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I have an unknown white powder, which is an ionic solid, and I need to identify what it is by using qualitative analysis. I've run quite a few tests, but I am having a lot of trouble identifying the anion. The data I've collected is below.

General Properties:

  • It is a white powder, and I know that it is an ionic solid.
  • It's density is around $\pu{2g/mL}$ (this is not an exact number)
  • It appears to be a hydrate (as it loses mass when heated over a bunsen burner)
  • It is fairly soluble in water. It definitely isn't very insoluble, but I have to add a significant amount of water in order to get it to dissolve.
  • The powder has a very strong smell (although it's hard to describe).
  • A solution of the solid is colorless.

Cation Test

  • A flame test produced a deep orange / orange-red flame

  • Adding $\ce{Na2CO3}$ to a solution of the solid results in the formation of a white precipitate

  • Adding a base, such as $\ce{NaOH}$ to a solution of the solid results in the formation of a white precipitate

Based on this data, I am fairly confident that the cation is calcium, but please correct me if there are any other metals that fit the pattern described above or if additional tests are necessary.

Anion Test

This is what has really been messing with me, because I can't seem to get a conclusive test for any anion.

  • Adding $\ce{AgNO3}$, $\ce{BaCl2}$, or $\ce{Pb(NO3)2}$ to a solution of the solid did not result in the formation of any precipitate, regardless of how much I added.
  • Adding $\ce{H2SO4}$ results in the formation of a white precipitate (and no bubbles or anything). However, I'm fairly sure that this precipitate is not related to the anion, but rather producing $\ce{CaSO4}$ (from the cation).
  • Adding $\ce{HCl}$ or $\ce{HNO3}$ to a solution of the solid did not result in any bubbles (neither did adding $\ce{H2O}$ directly to the pure solid)
  • Adding Ferric Nitrate to the solution just retained the orange color of the ferric nitrate. A dark red color or precipitate was not observed.
  • Adding $\ce{HCl}$ did not result in the smell of vinegar.
  • Adding $\ce{AgNO3}$ followed immediately by $\ce{HNO3}$ also did not result in a precipitate.
  • Adding ammonium molybdate and nitric acid to the solution and heating it caused the solution to turn yellow, but did not yield a precipitate (even after centrifuging). This test was done twice with different ratios of ammonium molybdate to nitric acid, and the exact same result was observed both times.
  • The brown ring test was conducted, and came back with a negative result. It is possible that this test was done wrong, because the acidified ferrous sulphate was added drop by drop on top of the solution, rather than by rolling it down the side of the test tubes, which may have interfered with test results.

Are there any tests for ions that I have missed? I've done almost every test I know trying to identify the anion, but no dice. The results for the ammonium molybdate test are also confusing to me, because I expected either a yellow precipitate (to confirm phosphate) or no reaction, but the test seems to have gone halfway by just turning yellow. Any guidance on what I should do next would be very much appreciated.

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  • $\begingroup$ If it were the only reasonably soluble phosphate of calcium, monocalcium phosphate, then the phosphate would have dropped out with the addition of Pb(NO3)2. Regarding solubility, do you know about how much water you had to add to how much solid? Did you ever notice a change in smell or the appearance of an oily layer upon adding acids to the solution? Did you observe any boiling, popping, melting when heated with bunsen burner? Maybe most importantly, did you quantitatively determine the weight loss on heating (masses before and after heating)? $\endgroup$ – airhuff May 24 '17 at 1:25
  • $\begingroup$ For the solubility, I do not have any quantitative measurements, but the volume of water I have to add is easily 10x the volume of the solid present. I can give more details in a few days. When adding acids to solution there is no change in appearance of the oily layer - it just seems to sink. I haven't smelled the solution after adding acid. When heating the solid, there is a white fume/vapor emitted that rises up, but there was no boiling, popping, or melting. The mass of the solid before heating was 0.26g, and after heating it was reduced to 0.21g (so a 0.05g difference). $\endgroup$ – L. R. J. May 24 '17 at 3:08
  • $\begingroup$ 1) Is it hydroscopic (i.e. adsorb water from air to form solution) ? 2) Could it loose oxygen on heating and not water? Or maybe some other gas? $\endgroup$ – permeakra May 24 '17 at 5:30
  • $\begingroup$ @permeakra 1) I'm not sure if it is hydroscopic, how would I test for that? One thing I did notice is that after a few days, the powder in the container started to slightly clump together. Would this indicate anything? 2) I don't believe it would lose any gas, because it didn't bubble or anything. Most likely, everything that was lost was water. $\endgroup$ – L. R. J. May 24 '17 at 5:57
  • $\begingroup$ @L.R.J. 1) leave a small quantity on open air overnight and see what happens 2)incorrect. thermal decomposition of perchlorates produces oxygen, but they won't produce gases when mixed with (diluted) acid. $\endgroup$ – permeakra May 24 '17 at 10:14
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What about a possible $\ce{F-}$ anion? While the nature of the cation might be guessed based on the flame test, since the $\ce{AgNO3}$ test gave no precipitate, it might have produced $\ce{AgF}$ (soluble in water).

This said, by looking at the densities of $\ce{M-F}$ systems with a cation with orange flame and density around $\pu{2 g ml-1}$, I'd say it's probably $\ce{NaF}$.

You might test for fluoride with a selective electrode, if you have one.

Otherwise, here are listed some other analytical methods for the identification of fluoride anions:PubChem: Sodium Fluoride-Analytical Methods

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    $\begingroup$ Maybe, but only if the OP is wrong (as you suggest may be true) about $\ce{Ca^2+}$ being the cation. What bothers me most is that $\ce{PbF2}$ is only soluble to about 600 ppm, which you'd think would have formed a precipitate. $\endgroup$ – airhuff May 24 '17 at 1:34
  • $\begingroup$ You are right, actually I'm not familiar with the lead test. Anyway, I excluded calcium simply because the density is quite different from the one provided (3.18 g/ml). $\endgroup$ – The_Vinz May 24 '17 at 3:07
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    $\begingroup$ What is the best way to conclusively determine if it is Calcium or Sodium? Would adding Sodium Oxalate do the trick? $\endgroup$ – L. R. J. May 24 '17 at 3:10
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    $\begingroup$ @L.R.J. Atomic Emission Spectroscopy. $\endgroup$ – CoffeeIsLife May 24 '17 at 5:07
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    $\begingroup$ @The_Vinz "I'd say it's probably NaF". Sodium gives an intense yellow color in flame test which contradicts OP's result which is deep orange to red color. $\endgroup$ – Nilay Ghosh Sep 23 '20 at 13:41
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Calcium formate is a white-to-yellow or off-white crystalline powder, has a density of $\pu{2.02 g/ml}$, solubility of $\pu{\sim 17\%}$, decomposes at $\pu{300 ^\circ C}$ . It irritates eyes severely and has a stinging taste, therefore could have an "hard to describe odor" or effect on the nose. One mental conflict I have is that I would expect that measuring the density of a powder with a density of $\pu{2.02 g/ml}$ would actually give a value substantially lower than 2, but maybe the OP's experimental technique was exceptionally good.

When heated, calcium formate decomposes to calcium carbonate (Ref 1). The remainder $\pu{0.21 / 0.26}$ from the original heating is $\pu{80.7 \%}$; whereas the mw of $\ce{CaCO3 (100)}$ divided by the mw of calcium formate ($\pu{138}$) is $\pu{76.9 \%}$. Close enough: perhaps not all of the material decomposed. If not all decomposed, the heating was probably not severe. A test: acidify the burned product with $\ce{HCl}$, if it is $\ce{CaCO3}$, it will bubble. Another test: severely heat the burned product, see if it goes to $\ce{CaO}$ by checking pH ($\ce{-> 13}$) and seeing a further weight loss to $\pu{56\%}$ of the original.

A pH test on the original material could be informative: one SDS reported $\pu{7.5}$ for calcium formate, whereas an SDS for calcium acetate reported $\pu{7.6}$. The difference is small, but suggests that if the pH is closer to neutral than $\pu{7.6}$, the anion would be of a stronger acid than acetic (which formic is).

Ref 1: http://www.geosc.com/Assets/Files/Products-Docs/P-C-Product-Docs/Trimet-Products/CAF-TDS-US-Format.pdf

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  • $\begingroup$ I had the same thought, but the solubility of Ca formate is higher than what OP observed. Calcium oxalate is a bit more dense and much less soluble, but should have given precipitate with barium, except that the low solubility of original material means very low anion in the solution to which barium was added $\endgroup$ – Andrew Sep 22 '20 at 12:34
  • $\begingroup$ @Andrew: The OP says: "For the solubility, I do not have any quantitative measurements,". Dissolution is often slow to get to the published maximums. "Easily 10x the volume" means 1cc of solid (2 g) in 10-15 cc H2O = 20 - 13.3%, av = 16.6%. Quite close to 17%. $\endgroup$ – James Gaidis Sep 22 '20 at 13:47
  • $\begingroup$ I do not expect CaO to give pH>13. Calcium hydroxide is not sufficiently soluble. Based on the solubility product for the hydroxide, actual pH with CaO would be between 12 and 13. $\endgroup$ – Oscar Lanzi Sep 23 '20 at 13:35
  • $\begingroup$ Hmm, if it was calcium formate, then it would be release formic acid vapor on reacting with sulfuric acid (from ref.), which would have a distinct smell but OP didn't report any vapor or bubbles or any smell from that reaction. But again, that white ppt could be CaSO4? $\endgroup$ – Nilay Ghosh Sep 23 '20 at 13:49
  • $\begingroup$ @Oscar Lanzi: You are absolutely correct. I rounded up for effect. $\endgroup$ – James Gaidis Sep 23 '20 at 20:34
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There are ways to quite reliably determine what this solid is. As mentioned above, the counter ion most probably should be $\ce{Ca^2+}$, however there are many organic or inorganic anions that still fit the criteria. The $\ce{AgNO3}$ tests provide insight that its a non-coordinating counter ion.

Ways to determine the anion - however this all takes infrastructure that may or may not be accessible:

  • Most straightforward way would be to acquire NMR of the sample in $\ce{D2O}$. This will reliably tell you, if there are H atoms, C atoms (and their symmetry), 19-F NMR will reliably tell the presence of a fluoride atom or other fluoride containing counter ions (such as $\ce{BF4-}$ etc.)
  • There are many ways to create adducts of organic anions for GCMS analysis.
  • A HPLCMS or UPLCMS analysis will provide You with several pieces of information:
  • UV-Vis spectra will provide insight whether aromatic structures are part of the counter ion, and the MS will provide the m/z of any anion that is heavier than 50. Fortunately almost every single weakly coordinating anion is fairly heavy and should show up.

These methods should reliably provide the answer for the question asked.

As for the simple answer, $\ce{Ca(BF4)2.xH2O}$ may be the answer:

  1. Tetrafluoroborate salts are soluble and known to form hydrates.
  2. Tetrafluoroborate salts are not the most stable and can actually release $\ce{HF}$ gas in low quantities resulting in strong, irritating odor.
  3. This salts is more consistent with everything examined.
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One idea is to retry the brown ring test with the proper procedure, as the question admits that this had not previously been used (acidified ferrous sulfate was added wrongly). The other evidence is consistent with calcium nitrate, so we need to get that test right.

Calcium nitrate also decomposes at 500°C and above, releasing nitrogen dioxide leaving the solid oxide behind (only alkali metal nitrates, and not all of those, give the nitrite with this reaction). You would then see the brownish nitrogen dioxide gas.

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  • $\begingroup$ I was also suspecting a nitrate salt given that the ring test wasn't performed properly but shouldn't that salt react with barium chloride to form calcium chloride + barium nitrate both of them of which are white ppt (?) or shouldn't it react with conc. sulfuric acid/hydrochloric acid(OP didn't mention the concentration though) to form nitric acid which is indicated by evolution of red nitric oxide fumes? OP didn't report any precipitate or bubble or fumes formation in his reaction process. $\endgroup$ – Nilay Ghosh Sep 23 '20 at 16:25
  • $\begingroup$ I do not know this. The question does not specify adding concentrated acids. $\endgroup$ – Oscar Lanzi Sep 23 '20 at 18:16
  • $\begingroup$ That's where OP missed critically. Did everything except testing it with conc. acids (we don't know, OP is not active). btw, calcium nitrate has a density 0f 2.5 g/ml as opposed to OP's salt which is 2 g/ml. Shouldn't that be a problem? $\endgroup$ – Nilay Ghosh Sep 24 '20 at 4:19
  • $\begingroup$ @nilay the density depends on whether the OP accounted for less than 100% packing in the powder. $\endgroup$ – Oscar Lanzi Sep 24 '20 at 9:39

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