Sciencemadness Discussion Board

Identifying an unknown

chemico - 23-5-2013 at 13:04

Good Day Chemistry Aficionados,

Please help me with this problem if you can. We have an unknown compound that we need to identify and a wide array of chemical and potential tests to identify it with.

The problem is not getting it identified -- the problem is time. We have a limited amount of the time and the quicker and more efficient we get the identification the better off we score.

We are asking you guys to let us know what you guys believe are the most valuable tests to perform...and any flow charts or materials we should consult before we officially start this experiment.

Any information or words of advice are greatly appreciated.

Thank you,

Andrew

Hexavalent - 23-5-2013 at 14:09

Do you have any information about your compound?
If it's inorganic, conventional wet techniques may be the way to go. If it's organic, ideally you'd want IR and NMR spectra and a melting point.

chemico - 23-5-2013 at 15:29

We know it is an inorganic compound, for sure. That is pretty much it. We have a list of a few hundred compounds that are all potentially it and all are inorganic,

AndersHoveland - 23-5-2013 at 15:37

The oxidation of ammonia by hydrogen peroxide is spontaneous, but the reaction is rather slow. Alkaline conditions favor faster reaction rate. Nitrites can certainly be formed from this reaction, but it should be remembered that the reaction rate for the oxidation of nitrite to nitrate is much faster than the oxidation of ammonia. And of course, boiling temperatures will just cause the nitrite to be reduced (by the ammonia) to nitrogen gas.

amazingchemistry - 23-5-2013 at 18:18

Well, what does it look like, feel like or even smell like? that shouldn't take more that 20 seconds to ascertain and would help you at least narrow it down a bit. I assume you don't have access to anything like ICP? If you do, there are SOPs for that (you'd have to start by dehydrating and possibly calcinating it). If that's not what you want to do with it you could start by grouping the "hundreds of possible compounds" into broad categories.

woelen - 23-5-2013 at 22:46

Without more information we cannot really help you. Color of the compound is a strong way to narrow down things. If it is white or colorless, then you can rule out nearly all transition metal salts (a few exceptions are zinc, cadmium, mercury, silver, titanium).

There are schemes for qualitative analysis, especially in older literature. These schemes in fact are a kind of decision tree, consisting of many tests, which lead you to the anion and the cation of the inorganic species. The use of sulfide ion, ammonia and weak acids, all in dilute solutions plays an important role in this decision tree for determining the cation in the tests.

phlogiston - 24-5-2013 at 00:05

Any bit of information you can share on the origin of your mystery compound will also help a lot. It should narrow the range of possible ions. (eg you would not expect to find a rubidium compound seeping out of the wall of your cellar unless you live in a strange place).

sonogashira - 24-5-2013 at 01:19

I would recommend that you Google (and download) Vogel's Qualitative Inorganic Analysis.

You will need to perform a chromatographic separation to make sure that you don't have a mixture of compounds, or a salt.

It's an interesting problem. It would be more interesting if it was completely unknown. In which case, I would perform a melting point experiment to crudely determine whether it is organic (charring) or inorganic (high melting point). If organic I would separate it on TLC (or confirm that it is pure by having it remain un-separated on TLC), then scan the plate with diode array.

adamsium - 24-5-2013 at 01:54

This sounds typical of the practical exams often given for university chemistry courses. I'm pretty sure that chemico is not attempting to identify an entirely unknown compound found laying around or the like. This means that it can be assumed that the sample is not a mixture and is quite pure (hopefully ;) ).

Usually, in order to maximise efficiency, what you want to do is to rule out entire groups of compounds quickly. An example of this would be testing for chlorides, bromides or iodides with a very quick silver nitrate test. If you're able to wipe out large portions of the list of potential compounds, you're well on your way. Essentially, it's often a good idea to begin by finding out what you don't have; then you can start working on deciding what you do have. Because you're given a list of possible compounds, this approach is entirely possible (and pretty much what they generally seem to expect students to do for these exams).

Also, as woelen mentioned, you can often garner some very useful information qualitatively. This can also help in ruling out large portions of the list of possible compounds. Woelen also mentioned decision trees, which is exactly what you need for this type of thing if you want to maximise efficiency.

sonogashira - 24-5-2013 at 02:35

Quote: Originally posted by adamsium  
. This means that it can be assumed that the sample is not a mixture and is quite pure.
One couldn't be satisfied that there is not a mixture without having tested for it. It's possible/probable that molecular geometry would allow for some sort of separation/identification (perhaps using electrophoresis - though inorganic analysis is not my area).

Simple chemical tests are all very well, but the smart tortoise picks the right analytical method, and a universal detector, and waits for the print-out whilst his colleagues are frantically trying to precipitate silver chloride!


[Edited on 24-5-2013 by sonogashira]

adamsium - 24-5-2013 at 02:42

Quote: Originally posted by sonogashira  
Quote: Originally posted by adamsium  
. This means that it can be assumed that the sample is not a mixture and is quite pure.
One couldn't be satisfied that there is not a mixture without having tested for it. It's possible/probable that molecular geometry would allow for some sort of separation/identification (perhaps using electrophoresis - though inorganic analysis is not my area).


No; the reason that it can be assumed that this is not a mixture is clearly that it's a laboratory exam. It's (extremely) highly unlikely that a mixture would be given for this exercise. Also, chemico states in the initial post that they need to identify "an unknown compound". It's safe to assume that, unless I am completely wrong and this isn't a university lab exam, it is quite reasonable to assume that it is not a mixture.

Edit to add: Talk of specialised tests and such is pointless in the context of an undergrad lab exam. It's all about 'simple chemical tests'. That's the whole point.

[Edited on 24-5-2013 by adamsium]

sonogashira - 24-5-2013 at 02:52

You may be right.

(If you get cisplatin with an anomalous melting point then don't say you weren't warned!)

phlogiston - 24-5-2013 at 03:06

Another question is how much of the compound do you have to sacrifice to analysis?
If you have many grams, that greatly increases the variety of tests you can do. If you have a few mg only, many of the already proposed methods are useless.
Quote:
We know it is an inorganic compound, for sure. That is pretty much it. We have a list of a few hundred compounds that are all potentially it and all are inorganic,


Please post that list.

BTW if this a school assignment feel free to ask for hints, but don't expect anyone to solve it. And ask the mods to move it to 'beginnings'.

[Edited on 24-5-2013 by phlogiston]

chemico - 24-5-2013 at 03:13

Thank you for all the replies and thoughts. The lab is NOT for using complicated/intricate tests like a spectrometer, or a comprehensive analytical method are unfortunately not allowed.

The compound is going to be pure, barring a freak contamination (very unlikely and unintentional). Once we get the unknown the first time I'll report back and let you guys know the basic prelim stuff and include a few pictures as well...the 5 senses (sans smell), silver nitrate test, perhaps a flame test, ammonia reactions, and then everyone will have a better idea of what it can be.

At that point I will also type out the list, as after those tests are done we should be able to get it down to a reasonable amount of options. Finally, we won't be starting this until Tuesday or Wednesday of next week.

sonogashira - 24-5-2013 at 03:16

If it is definitely one of a known-group of chemicals, then I would just do thin layer chromatography on all of them simultaneously (having found the best solvent system for the compound that you have). You could eliminate some based upon the information that you have gained from finding the best solvent system for your unknown (eg. if your unknown is non-polar, don't test any polar substances). You can narrow it down very easily, and it would take perhaps 30 minutes to run an initial and confirmation chromatogram, using the fact that that they behave identically to confirm that they are identical. (You could do 2D chromatography if you want to be extra sure). I think that this would be the fastest method by far.

adamsium - 24-5-2013 at 04:03

For these sorts of lab exams, you're generally not given samples of all the possible compounds, only the one they want you to identify. Even if samples of all few hundred possible compounds were provided, trying to perform TLC on every one of them, including perfecting solvent systems for each would take..... a very, very long time. It also probably (in fact, almost certainly) wouldn't be all that useful; Rf values are subject to variation, even from plate to plate in the same solvent system (and even on the same plate) and are a poor way to attempt identification and characterisation in most cases.

I think some people are missing the point. These lab exams are designed mostly to test a few skills: some (usually fairly simple) actual chemistry, general lab technique, and reasoning / logic ability.

sonogashira - 24-5-2013 at 04:16

You wouldn't have to perfect the solvent systems for each substance(!), just for the ONE that you are trying to identify. There is nothing wrong with the method, nor with the accuracy of the method. Comparative chromatography is the most accurate and specific test in the whole of analytical science.

woelen - 24-5-2013 at 04:34

Especially if the compounds on the list are salts (which is very likely if you are sure it is something inorganic) then having a list brings down the number of options very much. You can separately do the analysis for cations and for anions.

The list may have a few hundreds of compounds, but the number of anions and the number of cations probably is in the order of 20 each. With 20 anions and 20 cations in theory you can make 400 compounds, in practice probable a little less (some combinations of cations and anions may lead to unstable salts, such as ammonium nitrite or copper(II) iodide). But by splitting analysis of anions and cations in separate problems you can tremendously reduce your problem.

adamsium - 24-5-2013 at 04:34

Quote: Originally posted by sonogashira  
You wouldn't have to perfect the solvent systems for each substance(!), just for the ONE that you are trying to identify. There is nothing wrong with the method, nor with the accuracy of the method. Comparative chromatography is the most accurate and specific test in the whole of analytical science.


Fair point regarding perfection of the solvent system. However, even perfecting one solvent system can be very difficult and time consuming.

Also, TLC is only as useful as the resolution you get. It's just not a practical approach in this situation, even if they were to be provided with samples of all possible compounds (which they almost certainly won't be). Let's say you run your sample and get an Rf value of 0.51. Then you run the other (several hundred) known samples... when you finish a couple of days later ( :P ), how do you then decide if it's the one with an Rf of 0.50, or the one with an Rf of 0.49, or 0.52, or even 0.51? And, unless you use extremely long plates, you may well end up with more than one compound with the same Rf value. It's just not practical in this scenario.

That said, I am interested in knowing more about TLC for inorganic substances. I've only personally used it for organic compounds.

Fantasma4500 - 24-5-2013 at 05:03

i would say... first off colour.. then you can say alot of things its not
does it react giving off gas with acid? probably carbonate or something similar
is it soluble?
does it melt?
does it oxidize by any means (test with bad fuel then with really really good fuel.. magnesium super fine dust perhaps)
and can you make a compound with it that somehow have a specific colour? oxides are usually not looking very much alike each other, or well at least many oxides are different.. perhaps turn it into metal with thermite reaction?

blogfast25 - 24-5-2013 at 05:59

Quote: Originally posted by chemico  

Please help me with this problem if you can. We have an unknown compound that we need to identify and a wide array of chemical and potential tests to identify it with.



Problems where there is no a priori knowledge about the substance at all are very rare and extremely difficult to solve. You must 'know' something about the substance/compound before you even start testing, surely? Or did it come falling out of the blue sky onto your desk? :(


[Edited on 24-5-2013 by blogfast25]

binaryclock - 24-5-2013 at 06:14

Quote:
Or did it come falling out of the blue sky onto your desk?

They found it on a mirror in an abandoned hotel room.


[Edited on 24-5-2013 by binaryclock]

phlogiston - 24-5-2013 at 07:35

TLC for inorganic compounds would not be the first thing to try IMO.


1. note color
2. heat as dry powder. Observe carefully.
3. solubility in water (cold, then hot).
4. If insoluble: solubility in nitric acid.
5. If not soluble in nitric acid: test solubility in concentrated ammonia and sodium hydroxide solutions
6. If soluble in nitric acid: split the nitric acid solution and test for precipitations after adding solutions of:
sodium chloride
sodium carbonate
sodium sulphate
7. If you get a precipitate, heat to boiling, and allow to cool down.
8. In addition, if you got a precipitate upon adding chloride:
- allow to stand in light for a while
- add concentrated ammonia

Be careful when adding the bases to the nitric acid solution (add slowly and don't use highly concentrated solutions).

I think these tests should yield sufficient information to distinguish between many simple inorganic salts.

[Edited on 24-5-2013 by phlogiston]

binaryclock - 24-5-2013 at 07:37

Quote:


1. note color
3. heat as dry powder. Observe carefully.
4. solubility in water (cold, then hot).

..



all good suggestions, but why is #2 so secret?

[Edited on 24-5-2013 by binaryclock]

ScienceSquirrel - 24-5-2013 at 07:53

#2 could a be a flame test.
This will reveal the presence of sodium, potassium, etc.
Or if it burns, goes black an organic material.


[Edited on 24-5-2013 by ScienceSquirrel]

phlogiston - 24-5-2013 at 07:55

Why do you get the impression there is anything secret about it?

I said observe carefully because there are many things that can happen, some of them subtle but informative.
Perhaps a little condensation in the top of a test tube, indicating release of water.
Or a color change from white to faint yellow (eg. zinc oxide).

Things like that.

Eddygp - 24-5-2013 at 08:09

Or a dehydration, like CuSO4ยท5H2O to CuSO4... for example.

sonogashira - 24-5-2013 at 09:30

On second thoughts, the fastest and easiest thing to do would be to take a melting point (or decomposition point) and compare it with known values. That should narrow your "hundreds" of choices down to about 2 or 3. Then get your paper chromatography going!

[Edited on 24-5-2013 by sonogashira]

chemico - 24-5-2013 at 09:34

To clarify some points of confusion...

1. This is a chemistry lab for a class. The compound will be pure and is not some random chemical we just stumbled upon.
2. This is about the most efficient simple tests. Which order would most quickly determine the compound's identity. This is a great time for a discussion of the most valuable chemicals (AgNO3, sulfuric acid, ammonia, etc.)
3. We will have a small container full of solution w/ no refills. The bottle size is 3 oz and it is approx 1/2 way full.
4. We have not received the chemical yet! We will get it on Wednesday of next week.
5. We only have access to our unknown chemical not all of them.
6. I am not sure if heating to melting point is safest, fastest, or most accurate route to go. It would take a fair amount of time to heat to melting point and how would you get the most accurate temperature reading?

If anyone has any more questions feel free to ask. Again we won't get the compound until next Wednesday so we won't be able to provide specific information about it until then.

Any comprehensive flow charts for id'ing? Or anyone want to rank the best compounds to use to identify an unknown (eg. sulfuric, ammonia, silver nitrate)

sonogashira - 24-5-2013 at 09:37

Quote: Originally posted by chemico  

6. I am not sure if heating to melting point is safest, fastest, or most accurate route to go.
I assure you that it is. Don't you have access to a melting point apparatus?

blogfast25 - 24-5-2013 at 10:18

As far as didactical tools go this is pretty poor, IMHO. Unless you're not presenting it as it will be presented to you... I'd really like to know at what level of chemical education some tw*t has introduced this.

[Edited on 24-5-2013 by blogfast25]

Eddygp - 24-5-2013 at 10:38

Quote: Originally posted by blogfast25  
As far as didactical tools go this is pretty poor, IMHO. Unless you're not presenting it as it will be presented to you... I'd really like to know at what level of chemical education some tw*t has introduced this.

[Edited on 24-5-2013 by blogfast25]


Basically, I was going to say that.

[Edited on 24-5-2013 by Eddygp]

amazingchemistry - 24-5-2013 at 19:57

Normally, when identifying an unknown we look for specificity. In this case though, your best bet is to go for the most general tests possible as your first step. I echo the comments above about color. React a small amount of it with an acid. Does it fizzle? its probably a carbonate. Take advantage of solution chemistry. Nearly everything dissolves in nitric acid. If it doesn't then you have narrowed down your list quite a bit. If it does, then you start testing solubilities in other compounds. If after, going through your decision tree, your don't get a precipitate, try a flame test. Your compound probably contains sodium, potassium or ammonium ions. Always pay attention to your senses. Your eyes are the most obvious, but if on heating your solution starts smelling like cat pee, then you very probably have ammonium ion. Vogels book has already been mentioned as a reference, you should also have a spot plate handy and read Feigl's "Spot Tests in Inorganic Analysis." Good luck and keep us updated. I love these kinds of problems :)

adamsium - 25-5-2013 at 06:51

I'm just going to leave this here.... http://firstyear.chem.usyd.edu.au/SummerSchool/LabManual/E05...

chemico - 2-6-2013 at 08:37

Well we finished day one. Here are our findings and early thoughts. Feel free to chime in with recommended tests, questions, etc.

DAY 1 OBSERVATIONS
--White powder-like compound. Odorless or extremely faint odor.

-- Attempt to dissolve in water was a little bizarre. Not very soluble in water. Appeared like a little dissolved but most just stayed at bottom, water got cloudy color.

--Flame test orange

-- COMPOUND + HNO3 --> no reaction (rules out carbonates)

-- COMPOUND + AgNO3 --> cloudy weak/light ppt

-- COMPOUND + H2SO4 + MnO2 --> exothermic rxn but stayed black color of manganese dioxide.
** on heating, colorless, odorless gas emitted
** bleached/turned red the blue litmus paper
** suspected chloride

-- COMPOUND + HNO3 + HCl --> no ppt BUT addition of HCl cleared up solution's foggy color before (does this support our compound being a chloride?)


That is all we were able to do in ~20 minutes we had with it. We are planning on doing the following tests next lab day (Tues.):


And from our list of 300+ compounds here are a list of plausible ones:




Also, it's important to note that there are a lot of other plausible compounds. This is just a sampling to show you guys what we are dealing with.

Right now we have a hunch it may be mercury (i) chloride but it's just a hunch..

[Edited on 2-6-2013 by chemico]

Eddygp - 2-6-2013 at 09:35

Chemico, firstly, why do you rule out sodium, potassium and rubidium?? Secondly, it should contain chloride or bromide ions (silver precipitate?) and a less electronegative cation than Ag.

chemico - 2-6-2013 at 10:20

Eddygp, wouldn't virtually all sodium compounds readily dissolve in water? This was very slightly soluble in water. If anyone knows of any sodium compounds that don't really dissolve in water, I would be welcome to test further for sodium.

I agree with the second part. I'll take a second look at the list and refine at a later point.

amazingchemistry - 2-6-2013 at 10:29

Potassium and sodium should be ruled out at the end of the decision tree, as nearly all sodium compounds are soluble in aqueous sol. Excuse for the seemingly obvious question, but when you said, "no reaction" with HNO3, did you mean "no bubbling, but dissolution" or "inertness"? I'm inclined to think of the second one, as I'd be surprised if your compound was unreactive toward HNO3

chemico - 2-6-2013 at 10:32

Amazingchemistry, you are correct. By no reaction I meant no bubbling but dissolution.

Eddygp - 2-6-2013 at 11:42

I'm sorry chemico. I forgot that it was insoluble!

chemico - 2-6-2013 at 12:52

No problem, do you have any recommended tests we should perform? We really need a way to test for mercury or a test for lead.

Vargouille - 2-6-2013 at 13:48

Remember your flame test results: lead is reported to give a blue/white flame. You can try to reduce a mercury salt, and differentiate between similarly colored metals using your flame test results. You can also rule out sulfites, since you didn't note a sulfurous odor. The carbonates are gone too, not just unlikely, since it didn't off-gas with nitric acid. Many transition metal salts can be ruled out as well: a white salt is unlikely to be an iron, cobalt, or copper salt. You should really look through your salts and cross the colored ones off the list, as well.

amazingchemistry - 2-6-2013 at 18:03

It doesnt seem to be a carbonate. If im not mistaken, treating with HCl should tell you if you have a Pb compound. Flame tests are your surest bets for sodium, unless you have access to zinc uranyde acetate :P (uranium sodium salts are insoluble)

Finnnicus - 2-6-2013 at 19:59

What about sodium rhodizonate which tests for some heavy metals, IIRC.

amazingchemistry - 2-6-2013 at 22:42

Any decent text will tell you that if a precipitate is present when treating a solution containing your ions of interest (say, your nitric acid + compound solution) with HCl it can either be PbCl, Hg2Cl2, or AgCl. Pb is sufficiently soluble in hot water to test this wash by adding chromate ions (or some other ions whose Pb salts have a smaller water solubility than Cl) to get a precipitate. Hg2Cl2 should precipitate with ammonia. AgCl should dissolve and re-precipitate when acidifying. Given that this is standard procedure (that some cursory amount of reading would have uncovered) I'm starting to feel like we are just giving you the answers to what is supposed to be your puzzle. I apologize if this is a mistaken impression.

chornedsnorkack - 2-6-2013 at 23:01

Quote: Originally posted by amazingchemistry  
Any decent text will tell you that if a precipitate is present when treating a solution containing your ions of interest (say, your nitric acid + compound solution) with HCl it can either be PbCl, Hg2Cl2, or AgCl.

CuCl would get dismuted on solution, and so would AuCl. But how about TlCl?

blogfast25 - 3-6-2013 at 03:38

Quote: Originally posted by chornedsnorkack  
CuCl would get dismuted on solution, [...]


If you meant 'disproportionate' then no: in the absence of oxidisers CuCl is quite stable. It's also insoluble.

[Edited on 3-6-2013 by blogfast25]

phlogiston - 3-6-2013 at 05:22

You are not using all the information you already have.
On the basis of color alone you can already exclude all the chromium, copper, iron, cobalt compounds that you have listed and then some (lead (ii) iodide, mercury(ii)iodide for instance which are well known to be brightly colored), with the possible exception of copper(i)iodide which appears to be a bit off-white.

Did I understand correctly you observed complete dissolution upon adding HNO3?
An important hint! Oxides and hydroxides are likely candidates then.
The solubility of chlorides doesn't normally change upon addition of HNO3.

Also upon addition of HCl you would expect chlorides to become less soluble, not more. (look up "common-ion effect"), so again, your observations argue against a chloride.

Was there any noticeable heat effect upon mixing with water?
Solubility in hot water?
Effect of dry heating of the compound?
Given the flame test, perhaps you should include a few more calcium compounds on your list. So far, your observations are consistent with calcium hydroxide for instance, which isn't on your list.
You should repeat your flame test in the presence of chloride (dip wire in HCl, then in powder and heat in colorless flame). This will give you the well known green for barium/red for strontium/blue for copper if they are present.

So, there are a lot of simple things you can do to narrow down the list before resorting to very specific and complex tests for lead/mercury/chloride etc.

adamsium - 3-6-2013 at 05:49

I wouldn't get too carried away with looking for mercury and cadmium unless you really start ruling out pretty much everything else (in a methodical manner). The likelihood of them giving salts of mercury, cadmium, etc to undergrads (especially without telling you what it is... but even just at all) these days is probably infinitesimally minute. They used to just about throw that stuff around like table salt, but it's a pretty big 'no go' to use those things now, both for student safety and environmental reasons. I'm not saying that you should just totally assume it can't be that (although, really, part of doing science is using common sense, and this is common sense). Just know that the chances of it being anything like that are minimal, at most. Of course, I could be entirely wrong and it could be a mercury salt, but I'd be very, very surprised.

An orange flame test would perhaps indicate something like Ca2+. You also said that when you did the halide test, a gas was evolved, but it was odourless. Chlorine would certainly not be odourless. Don't discount the fact that there could just be halide contamination somewhere giving a slight cloudy precipitate.

"addition of HCl cleared up solution's foggy color before (does this support our compound being a chloride?)" No. I'd expect most chloride salts to be precipitated by addition of HCl, not dissolved, unless a complex is forming (which is unlikely, as you don't seem to have a d-block salt). Consider Le Chatelier's principle and, specifically, the common ion effect.

I'll leave it at that for now.

Edit: I started this and left it for some time. Looks like phlogiston was typing up much of what I was saying in the meantime :) Also, phlogiston, I was considering Ca(OH)2 as a very strong contender based on what has been presented, but stopped short of saying so... so I definitely agree with your assessment.

Honestly, forget about mercury (and anything else remotely nasty) for now. All that stuff is there simply to 'pad out' the list.


[Edited on 3-6-2013 by adamsium]

phlogiston - 3-6-2013 at 06:29

That's funny adamsium, we got nearly exactly the same points. We clearly think alike.

Vargouille - 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.

chemico - 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





[Edited on 4-6-2013 by chemico]

chemico - 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.

Vargouille - 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.

amazingchemistry - 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]

phlogiston - 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]

adamsium - 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.

phlogiston - 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.

adamsium - 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.

phlogiston - 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]

adamsium - 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.

phlogiston - 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]