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Author: Subject: Identifying an unknown
Vargouille
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[*] posted on 3-6-2013 at 07:06


Can you test the pH of a pure solution of the unknown? The only pH test I think you did was in regards to the gas that came off with the MnO2 test, though I could very well be incorrect.
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chemico
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[*] posted on 4-6-2013 at 13:57


@Vargouille The compound does not dissolve in water, so I am not sure how effective a pH test would be with litmus paper. If you still think it would be helpful to do one, please let me know.

DAY 2
*did not dissolve in ethanol
*when just the pure, unknown compound was heated a gas was created
*Na2CO3 + unknown --> no ppt
*Na2S + acidified solution (HCl) --> cloudyish color
***addition of NaOH w/ this acidified solution --> white ppt
*Test for sulfate: BaCl2 + unknown --> white ppt
*Test for nitrate: CuI + H2SO4 + unknown + heat --> colorless gas w/ pungent, choking odor
*NaBr + unknown --> no rxn, solid settled at bottom
*Na3PO4 + unknown --> no reaction, solid settled at bottom


I am currently revising the list of possibilities and will edit this post w/ it once I'm done.


EDIT: Revised list. No calcium or barium compounds due to no ppt w/ sodium carbonate. The white ppt w/ barium chloride also rules out chlorides, bromides, iodides...strong support for a sulfate



  • aluminum hydroxide monohydrate
  • aluminum phosphate
  • antimony sulfide
  • iron (ii) sulfate monohydrate
  • lithium phosphate
  • sodium sulfate decahydrate
  • zinc hydroxide
  • zinc phosphate
  • zinc sulfide
  • zinc sulfate dihydrate



[Edited on 4-6-2013 by chemico]
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[*] posted on 4-6-2013 at 14:14


As far as calcium compounds go...wouldn't a calcium compound react with sodium carbonate to make a white precipitate (CaCO3)? We had no reaction so that technically rules out all calcium, strontium, barium, and magnesium compounds.
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Vargouille
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[*] posted on 5-6-2013 at 06:16


How did you preform your tests? You should suspend your compound in water, filter, and then preform the tests on the solution, because you can miss slight precipitation otherwise. Nothing is completely insoluble, so a pH test would still be useful. Moreover, antimony sulfide is not white, and lithium gives a red flame test. Iron (II) sulfate monohydrate will absorb water and then dissolve. Zinc gives the wrong color flame test, and sources suggest that aluminum gives no color on a flame test. Sodium give bright yellow flame tests.
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amazingchemistry
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[*] posted on 5-6-2013 at 21:01


"When just the pure unknown was heated, a gas was created" Did this gas have any color or odor? You might have tested it with a lit wooden splint to see if a flame was produced from this gas. When you say you combined your compound with Na2CO3, do you mean you combined an aqueous solution of your compound with NaCO3? Or did you dissolve your compound in acid prior to attempting to react it with Na2CO3? How would sulfate compounds explain your previous results? Specifically, the results of the MnO2 test?

[Edited on 6-6-2013 by amazingchemistry]
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[*] posted on 6-6-2013 at 01:37


Quote: Originally posted by chemico  
@Vargouille DAY 2
*Na3PO4 + unknown --> no reaction, solid settled at bottom
[Edited on 4-6-2013 by chemico]


"solid settled at bottom"
suggests you are doing the precipitation reactions incorrectly.
You should be mixing two clear solutions and then look for any precipitate forming.
It is quite possible otherwise to miss an insoluble precipitate forming around the grains of the compound you added in solid form.

You generally leave us guessing a bit as to how you perform these tests exactly. I assume you are dissolving your unknown in HCl before the precipitation test, correct? Are the concentrations of your test solutions known?

It is a bit unusual (though not impossible) that you are getting a precipitate with NaOH and not with Na2CO3. If I can assume for the moment that you are indeed dissolving your unknown in HCl first, are you sure you added enough Na2CO3 (solution? solid?) to neutralize the HCl at all?

Quote:
when just the pure, unknown compound was heated a gas was created


This is why I said 'observe carefully' in my earlier post.

What exactly did you observe and how do you know a gas was created? Was there a residue left after the reaction? Any change in appearance? Smell? What were the properties of the gas? Color? Did you try to light it? Glowing splint? Burning splint?

[Edited on 6-6-2013 by phlogiston]




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[*] posted on 6-6-2013 at 02:57


Quote: Originally posted by chemico  
@Vargouille The compound does not dissolve in water, so I am not sure how effective a pH test would be with litmus paper. If you still think it would be helpful to do one, please let me know.

DAY 2
*did not dissolve in ethanol
*when just the pure, unknown compound was heated a gas was created
*Na2CO3 + unknown --> no ppt
*Na2S + acidified solution (HCl) --> cloudyish color
***addition of NaOH w/ this acidified solution --> white ppt
*Test for sulfate: BaCl2 + unknown --> white ppt
*Test for nitrate: CuI + H2SO4 + unknown + heat --> colorless gas w/ pungent, choking odor
*NaBr + unknown --> no rxn, solid settled at bottom
*Na3PO4 + unknown --> no reaction, solid settled at bottom


I am currently revising the list of possibilities and will edit this post w/ it once I'm done.


EDIT: Revised list. No calcium or barium compounds due to no ppt w/ sodium carbonate. The white ppt w/ barium chloride also rules out chlorides, bromides, iodides...strong support for a sulfate



  • aluminum hydroxide monohydrate
  • aluminum phosphate
  • antimony sulfide
  • iron (ii) sulfate monohydrate
  • lithium phosphate
  • sodium sulfate decahydrate
  • zinc hydroxide
  • zinc phosphate
  • zinc sulfide
  • zinc sulfate dihydrate



[Edited on 4-6-2013 by chemico]


If you're adding NaOH to an acidified solution of your unknown, you may well just be getting back a precipitate of your original unknown, which could be a hydroxide. I also am unsure how you're doing these tests and it sounds a bit like you're not quite doing some of them correctly. I think someone already explained how you need to do them, so, you may need to redo some.

Why are you testing for nitrate? Your compound is insoluble / sparingly soluble.
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[*] posted on 6-6-2013 at 03:42


None of the compounds on your list matches your test results.

aluminium hydroxide --> should give a precipitate with carbonate
aluminium phosphate --> does not dissolve in HCl
Antimony sulfide --> not white
iron(ii)sulfate--> not white, good solubility in water
lithium phosphate --> does not dissolve in HCl and even if it did would not give a precipitate with NaOH
sodium sulphate --> soluble in water, does not give precipitate with NaOH
zinc hydroxide --> should give precipitate with carbonate
zinc phosphate --> does not dissolve in HCl and even if it would give a precipitate with Na2CO3
zinc sulphide ---> solution in HCl would be smelly and should give precipitate with carbonate,
zinc sulphate --> soluble in water

As said, I doubt your Na2CO3 precipitate test result is correct.




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[*] posted on 6-6-2013 at 03:52


Quote: Originally posted by phlogiston  
None of the compounds on your list matches your test results.

aluminium hydroxide --> should give a precipitate with carbonate
aluminium phosphate --> does not dissolve in HCl
Antimony sulfide --> not white
iron(ii)sulfate--> not white, good solubility in water
lithium phosphate --> does not dissolve in HCl and even if it did would not give a precipitate with NaOH
sodium sulphate --> soluble in water, does not give precipitate with NaOH
zinc hydroxide --> should give precipitate with carbonate
zinc phosphate --> does not dissolve in HCl and even if it would give a precipitate with Na2CO3
zinc sulphide ---> solution in HCl would be smelly and should give precipitate with carbonate,
zinc sulphate --> soluble in water

As said, I doubt your Na2CO3 precipitate test result is correct.


I still think that the calcium hydroxide is looking pretty good here.

I think part of the problem, aside from the fact that the tests are likely being conducted / interpreted incorrectly, is that there seems to be a lack of methodicalness.
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[*] posted on 6-6-2013 at 04:21


Yes, exactly.
Actually, I think that is precisely the main lesson he is supposed to learn from this, apart from gaining a little basic experience doing experiments and making/interpreting observations. Woelen gave him good advice early on how to approach this problem in a systematic way. (Divide into cation/anion and make a decission tree. It is likely he would have identified the compound by now following that approach. I also still think calcium hydroxide is a likely candidate still, except that he got a precipitate upon adding barium chloride. Calcium sulfite is also consistent with many of his tests, including the BaCl2 test, but he should have noticed some unmistakable smells (SO2, H2S) in some of his experiments. He does not seem to report most details, however, or may have missed it completely (fumehood?).

[Edited on 6-6-2013 by phlogiston]




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[*] posted on 6-6-2013 at 04:44


Shouldn't there be a barium hydroxide precipitate with BaCl2 if it was calcium hydroxide, though? Or am I missing something?

I actually said something along those lines earlier, too, about what this sort of test at uni (I've done similar myself... I actually like them... and it's generally done in one lab session, so they expect you to be pretty organised and efficient) is supposed to be about. Also, if this uni is like mine, they generally seem to be more interested in the process than the end result. So, you can get entirely the wrong answer at the end, but if you went about it in a methodical and logical manner, perhaps misinterpreting or messing up one of the tests, you could still get a very good score... they don't seem to assign too many marks for the right answer at the end, it's mostly about the process and your approach.

I also linked to a document earlier... If you haven't had a look at it, chemico, you should definitely take a look. The flow chart arrangement is important.
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[*] posted on 6-6-2013 at 05:04


Barium hydroxide is more soluble than calcium hydroxide (3.89 g/100ml or 0.22M and 0.173 g/100ml or 0.023M) respectively at 20 deg C), so you would not expect a precipitate.
Also, he is not very clear about it but if I interpret his posts correctly he is doing his precipitation reactions with the solution of the unknown in hydrochloric acid. So then, you would not expect a precipitate either unless there is a cation that gives a precipitate with barium.

I also liked these kinds of experiments. Similarly, I also very much enjoyed the exams where you had to identify an organic compound on the basis of NMR, IR and MS spectra. Fun puzzle, and very satisfying to find you can extract pretty complex structures from that data.
At our uni the process was similar, and it should be. You should be taught how to think, learn and plan. You can always look up facts. (Solibilities in this case). The only exception was the analytical chemistry labcourse, where we got a penalty for getting a calibration curve with r2<0.9999 or if our measurement was off by more than 5%.

[Edited on 7-6-2013 by phlogiston]




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