingly pervasive as contaminants.

'Piranha' and Aqua Regia are largely irrelevant here. With nothing to oxidise, oxidisers don't work here.

deltaH - 25-7-2015 at 22:15

As a second year undergraduate a long time ago, I used to do vacation work at the analytical laboratory performing mineral assays for the mineral's processing group. The commonly used method that is absolutely not amateur friendly is sequentially boiling the sample in a series of acids including HF, perchloric and nitric, in a specially designed fume cupboard with a water scrubber on the exhaust.

Amazingly, this did not always work

Some samples high in chromites required something special. The 'special' was fusing the ore in cherry red hot sodium peroxide in a little zirconium crucible.

The latter simply required a blowtorch, a pair of tongs, a small Zr crucible (a bit pricey) and affordable Na2O2. The nice thing about the method is that it doesn't produce noxious fumes and is not ridiculously toxic like HF and works for nearly everything.

aga, please take a look at https://en.wikipedia.org/wiki/Peroxide_fusion for an introduction.

Here for a method:
http://www.perkinelmer.com/CMSResources/Images/44-130831APP_...
on page 3.

Quote:

In a zirconium crucible, precisely 0.2 g of finely ground sample was fused with 3 g of sodium peroxide (Na2O2) and 0.5 g of sodium carbonate (Na2CO3). The fused mixture was poured into a beaker containing 250 mL of a 20% acid mixture of 1:1 hydrochloric (HCl) and nitric (HNO3) acids.

Use face shield at all times and good PPE! Long-sleeved welder's/forgers thick leather gloves are also a must.

The following links to a more detailed description of the procedure, though it has the glaring mistake of talking about a ?1:1 HCl?, which is probably 1:1 HCl:HNO3 (conc. not mentioned). I'd use the quantities and concentrations mentioned in my quote above and just refer to this document for a description of the procedure's execution.

http://www.personal.psu.edu/hxg3/MCL/naperoxide.pdf

Zr crucibles come up on ebay from time to time for around $75-100 AFAIK. Important: you want 25-35ml (preferably 30ml) zirconium metal crucibles, NOT zirconium oxide crucibles!

[Edited on 26-7-2015 by deltaH]

aga - 26-7-2015 at 00:32

Thanks for all the references, especially the scary ones.

I never imagined that people had resorted to such extreme measures to get rocks to fully dissolve !

With 90% of the crust being silicates, i suppose that leaves a rather large 10% to more easily investigate before moving on to the silicates.

sample C looks like Kaolin?

It could well be.

It is now obvious that samples should be accompanied by a photo of them in their natural environment (noob error: my bad).

Colour me highly interested. This deserves to be a
sticky. So many practical applications.

Hmm. I think things go sticky after a while when they contain useful reference material !

As well as identifying rocks and minerals, it would be wonderful to know what they are composed of, and how to process them into useful chemicals.

A mini-dream is to have the ability to wander cross the face of this astounding planet with some depth of knowledge of what is underfoot.

j_sum1 - 26-7-2015 at 02:29

Hmm. I think things go sticky after a while when they contain useful reference material !

Of course you are right. and I did not mean that it should be stickied straight away. I just see potential for this thread to be a repository of a certain flavour of knowledge -- in particular, tests for ions in solution.

Of course the aluminosilicates are going to be the bane of the whole thing. Getting them to dissolve is going to be trouble. But that need not be a bad thing.
1. Acid selection can be part of the mineral identification process. (For example, dissolves in HCl with difficulty, AR easily, Sulfuric and nitric can't touch it. There fore it must be...) I would love to see that kind of information in this thread.
2. Following on from blogfast25's comment on the limitations of HF and AR on minerals... Oxidation is not the only possible process to employ. Given that silicon can be extracted from silica using a thermite reaction with Al (plus some sulfur), that opens up another possible route for analysis.
3. Part of the analysis involves what can be leached out. The fact that some undissolved material remains is neither here nor there.

Let me kick things off with some things I learned in the last 24 hours -- Credit to kecskesajt and gdflp

Testing for Al3+ in solution
Al3+ will precipitate Al(OH)3 as a gelatinous precipitate on the addition of NaOH. Further addition of NaOH will redissolve the precipitate as aluminate [Al(OH)4]-. Alizarin 0.1% solution added to the basic solution turns it purple. Addition of dilute acetic acid will cause the purple to disappear. If the aluminate ion is present then there will be a reddish precipitate remain -- a complex of the aluminate and the alizarin that does not dissolve in acetic acid.

Testing for Chromium 3+
Cr3+ in solution is a vivid green with perhaps a sense of grey-blue.
Addition of a base will cause a gel-like precipitate of Cr(OH)3.
Addition of a sufficiently strong oxidant (H2O2) will dissolve the hydroxide and oxidise the chromium to chromate (yellow). Basic conditions are required. Acidifying this will of course produce a nice orange dichromate.

Now, I feel as though I am attempting to tell people basics that they already know. ("Teaching Grandma to suck eggs" as they say back in the home country.) But, Hey, gotta start somewhere.

Now for some more useful ion tests. Or some tips on identifying minerals through acid digestion.

aga - 26-7-2015 at 07:02

I like the Cr3+ test - i have those reagents !

The fact that we may not be able to process the silicates 'properly' doesn't really detract from the idea of wandering the countryside with a small pick axe, then subjecting the collected rock samples to a variety of chemical tests.

It's sort of 'chemistry with a purpose' plus an active and healthy recreational activity (with a small pick axe).

Edit:
Photos of the places i got the samples from. A and C are from basically the same place, A being streak of hard white rock inbetween layers of C.

[Edited on 26-7-2015 by aga]

blogfast25 - 26-7-2015 at 07:08

Addition of dilute acetic acid will cause the purple to disappear. If the aluminate ion is present then there will be a reddish precipitate remain -- a complex of the aluminate and the alizarin that does not dissolve in acetic acid.

This is not entirely correct. I've done this test myself and it's quite finnicky.

The solution of Al (as Al3+ or aluminate) is carefully neutralised in the presence of Alizarin test reagent, Al(OH)3 then drops out. Al(OH)3 absorbs Alizarin, forming a pink so-called Alizarin Lake complex. The supernatant solution should become clear. Alizarin also forms a 'Lake' complex with Be(OH)2 and with Fe(OH)3, so the sample has to be iron-free for this Al identification method.

As for chromium, after oxidation to chromate/dichromate, add a Pb(+2) salt and yellow PbCrO4 precipitates.

[Edited on 26-7-2015 by blogfast25]

blogfast25 - 26-7-2015 at 07:13

I like the Cr3+ test - i have those reagents !

Unfortunately, chromium bearing streaks are rare. If you find one, you're probably trespassing on a mining company's property!

aga - 26-7-2015 at 08:40

Small pick axe and Running Shoes, say Nike for example ...

Do you mean places like these ?

[Edited on 26-7-2015 by aga]

aga - 26-7-2015 at 11:44

Did these sample dissolve? Any CO2?

Sample B was tested a few minutes ago with ~10w% HCl in a test tube with a cork/tube fitted so the gas could be led through another test tube containing a dilute calcium hydroxide solution (limewater).

The limewater went milky, so i guess that means it's CO₂

The reaction is fairly vigorous.

blogfast25 - 26-7-2015 at 12:26

Did these sample dissolve? Any CO2?

Did it dissolve more or less completely?

These so called carbonaceous rocks are about the only ones significantly attacked by mineral acids.

aga - 26-7-2015 at 12:33

I would have to say that it dissolved significantly.

There is not much left of the solid rock structure, just some dark grey mush with grains of presumably silicates in it.

Is this normal or did i just get lucky ?

Edit:

Easy way to find out : everyone go outside, grab a small rock and see if it dissolves in HCl.

Given the spread of where everyone lives, that should be a good test.

[Edited on 26-7-2015 by aga]

j_sum1 - 26-7-2015 at 14:46

I have no doubts that it is finickity . In the little reading that I have done it seems that it forms a wide range of complexes with different cations -- most in the purple to puce to crimson range of the spectrum with a few mustard yellows thrown in. It is also pH sensitive. This all leads to it not playing nice in mixtures: which is not what you want to hear if you are making rock soup.
Incidentally, one of the first hits I got while researching was your chemical supply company and its test for fluorides. Again, I don't suppose it will play nice if there are stray cations in the mix.
This video shows the Al test as well as Fe. The two precipitates do appear different in both texture and colour and so, I guess, could be differentiated in a careful test if only one was present. It also shows the purple complex with Mg.
I note that the most common scientific application seems to be the analysis of Ca in biological specimens. Good for analysing bone growth.

There is enough interesting here to elevate alizarin from something I have never heard of before to something that I should invest a couple of bucks on in my next chemical order. The first project would be to make a bunch of complexes with different metal ions and determine their colour. Then it could be used for some qualitative analysis.

[edit]
Quite right on the chromium and testing with a lead salt. A Ba salt can also be used with a similar result. I don't know how specific these tests are though. It seems that there is quite a lot out there that can make Pb and Ba precipitate and a yellowish hue would not be that unusual.

[Edited on 26-7-2015 by j_sum1]

blogfast25 - 26-7-2015 at 15:33

Here's some stuff I did with Gadolinite:

http://www.sciencemadness.org/talk/viewthread.php?tid=28071#...

This video shows the Al test as well as Fe. The two precipitates do appear different in both texture and colour and so, I guess, could be differentiated in a careful test if only one was present.

Yes, I've done the test with both Fe and Al and the Fe Lake Alizarin complex is distinguishable from the Al complex. But if both are present, Fe would mask the Al. Fe and Al are easy to separate though (Al's amphoterism).

[Edited on 26-7-2015 by blogfast25]

blogfast25 - 26-7-2015 at 15:41

Small pick axe and Running Shoes, say Nike for example ...

Do you mean places like these ?

No. Methinks I see lots of limestone and sandstone there.

[Edited on 26-7-2015 by blogfast25]

diddi - 26-7-2015 at 19:38

whose got a mass spec or AAS handy for all this?

j_sum1 - 26-7-2015 at 21:14

An AAS would be cool. Maybe Magpie could build one.

In the meantime, I am happy with wet chemistry.
And for the surprise of the day... Gadolinite contains no gadolinium.

I will happily do some work on the rock around here. Up on Australia's Great Dividing Range -- which is old volcanic. It is all that classic aussie red rock colour from the high iron content. (Soils in the area are classified as ferrosols.) It is quite ferromagnetic. I discovered this when I dropped a magnet in the back yard and it came up covered in fine dust. Further investigation revealed that magnetic separation and concentration of the iron compounds would not work. In the sample I worked with (which was devoid of organic material), all of the fine soil particles were slightly attracted to a magnet.

I had a student last year attempt to isolate the iron with a thermite reaction but he had little success. It seems that there is quite a bit of Si and Al in there that wouldn't budge. This all probably means that I am starting off with what might be one of the more difficult materials but we will see how that goes. I have some finely divided soil from the last bit of excavation done around here. I can also chip off a piece of the ancient lava flow that goes through the school. Results posted when I get them. I am however expecting problems with the digestion/leaching process.

Oscilllator - 26-7-2015 at 22:47

I will happily do some work on the rock around here. Up on Australia's Great Dividing Range -- which is old volcanic.

I live near the great dividing range (the blue mountains, to be precise) and it has always seemed to me that it was mostly sandstone. There is also a band of coal that runs under pretty much the entire blue mountain area, and there are a bunch of coal mines on the western side. I even had the opportunity once to do work experience in an underground one - one of the guys let me operate a continuous miner :cool:

[Edited on 27-7-2015 by Oscilllator]

j_sum1 - 26-7-2015 at 23:11

Ok. I am in QLD.
As I understand it the GDR is a volcanic hot spot that cooled down as the crust drifted over it. There are a string of volcanoes up north (Glasshouse mountains, Mt Warning etc) slowly becoming less dramatic as you progress south. I think it is a mistake to consider the range in terms of one kind of geology. But I believe its linearity is not accidental.

Diddi will undoubtedly be along later to correct me or fill in some more details.

blogfast25 - 27-7-2015 at 05:40

whose got a mass spec or AAS handy for all this?

XRF!

Texium - 27-7-2015 at 08:05

The rocks around where I live are primarily very white, pure limestone with little streaks of iron oxide running through it, giving it a mottled white and orange look.Occasionally there is some that looks more gray, but this seems to just be from a film of organic matter on the surface, as when broken they still have the characteristic white and orange look. I dissolved some in HCl about a year and a half ago, and the white carbonate part dissolved completely leaving behind most of the orange iron oxide. The solution was only slightly yellow, and boiling it down to dryness yielded off-white crystals of partially hydrated calcium chloride that worked pretty well for deicing my sidewalk (not needed during most of the year, but it was January).

aga - 27-7-2015 at 08:15

Surely the iron oxide would react to form FeCl₂ ?

blogfast25 - 27-7-2015 at 08:41

Surely the iron oxide would react to form FeCl₂ ?

Red iron oxide is ferric oxide and dissolves to FeCl3. But very old ferric oxide can be remarkably inert (not responsive to dilute acids).

aga - 27-7-2015 at 08:43

Really ?!?!

What is it about the Age that prevents it reacting ?

blogfast25 - 27-7-2015 at 09:04

Test it yourself. Prepare some fresh Fe(OH)3, filter and wash. Then 'age' it. As it ages and becomes drier it dissolves in acids less and less easily. That mineral stuff is millions (?) of years old.

aga - 28-7-2015 at 14:06

Heat. I bet it gets all jiggy when woken up with some heat.

Edit:

One way to find out ...

Wouyld some old rust (maybe 9 years old) be a candidate for 'old ' ?

[Edited on 28-7-2015 by aga]

blogfast25 - 28-7-2015 at 14:18

Heat. I bet it gets all jiggy when woken up with some heat.

Maybe, sometimes, also, too. I have some pottery grade Fe2O3 that doesn't dissolve in BOILING 98 % H2SO4. No dissolution whatsoever*. It all depends on the overall thermal history of the oxide vis a vis how it will react with acids/bases.

* But it makes great thermite! :cool:

[Edited on 28-7-2015 by blogfast25]

aga - 28-7-2015 at 14:24

Thermoperv.

So what is the mechanism by which the iron oxide looses interest in acids ?

Surely this has been investigated, studied and is well understood.

Are we going Quantum Iron now (QI) ?

[Edited on 28-7-2015 by aga]

Panic ! This has a horrible Groundhog Day feel to it.

[Edited on 28-7-2015 by aga]

blogfast25 - 28-7-2015 at 15:26

Freshly precipitated Fe(OH)3 is really mostly water. That structure is very reactive to dilute acids.

As the product dries already it loses water and slowly converts to Fe2O3, already a lot less reactive to acids.

Calcining further removes water and tightens crystalline structure more, rendering it even less responsive to acids.

diddi - 28-7-2015 at 17:33

re the iron: zts was looking at iron in sand recently here http://www.sciencemadness.org/talk/viewthread.php?tid=62710#...
the point here is that iron is everywhere and secondly that HCl strips it out. Also he is left with SiO2 unreacted.

blogfast25 - 28-7-2015 at 17:53

the point here is that iron is everywhere and secondly that HCl strips it out. Also he is left with SiO2 unreacted.

It did in that case (sand) but I can assure you that HCl doesn't always 'strip' iron out of minerals, at least not without some prior treatment. I suggest you try and dissolve a piece of mineral Hematite in HCl and report back to me on how well that went!

Many iron bearing minerals don't respond to HCl at all. Don't oversimplify by extrapolating from zts's particular experiment.

[Edited on 29-7-2015 by blogfast25]

Texium - 28-7-2015 at 18:32

Yep, I was actually rather surprised at how well it sucked the iron out of that sand, given how the deposits in the limestone that I dissolved barely reacted. I supoose it's due to a variety of factors including particle size, amount of time given to react, and the nature of the iron mineral. The sand was soaked over the course of a day and a half, which was a lot more time than the limestone spent in the HCl, and it was of course present as much smaller particles in the sand for the most part.

aga - 29-7-2015 at 09:02

Calcining further removes water and tightens crystalline structure more, rendering it even less responsive to acids.

Now there's an interesting thought, from an Energy perspective.

Fe₂O₃ in a crystal matrix 'tightening'.

Would it be that the matrix looses impurities, like water, and becomes more 'itself' thather than an impure version, or would it be that the bond lengths decrease/become more geometrically and electronically stable ?

The Volatile Chemist - 29-7-2015 at 09:56

"Friends don't let friends extrapolate" (Our school's stats teacher)
"But I love extrapolation..."
Stats is the voodoo of mathematics.
Anyways, I've got a sizable collection of those things earthy, and am going to southern Indiana this following week, which is where I like to go rock-hunting. I've only ever attempted to digest a few rocks, but I enjoy the chemistry related to it. It would be interesting (for me, at least) to chart up what color precipitates are made of different ratios of different ions when the ferricyanide and ferrocyanide indicators are used. I like them solely because I own about 20g of both.

aga - 29-7-2015 at 10:28

I like them cos the colours are pretty, tell you something about the composition, and always work as expected !

blogfast25 - 29-7-2015 at 11:16

Fluorite's two distinct colour variations:

https://www.google.co.uk/search?q=fluorite&rls=com.micro...

aga - 29-7-2015 at 12:54

Beautiful !

Dissolve them in acid at once.

A bit of rock molesting

aga - 29-7-2015 at 14:20

Samples A and C (see above) are basically the same, with A having a lot more silicate material, so sample A was discarded.

Added are samples D and E from the quarry (used to be a cement factory nearby) pictured above.
These are likely the same material too - they just looked different in situ.

3 portions of ~20g of each sample were put in 3 plastic 100ml pots (these are sold here for disposable party drinks very cheaply).

The sample pots were respectively treated with equal portions of approx 10% HCl Solution, ~48% H₂SO₄ solution and Distilled water.

The intention was to leave the samples to react for 24 hours. As usual, Life got in the way and they have been left for several days.

Here is what they looked like after being left to react.

The Top row is samples B,C,D,E in HCl solution.
Middle is with sulphuric acid.
Bottom row is with water which shows the rock samples more-or-less as they were found.

The liquor (or lixiviate - love that word) from each HCl dissolved sample was tested for iron using Ammonium Thiocyanate solution.

Suprisingly samples D and E showed NO iron cations !

Sample C is so heavy with iron that the test went 'blood red' as described in the textbooks.

Next the lixiviate of each H₂SO₄ sample was tested for Fe3+ in the same way.

Suspicious of the positive results in samples D and E, a sample of the acid used was also tested as a control.

The photo may not show it well, but there is a very very faint red tinge to the control sample (right hand side test tube) indicating some traces of iron.

Notwithstanding, samples D and E have more Fe3+ as evidenced by the more pronounced colour.

The Water samples were tested, and none showed any sign of iron in this test.

Next a 6 [M] solution of NaOH was prepared by dissolving 12g of NaOH in 50ml of water.

For the following, each sample was taken with a clean plastic dropper-pipette, and 3 drops of each lixiviate was taken and added to a test tube.

The pipette was washed 3 times with 1.5ml clean DIW with the wash water added to the test tube, making approximately 5ml of dilute solution in the test tube.

Drops of the 6 [M] NaOH solution were added until the solution showed as Basic with a universal pH test paper.

HCl lixiviate test results :

Sample B has a definite solid white powdery precipitate.
C has an indistinct Brown precipitate distributed throughout the liquid.
D and E (probably the same material) have a gelatinous white suspension.

The same test was performed on the Water samples, and no reaction was seen.

H₂SO₄ lixiviate with NaOH test results :

The results are similar to the HCl results, although sample B shows no reaction at all and Sample C has a very different character to the precipitate.

A surpring test was the pH of the resulting HCl solutions.

Samples D and E consumed ALL of the acid !

Not pictured are the tests with BaCl₂ solution which showed no white precipitate with the HCl or Water portions, and as a control, immediately showed a white precipitate with the H₂SO₄ solutions.

If i were a Proper Chemist, i would be able to work something out from all this.

blogfast25 - 29-7-2015 at 15:13

aga:

The slightly surprising results here are the white precipitates (after NaOH), almost certainly a metal hydroxide and likely to be Al(OH)3 (it certainly looks like hydrated alumina).

If Al(OH)3 it will dissolve in an excess of NaOH but not in strong NH3. To test this properly filter, wash and collect filter cake. Then subject to 6 M NaOH and NH3 (aq).

During the acid treatments, did you see any CO2 evolve?

[Edited on 29-7-2015 by blogfast25]

aga - 29-7-2015 at 15:38

Gas was evolved in all samples, all acids.

Most vigorous with HCl, tested +ve for CO₂.

I would go get the notes, but there's a snake in the shed again at the moment.

When they're happy and at ease they move slowly.

If Startled, they could easily smash up my glassware.

Usually they are gone in the morning.

blogfast25 - 29-7-2015 at 15:44

CO2 means most likely CaCO3 and/or MgCO3. It also means your HCl lixiviates contain Ca2+ and/or Mg2+.

[Edited on 29-7-2015 by blogfast25]

aga - 29-7-2015 at 15:47

So if i carefully neutralise the acid with NaOH and then chuck in some freshly scraped Mg, if it reacts (displacement), then it was CaCO₃ ?

Edit:

Doh !

Other way round, and i have no Ca metal ...

[Edited on 29-7-2015 by aga]

blogfast25 - 29-7-2015 at 16:47

Mg and Ca are both electropositive elements and cannot be displaced from water. So that doesn't work at all.

Hint for a Ca2+ test: CaSO4 is basically insoluble. MgSO4 is highly soluble.

[Edited on 30-7-2015 by blogfast25]

diddi - 29-7-2015 at 22:08

you have been busy aga. might have to get involved in a bit of molestation myself, so to speak

aga - 29-7-2015 at 23:11

Add Ca metal to water ? Tsk Tsk Tsk.

Note to self: do not post really dumb things after midnight.

Hint for a Ca2+ test: CaSO4 is basically insoluble. MgSO4 is highly soluble.

So if i add MgSO₄ solution and get a precipitate, then the original material was CaCO₃.

Thanks !

Edit:

Found this useful list of some rock compositions :
http://www.physicalgeography.net/fundamentals/10d.html

[Edited on 30-7-2015 by aga]

diddi - 30-7-2015 at 01:39

heres some chemical compositions of minerals I have researched so far.

Attachment: Minerals.xlsx (116kB)
This file has been downloaded 653 times

blogfast25 - 30-7-2015 at 05:52

Hint for a Ca2+ test: CaSO4 is basically insoluble. MgSO4 is highly soluble.

So if i add MgSO₄ solution and get a precipitate, then the original material was CaCO₃.

Correction: "So if i add MgSO₄ solution and get a precipitate, then the original material contained CaCO₃". Your stuff may not be a pure or simple mineral composition.

If you're rock is mainly limestone your sample should dissolve completely in HCl (provided there's enough of it, of course).

[Edited on 30-7-2015 by blogfast25]

aga - 30-7-2015 at 08:30

A. Yes. Good point.

The rocks are a mix mish-mash of all sorts of stuff, a bit like the inside of my head.

Edit:

Snake's gone, so i checked the notes.

Turns out that i only tested the HCl digested samples for CO2 gas, not 'all acids'.

[Edited on 30-7-2015 by aga]

blogfast25 - 30-7-2015 at 11:19

I would take a small amount (a few g) of B and of C. Then treat those samples with HCl and observe (dissolves? Gas?)

Then filter and carefully neutralise until precipitates appear.

Then filter off and wash precipitates and test these with NH3 and NaOH.

[Edited on 30-7-2015 by blogfast25]

Texium - 30-7-2015 at 19:30

Here's a pic of a chunk of limestone from my back yard. It's nice and clean because that face is from the inside of a large rock that I broke up with a hammer. It's very glittery when you move it about in the light, which I think is due to little quartz crystals embedded in it. The brown spot on the right appears to be a larger silicate inclusion. You can also see some orange bands of iron oxide running through the upper part of the rock.

It looks more white in person than it does in the picture.

diddi - 30-7-2015 at 19:48

quartz or mica? the facets look quite flat on the sparkly bits.

aga - 2-8-2015 at 12:48

I would take a small amount (a few g) of B and of C. Then treat those samples with HCl and observe (dissolves? Gas?)

Then filter and carefully neutralise until precipitates appear.

Then filter off and wash precipitates and test these with NH3 and NaOH.

If you would, then it is likely worth copying, so i did.

All prepared as suggested (minus the careful bit) and are drying overnight.

Samples D and E were taken from the same spot, and are behaving the same, so likely the same material.

Eliminating E leaves samples B, C and D

B happily dissolves like crazy in HCl.
C just makes the water brown - no bubbles evolved at all
D Also goes like crazy in HCl, evolving CO₂ gas.

(if you smoke, just get the fag going and stick it in the beaker : the redness on the end immediately disappears)

D in 96w% sulphuric acid goes slowly, unless you heat it, in which case it evolves gas like crazy.

Each dissolved sample was filtered, and the resulting acidic liquor neutralised with a 6 [M] NaOH solution.

In all cases the point of neutrality was missed, despite using a 3ml dropper pipette, NaOH added drop by drop, and the resulting liquors all tested to be > pH 8.

Sample B in HCl gave a grainy white precipitate which rapidly settled to the bottom of the beaker.

Sample C in HCl gave a Brown flocculous precipitate.
Watching it form, it's like a series of tiny brown avalanches running down a brown mountain.

Sample D in HCl gave a creamy white precipitate.

Sample D in sulphuric acid was odd.

A White flocculating precipitate formed during heating with H₂SO₄ (before the acid was neutralised).

This was filtered out (is drying) and kept for later testing.

All of the HCl stuff in the filter papers was washed once with distilled water.

The liquid leaving the filter (filtrant ?) was Clear like water in each case.

Tomorrow some dissolving in NH₃ and NaOH.

All and Any pointers on best (most useful) procedure most welcome.

[Edited on 2-8-2015 by aga]

blogfast25 - 2-8-2015 at 14:50

My money is on B and D being mainly limestone. C 'sounds' the most interesting of the three.

The Volatile Chemist - 2-8-2015 at 16:08

Indeed, do you have pictures of the original rocks un-digested?
Nice spreadsheet diddi!
I'm going a'rock hunting this week, and will show some pics when I'm back from vacation.

diddi - 3-8-2015 at 02:24

thx.

aga - 3-8-2015 at 05:13

The filtered material was dry today, apart from D which is more like a toothpaste consistency.

A small quantity of each was added to 2 test tubes, placed in a rack side by so that tubes 1 & 2 held sample B, 3 & 4 sample C, 5 & 6 sample D. Tube #1 is on the left.

~0.5ml of water added to all tubes. Unsurprisingly none dissolved in water.

6 [M] NaOH solution was added to tubes 1, 3 and 5 in 1ml portions with vigorous shaking between additions.

33w% NH₃ solution was added to tubes 2, 4 and 6 in the same manner.

Here are the results :-

It appears that C does not dissolve significantly in either.

B and D appear to slightly dissolve in both NaOH and NH₃, so would it be said that they Do dissolve or Do Not dissolve ?

blogfast25 - 3-8-2015 at 05:21

B and D appear to slightly dissolve in both NaOH and NH₃, so would it be said that they Do dissolve or Do Not dissolve ?

Considering some of it clearly hasn't dissolved, I would say these precipitates are insoluble (in NaOH and NH3(aq)). That points to Ca(OH)2 or Mg(OH)2.

If you redissolve them in a small amount of HCl, then add some H2SO4, in the case of Ca you should get poorly soluble CaSO4.

[Edited on 3-8-2015 by blogfast25]

aga - 3-8-2015 at 06:23

CaSO₄ it is then !

blogfast25 - 3-8-2015 at 06:41

So what you've got with B and D is (probably) mostly limestone, which is common as muck.

C is therefore your most interesting rock so far. Can you put up a photo of C?

[Edited on 3-8-2015 by blogfast25]

aga - 3-8-2015 at 08:43

B, C and D in the nude :

Close-up of C (which is basically the earth around here) :

[Edited on 3-8-2015 by aga]

aga - 3-8-2015 at 11:52

The C sample after dissolution in HCl seems to have some Sulphur deposited on it, as there is a dusting of Sulphur Yellow on it, collected in nooks and crannies.

Might zap it with Cl₂ and see if it dissolves away.

Is there a simpler (less toxic) test for sulphur ?

[Edited on 3-8-2015 by aga]

blogfast25 - 3-8-2015 at 12:28

C also yielded bubbles on dissolution in HCl, correct?

Try heating a sample in a Bunsen or such like: the smell of SO2 is quite distinctive (think yellow matches!) Taking chlorine to it seems overkill. I think your powder is a bit of gangue material.

aga - 3-8-2015 at 12:37

Actually (now that i have the notes on hand) sample C did Not evolve bubbles when HCl was added.

There were a very few (30 or so) bubbles when the HCl was added, which could easily have been air escaping.

My earlier statement was wrong.

Distrusting even my own scribblings i tried it again just now with NaOH, HCl, H₂SO₄ and HNO₃.

Sample C give no bubbles, or even any visible reaction, apart from the HCl solution going pale green.

The Yellow stuff may well be gangue/SOS as there was no SO₂ smell when the rock was burnt.

[Edited on 3-8-2015 by aga]

[Edited on 3-8-2015 by aga]

blogfast25 - 3-8-2015 at 17:12

Well, at least C appears more interesting than ordinary carbonaceous crap!

Texium - 3-8-2015 at 17:18

Well, at least C appears more interesting than ordinary carbonaceous crap!

I wish that I had more than just ordinary carbonaceous crap here.

Just west of where I live everything is made of granite, the pinkish kind that they used for building the Texas State Capitol. That stuff would be even less fun to play with though, being as it's primarily silicates and alumina and will likely yield to nothing. At least limestone can be dissolved.

[Edited on 8-4-2015 by zts16]

blogfast25 - 3-8-2015 at 18:48

Quote: Originally posted by zts16

At least limestone can be dissolved.

Sure. And you get CaCl2. Exciting! (NOT!)

Not meaning to rain on anyone's parade but so far I haven't seen much 'rock molesting' here...

[Edited on 4-8-2015 by blogfast25]

diddi - 3-8-2015 at 19:04

Well I have started a very small scale examination of Lazurite (Na,Ca)8[(S,Cl,SO4,OH)2|(Al6Si6O24)]. it is the blue mineral in Lapis Lazuli. It is a cyclosilicate with S3- ions responsible for the distinctive blue colour.

Small whole pieces have been added to NHO3, HCl, H2SO4 and NaOH with the following observations after 5 days.
All 3 acids are capable of decolourising the material, with nitric acid acting the fastest.
Whilst small bubbles are slowly evolved from each, Sulfuric acid was most noticeable with bubbles surrounding the stone in a couple of hours
All acid samples have resulted in a collapse of the solid structure of the stones into a gritty mush. none of the samples have been agitated.
The NaOH has no effect on lazurite at all. more to follow, and maybe even a proper write up if anything interesting happens

aga - 3-8-2015 at 23:52

Cool !

"We're not just bashing rocks together here .."
- we're dissolving them in acid !

diddi - 6-8-2015 at 18:05

Further to the lazurite observations:

After about 10 days there has been an interesting turn of events.
Of the 3 acids, only the dilute (about 5M) H2SO4 has had any further activity. The material has almost completely dissolved to a very fine grit which when swirled takes an hour or so to settle. The particulate matter is white and there is no evidence of the blue colour. what is interesting to note is that an amount of black/gold grit has appeared which is considerably denser than the white solid. under microscope I confirmed that the grit is tiny pyrite cubes, barely .1mm across!

So I am now at the stage where I will do a quantitative analysis using H2SO4 to determine the mass of solids remaining after digestion of the lazulite. any suggestions welcome

[Edited on 7-8-2015 by diddi]

aga - 7-8-2015 at 13:05

Get on with it !

People tend to prefer to criticise before they Enlighten.

The Volatile Chemist - 9-8-2015 at 14:45

I just cleaned a geode with 100mL HCl. Maybe later I'll provide before and after. It had a calcite inclusion, which made washing it tedious, but it worked for the part I could immerse without ruining the inclusion. 18-pounder, the bigger half was upon breaking.

aga - 9-8-2015 at 15:06

What ?

You're rubbing it with HCl ?!?!

Why ? (Fetishes are allowed.)

j_sum1 - 9-8-2015 at 15:29

I have a few rocks in storage that I might have a go at. Interestingly, in doing a quick wikipedia search I found a photo of the location where I got them from: https://en.wikipedia.org/wiki/File:Multi-coloured_scoria_in_...

My intention was to make a chess set from different coloured rocks that I collected from the crater of Mt Tarawera.
https://en.wikipedia.org/wiki/Mount_Tarawera The crater is a rift 16km long.

Anyway, I have samples of both red scoria and black scoria and also some white rock that I believe is ignimbrite. I rather expect the scoria samples to be unfruitful -- very dense and containing a lot of silicates. I wouldn't be surprised if the ignimbrite throws up some interesting things though. It will probably be at least a week before I get to these however. I'll do some local rock and soil at the same time.

diddi - 9-8-2015 at 15:38

@TVC - where is your geode from?

@j_sum - wow the white one almost looks like pumice, but whiter. I bet it would have beaut microstructure

j_sum1 - 9-8-2015 at 15:42

My white rocks are different from that one. They are quite different from the scoria. They are not porous, have no gas bubbles and are relatively soft -- almost chalky. Quite large sparlky crystals are visible in the right light -- maybe greater than 0.5mm. I will have to dig them out and have another look soon. When I do, I will take photos.

Oh. I do have some scoria pieces that were cut into thin strips using a tile-saw. Might be good for a look under a microscope. I can send you some if you are interested.
J.

The Volatile Chemist - 12-8-2015 at 08:10

My geode is from southern Indiana. I rubbed it in HCl to clean the rust and dirt from the interior. Looks great, except now there's ferric chloride on the outside of the rock

Texium - 12-8-2015 at 10:06

I have a whole bunch of thundereggs ranging from grape size to softball size that I dug up with my cousins in eastern Oregon a few years ago. I've considered trying to clean some of them up, but I'm not sure how exactly to go about it, and I'm nervous about possibly damaging them.

aga - 12-8-2015 at 14:55

Thundereggs ?

Do what ?

j_sum1 - 12-8-2015 at 15:34

thundereggs = geodes
https://en.wikipedia.org/wiki/Geode
https://www.google.com.au/search?q=geodes&safe=strict&am...

[Edited on 12-8-2015 by j_sum1]

Texium - 12-8-2015 at 15:38

thundereggs = geodes
https://en.wikipedia.org/wiki/Geode
https://www.google.com.au/search?q=geodes&safe=strict&am...

[Edited on 12-8-2015 by j_sum1]

Not exactly https://en.wikipedia.org/wiki/Thunderegg
Although I'm not 100% sure what the difference is. Maybe that thundereggs are usually not hollow?

j_sum1 - 12-8-2015 at 15:52

I thought they were synonyms.
It seems that that particular wiki article is splitting hairs on the definition but lacks the language precision to make matters clear.
Quote:
Thundereggs are found globally wherever conditions are right. /quote
Well, duh!

Texium - 12-8-2015 at 15:58

Yeah, it's certainly not a great article. But if you look at pictures of thundereggs compared to pictures of normal geodes, geodes seem to usually be hollow and contain well defined crystals, while thundereggs are generally solid and less "crystally," which is how I expect mine to be inside, so I guess that's the difference.

diddi - 12-8-2015 at 20:17

no difference. geode is the more geologically accepted name. thunderegg is likely a trade/invented name. they have nothing to do with thunder. they are volcanic

Texium - 12-8-2015 at 20:21