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elementcollector1
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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]
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elementcollector1
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Does anyone know the decomposition temperature for Cr(OH)3 to Cr2O3 and H2O? I suspect around 200-400 C.
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bhattshivamm
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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 !!! ]
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DraconicAcid
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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.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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bhattshivamm
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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]
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blogfast25
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| 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]
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elementcollector1
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Can chromate be reduced to the Cr(III) ion effectively, or does it have to be dichromate?
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DraconicAcid
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Chromate can be reduced; it's just easier to reduce in acidic solution, and in strongly acidic solution, it converts to dichromate.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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bhattshivamm
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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 
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Zyklon-A
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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.
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elementcollector1
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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]
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blogfast25
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| 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+</sup> as
hydroxide from alcohol reduced dichromate. It's possible it still contained sulphate ions but on calcination these are probably driven off.
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Zyklon-A
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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.
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blogfast25
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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]
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Zyklon-A
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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...
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blogfast25
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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.
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Amos
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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?
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blogfast25
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| 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]
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Amos
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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!
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blogfast25
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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)<sub>3</sub>(H<sub>2</sub>O)<sub>3</sub>, 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
M<sup>3+</sup>(H<sub>2</sub>O)<sub>n</sub> (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 CO<sub>2</sub> (decomposition of carbonic acid). When
M<sup>3+</sup>(H<sub>2</sub>O)<sub>n</sub> has lost 3 protons, M(OH)<sub>3</sub> 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]
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Amos
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| 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.
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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. |
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blogfast25
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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)<sub>3</sub> 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]
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woelen
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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.
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Amos
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| 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.
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