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woelen
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Solubility of elements without chemical reaction
I am working on a periodic table display of the elements and for this I prepare samples of the elements in different forms. They, however, must remain
present as elements, not as chemical compounds. For this purpose I also prepare some solutions of elements, which display the element's properties
very nicely (e.g. a nearly saturated solution of Cl2 in CCl4, which very nicely demonstrates the color of Cl2 at high concentration and which remains
stable indefinitely).
Now I wonder, for which elements I can make such display-solutions, which contain the element as free element and which are stable over long times. I
identified the following:
+ Cl2 in CCl4
+ Br2 in CCl4
+ I2 in CCl4
+ I2 in CH3CH2OCH2CH3
- S in CS2
- Se in CS2 (??)
- P4 in CS2
+ C60 in CS2
The ones, listed with a "+" I actually made and ampouled.
This is all which I can imagine. Any other ideas? Metals, alloyed with other metals, do not count. I suspect this only works for non-metal elements.
If you know of any other element which can be dissolved in a suitable solvent, then I would like to know.
I don't think I'll make P4 in CS2. This is just a colorless liquid.
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DraconicAcid
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Oxygen is very soluble in some freons. I wonder if ozone would similarly dissolve (and might be more blue than colourless).
I don't expect you're going to dissolve metals in anything but mercury, unless a suspension of colloidal silver or gold would count.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Magpie
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I don't suppose Na dissolved in liquid ammonia would qualify?
The single most important condition for a successful synthesis is good mixing - Nicodem
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Eddygp
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Na in liquid ammonia technically loses an
electron so is not the "pure" element.
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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blogfast25
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And in those 'Fluorinert' things: https://en.wikipedia.org/wiki/Fluorinert
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woelen
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No, this does not count for me, the color of
these solutions is not due to the element dissolved, but due to the solvated electrons. Some time ago I made a deep blue solution like this with
lithium metal in ethylenediamine.
--------------------------------------------------------------------------------------
I have done some research though and found that selenium and tellurium can be dissolved in a mix of ethanethiol and ethylene diamine:
http://pubs.rsc.org/en/Content/ArticleLanding/2014/SC/C4SC00...
On evaporation, the element appears again in crystalline form.
Ethylenediamine alone does not dissolve selenium, I tried that.
I, however, do not have ethanethiol, and I wonder whether I do want this compound in my lab (boiling point just 35 C while its odour threshold is
appr. 0.5 ppb). I can buy it (appr. EUR 60 per 500 ml).
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DraconicAcid
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If you don't want ethanethiol in the lab, I suspect other thiols will have similar solvent properties.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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careysub
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C70 fullerene solutions are a different hue from C60. Showing that difference would be interesting.
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careysub
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I would say that you do not want ethanethiol in your lab. I was once affected by an ethanethiol spill evacuation in Long Beach, CA around 1999. ~100
grams evacuated a square mile area, and the place smelled awful for week after they let us come back.
Also the odor of ethanthiol from a wood pulp plant at Port St. Joe, FL can be detected on occasion at Ft. Walton Beach 100 miles away.
To significantly reduce the thiol odor problem you need a higher thiol, octanethiol perhaps (odor is described as "mild").
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woelen
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I know
of the differences. Apparently, solutions of C70 in CS2 are bright red, while solutions of C60 in CS2 are purple with a bluish hue. I have 200 mg of
C60, purchased that for EUR 13 or so. C70 is much more expensive though. I found a seller, who sells it at EUR 40 or so per 50 mg, which is too
expensive to my taste.
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deltaH
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About the solutions of metals. I don't think having the metal in solution in an ionic form betrays much if you call this a 'solution periodic table'
and display them in a standardised way. For example, metal oxides can be dissolved in anhydrous choline chloride/urea to show their colour and this
would make a great table if you use the most stable metal oxide for each metallic element:
See how nice this works from this picture from p52 of the presentation:
https://wet.kuleuven.be/english/summerschools/ionicliquids/l...
I suspect this would make something very special when done with all the metals 
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woelen
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This indeed looks very nice, but to my opinion it is quite arbitrary how this looks like. If I use aqueous solutions of ionic compounds of metals,
then the set of colors becomes quite different. Then there is the issue of different oxidation states. There can be coordinating agents in the
solution, which again change the color.
If I do such a thing, then I will do this for the transition metals in all their common oxidation states, dissolved in water, assuring that no ligands
other than water are available in the solutions, but I am still pondering on whether I will do that or not. The number of samples will be large, and
still there is the feeling of randomness (e.g. for chromium, should I only display sky blue Cr(2+) and grey/violet Cr(3+) solutions, or should I add
yellow CrO4(2-) and orange Cr2O7(2-) as well?). Also, the space, available in my display, is limited. Only appr. 7 cm per element and a height of
appr. 10 cm. Having many element boxes crammed full with vials is not attractive.
[Edited on 11-2-16 by woelen]
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j_sum1
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Interesting deltaH. But it is definitely a different kind of table. To me, this has more appeal than compound collecting or acid collecting that
have been recently suggested. I have long been fascinated by transition metals and their assorted colours. A display showing all stable oxidation
states would be really cool. I am reminded of strontiumred's work. Not complete but a really good start with some nice preparations to follow. (Not dissimilar to some sections of your pages
either, woelen. I am thinking of where you show different oxidation states of n and V and others.)
I intend to play with some Ti at some stage soon. Beginning with dissolving in HCl and then producing the red peroxide complex.
[edit]
I love SR's gorgeous photo of Mn(V) that he uses as a slide at the top of his pages. That's on the to do list too.
Same photo in the thread below.
http://www.sciencemadness.org/talk/viewthread.php?tid=14644&...
[Edited on 11-2-2016 by j_sum1]
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MeshPL
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I'm surprised nobody mentioned amalgams! Plus there are some other alloys liquid in room temperature. Most alloys are not compounds. Technicaly brass,
bronze, cast iron (carbon solution , who needs fullerenes?), all the jewlery and
some on, are solutions.
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DraconicAcid
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Quote: Originally posted by MeshPL  | I'm surprised nobody mentioned amalgams! Plus there are some other alloys liquid in room temperature. Most alloys are not compounds. Technicaly brass,
bronze, cast iron (carbon solution, who needs fullerenes?), all the jewlery and
some on, are solutions.
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That's because alloys were ruled out in the original post.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Texium
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Regarding the choline chloride/urea eutectics... unfortunately the picture deltaH used upthread was actually depicting solutions of metal oxides in
choline chloride/malonic acid. The power point presentation that it is from mistakenly labeled it as choline chloride/urea. I have made the choline
chloride/urea eutectic before, and it does not dissolve a significant amount of any metal oxides. Claims about it in that presentation are
unfortunately false.
See this other presentation for a comparison of the solubilities of several metal oxides in choline chloride/urea versus choline chloride/malonic
acid: http://www.erc.arizona.edu/seminar/Current-2011/DineshThanu_...
[Edited on 2-12-2016 by zts16]
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MeshPL
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Ok, sorry, alloys were ruled out , but only as metal-metal mixtures.
| Quote: |
Metals, alloyed with other metals, do not count.
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But still, you can alloy many non-metals with metals, although they will kinda form compounds. Not sure if in all cases though. Like do boron, silicon
or arsenic containing alloys contain compounds of respective element?
Also, do gaseous non-metals liquified under pressured and mixed together count? Like I'm pretty much sure liquid nitrogen, oxygen, helium, neonium,
argonium, kryptonium and xenonium will form liquid mixtures under certain conditions.
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careysub
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| Quote: Originally posted by woelen | | I know
of the differences. Apparently, solutions of C70 in CS2 are bright red, while solutions of C60 in CS2 are purple with a bluish hue. I have 200 mg of
C60, purchased that for EUR 13 or so. C70 is much more expensive though. I found a seller, who sells it at EUR 40 or so per 50 mg, which is too
expensive to my taste. |
Onyxmet.com has it a bit cheaper: EUR 30 for 100 mg.
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deltaH
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| Quote: Originally posted by zts16 | Regarding the choline chloride/urea eutectics... unfortunately the picture deltaH used upthread was actually depicting solutions of metal oxides in
choline chloride/malonic acid. The power point presentation that it is from mistakenly labeled it as choline chloride/urea. I have made the choline
chloride/urea eutectic before, and it does not dissolve a significant amount of any metal oxides. Claims about it in that presentation are
unfortunately false.
See this other presentation for a comparison of the solubilities of several metal oxides in choline chloride/urea versus choline chloride/malonic
acid: http://www.erc.arizona.edu/seminar/Current-2011/DineshThanu_...
[Edited on 2-12-2016 by zts16] |
Excellent moderation zts16! I didn't understand how metal oxides dissolved in ChCl/urea, but an acid makes perfect sense! So
these are actually nearly anhydrous solutions of metal malonates. Well that's not so interesting, metal oxides dissolve much better in conc. HCl and
it's OTC
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Texium
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Yeah, it's very disappointing. 
Although, choline chloride is a pain to work with anyway. It's so hygroscopic that even the reagent grade stuff from Fischer that my school purchased
was quite wet before the jar was even opened! I'm kinda glad I've moved onto other projects.
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woelen
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If you want an interesting set of metal ions in solution, then I would not use HCl, but preferrably HClO4, or dilute HNO3. These do not coordinate to
the metal ions and then you get solutions which really represent the color of aquated ions, and not the color of ions, coordinated to some anion.
E.g. an iron(III) solution in dilute HClO4 or dilute HNO3 is nearly colorless (actually very faint violet), while in hydrochloric acid such a solution
is bright yellow with a brownish hue (due to formation of the FeCl4(-) complex). Copper(II) tends to form green solutions when chloride is present at
moderate concentrations, while it is pure sky-blue in dilute HClO3 or dilute HClO4.
Such a display would indeed be nice. The first row of transition metals has the following simple aqueous ions, of which Ti(3+), Cr(2+) and FeO4(2-)
are borderline stable and I doubt whether solutions of these can be kept around in ampoules for years.
Sc(3+)
Ti(3+) / TiO(2+)
V(2+) / V(3+) / VO(2+) / VO2(+)
Cr(2+) / Cr(3+) / Cr2O7(2-) / CrO4(2-)
Mn(2+) / Mn(3+) / MnO4(2-) / MnO4(-)
Fe(2+) / Fe(3+) / FeO4(2-)
Co(2+)
Ni(2+)
Cu(2+)
Zn(2+)
You could also argue about whether you want stuff like Cr2O7(2-) or MnO4(-) in the display.
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blogfast25
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| Quote: Originally posted by woelen |
E.g. an iron(III) solution in dilute HClO4 or dilute HNO3 is nearly colorless (actually very faint violet), while in hydrochloric acid such a solution
is bright yellow with a brownish hue (due to formation of the FeCl4(-) complex). |
Interesting. I always believed the 'yellow with brownish' hue was due to [Fe(OH)(H2O)5]<sup>2+</sup> but I'm
probably conflating.
Thinking about it, dilute solutions of ferric alum don't appear yellow, only slightly brownish, depending on concentration, of course. And they're a
bit thermochromic: darken on heating, presumably the equilibrium in pushed to the hydroxy complex at higher T.
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woelen
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If you dissolve an iron salt like Fe(ClO4)3 or Fe(NO3)3, then you get brown solutions. This brown color indeed is due to [Fe(OH)(H2O)5](2+) and
possibly further deprotonated ions. When chloride ion is added to the game, then the color shifts more to yellow due to ligand replacement by chlorine
ions. At low pH, you get FeCl4(-).
Pure [Fe(H2O)6](3+) is nearly colorless. Try dissolving some Fe(NO3)3·9H2O or Fe(NH4)(SO4)2·12H2O in dilute HNO3 or dilute HClO4. You get a nearly
colorless solution, not brown, nor yellow. Next, add a little HCl or NaCl and the liquid at once turns bright yellow.
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Bezaleel
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Sulphur is reported to be soluble in toluene, and can be recrystallised from it. Doesn't add a new element to the list, though.
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chornedsnorkack
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| Quote: Originally posted by woelen |
Such a display would indeed be nice. The first row of transition metals has the following simple aqueous ions, of which Ti(3+), Cr(2+) and FeO4(2-)
are borderline stable and I doubt whether solutions of these can be kept around in ampoules for years.
Sc(3+)
Ti(3+) / TiO(2+)
V(2+) / V(3+) / VO(2+) / VO2(+)
Cr(2+) / Cr(3+) / Cr2O7(2-) / CrO4(2-)
Mn(2+) / Mn(3+) / MnO4(2-) / MnO4(-)
Fe(2+) / Fe(3+) / FeO4(2-)
Co(2+)
Ni(2+)
Cu(2+)
Zn(2+)
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I´m also unsure about V2+ and Mn3+.
While nitrate is a poorly complexing anion, it also tends to decay to NO2, which is yellow. Dilute perchlorate is more stable, but
Cl2 and ClO2 are also yellow.
How complexing is sulphate? E. g, can iron(III) be observed to turn from brown although still clear in a partially hydrolyzed solution to clear as
sulphuric acid excess is added?
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