Sciencemadness Discussion Board - Precipitation of Cromium(III)

Precipitation of Cromium(III)

bhattshivamm - 28-2-2013 at 07:58

i have potassium dichromate... after reducing with sodium metabisulfite, i got a nice and clear green solution [Cr(III)] as I should. but i am interested in precipitating Cr(III) out of the solution... i tried sodium hydroxide, potassium hydroxide to precipitate Cr(III) as chromium hydroxide, but I couldn't. then i tried with oxalic acid, sodium oxalate and potassium oxalate, but no luck. then i tried to precipitate it as chromium carbonate by sodium carbonate and sodium carbonate, but again, i got nothing out of it. finally i tried to precipitate it in the form of benzoate using saturated solution of sodium benzoate, but.... but.....

i know that chromium tends to form various complexes in solution by reacting with excess OH- and ions like that. but i tried every possible concentrations of the solutions....not a single reaction could precipitate Cr(III).... plz help me out.

Endo - 28-2-2013 at 09:07

Have you tried ammonia?

Cr3+(aq) + 3NH3(aq) + 3H2O(l) <==> Cr(OH)3(s) + 3NH4+(aq)

I believe that if you use an excess amount amount of ammonia the solubility of the precipitate decreases.

elementcollector1 - 28-2-2013 at 09:34

If you use an excess of ammonia, the Cr(OH)3 redissolves into ammonium chromate. You have to use just the right amounts of a weak base, such as Na2CO3 or NaHCO3.

DraconicAcid - 28-2-2013 at 09:46

If you use an excess of ammonia, you will get complex ions such as [Cr(NH₃)₆]³⁺. Add the sodium carbonate and give it time. Chromium(III) compounds are notoriously slow to react with a lot of things.

[Edited on 28-2-2013 by DraconicAcid]

blogfast25 - 28-2-2013 at 14:06

Quote: Originally posted by elementcollector1

If you use an excess of ammonia, the Cr(OH)3 redissolves into ammonium chromate. You have to use just the right amounts of a weak base, such as Na2CO3 or NaHCO3.

That's unlikely: NH3 is too weak an alkali to dissolve Cr(OH)3. But Cr (III) does form an ammonia complex: Cr(NH3)6(3+). It would take strong NH3 and time and a lot of shaking to get Cr(OH)3 to dissolve in it though...

[Edited on 28-2-2013 by blogfast25]

Boffis - 28-2-2013 at 17:38

Did you heat the solution? If you boil chrome alum solution it will turn green and not precipitate with most reagent. It slowly turns purple again in the cold and will the precipitate with NaOH etc. A precipitate will form slowly.

This is why when preparing many Cr3+ complexes it is first reduced to Cr2+ and then allowed to oxidize in the presence of an excess of the new ligand (check out the chromium complex preparation in Brauer (forum library))

bhattshivamm - 1-3-2013 at 01:24

actually, I am running out of ammonia right now, I would try it in a few days...

woelen - 1-3-2013 at 01:35

Do not use sulfite or bisulfite as reductor for dichromate to chromium(III). You get a green solution and this green complex is a remarkably stable anionic sulfato-complex of chromium. Once you have this complex you need to wait for weeks (!!) to get free chromium(III) again.

If you use ethanol as reductor and assure that the liquid remains cold, then you get a purplish/greyish solution of hydrated Cr(3+) ions and from that solution you easily can precipitate chromium(III) with just enough NaOH or with slight excess of ammonia. The reduction with ethanol must be done with dichromate, dissolved in dilute sulphuric acid. If the temperature remains low, then you don't get the green sulfato complex. Even better is to dissolve the dichromate in dilute (e.g. 5%) nitric acid. If you use dilute nitric acid, then you get no complex and the solution becomes purplish/grey even if the liquid becomes somewhat warmer. If the nitric acid is dilute, then the acid itself does not act as oxidizer, it just acts as acid.

blogfast25 - 1-3-2013 at 05:22

Woelen: hydrochloric acid should work, no? In dilute solution of course.

elementcollector1 - 1-3-2013 at 08:46

I'm assuming sulfuric acid works the same way?
That would explain a LOT, actually.
I guess I'll distill my hydrochloric to get rid of the iron contamination, and reduce my chromate with that instead. Then probably boil down to get rid of excess hydrochloric, and precipitate with baking soda...
The only problem with this synth is, filtering has taken weeks due to the sticky, muddy ferric hydroxide I've had to separate out, as well as the fact that I now have 3-4 liters of chromate solution. Fun.

woelen - 1-3-2013 at 13:11

Quote: Originally posted by blogfast25

Woelen: hydrochloric acid should work, no? In dilute solution of course.

No, hydrochloric acid does not work. The chloride does not reduce the dichromate, but when it is present while dischromate is being reduced, then it imemdiately attaches to the formed Cr(3+) and forms a chloro-complex, which again is very stable and once the chloride is attached to the chromium, you have to wait a LONG time before you have free aqueous Cr(3+) again.

Yes, chromium(III) is an interesting, but also quite annoying ion. It coordinates to almost everything very quickly during the short period of its formation, but once it has reached a stable coordination-state, it hardly is willing to exchange ligands. For this reason, synthesis of special chromium(III) complexes usually is not done with chromium(III) salts as starting point, but with hexavalent chromium as starting point, or chromium(II) as starting point.

elementcollector1 - 1-3-2013 at 13:14

What acid does work to obtain Cr(III) ions, then? Remaining possibilities are nitric and possibly oxalic (sulfuric is discounted, as mentioned earlier).

ScienceSquirrel - 1-3-2013 at 13:36

The best approach from a dichromate salt is to dissolve it in sulphuric acid and then add ethanol dropwise keeping the temperature low.
You can use any ethanol source that you like, purple methylated spirits works well, the important thing is to keep it cool, this will stop the sulphate moving in as a ligand.
It is the classic way to make chrome aluim and was a standard lab preaparation to teach technique for years.
Aternatively you can use 30% hydrogen peroxide as the reducing agent, this will run at 10 C and give you a simple uncomplexed chromium III solution.
It goes through a blue chromium peroxide intermediate that forms a complex with and extracts into diethyl ether.
It will easily form brilliant blue needles on evaporation of the ether.
http://en.wikipedia.org/wiki/Chromium_peroxide
The colour changes are nice.
Pale yellow dichromate goes to darker yellow chromate on addition of sulphuric acid, then blue on addition of hdrogen peroxide, oxygen effervesces and the solution turns green.
Take care, do it on a small scale and impress the socks off the next generation of chemists.

woelen - 1-3-2013 at 13:56

Quote: Originally posted by elementcollector1

What acid does work to obtain Cr(III) ions, then? Remaining possibilities are nitric and possibly oxalic (sulfuric is discounted, as mentioned earlier).

I did not discount sulphuric acid, but you have to keep the reaction vessel cool, as ScienceSquirrel also mentions. I discounted sulfite as reductor.

Oxalic acid is totally useless for the purpose of making free aqueous chromium(III). If you use oxalate or oxalic acid then the reduction of dichromate leads to formation of an intenely colored deep purple oxalato complex of chromium(III). A rather interesting experiment is to add oxalic acid to a dilute solution of sodium dichromate or potassium dichromate. The result is a very dark purple solution.

The only acids which can be used for making free aquated chromium(III) ions are:
- dilute sulphuric acid (not above 40 C, otherwise sulfato complex formation)
- dilute nitric acid
- dilute perchloric acid
- dilute tetrafluoroboric acid

The latter two acids are not very common, especially HBF4 is not something you can easily find as a home chemist.

elementcollector1 - 1-3-2013 at 14:09

Dilute sulfuric it is, then. How dilute? <20%?

blogfast25 - 1-3-2013 at 14:52

Quote: Originally posted by elementcollector1

Dilute sulfuric it is, then. How dilute? <20%?

You can work out the amount of H2SO4 needed from the stoichiometric equation to obtain Cr2(SO4)3.

C2H5OH + H2O === > CH3COOH + 4 H+ + 4 e- (oxidation of ethanol)

Cr2O7(2-) + 14 H+ + 6 e- === > 2 Cr3+ + 7 H2O (reduction of dichromate)

Now balance and complete! (5 marks)

elementcollector1 - 1-3-2013 at 15:45

Gah, I hate these. I much prefer full reactions, anyway. Net ionic and half-reactions just don't click for me, full reactions do. (In fact, on our recent chemistry test, this was the only type of question I failed. Fortunately, there was only one of them.)
But enough whining, let's give this a shot...
7 C2H5OH + 7 H2O -> CH3COOH + 28H+ + 28e-
2 Cr2O7(2-) + 28 H+ + 12 e- -> 4 Cr(3+) + 14H2O
Those darn electrons.
3 C2H5OH + 3 H2O -> 3 CH3COOH + 12H+ + 12e-
2 Cr2O7(2-) + 28 H+ + 12e- -> 4 Cr(3+) + 14 H2O
Those darn hydrogen ions.
...
Well, I'll just use regular stoichiometry, then. Replaced ethanol with isopropanol for availability issues.
Na2Cr2O7 + 4 H2SO4 + 3 C3H8O -> 3 C3H6O + Cr2(SO4)3 + Na2SO4 + 7 H2O
Huh. This looks familiar. A Cr2O7(2-) ion, 4 H+'s, 7 H2O's...

ScienceSquirrel - 1-3-2013 at 18:28

There is a preparation here

http://books.google.com/books?id=v5xLjrMEZ1QC&pg=PA208&a...

blogfast25 - 2-3-2013 at 12:39

Quote: Originally posted by elementcollector1

Gah, I hate these.

[snip]

Huh. This looks familiar. A Cr2O7(2-) ion, 4 H+'s, 7 H2O's...

I'm too lazy to fully check it but your result looks perfect to me, on the face of it.

elementcollector1 - 2-3-2013 at 19:30

I tried precipitation of Cr(OH)3 using a pure potassium alum that I got recently. I got an initial precipitate with KOH, but it seemed to 'disappear'! The solution was initially purple, but switched to green over time. Addition of baking soda caused two layers to form, with the extremely dark chromium layer on bottom and the baking soda on top, with a purple-seeming precipitate in between. Odd... And it doesn't bode well for my chromate, as if I can't get the pure stuff to do what I want...

S.C. Wack - 2-3-2013 at 19:46

If the green chromite solution is boiled, there is hydrolysis or something, and hydroxide reprecipitates.

elementcollector1 - 2-3-2013 at 20:15

Seriously? I'll give that a go, then. Unfortunately, it's late. Will post results and possibly pictures tomorrow.

blogfast25 - 3-3-2013 at 05:48

Quote: Originally posted by elementcollector1

Seriously? I'll give that a go, then. Unfortunately, it's late. Will post results and possibly pictures tomorrow.

Beware of temperature, as SS mentioned. Above a certain temperature stable Cr(III)-sulphato complexes form that do no behave like 'naked' Cr3+. It happened to me. The sulphato complex can take months to disintegrate back to 'normal' Cr3+, from which Cr(OH)3 can be precipitated.

That preparation referred to by Squirrel looks very sound.

[Edited on 3-3-2013 by blogfast25]

AJKOER - 3-3-2013 at 08:29

Here is a good internet available source, Atomistry (link: http://chromium.atomistry.com/chromic_hydroxide.html ). To quote some interesting parts:

"Chromic Hydroxide, Cr2O3.Aq., is obtained by precipitation of a solution of a chromic salt by means of potassium hydroxide; if excess is used, the precipitate dissolves, forming a green solution from which the hydroxide is again precipitated on keeping or boiling. Such precipitate retains alkali which cannot be removed by repeated washing with hot water; it is usual, therefore, to employ ammonium hydroxide as the precipitating agent. Even in this case, excess of ammonia dissolves chromic hydroxide, yielding a reddish-violet solution: methyl-amine behaves similarly, but di- and tri-methylamine at once precipitate chromic hydroxide completely. When freshly formed, the precipitate appears to be a well-defined chemical compound, the solubility product of which, according to Bjerrum, is 4.2×10-16 at 0° C. and 54×10-16 at 17° C. in 0.0001 molar units. According to Weiser, however, the precipitate obtained by the addition of alkali to solutions of chromic salts does not contain any definite hydrate. In the cold, the freshly formed precipitate readily dissolves in acids, but becomes insoluble on keeping or heating; between these two extremes of solubility an indefinite number of hydrous oxides exists. By precipitating at temperatures ranging from 0° C. to 225° C., products have been obtained varying in colour from greyish blue to bright green. Since most of the chromic salts exist in two distinct modifications, the violet and the green, it has long been assumed that there must be two isomeric chromic hydroxides corresponding to these two series of salts. Such isomerides have not, however, been isolated. The properties of chromic hydroxide vary considerably with age, especially as regards solubility in acids and alkalies, and the "ageing" is accelerated if the precipitate is allowed to remain under alkaline solution; the rate appears to increase with hydroxyl-ion concentration and also with increase of temperature, and the change in properties appears to be due to change in the size of the particles."

On Colloidal Chromium Hydroxide, to quote:

"The hydrosol is obtained as a deep green solution by the peptisation of the hydroxide by means of chromic chloride, or by a solution of copper oxide in ammonia. As already stated, the freshly precipitated hydroxide forms an apparently clear green solution with excess of an alkali hydroxide. That the chromic hydroxide is peptised and not dissolved is shown by the fact that it can be completely filtered out by means of a collodion filter, leaving a colourless filtrate. The colloidal solution is stable while hot, but slowly yields a gel on keeping at ordinary temperatures."

And also, "When potassium hydroxide is added to solutions containing ferric chloride and chromium sulphate in varying proportions, the iron is not precipitated so long as the chromium is present in excess."....

Of possible value note that if the FeCl3 is in excess, it is also precipitated along with any adsorbed chromium.

Interestingly, both positively and negatively charged colloids have been prepared. Positive charged when the hydrated oxide is peptised with chromic chloride, or by hydrolysis of the chloride or nitrate. The negatively charged has been created by peptising the hydrous oxide with sodium or potassium hydroxide.

Now, with respect to Chromic oxide jellies, they can be created, to quote "by adding sodium or potassium hydroxide or ammonia to a solution of chromic sulphate or chloride containing sodium acetate; or by adding sodium or potassium hydroxide, but not ammonia, to a solution of chrome alum. The jelly is violet if prepared by the addition of ammonia or of a slight excess of the alkali metal hydroxide; if the latter is added in larger quantity the jelly is green. The jellies dissolve in hydrochloric acid, but re-form on neutralising the solution if sufficient sodium acetate is present."
-------------------------------------------------------------

My take on this with respect to the preparation and use of Cr(OH)3:

1. Precipitate from a chromic salt by means ammonium hydroxide (avoid KOH), but not an excess of ammonia as it dissolves chromic hydroxide, yielding a reddish-violet solution.

2. Dissolve the freshly formed precipitate in an acid, but employ as soon as possible as the Cr(OH)3 becomes insoluble on keeping or heating.

For a hydrosol (a deep green solution) can be prepared by the peptisation of Cr(OH)3 with a solution of copper oxide in ammonia. Note, a cited creation of a colloidal suspensions of copper oxide (CuO) nanoparticles is via an alcothermal method with reaction of copper acetate and sodium hydroxide in the presence of acetic acid in ethanol at 78°C (see http://onlinelibrary.wiley.com/doi/10.1111/j.1551-2916.2006.... ). Perhaps substituting aqueous ammonia (not in excess) for NaOH, resulting in a larger nanoparticle, may be acceptable.

[Edited on 3-3-2013 by AJKOER]

blogfast25 - 3-3-2013 at 09:30

"Even in this case, excess of ammonia dissolves chromic hydroxide, yielding a reddish-violet solution: methyl-amine behaves similarly, but di- and tri-methylamine at once precipitate chromic hydroxide completely."

The 'reddish-violet' is the hexa-ammonia complex. Obviously the methyl amines do not form such complexes.

elementcollector1 - 3-3-2013 at 12:41

I'm happy to report a success with boiling: viridian green Cr(OH)3 has precipitated out, leaving a pale green solution behind (could probably boil that to get a clear solution and a little more precipitate).

EDIT: Viridian is the actual name for the pigment derived from Cr(OH)3. Ironic.

[Edited on 4-3-2013 by elementcollector1]

elementcollector1 - 3-3-2013 at 22:46

Does anyone know the decomposition temperature for Cr(OH)3 to Cr2O3 and H2O? I suspect around 200-400 C.

bhattshivamm - 4-3-2013 at 09:25

As woelen says, Cr(III) forms a stable sulfato-complex which then bothers in precipitating Cr(OH)3... so what about reducing K2Cr2O3 by nitrite ???
does that work ? or it also forms any sort of 'stable' complex with it ??
[ i can't use ethanol to reduce dichromate because i can't buy or get ethanol in my state legally (Gujarat,India). even no alcoholic drinks. no way to get ethanol anyhow !!! ]

DraconicAcid - 4-3-2013 at 09:51

Quote: Originally posted by bhattshivamm

As woelen says, Cr(III) forms a stable sulfato-complex which then bothers in precipitating Cr(OH)3... so what about reducing K2Cr2O3 by nitrite ???
does that work ? or it also forms any sort of 'stable' complex with it ??
[ i can't use ethanol to reduce dichromate because i can't buy or get ethanol in my state legally (Gujarat,India). even no alcoholic drinks. no way to get ethanol anyhow !!! ]

Can you get methanol or formaldehyde? Those will also work.

bhattshivamm - 4-3-2013 at 23:30

yeah, i have 250 ml methanol, but its quite expensive here...
can i use NaNO₂ ? will Cr(OH)₃ precipitate after reducing Cr(VI) with nitrite ? other wise i'll have to use methanol which i don't want to use...

[Edited on 5-3-2013 by bhattshivamm]

blogfast25 - 5-3-2013 at 04:54

Quote: Originally posted by bhattshivamm

yeah, i have 250 ml methanol, but its quite expensive here...
can i use NaNO₂ ? will Cr(OH)₃ precipitate after reducing Cr(VI) with nitrite ? other wise i'll have to use methanol which i don't want to use...

[Edited on 5-3-2013 by bhattshivamm]

Try 'methylated spirits' aka 'denaturated alcohol'. Cheap as chips.

Assuming even it works, nitrite will be far more expensive...

[Edited on 5-3-2013 by blogfast25]

elementcollector1 - 5-3-2013 at 09:54

Can chromate be reduced to the Cr(III) ion effectively, or does it have to be dichromate?

DraconicAcid - 5-3-2013 at 09:56

Quote: Originally posted by elementcollector1

Can chromate be reduced to the Cr(III) ion effectively, or does it have to be dichromate?

Chromate can be reduced; it's just easier to reduce in acidic solution, and in strongly acidic solution, it converts to dichromate.

bhattshivamm - 6-3-2013 at 20:24

I reduced dichromate by nitrite (thanks to woelen)... i got nice slurry of Cr(OH)3 after reacting Cr(III) with aq. NaOH... but I couldn't post its images as the memory-card of my camera is lost

Zyklon-A - 21-4-2014 at 21:25

I want to make some chromium (III) oxide today, elementcollector1 said isopropanol can be used:

Quote: Originally posted by elementcollector1

[...]Replaced ethanol with isopropanol for availability issues.
Na2Cr2O7 + 4 H2SO4 + 3 C3H8O -> 3 C3H6O + Cr2(SO4)3 + Na2SO4 + 7 H2O
Huh. This looks familiar. A Cr2O7(2-) ion, 4 H+'s, 7 H2O's...

Anyone know if this works? He just balanced the equation, didn't provide any references.

elementcollector1 - 21-4-2014 at 21:42

It should work, by all means - and yet, chromium(III) laughs at "should".
I'll warn you - Cr(OH)₃ is easily one of the trickiest substances I've ever encountered. You'll think you've got Cr³⁺, and your base is ready to go, when in reality... the chromium has already complexed itself away. I wanted to do this from stainless steel, but I ended up caving in and using pure alum. Even then, it was hard.

[Edited on 4-22-2014 by elementcollector1]

blogfast25 - 22-4-2014 at 04:11

Quote: Originally posted by elementcollector1

[...] the chromium has already complexed itself away. I wanted to do this from stainless steel, but I ended up caving in and using pure alum. Even then, it was hard.

[Edited on 4-22-2014 by elementcollector1]

The sulphato complex appears to be the weaker of the series, from what I've done. I didn't have any problem precipitating Cr<sup?3+ as hydroxide from alcohol reduced dichromate. It's possible it still contained sulphate ions but on calcination these are probably driven off.

Zyklon-A - 22-4-2014 at 05:51

My dichromate is pure, not from SS, I bought it. The end product is elemental chromium for my collection. I'll reduce dichromate, precipitate hydroxide, decompose to Cr₂O₃ then thermite with aluminum powder for elemental chromium. I'll try out isopropanol first, then methanol or ethanol next.

blogfast25 - 22-4-2014 at 13:04

Zb:

Quite a laborious and expensive way to prepare Cr (III) oxide, considering how cheap the pottery version is....

[Edited on 22-4-2014 by blogfast25]

Zyklon-A - 22-4-2014 at 13:42

True, but I don't have any money at the moment, I bought the dichromate because it's awesome, I got a good deal, decent source of chromium, and I wan'ted to to grow some crystals of it. I have over 300 grams and using enough to make several grams of chromium will be easily worth it. But yeah, it's more difficult to precipitate than I thought...

blogfast25 - 23-4-2014 at 08:14

Quote: Originally posted by Zyklonb

But yeah, it's more difficult to precipitate than I thought...

I don't think it's that difficult. Make sure the reduction is complete, use only small excesses of reagents and leave to stand overnight to complete reaction.

Then add the right amount of conc. NH3 to neutralise everything only just (excess NH3 will also complex the Cr (III) - if you use NaOH, excess NaOH can form chromite). Filter and wash profusely. Calcine the Cr(OH)3 to as high as you can go.

Amos - 17-4-2015 at 17:27

So tonight I decided to try preparing chromium(III) sulfate by reducing sodium dichromate with sulfuric acid and ethanol. I used excesses of both sulfuric acid and ethanol, and kept the reaction flask very cold in an ice bath while adding ethanol dropwise. The solution never got warm enough for acetaldehyde produced by the oxidation of ethanol to escape the flask, and the end product was a midnight blue solution with some hints of green, but nowhere near as green as the sulfato-complex I have produced previously by aggressively boiling a slurry of chromium(III) oxide in sulfuric acid.

I then added the solution, still cold, to a chilled solution of saturated sodium bicarbonate, and obtained a precipitate with a handsome muted blue-violet coloration. I was expecting this to be chromium(III) hydroxide, as I've previously been told (I believe by blogfast25, but I can't be sure) that chromium(III) carbonate is unstable and quickly gives off carbon dioxide, converting to the hydroxide.

The problem is, everything I seem to encounter (the wikipedia page, google images, and this page) seem to indicate that chromium hydroxide is a green color. Do I have hydroxide, a basic sulfate, or something else entirely? Can anyone shed some light on what this precipitate is?

blogfast25 - 17-4-2015 at 18:12

Quote: Originally posted by Amos

Do I have hydroxide, a basic sulfate, or something else entirely? Can anyone shed some light on what this precipitate is?

It's likely to be the hydroxide, colour can very acc. precipitation conditions. What I precipitated I wouldn't have described as green either.

Check for sulphate by washing the precipitate very carefully, then dissolve a portion in sulphate-free HCl (or any other acid). Test that solution for sulphate with Ba chloride or nitrate solution.

Also, try and calcine a bit of the washed precipitate, see what colour you get.

[Edited on 18-4-2015 by blogfast25]

Amos - 17-4-2015 at 18:58

blogfast25,

Since woelen mentioned earlier in the thread that hydrochloric acid immediately forms a chloro complex with chromium(III), I elected to dissolve me precipitate in dilute nitric acid, which took some time but eventually gave me a deep blue(but again, slightly green) solution. The addition of aqueous calcium nitrate gave no precipitate.

Sounds like I may be able to make chromium(III) sulfate after all!

blogfast25 - 18-4-2015 at 05:54

Amos:

The green Cr(III) sulphato complex you mentioned only seems to form at higher temperature. With your ‘cold’ reduction reaction you probably obtained the blue Cr(III) hexaaqua ion. With a base that gives blueish Cr(OH)3(H2O)3, in accordance with (scroll down a bit):

http://www.chemguide.co.uk/inorganic/transition/chromium.htm...

This also means that if you want to prepare some Cr(III) sulphate, it's advisable to dissolve the Cr(III) hydroxide in COLD sulphuric acid, which should yield a blue solution. Boiling this will cause the colour change to green, due to the sulphato complex. Unfortunately it will revert back to blue only very, very slowly (weeks, rather than days).

****************

What you wrote above with regard to ‘Cr(III) carbonate’ is not entirely correct:

Quote: Originally posted by Amos

I then added the solution, still cold, to a chilled solution of saturated sodium bicarbonate, and obtained a precipitate with a handsome muted blue-violet coloration. I was expecting this to be chromium(III) hydroxide, as I've previously been told (I believe by blogfast25, but I can't be sure) that chromium(III) carbonate is unstable and quickly gives off carbon dioxide, converting to the hydroxide.

In fact the Cr(III) carbonate simply never forms. The strong central electrical field of M3+(H2O)n (M = Al, Fe, Cr…) causes these cations to be (weakly) acidic through deprotonation of the coordinated water molecules (three possible, subsequent deprotonations). The resulting oxonium ions neutralise the bicarbonate ions with release of CO2 (decomposition of carbonic acid). When M3+(H2O)n has lost 3 protons, M(OH)3 is formed. No ‘M carbonate’ is ever on the horizon.

This theory explains also why ferric solutions require low pH (3 - 4, depending on ferric concentration) to avoid hydrolysis: high oxonium concentration suppresses the deprotonations.

[Edited on 18-4-2015 by blogfast25]

Amos - 18-4-2015 at 07:52

Quote:

Quote: Originally posted by blogfast25

Amos:

This also means that if you want to prepare some Cr(III) sulphate, it's advisable to dissolve the Cr(III) hydroxide in COLD sulphuric acid, which should yield a blue solution.

That's precisely what I had planned, I'm glad to get the confirmation.

And about the chromium(III) carbonate that doesn't exist, I remember you telling me that it never formed now; I had just been going off memory and misquoted you. The explanation for why is quite cool, actually. It seems the more chemistry I learn, the less I know.

blogfast25 - 18-4-2015 at 07:58

Quote: Originally posted by Amos

It seems the more chemistry I learn, the less I know.

Wise men know just how little they know.

Sadly, ignorant barstools on the other hand think they know it all...

Incidentally:

Quote: Originally posted by Amos

[…], I elected to dissolve me precipitate in dilute nitric acid, which took some time but eventually gave me a deep blue (but again, slightly green) solution.

… was interesting and the right choice of acid. Nitrates are the least likely ligands, so the blue colour obtained is essentially the ‘true’ colour of the Cr(III) hexaaqua cation. No great surprise then that Cr(OH)3 precipitated from Cr(III) nitrate is also broadly blue.

***********************

As regards the ‘true colour’ of Cr(III) hydroxide I wonder, very tentatively, whether different colours can be obtained due to often overlooked isomers of Cr(III) complexes, see Wikipedia on that subject:

http://en.wikipedia.org/wiki/Chromium#Chromium.28III.29

Quote:

Chromium(III) ions tend to form octahedral complexes. The colors of these complexes is determined by the ligands attached to the Cr center. The commercially available chromium(III) chloride hydrate is the dark green complex [CrCl2(H2O)4]Cl. Closely related compounds have different colors: pale green [CrCl(H2O)5]Cl2 and the violet [Cr(H2O)6]Cl3. If water-free green chromium(III) chloride is dissolved in water then the green solution turns violet after some time, due to the substitution of water by chloride in the inner coordination sphere. This kind of reaction is also observed with solutions of chrome alum and other water-soluble chromium(III) salts.

Similar isomers exist also with other Cr(III) complexes.

Could ‘Cr(OH)3.nH2O’ vary in colour, depending on what kind of isomer it was precipitated from?

[Edited on 18-4-2015 by blogfast25]

woelen - 20-4-2015 at 12:35

I made Cr(OH)3 from purplish/blue/grey solutions of aqueous chromium(III) and this indeed cannot be described as green. I would say that the precipitate is grey with a blue/green tinge. The color also depends somewhat on the type of lighting used. I find it grey with a hue, depending on type of light. The hue can vary from dull green/blue to dull blue/violet, but the basic color I would describe as grey.

Amos - 20-4-2015 at 17:30

Quote: Originally posted by woelen

I made Cr(OH)3 from purplish/blue/grey solutions of aqueous chromium(III) and this indeed cannot be described as green. I would say that the precipitate is grey with a blue/green tinge. The color also depends somewhat on the type of lighting used. I find it grey with a hue, depending on type of light. The hue can vary from dull green/blue to dull blue/violet, but the basic color I would describe as grey.

The lighting definitely does seem to effect it. Also, I realized why I kept thinking my Cr(III) solutions looked so different than the images I found online. It seems that once again, as happens to a lot of highly saturated colors but especially to blue-green ones, digital cameras interpret the color much differently than does the human eye. Looking at the solution through my iPhone gives me a more decidedly blue image, sometimes with a bit of violet.