Darkblade48
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
Reactions of dichromates
Normally, I wouldn't play around with (hexavalent) dichromates due to their carcinogenity, but I was curious as to the various colours chromium (and
other transition metals) have in different conditions.
Firstly, I added some potassium dichromate to a solution of sodium sulfite, and the solution went from a bright orange colour to a dull green colour
(assuming the +6 to the +3 state, which would explain the bright green colour).
However, is the colour simply due to the Cr+3 ions in water? Or are they perhaps complexed with the water molecules to give this green colour?
Is there any possibility that I may have formed chromium (III) oxide (Cr2O3?)
|
|
|
Mephisto
Chemicus Diabolicus
Posts: 294
Registered: 24-8-2002
Location: Germany
Member Is Offline
Mood: swinging
|
|
You've produced Cr(III) as hydrate complex. I think the heat of the redox reaction might cause the formation of a green hydrate isomer. This isomerism
of Cr(III)-hydrate complexes is known from chrome alum solutions, which are turning from violet to green when heated above 70 °C. It takes some days
until the violet hydrate isomer is formed again. When using sulfuric acid and ethanol instead of the sodium sulfite for reducing Cr(VI) to Cr(III),
you would get pure chrome alum (chromic potassium sulfate) after removing the aldehyde.
~Mephisto
Edit: No, I don't think there was any chromium(III) oxide formed.
[Edited on 11-12-2005 by Mephisto]
|
|
|
Darkblade48
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by Mephisto
When using sulfuric acid and ethanol instead of the sodium sulfite for reducing Cr(VI) to Cr(III), you would get pure chrome alum (chromic potassium
sulfate) after removing the aldehyde.
|
So in this case, using an acidified solution of ethanol would reduce the Cr(VI) to Cr(III), and I would get the pure chrome alum (which is supposed to
be violet in colour?). Also you mention an aldehyde, would the OH group not be oxidized completely to the COOH functional group? I know that Jones
reagent (CrO3) has the Cr in the +6 state, and this seems to oxidize the OH groups to COOH groups.
| Quote: | Originally posted by Mephisto
Edit: No, I don't think there was any chromium(III) oxide formed.
|
On a totally different note, is it even remotely possible to get chromic oxide from the (potassium) dichromate, or is it only possible through an
ammonium dichromate (thermal) decomposition?
|
|
|
Mephisto
Chemicus Diabolicus
Posts: 294
Registered: 24-8-2002
Location: Germany
Member Is Offline
Mood: swinging
|
|
| Quote: | Originally posted by Darkblade48
Also you mention an aldehyde, would the OH group not be oxidized completely to the COOH functional group? I know that Jones reagent (CrO3) has the Cr
in the +6 state, and this seems to oxidize the OH groups to COOH groups.
|
You're right. I've read the thing with the aldehyde in a book, in which it was the goal to produce chrome alum. Therefore the writer didn't care about
the further oxidation of the acetaldehyde to acetic acid. The formed chrome alum (which crystallize from that solution) would be pure enough with only
little traces of the by-products in its crystalline structure.
If you want do oxidise a primary alcohol just to the aldehyde, the Jones reagent isn't selective enough. Except for the case, in which the formed
aldehyde is immediately distilled off from the reaction flask, to prevent a further oxidation.
| Quote: | Originally posted by Darkblade48
On a totally different note, is it even remotely possible to get chromic oxide from the (potassium) dichromate, or is it only possible through an
ammonium dichromate (thermal) decomposition?
|
Yes, potassium dichromate decomposes at 400 °C to chromium(III) oxide, potassium chromate and oxygen.
4 K2Cr2O7 → 2 Cr2O3 + 4 K2CrO4 + 3 O2
Additionally the Merck Index says, that chromium(III) oxide can be formed by the reaction of sodium dichromate or chromate with sulfur.
~Mephisto
|
|
|
Darkblade48
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
Ah, an interesting reaction I noted today. Mixing a solution of potassium dichromate with a solution of oxalic acid yielded (what I believe to be) a
reduction of the Cr(VI) to a Cr(III) state, as the solution went from a bright orange to a dark, muddy brown colour.
However, I noticed the evolution of gas, would this be O2?
|
|
|
Mephisto
Chemicus Diabolicus
Posts: 294
Registered: 24-8-2002
Location: Germany
Member Is Offline
Mood: swinging
|
|
With oxalic acid and potassium dichromate you get an tris(oxalato)chromate(III) complex. The evolved gas is CO2. With additional KOH you will get
tripotassium trioxalatochromate(III), which is a nice substance for crystal growing. I tried this myself some years ago.
~Mephisto
|
|
|
Darkblade48
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
Most interesting. Is it possible to use a base other than KOH (i.e. NaOH) and get the respective sodium salt instead?
If not, how would I go about making the tripotassium trioxalatochromate (III) solution for crystal growing? Do I simply need to mix in the oxalic acid
with the dichromate and wait for the reaction to occur before adding in the KOH? I assume then allowing the solution to stand and evaporate will allow
crystals to form?
|
|
|
Mephisto
Chemicus Diabolicus
Posts: 294
Registered: 24-8-2002
Location: Germany
Member Is Offline
Mood: swinging
|
|
| Quote: | Originally posted by Darkblade48
Most interesting. Is it possible to use a base other than KOH (i.e. NaOH) and get the respective sodium salt instead?
|
So you would get a double salt like KNa2[Cr(C2O4)3]3 with 1 K+ from the potassium dichromate and 2 Na+ from NaOH. Okay, but this may influence the
crystall shape in an unwanted manner. Maybe you are able to find simply potassium carbonate instead of KOH.
You can calculate the needed amounts from the following equitation:
4 KOH + K2Cr207 + 9 C2O4H2 · 2 H20 → 2 K3[Cr(C204)3]3 · 3 H2O + 6 CO2 + 23 H2O
The preparation is very simple. I've wrote down a detailed description in German language on www.LambdaSyn.com/synfiles/kaliumtrioxalatochromat.htm. Here is a quick translation:
284 g potassium carbonate or an equivalent quantity of potassium hydroxide are dissolved in 700-1000 ml water. 1165 g oxalic acid dihydrate are added
slowly in small portions to the solution (CO2-formation ...). The solution is heated up to 80 °C. Under strong agitating 302 g potassium dichromate
are added in smallest portions. The heating source is meanwhile switched off. It is important to add the potassium dichromate slowly and in small
quantities, as the carbon dioxide evolvement can be very violent. This reaction is exothermic, too. After all potassium dichromate was added, the
solution is boiled for one hour. It is hot filtered into container and left there for one week at room temperature. After this time large black
crystals of tripotassium trioxalatochromate(III) have formed.
The fine spray, which escapes during the CO2-formation contains chromium. The preparation should be therefore done outside or in a fume hood.
~Mephisto
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
There is still more
1. I read on a webpage about synthesizing of pigments that you can create a hydrated chromium oxide from potassium dichromate and boric acid. You melt
them together at red heat for an hour or so then extract the potassium borate with boiling water.
2. A process simular to the oxalate complex is reduction of the dichromate with six moles of thiocyanate to form the crystalline Potassium
Trithiocyanchromate.
3. You can neutralize with KOH to get the Chromate and add strontium chloride and ammonia to precipitate strontium chromate. Strontium Chromate is
slightly soluble in boiling water so you can make a suspension of the strontium chromate in the equivalent amount of sodium bisulfate to get sodium
dichromate and strontium sulfate.
Fellow molecular manipulator
|
|
|
olmpiad
Harmless
Posts: 29
Registered: 2-6-2006
Location: this thing known as earth
Member Is Offline
Mood: Dandy
|
|
Could the procedure mentioned by Mephisto be done with the Na salts rather than K, so Na2Cr2O7, NaOH, and Oxalic Acid, to make Sodium
Tris(Oxalato)Chromate (III)?
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Most likely it will work, but I doubt that you get nice crystalline results. Frequently, the sodium salts do not crystallize well and are very
hygroscopic. So, isolation of the salt may be a real pain.
|
|
|
olmpiad
Harmless
Posts: 29
Registered: 2-6-2006
Location: this thing known as earth
Member Is Offline
Mood: Dandy
|
|
Actually, I tried this myself last night, It formed crystals overnight, and for once, they are not hygroscopic!
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Are you sure these crystals are of the complex? I also did experiments with Na-salts and oxalato complexes, but I usually ended up with crystals of
sodium oxalate, which were colored, due to impurities of the transition metals. Sodium oxalate is only sparingly soluble, IIRC 3 gram / 100 ml.
|
|
|
olmpiad
Harmless
Posts: 29
Registered: 2-6-2006
Location: this thing known as earth
Member Is Offline
Mood: Dandy
|
|
This dissolves fairly well in water, I cannot say the exact solubility yet as I am waiting to let crystals finish drying. The crystals are just
extremely dark, I do not think that they are Na oxalate. I will post a picture of them when I am finished however, and let you be the judge, as you
seem to have done quite a lot with chromates and dichromates,
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Yes, if you have some pictures, that would be nice. And of course, I also can learn from this 
It is my experience that Na-salts usually are much more hygroscopic than K-salts, so if this is a nice exception, then I am always willing to adjust
my ideas about this. I'll wait for your pics.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
