Sciencemadness Discussion Board - Identifying Zn2+/Al3+ ions

Identifying Zn2+/Al3+ ions

AWLB - 27-7-2015 at 02:31

Does anyone know how Zn2+/Al3+ ions can be identified in solution. I recently found a sample of metal which I assumed to be either zinc or aluminum; which I dissolved in conc. hydrochloric acid- the salt formed is of course ZnCl2 or AlCl3, is there a way to determine which cation (Zn2+/Al3+) is in solution by a simple test?

Thanks in advance.

unionised - 27-7-2015 at 02:35

Add excess ammonia. Al(OH3) won't dissolve Zn(OH)2 will.

Of course, it could be an alloy.
If there's a precipitate that doesn't dissolve, filter it off or wait for it to settle then pour off the clear solution. Add a dilute acid slowly- if a precipitate forms whrn the excess ammonia in neutralised then there's probably zinc in there as well as aluminium.

AWLB - 27-7-2015 at 02:38

Thank you very much for replying, that seems like a good way to test for the ions. Why does the Zn(OH)2 dissolve with excess ammonia, does it form a complex?

Thank you for your help.

j_sum1 - 27-7-2015 at 03:11

Quote: Originally posted by unionised

Add excess ammonia. Al(OH3) won't dissolve Zn(OH)2 will.

Of course, it could be an alloy.
If there's a precipitate that doesn't dissolve, filter it off or wait for it to settle then pour off the clear solution. Add a dilute acid slowly- if a precipitate forms whrn the excess ammonia in neutralised then there's probably zinc in there as well as aluminium.

It ain't quite that easy.
Zn and Al behave in a very similar manner.
https://www.youtube.com/watch?v=KNvPfGBnN94

(I only know this because I have been looking at it recently. This is the third time recently I have had occasion to post this video -- which I believe shows a method for differentiating between these two ions. https://www.youtube.com/watch?v=IFlIi_Xk3FE)

AWLB - 27-7-2015 at 03:37

Thank you for posting those informative videos j_sum1, they are very helpful. I believe I have now identified the cation using ammonia- it appears to be aluminum, though as you said they are almost indistinguishable(though the precipitate does not dissolve in excess NH4OH); the metal also is also not dense enough to be zinc. Thank you very much (is that your YouTube channel ? Your videos are very good!)

j_sum1 - 27-7-2015 at 03:52

Not my channel.
Agreed, they are presented to a high standard. I have bookmarked it because I have lots to learn.

AWLB - 27-7-2015 at 03:55

Thank you nonetheless for helping me j_sum1, I appreciate it.

I also have a lot to learn about chemistry, but learning is perhaps the greatest enjoyment of this hobby, I am still at school so these videos, this blog and home chemistry are very helpful.

[Edited on 27-7-2015 by AWLB]

blogfast25 - 27-7-2015 at 05:59

j_sum1:

Zinc does form a tetrammine complex and Al does not. But redissolving the firstly formed Zn(OH)2 into an excess of NH3(aq) takes time and agitation.

Another difference between Al and Zn is that when a carbonate solution is added Al(OH)3 precipitates (with CO2 escaping), where as with Zn the carbonate or basic carbonate precipitates.

[Edited on 27-7-2015 by blogfast25]

AWLB - 27-7-2015 at 06:23

Thank you blogfst25 for providing an alternate way to distinguish between Zn2+ and Al3+ ions in solution.

unionised - 27-7-2015 at 06:34

Quote: Originally posted by blogfast25

j_sum1:

Zinc does form a tetrammine complex and Al does not. But redissolving the firstly formed Zn(OH)2 into an excess of NH3(aq) takes time and agitation.

Another difference between Al and Zn is that when a carbonate solution is added Al(OH)3 precipitates (with CO2 escaping), where as with Zn the carbonate or basic carbonate precipitates.

[Edited on 27-7-2015 by blogfast25]

Tell me, if the zinc forms a basic carbonate, where does the CO2 go?
Also, if the solution is already acidic 9and it will be- remember -, it was made by dissolving the metal in acid- then the zinc solution and the aluminium solutions will both fizz and give a white solid.
Not quintessentially specific.

unionised - 27-7-2015 at 06:41

Quote: Originally posted by j_sum1

Quote: Originally posted by unionised

Did you actually watch the video?
At about 5:10 you can see that the Zn(OH)2 dissolves.
At about 4:00 they confirm that Al(OH)3 doesn't dissolve in excess
The reactions are fast ( it takes a few seconds to mix the solutions properly in the tube.)

blogfast25 - 27-7-2015 at 06:57

Quote: Originally posted by unionised

Actually the first caption (a propopos Zn2+) says: 'A white precipitate forms that does not redissolve in excess ammonia'.

Problem is that they've got their proportions all wrong: far too much ZnCl2 solution. With about 1 ml of solution, redissolving the formed Zn(OH)2 into excess NH3(aq) would be quick and conclusive.

blogfast25 - 27-7-2015 at 07:02

Quote: Originally posted by unionised

I'm not claiming it's 'quintessentially specific'. But one can carefully wash the precipitate and subject it to acid again. The Al precipitate will not fizz, the Zn one will. That's distinctive: the first one's a hydroxide, the second one's a carbonate.

If a basic zinc carbonate forms then of course CO2 also evades. It's not clear to me in which circumstances 'straight' or basic carbonate forms though.

A mineral basic zinc carbonate - Hydrozincite:

Zn5(CO3)2(OH)6 = 2ZnCO3.3Zn(OH)2

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

[Edited on 27-7-2015 by blogfast25]

unionised - 27-7-2015 at 07:17

Quote: Originally posted by blogfast25

Quote: Originally posted by unionised

The first caption is wrong; the first solution is Mg++

blogfast25 - 27-7-2015 at 07:27

According to this somewhat fluffy paper synthetic basic zinc carbonates do exist:

http://rruff.info/uploads/CM8_92.pdf

unionised - 27-7-2015 at 07:31

Quote: Originally posted by blogfast25

Quote: Originally posted by unionised

indeed, but that's not what you wrote at first is it?

It's not always possible to produce straight carbonates. Using bicarbonate as the precipitant sometimes helps (slightly lower pH an thus slightly less OH- about the place) and working under a high pressure of CO2 canalso work.

blogfast25 - 27-7-2015 at 07:33

Quote: Originally posted by unionised

The first caption is wrong; the first solution is Mg++

Yep. It's a shame about the caption mix-up: in the third experiment (with Zn2+) the precipitate clearly redissolves quite quickly in NH3(aq).

[Edited on 27-7-2015 by blogfast25]

blogfast25 - 27-7-2015 at 07:36

Quote: Originally posted by unionised

indeed, but that's not what you wrote at first is it?

'Pffft' and 'sigh' spring to mind.

AJKOER - 27-7-2015 at 10:11

Here is an interesting way based on the observation that adding Na2CO3 or NaHCO3 to say ZnCl2 creates a precipitate of ZnCO3 that can be partially redissolved upon the addition of more Na2CO3 or NaHCO3. Reference: see http://pubs.acs.org/doi/abs/10.1021/ic50190a004 and a very dated observational account at https://books.google.com/books?id=FSxGAQAAMAAJ&pg=RA2-PA... of Zinc bicarbonate's stability/solubility in natural streams.

So add a small amount of NaHCO3 to aqueous ZnCl2 in a long thin tube. Shake and poor half of the suspension into an identical tube. Now, add more solid Na2CO3 to just one tube and stir to dissolve. If the relative height of the precipitate is lower in the treated tube on standing (which is the best technique but possibly sooner, if comparisons are taken at the same time when both tubes were shaken together to dissolve the added carbonate in one tube assuming vicosity differences can be adjusted or not too material) , than some of the treated ZnCO3 is now a soluble complex (perhaps as Zn(HCO3)2 ). Add more Na2CO3 and see if a further relative decline is visible. This property is not shared by Aluminum unless the Na2CO3 /Al(OH)3 mix is boiled, at which point an aluminate is formed.

I am assuming we known the weight of the mystery metal and max % Zn so that a suitable amount of Na2CO3 can be added to precipitate all the Zinc as ZnCO3.

[Edit] Note, this test works in alloy situations as more ZnCO3 relatively dissolves. Also, boiling the carbonate rich tube reveals the presence of Aluminum.

[Edited on 27-7-2015 by AJKOER]

blogfast25 - 27-7-2015 at 13:21

Quote: Originally posted by AJKOER

So add a small amount of NaHCO3 to aqueous ZnCl2 in a long thin tube. Shake and poor half of the suspension into an identical tube. Now, add more solid Na2CO3 to just one tube and stir to dissolve. If the relative height of the precipitate is lower in the treated tube on standing (which is the best technique but possibly sooner, if comparisons are taken at the same time when both tubes were shaken together to dissolve the added carbonate in one tube assuming vicosity differences can be adjusted or not too material) , than some of the treated ZnCO3 is now a soluble complex (perhaps as Zn(HCO3)2 ). Add more Na2CO3 and see if a further relative decline is visible.

Interesting as all this sounds, but have you tested this?

And this bit:

Quote:

This property is not shared by Aluminum unless the Na2CO3 /Al(OH)3 mix is boiled, at which point an aluminate is formed.

... is, despite your past announcements, NOT true. Carbonate solutions are not alkaline enough to form aluminate in any significant amounts. Instead, carbonate solutions can be used to precipitate Al(OH)3 (although NH3 is preferred).

[Edited on 27-7-2015 by blogfast25]

Brain&Force - 27-7-2015 at 14:52

Really easy test:

Run a current through the solution. Zinc plates out as the metal, aluminum makes a gummy gunk.

j_sum1 - 27-7-2015 at 15:37

Glad to have provoked some discussion here.
@unionised -- Yes I did watch the video. But did not immediately review before posting. I was actually recalling in my mind Al behaviour with NaOH in which it does redissolve with excess. My bad. In my defense I had been watching a gazillion YT vids and... well, details blur.

@zts16 -- Of course. If one has narrowed down to Al or Zn, then I think that is by far the easiest test.

[edit] [smacks forehead]
@brain and force -- Of course. If one has narrowed down to Al or Zn, then I think that is by far the easiest test.
(Thanks blogfast)

[Edited on 28-7-2015 by j_sum1]

blogfast25 - 27-7-2015 at 15:57

Quote: Originally posted by j_sum1

@zts16 -- Of course. If one has narrowed down to Al or Zn, then I think that is by far the easiest test.

I didn't see zts16 comment on this thread?

AJKOER - 27-7-2015 at 16:59

Blogfast:

You may be right that a fresh precipitate of Al(OH)3 does not dissolve in boiling concentrated aqueous Na2CO3.

Honestly, I am not completely sure, but please explain how Aluminum metal (naturally coated with Al2O3, certainly less reactive than Al(OH)3 ) manages to entirely dissolve, nevertheless, in boiling concentrated Na2CO3 (which I have performed on several occasions in a microwave) and previously reported on SM as a path to sodium aluminate.

blogfast25 - 27-7-2015 at 17:31

AJ:

Firstly, you're talking small amounts of Al with CONCENTRATED, BOILING Na2CO3. Chances are that on cooling Al(OH)3 dropped out anyway. Carbonate solutions are ORDERS of MAGNITUDE less alkaline than strong (even dilute) alkali (NaOH e.g.). Aluminate is only stable at very high pH, although at say 11 very limited amounts of Al(OH)3 can stay in solution.

In the context of what is being discussed here, you're peddling misinformation.

If you still believe Al(OH)3 can be dissolved in boiling, concentrated Na2CO3, then prove it but I'll remind you that once not so long ago you believed NH3 can do that too.

AJKOER - 28-7-2015 at 06:07

Blogfast:

Assume we heat up aqueous Na2CO3 to boiling in a wide mouth vessel in air, then we have the reactions:

2 Na2CO3 + 2 H2O = 2 NaHCO3 + 2 NaOH

And possibly to a limited extent as well the breakdown of the aqueous Sodium bicarbonate on boiling (although the supplied reference claims completion on boiling):

2 NaHCO3 + H2O ---Boiling → CO2(g) + 2 H2O + Na2CO3 (see http://antoine.frostburg.edu/chem/senese/101/inorganic/faq/c... )

Or, on net, an optimistic, but likely inaccurate assessment of the max possible NaOH formation:

Na2CO3 + H2O ---Boiling → CO2(g) + 2 NaOH

And indeed a limiting one as the following reaction has come to my attention in the presence of say Al(OH)3, the NaHCO3 is possibly consumed. To quote https://books.google.com/books?id=FijWBgAAQBAJ&pg=PA149&...

Al(OH)3 + NaHCO3 = NaAl(CO3)(OH)2 + H2O

Now, we bring into play some electrochemistry, to quote Wikipedia on the Aluminum-air battery (https://en.m.wikipedia.org/wiki/Aluminium–air_battery#Elec... :

"The anode oxidation half-reaction is Al + 3OH− → Al(OH)3 + 3e− +2.31 V.

The cathode reduction half-reaction is O2 + 2H2O + 4e− → 4OH− +0.40 V.

The total reaction is 4Al + 3O2 + 6H2O → 4Al(OH)3 + 2.71 V.

About 1.2 volts potential difference is created by these reactions, and is achievable in practice when potassium hydroxide is used as the electrolyte. Saltwater electrolyte achieves approximately 0.7 volts per cell."

Note the overall consumption of water which could important implications for pH and solubility.

Now, in our case upon adding Aluminum, to boiled Na2CO3 in the presence of oxygen (or an oxygen source), we have a better electrolyte than saltwater but not as effective as the KOH. Note, it is O2 and H2O attacking the Al and not the electrolyte!

At cooler temperatures (under 50 C) and agreeable pH, some Sodium aluminate could form via the equilibrium reaction:

Al(OH)3(s) + NaOH = NaAl(OH)4

With respect as to whether the Sodium aluminate is stable see https://www.google.com/url?sa=t&source=web&rct=j&...

A pH related comment on the article page number 195 is, to quote:

"The optimum pH during precipitation of aluminum hydroxide is within the range 8.7–9.3 [5,6]. "

at least for the purposes of that article.

To speed things up in the electrochemical half reactions, one could use dilute H2O2 as an oxygen source and I suspect microwave heating will also be benefical.
------------------------

I have since run an experiment and for courtesy, have decided to place the results with pictures in a soon to be created new thread (exists link: http://www.sciencemadness.org/talk/viewthread.php?tid=63150 ).

[Edited on 29-7-2015 by AJKOER]

blogfast25 - 28-7-2015 at 11:28

AJ:

Your first ‘reaction; is nothing but the first hydrolysis of the carbonate ion:

CO32-(aq) + H2O(aq) < === > HCO3-(aq) + OH-(aq)

pKb = 3.67

Approximately:

pOH = ½(pKb + pC)

So for carbonate concentration of 1 M:

pOH = 1.8, pH = 12.2

And [HCO3-] = 0.015 M

For higher concentrations of carbonate, this increases approx. with the square root of that concentration.

Sure, bicarbonate slowly loses CO2 but boiling a carbonate solution is a near infinitely slow way of preparing an NaOH solution!

As regards the oxidation of Al by water and oxygen to Al(OH)3, bear in mind that gases like oxygen are very poorly soluble in boiling water.

[Edited on 28-7-2015 by blogfast25]

unionised - 28-7-2015 at 12:09

Ajoker, to put it simply, this reaction
2 Na2CO3 + 2 H2O = 2 NaHCO3 + 2 NaOH
does not happen. In fact the opposite reaction goes practically to completion.

And this reaction
Na2CO3 + H2O ---Boiling → CO2(g) + 2 NaOH
is even less plausible.

However at very high concentrations sodium carbonate will dissolve a little aluminium hydroxide.
Who cares?
It's a lousy way to test for Al in the presence of Zn.

AJKOER - 28-7-2015 at 19:00

Unionised:

On boiling, Aluminum in hot Washing soda is a known hazard for Aluminum cookware!

As Na2CO3 is a common ingredient in dishwasher detergent, an ongoing one at that.

It literally dissolves the Al, so please verify it if you don't believe.

The simplest chemical argument is that the reaction is moved to the right with the NaOH attacking the Aluminum.

Na2CO3 + H2O --boiling--) NaHCO3 + NaOH

It is a vigorous reaction, but if you want to see insane, replace H2O with dilute H2O2, heat treat your Al foil in a flame before employing, and jump start the reaction in a microwave. For more details on the chemistry, see my latest new thread on Dawsonite.

I have decided to edit the second reaction to make it clearer it is likely an inaccurate expression giving limited guidance as to the max amount of NaOH possibly formed.

Also, Al(OH)3 apparently can react with NaHCO3, especially in the presence of Na2CO3, to form a mixed carbonate hydroxide (there is a Patent on Dawsonite that so claims), which limits the effective aqueous decomposition of heated NaHCO3 here.

[Edited on 29-7-2015 by AJKOER]

unionised - 29-7-2015 at 12:59

Why do you think it matters that washing soda attacks aluminium?
How much of the metal dissolves and how much is converted into a suspension of hydrated Al2O3?

Anyway, while you are rewriting your post you might want to take out the stuff about electrochemistry- it's got nothing much to do with the subject.

AJKOER - 29-7-2015 at 16:28

Unionised:

A quote:

"So the alkaline sodium carbonate in the dishwasher detergent does indeed attack aluminum, at the very least eating deeply enough into the surface to make it dull and pewter-gray with aluminum compounds. For this reason most manufacturers of quality aluminum cookware advise against putting it in the dishwasher."

Link: http://superbeefy.com/why-does-aluminum-cookware-corrode-and...

blogfast25 - 29-7-2015 at 18:39

AJ:

NOBODY disputes sodium carbonate solutions corrode Al objects. But that is not an argument that proves sodium carbonate can dissolve SIGNIFICANT amounts of aluminium. Mild surface corrosion does simply not equate to significant dissolution. Corrosion is also affected by other factors. To equate corrosion so insistently with dissolution as you do is simply silly.

AJKOER - 30-7-2015 at 08:04

Yes, a fair question.

Perhaps looking at the starting to finish pictures (sorry, not the best quality) in my new thread (link: http://www.sciencemadness.org/talk/viewthread.php?tid=63150 ) where to the heat treated Aluminum foil (albeit more reactive) and washing soda, some dilute H2O2 (again better than air) was also added (that is, in place of just boiling aqueous Na2CO3 in an open vessel in air) may still give one cause to be astonished given that it was accomplished in around 10 minutes with only 30 seconds in the microwave!

If the Al foil was more heat treated via a methane flame a little more consistently (that is, made to glow red creating the weak gamma Al2O3 over the Aluminum), I would expect near total destruction in short order. Otherwise, the process is slower for sure, but the ultimate result is the same.

[Edit] Also interesting is the Aluminum-air battery electrochemistry involved (discussed at the top of this page), implying that Na2CO3 acts in one role as an electrolyte, which is a catalytic role only! So as long as there is water and an oxygen source, the Aluminum could continue to be attacked!

[Edited on 30-7-2015 by AJKOER]

unionised - 30-7-2015 at 10:57

Ajkoer,
That piece of aluminium you were looking at as it dissolved.
Do you know that the metal was made by electrolytic reduction of Al2O3 dissolved in molten cryolite?
I guess you did.
Do you know where the Al2O3 comes from?
Most of it is from an ore called bauxite.
i guess you knew that too.
The Bauxite isn't very pure- it has quite a lot of Fe2O3 in it.
You may well have known that too.
Do you know how they separate the Alumina from the iron oxides?
They dissolve the alumina by digesting it with hot NaOH solution then filter off the "red mud" of (mainly) Fe2O3.
Do you know how they get the alumina back again?
There are a couple of ways.
One is to add CO2 to the solution of sodium aluminate.

So, the fact that you can find a bit of cheap aluminium to dissolve is a consequence of the fact that the reaction you are relying on goes backwards.

A multi million dollar industry says one thing and you say the other .
Which one do you think I should believe?