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blogfast25
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I wonder to try the reduction of Cr (III) to Cr (II) with the more powerful reducing agent Al powder: Al (0) === > Al<sup>3+</sup> + 3e
… E<sub>ox</sub> = + 1.662 V. The Cr (III) solution could be buffered to pH ≈ 5 with acetate buffer, to ‘eliminate’ H3O+ as a
player?
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DerAlte
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@ blogfast
I have carefully reviewed my calculations and am using new numbers from CRC rather than the earlier ones – they are not much different. Eo values
vary according to source.
My conclusion is that you erred in your value of K – it’s much worse, a lot higher!
Incidentally as to your earlier point about entropy – all we bother about is the change in energy dG and the relative potentials Eo, all at STP 1N
etc. There is no change of state – all is in the aqueous environment.
OK, since you are a dG man and I am a dE man by virtue of being (once) an EE, let’s work it both ways, for standard conditions. First, repeating the
equations,
Cr3+ + e− → Cr2+ dEo = - 0.407 V …. (1) n=1 electron transfer
Zn(s) -- > Zn++(aq) + 2 e- dEo = 0.7618 V …..(2) n=2
Add 2x(1) to 1x(2) to get rid of electrons
2Cr+++(aq) + Zn(s) < --> 2Cr++(aq) + Zn++(aq) d Eo = dE(3) …(3)
dG(1) = -nF.dE(1) = 1* 96500*0.407 = 39276 j/mol
dG(2) = 2*96500*0.7618 = 147027 j/mol
So dG(3) for reaction (3) as written is 147027-2*39276 = 68475 j/mol
If we want dEo(3), it is 74495/(96500*2) for a 2-electron transfer = 0.3548 V, a cross check that I haven’t screwed up yet.
K for reaction (3) is then exp(nF*dE/RT) = exp(dG/RT) = e^27.66 = 1.03E12 – far larger than the 30 you suggested, which only makes our problem
worse. While not explosive, this is a pretty brisk reaction.
Now, let me throw you a curly one:
The overvoltage of hydrogen on zinc is around 0.6 – 0.7 volts, whether acid or alkali.
.
And next a googly: this is often explained as either a hydride production, ZnH2 (see WIKI on this) or due to occluded hydrogen. In other words, zinc
is ‘passivated’ by H2.
A personal experience is that the zinc becomes dull in acid or a Cr+++ solution. Taking blackened zinc from a Zn/acid reduction, washing well
and reacting with acid shows no sign of Chromium with H2O2 (perchromate test, blue coloration, very sensitive.) The black parts react slowly. Hence
the blackening is not Cr metal (very unlikely) but a surface condition of the Zn, possible hydride.
Hence Bezaleel’s comment above looks a possible candidate. This explanation was popular in the 1930s.
Added, on seeing your last posting: Al would definitely be overkill - it might even produce Cr metal.
Regards,
Der Alte
[Edited on 23-12-2011 by DerAlte]
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blogfast25
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Der Alte:
Ooops, my error was not inverting the lnK (= 27.7) to get K = e<sup>27.7</sup>, an exceedingly silly mistake. Dunce cap on! 
Entropy? I didn’t mention it. Only that G is a State Function (see Hess) and that ΔG isn’t therefore affected by any catalysis (which only
affects the reaction path, not the end state). Rate: yes, overall ΔG: NO!
The blackening of metals when acids react on them could just as well be a micro coarsening of the ‘smooth’ surface. Most metals show this
behaviour. Fairly weak evidence for a hydride, IMHO.
You may be right about Al. I’ve got Ered = - 0.744 V for Cr<sup>3+</sup> + 3e === > Cr (0), so + 1.662 V covers that amply.
I’ll give that a go tomorrow.
Don’t ask me about over potentials: I really leave that to the dedicated folk at the EC department! 
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DerAlte
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Agree about the black surface being possibly pitted. As I say, it's just a strawman - by all means burn! It's most unlikely we can solve the mystery
here. I would like to hear of any super modern explanations if they exist.
Regards
Der Alte
[Edited on 23-12-2011 by DerAlte]
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blogfast25
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Truest words written on the last couple of posts. I'm now favouring the (old) statu nascendi thingy:
Zn + 2H3O+ === > Zn2+ 2 H* + 2H2O
2 x [Cr3+ + H* === > Cr2+ + H+]
2 x [H+ + H2O === > H3O+]
============================
Zn + 2 Cr3+ === > Zn2+ + 2 Cr2+
Appealing because H3O+/H* would be a catalytic system, not affecting Delta G or Ecell, but overcoming some unknown kinetic barrier of the 'naked'
(overall) reaction. Hydrogen would evolve because the catalytic reaction would compete with 2 H* === > H2 (H* a hydrogen radical).
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DerAlte
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Let me lob a couple of thoughts your way.
In the course of the Cr experiments described in this thread, I also tried the effect of making a Zn/Cu couple instead of straight Zn/acid. The
hydrogen then comes off the Cu as is well known. Hence here H2 and Zn are separated. The couple reduced the Cr+++ the same way as Zn alone, but what
did it, the H, H*, H2 or the Zn? I suppose you could also ask what electrolysis does, a very good reducing agent that keeps the Cr++ clean of Zn. Have
done that also.
This prompts the idea of a cell with a barrier to answer this question, connected electrically and by salt bridge. Cr+++ ions in both cells, and H+
ions in both, using Zn and Cu. Got to mull that one over a bit. In which cell – or both – would reduction take place?
Tried Al today in the form of turnings with dilute HCl (~15%). The oxide passivation does not allow this to start easily so I usually add a bit of Cu
foil (which reminded me of above). However, once contact is made the reaction takes off with a vengeance. And of course, reduces the Cr+++. When the
acid is done the Al has a thick coat of a black substance. Cr metal? TBD…
Hoist a jar with me this Xmas Eve,
Der Alte
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blogfast25
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Der Alte:
In the case of the Cu/Zn couple, you could check for Zn2+ post reaction, to exclude some things. Ditto Cu2+. Strictly speaking, for the ‘old’
proposed catalysis to work (based on H*) contact between the Zn and H isn’t required I’d imagine.
The black stuff on the Al must be Cr metal: did the solution clear up fully? Dissolve the Al/Cr (presumed) in acid to check for Cr, the usual way. But
with a cell potential for Cr3+/Al as noted above the reduction’s gotta go all the way down to Cr (0), as you suggested.
Can we think of a metal that would generate H2 (and thus H* intermediately) but isn't capable of reducing Cr (III) to Cr (0)? Tin for instance: Sn
(0) === > Sn2+ + 2e, Eox = + 0.1375 V doesn’t cover the Cr (III) to (II) reduction. Tin metal doesn’t react very enthusiastically with strong
HCl but it does fairly slowly and it is a method for synthing SnCl2 hydrate. Hottish it’s fairly brisk: first hand experience. Tin powder should be
reasonable fast…
Mine’s a pint of Old Speckled! Merry X-mas. (Dare I say Happy Winterval? ;-) )
[Edited on 25-12-2011 by blogfast25]
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Bezaleel
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Blogfast25, your suggestion with Sn is indeed a great test. I doubt it will indeed reduce the Cr+++, though. (I don't believe H* to be sufficiently
stable to get to the Cr+++.) I'm much looking forward to your experimental outcome therefore.
Quote: Originally posted by DerAlte  | Tried Al today in the form of turnings with dilute HCl (~15%). The oxide passivation does not allow this to start easily so I usually add a bit of Cu
foil (which reminded me of above). However, once contact is made the reaction takes off with a vengeance. And of course, reduces the Cr+++. When the
acid is done the Al has a thick coat of a black substance. Cr metal? TBD…
Hoist a jar with me this Xmas Eve,
Der Alte |
Cheers! The Saviour was born!
Since you speak of a thick layer, it indeed must be Cr metal. Otherwise it might well be a remnant of the Al foil. Dissolving a piece of
Al-foil in HCl solution, also leaves behind some blackish mass, that only dissolves on longer standing, and has been shown to be some mixture of Al
and Al2O3.
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DerAlte
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Bezaleel,
A thick layer was probably an exaggeration. I did not use foil - it's quite useless and just floats around lifted by hydrogen bubbles. This was about
1mm thick pieces of Al.
However, on returning home after 2 days away I found that the Al 'thick coating' had vanished leaving dull Al, so perhaps as you say, it's just a
mixture of Al and Al2O3.
Also, the provenance of the Al was unknown, except it was soft. Might have been an alloy.
Der Alte
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DerAlte
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Chromium Revisited – Part IIIa
(5) Chromium II Compounds.
(5A) Reduction with zinc/acid.
This was mentioned above. To see the color of the Cr++ ion (hexahydrated) is quite simple. Take a long thin test-tube (longer and thinner the better)
and load it half high with cleaned zinc foil or sheet (hammer out if necessary and use dilute acid to dissolve of oxide). The aim is to maximize the
surface area of the zinc and minimize that of the liquid surface when filled. Dissolve some Cr[III] salt – chloride or sulphate {not saturated! Say
25g/100g}, chrome alum (saturated) etc. – in water to produce a deep green or violet solution. Fill the tube to about the top of the zinc. Little
happens. Add conc. HCl (If using chloride) or 40% sulphuric acid in sufficient quantity to react with part of the zinc but not all (difficult to judge
– add in stages). That way the final solution is not heavily acidic.
The usual zinc/acid reaction occurs. Plug the top of the tube with tightly rolled tissue paper, cotton wool or fine glass wool. This minimizes
convected air entering the tube and tends to keep a hydrogen atmosphere above the liquid which improves efficiency. The liquid begins to get paler at
first and violet (or green) turns to a clear sky to darker blue that is quite different as the Cr[II] develops. The color change is quite striking,
especially with the green. Equations shown earlier.
The solution containing the Cr[II] can be kept for some time if poured off into a smaller test tube nearly full and tightly corked with a rubber bung,
especially if refrigerated. I never managed to crystallize the Cr[II] salt out because the potassium sulphate or the zinc sulphate come out of
solution first and you cannot boil to evaporate! In order to get the chromous salt in a pure state electrolysis of chromic salt (preferable sulphate)
in a divided cell with Cr[III] in both anode and cathode compartments (when one can simultaneously produce (di)chromate at the anode!)
(5B) Instability of Cr[II] salts WRT water.
Take some of the solution from (5A) and heat in a test tube. Once it gets hot bubbles of H2 appear and on boiling the blue color reverts to green,
showing that Cr[III] is produced by reduction of water.. Equations shown earlier.
(5C) Instability of Cr[II] salts WRT Oxygen (air).
Pour some of the liquid into a 50ml beaker and leave for a day or so. The solution reverts to the violet form of Cr[III].
(5D) Chromium amalgam, which can be made by electrolysis at an Hg cathode, is said to produce black CrO on oxidation in air. Have not tried this due
to very short Hg supplies!
(5E) Chromous Acetate, Cr2(CH3CO2)4(H2O)2:
It can be made as a red solution by adding a reasonably strong solution of Zn reduced Cr[III] sulphate to a saturated solution of sodium acetate.
With care, small red crystals can be made and kept. I used the following method:
Take a glass tube, 1 to 2 cms ID X 25 cms long or thereabouts and fit the lower end with a rubber bung and a short glass tube. Attach a short
rubber/neoprene tube with a Mohr’s clip or a pinch cock to a longer glass tube. Prepare a saturated solution of sodium acetate in water and cool to
around 5C in a refrigerator.
Using a suitable support, place the long tube vertically with a mat of glass wool next to the bung and fill with zinc pieces to a height of around
10-15 cms. Add a moderately strong solution of Cr2(SO4)3 (this is very soluble, don’t use a too concentrated solution) to the zinc with the Mohr’s
clip in position. Add 40% or so H2SO4 and proceed with the reduction.
In this preparation chrome alum is undesirable because of the potassium content; potassium sulphate may crystallize out at a later stage. {If all your
have is the alum, convert it to hydroxide and dissolve in sulphuric acid, which is the way I did it.}
Once the acid is consumed (there should be zinc remaining), run the solution into a Erlenmeyer flask containing the cooled sodium acetate by releasing
the Mohr’s clip. The end of the lower glass tube should be under the surface of the acetate solution. The quantities should be such that the flask
is nearly filled to the top with the incoming solution – a bit of prior experimentation can ensure this. The acetate should be in excess.
The blue solution entering should produce a yellow then a red coloration. Remove the glass tube and put a rubber bung on the flask. Crystals of the
Cr[II] acetate may form immediately; if not they should on cooling in an ice mixture in a refrigerator. The color is brick red. After cooling for 12
Hrs, decant and rapidly dry between good quality filter papers. The red crystals can be kept some time in a small tube but will turn darker, then
brownish and finally greenish due to oxidation and/or water reduction. If you have a dessicator, dry in that first (over conc. H2SO4).
(5F) Chromous Oxalate can be made by the same process and is also red. It is said to be less soluble and more stable than the acetate. In this case
use the potassium salt because sodium oxalate is poorly soluble.
(5G) Cr[II] Hydroxide.
In the presence of the zinc from the reduction, addition of sodium hydroxide precipitates white zinc hydroxide which masks the yellow color of the
Cr[II] compound. However, the production of Cr(OH)2 can be inferred by the fact that the solution becomes colorless and the precipitate is not pure
white but yellowish. On standing the gray-green color of the Cr[III] hydroxide appears.
Discussion of Cr[II] compounds.
All these compounds are strong reducing agents and deteriorate in air due to oxidation. Nor are they stable in solution due to the ability to reduce
water to H2 gas. As such they are laboratory curiosities and cannot be kept for long. They cannot be concentrated in solution by boiling since this
accelerates the reduction. Apart from the acetate I have not managed to crystallize out any Cr[II] compound. The Ac is atypical with its red
coloration and is both dimeric, covalent co-ordinated and non ionic (see Wiki). Most ionic acetates are very soluble but Cr[II]Ac2.H2O is sparingly
so, although soluble in hot water (hence the cooling required above).
Anhydrous Cr[II] compounds can be made by the action of eg chlorine on Cr metal or oxide and carbon. They are white according to the texts, but the
processes require very high temperatures.
The catalytic action of Cr[II] ions on insoluble Cr[III] compounds never fails to be mentioned in texts on Cr compounds so I might as well join the
club. The explanation given is as follows: The Cr[III] compounds dissolve rapidly with a trace of Cr[II] ions by the following reduction chain
reaction:
Cr[III]X3(s) + Cr++(aq) -- > Cr++(aq) + 3X- + Cr+++(aq):
The Cr++ ion reduces the insoluble covalent Cr[III] compound to a new Cr++ ion, also soluble, and is oxidized to Cr+++, and so on. I guess the
thermodynamics allows this.
Next: Other valences of Cr.
Der Alte
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DerAlte
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Chromium Revisited Part IV
The current version of the Periodic table is in some respects less genuinely chemical in flavor than Mendeleef’s original, which he based on
‘higher saline oxides’ of the elements as representing the maximum valence state. The current version (IUPAC) stretches across 18 elements. If it
included the 4f block [rare earths] it would stretch all the way across your wall! Elements exhibiting similarities due to valence are now widely
separated.
Older versions placed oxygen and sulphur etc, as Group VI elements along with Cr and Mo, etc., in VIa and VIb sub-groups. Between S and Cr there are,
in fact, distinct similarities such as SO2Cl2 and CrO2Cl2, or chromates and sulphates, CrO3 and SO3 and so on. Mendeleef also noted the sideways
similarities such as those between V, Cr and Mn. By such means he was able to predict unknown elements yet to be discovered.
Chromates
Comparing the lower oxidation states of Mn and Cr, we can note that with Mn the Mn[III] compounds are powerful oxidants whereas Mn[II] is stable. With
Cr, Cr[II] is a good reducing agent whilst Cr[III] is the stable state. Permanganates are considerably more powerful oxidants than Chromates:
Acidic at pH=0: MnO4- -- > Mn++ , Eo = 1.51v; Cr2O7- - -- > Cr3+++, Eo=1.38V;
at pH=14, MnO4- -- > Mn(OH)2, Eo= 0.34V and CrO4- - -- > Cr(OH)3, Eo= -0.11V
and if the hydroxide is ‘redissolved’ in the alkali, forming the chromite ion,
CrO4- - -- > Cr(OH)4- with Eo = -0.72V
As a consequence chromates can be easily made in the wet way unlike permanganates (see the long permanganate thread). Suitable oxidants to take the Cr
from [III] to Cr[VI] state are hydrogen peroxide (acidic or alkaline) or NaOCl (alkaline only).
(6a) Under approximately neutral conditions:
Starting from a Cr[III] salt, precipitate Cr(OH)3 as above. Add H2O2 (3% is OK, it is around 1N). The greenish precipitate dissolves and the solution
turns yellowish due to formation of chromate ion:
2Cr(OH)3 + 3H2O2 + 4OH- -- > 2CrO4- - + 8 H2O
Notice that for neutrality 4 positive ions must be present in the above. If only water is present initially then a solution with H+ ions would have
to result, implying a solution of chromic acid. However, if instead we make the solution alkaline with , say, NaOH, then sodium chromate results, so
long as the final solution is near neutrality or alkaline. Which leads us directly to:
(6b) Produce a solution of the green ‘chromite’ by carefully redissolving the precipitate in NaOH. Add H2O2 (or NaOCl solution) slowly until the
solution is bright yellow:
Cr(OH)3(s) + OH- -- > Cr(OH)4-(aq) ….. green
Cr(OH)4-(aq) + 3H2O2 + 2OH- -- > CrO4- - (aq) + 8H2O …… yellow
(6c) Alkaline hypochlorite (bleach is about 0.8N at 6%) can be used to make sodium chromate equally well:
CrCl3 + 2NaOCl + 6NaOH -- > Na2CrO4 + 3H2O + 5 NaCl
Any Cr[III] salt can be used in place of the chloride, and H2O2 used instead of hypochlorite, or potassium salts used instead. Na2CrO4 is quite
soluble (85g/100g aq at RT) but using the above method does make the extraction and purification somewhat difficult due to the preponderance of NaCl.
Cr2(SO4)3 and H2O2 as oxidant is a bit easier (I’ll leave it to you to work out the equation!).
(6d) Neutralize the yellow solution produced in (6c) with acid. The solution turns orange due to formation of dichromate:
2CrO4-- + 2H+ -- > Cr2O7- - + H2O; at pH ~ < 5.9 and this can be reversed:
Cr2O7- - + 2OH- -- > 2CrO4- - + H2O at pH > ~ 6.7
Dichromates are still in oxidation state [VI]; this is a distinct difference from manganates, where the oxidation state varies from [V] in the hypo- ,
[VI] in the manganate and [VII] in the permanganates.
The structure of the chromate ion is (-O)(O)Cr(O)(O-) and that of the dichromate can be written as (-O3)Cr-O-Cr(O3-). Further addition of CrO3
groups is possible to produce polychromates M2CrO4(CrO3)n , all in oxidation state Cr[VI], which tend to be marginally stable. Treating K2Cr2O7 with
concentrated nitric acid is said to produce a trichromate K2Cr3O(10) - (haven’t tried this).
Many chromates are insoluble. K, Na, Mg and Ca are all soluble to some extent. Most chromates are yellow, but Ag, Hg are red; most dichromates red to
reddish brown and tend to be more soluble than the chromates.
Treating a chromate or dichromate with H2O2 causes a transient blue solution (due to the formation of CrO(O2)2 (Wiki)). This is a very sensitive test
for chromate or dichromate.
A final reminder that all chromates and Cr[VI] compounds are fairly poisonous. Although Cr[III] compounds are only marginally so, chromates are
because of their oxidation capabilities and are also carcinogenic; they stain and can be absorbed through the skin. Wear gloves and watch eyes. They
are not usually an environmental hazard because of ready conversion to Cr[III] state, except in very acidic waters.
To be continued….
Der Alte
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ScienceSquirrel
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The blue chromium (VI) oxide peroxide is stable as an ether complex in which it forms a deep blue solution.
The solution can be carefully evaporated to form dark blue needles.
http://en.wikipedia.org/wiki/Chromium_peroxide
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woelen
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I tried the addition of hydrogen peroxide to an alkaline solution of chromium(III). It indeed works, but the product obtained is not a purely nice
yellow chromate solution. The final product tends to be yellow/brown, due to formation of peroxo complexes. When most of the green chromium(III) is
converted to chromium(VI), then addition of more hydrogen peroxide causes partial conversion of the chromate to tetraperoxochromate(V). The latter is
a deep brown ion with formula [Cr(O2)4](3-).
A very nice oxidizer for converting chromium(III) to chromium(VI) in acidic solution is peroxodisulfate. The reaction is slow, but on heating the pure
orange color of dichromate ion is produced. If a pinch of silver nitrate is added, then the reaction proceeds MUCH faster, the silver(I) ion is a
strong catalyst of the reaction. It is oxidized very quickly to silver(III) by the peroxodisulfate and the silver(III) in turn oxidizes the
chromium(III) to chromium(VI) very quickly.
Chromium(VI) has a beautiful and colorful chemistry. The solid compounds also are very nice and brighly colored compounds. Sad that these compounds
are so toxic.
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DerAlte
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Thanks, gentlemen, for your comments.
@ Woelen
I agree 100% that the peroxide reactions with Cr[III] salts can produce all sorts of anomalies. (I have something re this in a later section). It
seems to depend critically upon pH and concentration. The reaction is often slow. I had thought earlier that H2O2 would be the ideal oxidant for the
Cr+++ to chromate reaction because the product is H2O but I now prefer my old favorite NaOCl in alkaline solution as the oxidant, regardless of the
baggage it carries.
You can get the blue peroxy color from excess H2O2; the oxidation can appear to not occur but later will have completed itself to yellow chromate or
orange dichromate. Another anomalous result is the production of a deep reddish color by adding the peroxide to the Cr+++ first and then the
hydroxide. This is too deep for dichromate and occurs in alkaline solution. Any ideas?
Addendum Note: the precipitate obtained from using Na2CO3 on a Cr+++ solution has a bluish-purple color that is persistent. I had a test tube with
some in left for several weeks. On adding acid it fizzled as if some sort of carbonate or a carbonato- complex had been formed. But I may not have
washed it sufficiently…
Correction: The equation in the last part for hypochlorite oxidation of Cr+++ should read
2CrCl3 + 3NaOCl + 10NaOH -- >2Na2CrO4 + 5H2O + 9NaCl - a balancing error.
Der Alte
[Edited on 7-1-2012 by DerAlte]
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ScienceSquirrel
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I think that chromium chemistry is fairly safe if you do it on an appropriate scale and handle the compounds with reasonable care eg gloves, etc and
contain or avoid making areosols.
Potassium dichromate is a great starting material and you can make a lot of interesting compounds from it. 100g will go a long way and you can make
chrome alum, chromium II acetate and some nice complexes like the blue peroxo complex mentioned above.
Chromium VI quickly turns in to chromium III in the environment and it is non toxic at low levels as far as we know.
One or two grams of chromium salts down the sink will not make any difference by the time they reach the sewage farm.
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DerAlte
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More on Chromates I
(7a) Preparation of Chromates from Cr+++ salts.
Unlike with Manganates, the Chromates all represent the single oxidation state Cr[VI]. Dichromate is stable only as solid or in acid solution and can
be considered as a condensation product of the chromate ion. Further condensations to the poly chromates are possible. Neither of the acids H2CrO4 or
H2CrO7 can be isolated although there is good evidence for the existence of the ion HCrO4-; however, hydrogen acid salts of Cr also do not stably
exist.
As shown before, chromates are readily converted to dichromates and hence only one species need be considered. If the other is wanted then either a
suitable acid or alkali can be used to make it simply.
Remember that all Cr[VI] compounds are poisonous. Wear gloves; do not breath in vapors when evaporating
For several reasons the oxidant H2O2, which might seem ideal as its only product is water, was found to be unsuited to this purpose. For a start it
has a tendency to decompose to O2 in the presence of heavy metal ions; and Cr+++ is no exception. In acid solutions with Cr+++ ions present it has a
tendency to produce peroxy compounds which are unstable; these decompose with evolution of O2. In ‘neutral’ (unacidified) solutions O2 is also
produced (the solutions are actually acidic). In alkaline solutions, depending upon the order of mixing of components, the hydroxide may be
precipitated and the reaction is then slow; also peculiar products seem to be formed as intermediates (See later on this). The net result is that no
sensible stoichimetric equation can be written for the action of H2O2 on the Cr+++ ion in solution; it depends upon exact conditions such as
concentration, temperature, pH and just possibly phase of the moon! It is inefficient. Another Oxidant is needed. I wasted a lot of effort on this
possible method, concentrating H2O2 etc. (There’s a good thread in SciMad on H2O2 concentration, should you want to do it).
So I turned to the old favorite (and cheap!) sodium hypochlorite. This is unstable in acidic solution, producing chlorine, oxygen, or a combination,
depending on acid. Hence attention is directed to alkaline solutions where chromate will be the product.
A putative oxidation in ionic terms might be written:
2Cr+++(aq) + 3OCl-(aq) + 10OH-(aq) -- > 2CrO4- -(aq) + 3Cl-(aq) + 5H2O …(1)
Violet (or green) -- > Yellow
However, if solutions containing these ions were mixed, since ionic reactions are (nearly) always faster than oxidations because the bond in the
oxidant has to be broken, and the hydroxide is predominant, the likely result would be a large amount of gelatinous hydroxide, especially if
concentrated solutions are involved, giving a modified equation in which the precipitated hydroxide is oxidized:
2Cr(OH)3(s) + 3OCl-(aq) + 4OH- (aq) -- > 2CrO4- -(aq) + 3Cl-(aq) + 5H2O …(2)
Green-gray ppt. + Colorless liquid -- > Yellow liquid
It is best to slowly add a combined solution of hypochlorite and hydroxide gradually to the Cr+++ ionic solution stirring if necessary to avoid a
mess of Cr(OH)3 forming. An alternative method (one I favor) is to pre-precipitate the hydroxide and just redissolve it, then adding the NaOH
dissolved in the NaClO solution:
2Cr(OH)4-(aq) + 3OCl-(aq) + 3OH- (aq) -- > 2CrO4- -(aq) + 3Cl-(aq) + 5H2O …(3)
Green -- > Yellow.
The overall reaction using Na+ and SO4- - as spectator ions can then be indicated as:
Cr2(SO4)3 + 3NaOCl + 10NaOH -- > 2Na2CrO4 + 3Na2SO4 + 3NaCl + 5H2O …(4)
I prefer this on principle to using Cr[III] chloride for reasons discussed below. A very small excess of NaOH is desirable to keep the solution
alkaline (> ~pH 8). The color change serves as a weak indicator (the reaction is not reversible).
Stirring helps during the additions; certain peculiar side reactions which do not seem to radically affect the product may also occur. Do not use
highly concentrated solutions – the chromic sulphate is extremely soluble – or you will land up with an intractable mess. 5% Hypochlorite is
a good strength (e.g., bleach).
If chromic chloride is used, 9mols of NaCl is produced for every 2 of chromate which renders extraction difficult. Do not use KOH in place of NaOH or
it will confuse the products more than necessary.
The solubilities of the various products (g/100g Aq.) are ~ as follows: (Wiki)
Compound Temp C = 0 30 60 100
Na2CrO4 32 88 115 126
Na2SO4 5 41 45 43
NaCl 36 36 37 39
The relative weights of expected products are Chromate, 1.00; Sulphate, 1.58; Chloride, 0.65. Because of the common ion effect (Na+), actual
solubilities will be somewhat less than above; how much less we cannot calculate. What strategy should be used to isolate the chromate? First estimate
how much water would dissolve each component, assuming no common ion. This is proportional to mass and inversely proportional to solubility; in the
same order as above, the ratios are @100C are roughly, 0.793 : 3.67 : 1.67 . At 0C ratios are: 3.13 : 31.6 : 1.81.
The strategy becomes obvious. (a) Evaporate @ ~ 100C until crystals appear. At this point the sulphate is saturated. On cooling to 0C most of the
sulphate will precipitate. (Be careful about supersaturation of the sulphate; it’s about the worst offender I know).
Those familiar with fractional crystallizations know how to proceed from then on. Some hate the process; I revel in it. It takes a lot of time and
patience. To avoid the massive task of actually crystallizing out the chromate, after eliminating as much NaCl as possible (important!), we pull an
ace out of our sleeve.
We convert to dichromate by adding one equivalent of sulphuric acid:
2Na2CrO4 + H2SO4 -- > Na2Cr2O7 + Na2SO4 + H2O
Now the sodium dichromate is incredibly soluble – 163g/100 aq. @ 0C and 417 g/100g @100C – no way that is going to precipitate while you get rid
of more sulphate by the same process as the chromate. So you land up with a strong solution of Na2Cr2O7 plus minor amounts of chloride and sulphate.
The reason for keeping the chloride level as low as possible is that HCl or Cl2 may be generated if you let the acid level get too high. If you use
appropriate amounts this will not happen, but it requires careful measurements. With strong acid dichromate may form chlorine or, worse, CrO2Cl2. The
latter is unlikely and even if formed will hydrolyze. See later.
Finally, the dichromate is precipitated by the addition of a potassium salt (but not sulphate).
The solubility of K2Cr2O7 is far less than the sodium salt and it separates on cooling to 0C. Solubility @0C, about 5g/100; @ 30C, 18g/100; @ 60C
46g/100; and @ 100C, ~80g/100
The ammonium salt is also far less soluble and can be similarly produced. Solubility is around 18g/100g @0C. Do not boil a solution containing
ammonium dichromate, it will decompose. The color of this dichromate is more orange.
The above wet method is a bit time consuming. The following fusion method is easy and recommended. It is much easier than the corresponding
permanganate production.
(7a) Preparation of Chromate from Cr2O3 or Cr(OH)3.3H2O by fusion
Warning: Take care with fused hydroxides and nitrates. Use stable apparatus and do not perform in a dwelling area.
Use a nickel, iron or SS crucible, small cup or bowl for this process. Do not use porcelain. The generalized reaction sought is as follows,
Cr2O3 + 4MOH + 3[O] -- > 2M2CrO4 + 2H2O ;
M = Na or K; the [O] to be provided by a suitable oxidant.
O2 from air can be used but takes a very long time unless sparged through the melt. The Cr2O3 can be replaced by dried hydroxide Cr(OH)3.3H2O which
decomposes to Cr2O3 at around the temperature of the melt. Suitable oxidants are nitrate, chlorate or perchlorate, the last needing a slightly higher
temperature. Chromate is stable up to around 900C AFAIK so heating to bright red heat does not destroy it, although dichromate will not survive.
If MNO3 is used as oxidant, a proportion Cr2O3 : MOH : MNO3 of 1:1:1 by weight is good for M=K and 3:2:2 for Na, but add a little extra MOH to account
for impurities such as carbonate.
The nitrate/hydroxide mixture melts at temps. between 200-300C, lower for the K salts. The time taken to complete the reaction depends on the nature
of the Cr2O3 – if gritty it takes longer. If chromic hydroxide is used it must be dried and/or at least partially decomposed to oxide, to avoid
splattering as it is added to the fused mixture. Pre-recalescent dried oxide/hydroxide reacts very much faster than technical Cr2O3 from pottery
stores or oxide that has been strongly heated.
The reaction hoped for is 5Cr2O3 + 14K0H + 6KNO3 -- > 10K2CrO4 + 7H2O + 3N2
Steam and gas are emitted and the material may bubble so allow room for this in the crucible. The reactants will go pasty after ~15 minutes and should
be stirred with a thick iron wire like a coat hanger. When the product becomes almost solid, increase the heat to a low red heat preferably with a
cover over the crucible. (The nitrite first formed tends to decrepitate). Chromate is stable to about 900C so heating will not destroy it.
About 45 mins at a red heat is sufficient to complete the reaction. It is far easier than the manganate fusions and the yield is generally very good.
Leach out the contents of the crucible after cooling. The main impurity will be KOH – the nitrate should have been fully decomposed. K2CrO4 will
easily crystallize out as small anhydrous yellow crystals on filtering the leach water and evaporation.
You can use ordinary filter paper with chromate but not with acidified dichromate. Sodium chromate will contain water of crystallization. Both
chromates are hygroscopic or deliquescent.
It is said that K2CO3 can be used in place of KOH but I haven’t tried it. The carbonate is not easily fused, nor does it form a useful eutectic with
nitrate.
(7c)Dichromate to Chromate conversion.
This can be conveniently done using carbonates instead of hydroxide. Boil to remove the CO2:
M2Cr2O7 + M2CO3 -- >2M2CrO4 + CO2; M=Na or K
To be continued...
Der Alte
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DerAlte
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More on Chromates II
(7d)Cr2O7- - to Cr+++ Reduction using alcohols
Since Potassium Dichromate is a more common compound than any Cr[III] salt, the experiments with Cr+++ described can be done in reverse by reducing
dichromate to produce the sulphate (as chrome alum). This demonstrates the oxidation power of acidified dichromate, which is moderate. Chromate is a
poor oxidant.
You can use ethanol or isopropyl alcohol. With C2H5OH, acetaldehyde is first produced but this may be further oxidized to acetic acid unless the
aldehyde is expelled by heating. Isopropyl alcohol, being a secondary alcohol, produces a ketone, acetone, resistant to further oxidation. (Tertiary
alcohols cannot be oxidized by dichromate). Hence, since we don’t want involatile acetates, use isopropyl and boil off the acetone (or condense it
and use it!).
2(CH3)(CH3)COH +{O} -- > 2CH3(CO)CH3 + H2O
CR2O7- - + 8H+ -- > 2Cr+++ + 4H2O + 3{O}
The H+ comes from the sulphuric acid which also provides sulphate ions; the dichromate provides three {O} moieties for oxidation of the alcohol, and
the potassium provides the other base in the alum KCr(SO4)2.12H2O as the final product. The Cr can be isolated as hydroxide if so desired and used to
produce any Cr+++ salt with a suitable acid.
Dichromate is frequently used in organic chemistry as a moderate oxidant, but has no use in pyrotechnics AFAIK. Potassium dichromate crystallizes well
giving large deep red or garnet crystals and is a favorite with crystal enthusiasts.
Ammonium Dichromate auto-reduces to a lovely green form of Cr2O3. It can be set off by a red hot wire or a match.
(NH4)2Cr2O7 (heat>200C) -- > Cr2O3 + 4H2O + N2.
Heating Potassium dichromate with sulphur also gives the same product:
K2Cr2O7 + S + heat -- > Cr2O3 + K2SO4 (or something like that).
……..
To be continued...
Der Alte
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Bezaleel
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| Quote: Originally posted by DerAlte | | You can get the blue peroxy color from excess H2O2; the oxidation can appear to not occur but later will have completed itself to yellow chromate or
orange dichromate. Another anomalous result is the production of a deep reddish color by adding the peroxide to the Cr+++ first and then the
hydroxide. This is too deep for dichromate and occurs in alkaline solution. Any ideas? |
H2O2 may behave as a reductor in alkaline solution. The pH needs to be quite high for that. So the reddish colour may be due to the formation of Cr++
(probably mixed with some Cr+++). I'm only theorizing here, not experimenting.
Edit: I much like this thread, Der Alte. Keep up the great work!
[Edited on 13-1-2012 by Bezaleel]
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DerAlte
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More on Chromium[VI]
The following experiments are potentially DANGEROUS. Take note of the hazards and act accordingly.
(7e) Chromyl Chloride, CrO2Cl2
{Not for k3wls; or even intelligent but over enthusiastic teens below the age of reason without knowledgeable adult supervision, or the
cack-handed or panic prone. All those attempting ought to have had experience managing at least one previous disaster.}
This liquid (bp 117C) is the acid anhydride of chromic acid, a Cr[VI] compound. It is an example of a volatile chromium compound at near RT. CrO2Cl2
is relatively easy to prepare. It is a deep red liquid, fumes in moist air due to hydrolysis, and has a vapor pressure of about 25 mm Hg at 25C
(IIRC). It is definitely quite volatile. It has a very nasty acrid smell, vaguely like Cl2 or Br; the liquid is quite similar to Br in appearance.
HAZARD: very poisonous vapor, about as bad as phosgene. As a bonus, it not only destroys lung tissue but deposits carcinogenic Cr if you
survive. However, it stinks so badly you will know it is deadly and receive a prior warning!
Do not perform this within the confines of any dwelling house, or in any apartment block. Escaped vapor can hang around for a long time and leave its
odor. The only safe conditions are in a proper lab under a fume hood with a good ventilation, or outside. Outside, use a fan to keep
up a current of air toward the apparatus and away from the operator or any spectator. It also involves the use of concentrated H2SO4; familiarity of
handling this is essential.
All glass apparatus is virtually essential. {Over 50 years ago my father and I used a retort with a long neck thrust well into an RB flask cooled by
flowing water at a few degrees above freezing. Even so the smell of the stuff hung around our small private lab for days.}
Use a RB flask (a distilling flask with side tube would be best) with a dropping funnel to pour in concentrated acid slowly and attach a condenser to
the tube from the flask. Lead into an ice cooled receiver flask with an arrangement to vent gases (HCl) and prevent ingress of moisture. Use no rubber
or plastic tube connections except PTFE.
Do not attempt in an open test tube! I did do a very small scale effort to remind me last year, mixing by containing the acid in a small tube lowered
carefully into a larger one, but I suggest you not try this.
Grind together about 5 parts common salt and 8 parts of potassium dichromate. The salt should be dried first by heating. Place in the flask with a
dropping funnel containing a shut off e.g. a separating funnel. Fill the latter with concentrated acid (c 98%). Allow slow entry of acid, drop by
drop. Excess acid beyond the stoichimetric proportion is needed as a dehydrating agent to prevent hydrolysis.
The reaction is firstly between the sulphuric acid and the salt to produce HCl:
NaCl + H2SO4 -- > NaHSO4 + HCl
The reaction is immediate and rapid at normal RT. Reddish fumes appear as the dichromate undergoes partial reduction to Chromyl Chloride as the HCl is
oxidized:
K2Cr2O7 + 4HCl + H2SO4 -- > 2CrO2Cl2 + 3H2O + K2SO4
Excess H2SO4 prevents hydrolysis. Heat is evolved but in the latter stages a gentle heat will be needed to distill off the remaining Chromyl Chloride.
The product in the receiving flask is a very dark red liquid. It will probably contain dissolved HCl, removable by further distillation (if you so
desire).
It is not of much use and potentially dangerous. Although it could be kept as a trophy chemical by those so inclined, I am not sure of its stability
long term for inclusion in an sealed ampoule. Sunlight (UV) is said to decompose it.
CrO2Cl2 hydrolyzes readily in water, and hence fumes in moist air, forming HCl and a solution of H2CrO4 = CrO3 + H2O: CrO2Cl2 + 2H2O <
-- > 2HCl + H2CrO4
It is a strong oxidant and chlorinating agent. Great care must be taken not to let it get on the skin or serious burns result.
(7e) Potassium Chlorochromate, KCrO2Cl
A stable salt, potassium chlorochromate, KCrO2Cl can be made by adding concentrated KCl solution to chromyl chloride. It has a bright orange color
& is stable if kept dry.
CrO2Cl2 + KCl + H2O -- > KCrO3Cl + 2HCl.
Alternately this salt can be made without chromyl chloride as follows:
Grind up 10g potassium dichromate. Heat gently if necessary until dissolved by digesting with about 25 ml of 38% HCl solution. Filter through glass
wool. Large red prisms of the chlorochromate develop slowly on standing, preferably at near 0C; keep in a refrigerator in ice/water for a day or so.
(NOT, of course, one used for food!).
Pour off the liquid and dry crystals on a porous plate (eg, terra cotta from gardening supplies; it is an strong oxidizing/chlorinating agent and
destroys filter paper). It can be dried carefully in a stream of heated dry air, but too much heat causes release of chlorine. It hydrolyzes slowly if
damp and is hygroscopic:
K2Cr2O7 + 2HCl < -- > 2KCrO3Cl + H2O
(7e) Chromium Trioxide CrO3
Hazard: Hot concentrated Sulphuric Acid. Powerful Oxidizing agent
This oxide can be produced by, essentially, dehydrating a solution of chromic acid with concentrated sulphuric acid:
K2Cr2O7 + H2SO4 --> K2SO4 + ‘H2CrO4’ -- > K2SO4 + H2O + CrO3
The operation is rather tricky, in my experience, and yields tend to be disappointing. The problem, on investigation of the literature, seems to be
due to the fact that CrO3 is least soluble in about 75% acid and one should aim for this percentage as an end point for good results.
Weaker or stronger acid dissolves more CrO3. Unfortunately a diligent search of the web yielded no more precise data than this. Various sources quote
the optimum as 66-85%. What the actual solubility is I do not know.
Further problems are the precipitation of K2SO4 at the same time; the extreme hygroscopic nature of CrO3 once precipitated and its strong oxidizing
power. The CrO3 cannot be dried with any of the usual organic solvents – for instance, it will set fire to ethanol with incandescence. Since the
solvent is sulphuric acid, the usual desiccators are useless. Try drying on a porous tile. Suction from A Buchner funnel on a glass mat is the best
that can be done. However, it can be redissolved in water and purified by careful evaporation/crystallization.. The solubility is about 62 to 65% by
weight (185g/100g aq) over 0-90C.
At the level of acid needed the K2SO4 will exist as KHSO4: 75% acid is about H2SO4 + 2H2O. The solubility is then less than 10g/100 solution at 0C.
With care, KHSO4 can be precipitated before CrO3.
The following method is copied from an old source: (quantities reduced):
“Dissolve 100g K2Cr2O7 in 160ml of hot water (temp has to be >80C for this) and add 135ml conc. H2SO4 very slowly and cautiously with active
stirring. Stand over night in a cool place and next day decant solution off the potassium sulphate precipitate formed. Heat solution to 90C and add
50ml H2SO4; if a ppt. of CrO3 forms, add just enough water to re-dissolve. Evaporate till crystals appear and let stand for a day in a cool place.
Collect and evaporate the solution for a further crop. Drain and wash with 15ml HNO3 (sg 1.46), then with a further 7.5ml. Dry on a sand bath.”
I used about the same procedure except for the HNO3, which I didn’t possess, re-crystallizing instead. It works but yield is only around 50% after
re-crystallization.
I have not tried sodium dichromate but its easy solubility makes the following process sounds attractive. Adding conc. acid to a solution near 90C as
above is a bit hairy!
“CHROMIC ANHYDRIDE, CrO3
The addition of sulphuric acid to a solution of either a chromate or a dichromate liberates chromic acid which is very soluble and can exist in
solution in the different forms, H2Cr04, H2Cr207 and CrO3, in equilibrium with each other. With the addition of a
large excess of concentrated H2SO4, water is withdrawn from the hydrated forms and the anhydride separates in the shape of red needles.
Materials: sodium dichromate, Na2Cr2O7-2H2O, 100 grams = 0.33 F.W.
36N H2SO4, 400 cc
Apparatus: 8-inch porcelain dish. Glass plate to cover the 8-inch dish, suction filter with glass marble, glass-stoppered sample bottle, tripod.
Bunsen burner.
Procedure: Dissolve the 100 grams of sodium dichromate in 250 cc. of water and filter from any sediment. Add rather slowly with constant stirring
about half of the concentrated sulphuric acid until a slight permanent precipitate of CrO3 is formed. Let the mixture cool for half an hour or longer,
then add slowly, while stirring, the rest of the sulphuric acid. Let the mixture stand over night covered with a glass plate in order that the crystal
meal may become somewhat coarser. In such a crystal meal standing in its saturated solution, the smaller grains dissolve and their material deposits
out on the larger crystals. But even now the crystal meal will be rather fine and it will at first run through the filter; if, however, while waiting,
the mixture is heated with stirring to 100° and allowed to cool slowly, and this process is repeated once or twice, a more satisfactory product will
be obtained.
To collect the crystals, use a suction filter, but place a small glass marble in the funnel instead of the usual plate and paper. If the red crystals
at first run past the sides of the marble, pour the liquid in the bottle repeatedly back on to the filter until finally the filtrate runs clear (see
last sentence of Note 3 on page 5). After draining the crystals completely and pressing the surface with the round end of a test tube, stop the
suction and
pour 15 cc. of 16 N HN03 so as to wash down the sides of the funnel and cover the surface of the product. Stir up the product with this washing fluid
for a depth of about 1/2 inch. Suck dry and repeat the operation twice with 10 cc. of nitric acid each time. Finally drain the red crystals as free of
liquid as possible, transfer the crystals to a dry 8-inch evaporating dish and place this on a hot plate to let the nitric acid evaporate. When the
product is dry and no longer gives off vapors of nitric acid place it in the glass-stoppered sample bottle.”
(Blanchard, Phelan, & Davis, Synthetic Inorganic Chemistry; fifth edition 1936)
Re CrO3 also see http://www.sciencemadness.org/talk/viewthread.php?tid=6116 esp. the post by S.C.Wack.
The similarity with SO3 is marked. Both tend to polymerize in long strings and produce wooly strands. The normal form of CrO3 is rhombic needles.
Compare SO2Cl2 to CrO2Cl2 and the chromic acids H2Cr2O7/H2CrO4 with H2S2O7/H2SO4. The Cr2O7- - ion predominates at low pH; Ka for chromic acid is
about 0.7 but the acid(s) are not isolable. As I said above, all this gets lost in the current periodic table which reflects physics rather than
chemistry.
To be continued…
Der Alte
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DerAlte
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@Bezaleel
Der Alte wrote:
Another anomalous result is the production of a deep reddish color by adding the peroxide to the Cr+++ first and then the hydroxide. This is too deep
for dichromate and occurs in alkaline solution. Any ideas?
Bezaleel wrote:
H2O2 may behave as a reductor in alkaline solution. The pH needs to be quite high for that. So the reddish colour may be due to the formation of Cr++
(probably mixed with some Cr+++). I'm only theorizing here, not experimenting.
……..
I forgot to add that O2 is also given off when this occurs.
H2O2 in alkaline pH=14 (1N): H2O2 is a very weak acid with pKa1 ~11.7, but at this alkalinity the dissociation H2O2 < -- > HO2 - + H+ should
lie well to the right because [H+] =10^-14. HO2- thus represents the oxidation state -1 for oxygen at pH=14
The O2 may be due to the auto-oxidation leaving the oxygen in the 0 oxidation state.
HO2- + HO2- -- > O2 + 2e- + H2O
The deep orange color cannot be due to dichromate or polychromates because they do not exist at pH 14. Nor do I think it could be due to Cr[II]
compounds due to the weak reductive power of H2O2 – and Cr[II] compounds are blue, if ionic.
However, under the conditions of “adding the peroxide to the Cr+++ first and then the hydroxide” , the pH is actually about 4 on entry of
hydroxide. The H2O2 will exist as H2O2 in such a solution. It is possible then for dichromate to exist too, transiently as the acidity is neutralized.
But the color looks too orange for that…
But, as you say, I too am only theorizing here, not experimenting!
Regards
Der Alte
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DerAlte
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Other Oxidation States of Chromium and comments
This is a final wrap-up of this series (did I hear a sigh of relief?). Added for completeness, it is mainly a compilation of the literature searches I
have done during the course of my Chromium madness phase. If I have actually done an experiment I have added an asterisk on this page.
Correction: Cr[II] compounds: Cr[II] oxalate, CrC2O4.H3O is a yellowish-green powder, not red as stated previously, according to Brauer and CRC.
Examples of other uncommon valences:
Cr{0}: Chromium hexacarbonyl, Cr(CO)6, is a Cr[0] compound is solid at RT but potentially explosive on heating as are all carbonyls.
Not made directly by passing Co over metal: difficult to make.
Brauer also cites a few metalo-organics of Cr[0], like Dibenzenechromium(0).
Brauer represents Cr[I] by Dibenzenechromium(I) Iodide, [(C6H6)2Cr]I, which I guess is valid by the usual rules.
Cr[IV]: Barium Orthochromate (IV), Ba2CrO4; obtained by a high temperature fusion (See Brauer) : BaCrO4 + Cr2O3 + 5Ba(OH)2 = >
3Ba2CrO4 + 5 H2O
Cr[V]: Potassium Tetraperoxochromate(V), K3CrO8, prepared from KOH, CrO3 and H2O2. Doable by an amateur. (See Brauer and Wiki)
Another high temp fusion requiring especial conditions: Barium Chromate(V) (See Brauer)
2BaCrO4 + BaCO3 => Ba3(CrO4)2 + CO2 + BaO2
A side comment: The essential usefulness of Barium compounds means that every one remotely interested in experimental amateur inorganic chemistry
should always have some on hand. It has quite unique properties for an alkaline earth element; eg. it will fix atmospheric oxygen as peroxide; its
sulphate is about as insoluble as you can get; it has a fairly soluble hydroxide and it’s the only way I know of keeping a manganate ion
stable.
The Oxides of Chromium
Some are covered previously. (CrO, Cr2O3, CrO3): vide supra.
Chromium[IV] Oxide (CrO2) and others
All Chromium compounds AFAIK are paramagnetic due to unpaired 3d orbitals but CrO2 is actually ferromagnetic (consult a physics text for the
difference, if necessary!). Hence you can use a physical test, since its magnetic susceptibility is very high. It was used for high quality audio
recording prior to the digital era (remember Chromium tape?). Various ways of making it exist, most probably impure.
(1) Careful heating of CrO3 is said to produce distinct steps as oxygen is disengaged: (See Mellor on Chromium).
The following have been reported but not all verified:
CrO3 -- > Cr5O9 -- > Cr3O5 -- > Cr3O6-- > Cr5O13 -- > Cr5O12 -- > Cr6O15 -- > Cr2O3 as the temperature is raised from RT to
450C. (Note that Cr3O6 = 3CrO2).
Most of these – if they really exist – can be related to mixed oxides of the type [xCrO3, yCr2O3] or [xCrO2, yCr2O3] etc., i.e. Cr{III,VI} etc,,
mixed oxides.
Commercially Dupont produced the dioxide by decomposing chromium trioxide in the presence of water at a temperature of 800 K and a pressure of 200 MPa
(Wiki) – obviously well outside amateur capability.
(2)*It is alleged to be produced by heating chromic nitrate until all the nitric oxides are evolved. Since the nitrate can be easily be made from
Cr[III] hydroxide plus dilute nitric acid, I tried evaporating a solution and gently heating it, whereupon a black solid remains that is, in fact,
attracted to a magnet.
*Another process I have tried is a sort of parallel to the disproportionation exhibited by manganese compounds – specifically the reaction of MnO4-
with Mn++ ions which produces MnO2. For Chromium, any reaction between a chromate and a chromic or chromous ion sulphate might give what could be
called potential Chrom(ous)(ic) (di)chromates. Such possible reactions total four, giving
Cr[II].Cr2O7 = Cr3O7; Cr[II].CrO4 = Cr2O4=2CrO2;
Cr[III]2.(Cr2O7)3 = Cr3O7; Cr[III]2.(CrO4)3 = Cr5O12
I did not try all of these but Cr[II] and CrO4- did give some blackish ppt., along with a heap of zinc hydroxide which obscured everything (from the
reduction with Zinc). Results indeterminate!
Fluorides – Cr gives every valence from 2 to 6, as might be expected.
Sulphides are not made with H2S
……
Enough; labor meus finitus est: I am done.
Der Alte
>>>>>>>
Chromium - Epilogue
Don’t have any Chromium salts? Then use this CRUD method.
(CRUD = Chemical Reagent from Utter Dross, the art of producing reagents from household trash).
Get some scrap stainless steel and dissolve in moderately concentrated HCl (not H2SO4 or HNO3, they will probably just passivate the SS), keeping
metal in excess. Evaporate to dryness but don’t overheat. Just dissolve the chlorides in water and add slowly a solution of NaOH in bleach (5-6%
NaOCl) – the Ni and Fe precipitate as oxides or hydroxides and the Cr gets converted to chromate. Let stand for a day, boil and decant and filter.
Evaporate and crystallize out the Na chromate.
What? Don’t have any HCl? Then try the following, a true MadScience adventure into the absurd, a Der Alte special CRUD encore.
Fix up a crude electrolytic cell with a sacrificial SS anode and a cathode of carbon rods from old drycells. Add common salt, water and electrolyze.
See if you can work out the rest…
Regards
Der Alte
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| Quote: Originally posted by DerAlte | More on Chromates II
Dichromate is frequently used in organic chemistry as a moderate oxidant, but has no use in pyrotechnics AFAIK.
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This is not correct, Potassium dichromate has use in pyro.
1.For Coating metals
2. For use in strobes, strobe rockets, and other strobing devices.
[Edited on 11-4-2018 by AlChemicalLife]
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JJay
International Hazard
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I believe it is also used in thermite boosters, especially for chromium trioxide thermites. It is easy to purify and non-hygroscopic.
Sodium dichromate is more generally useful due to its higher solubility at room temperature, but it's deliquescent.
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AJKOER
Radically Dubious
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The issue of why chromium salts can lead to DNA damage is likely complex owing to multiple valence states feeding possible fenton-type redox reactions
in the presence of H2O2 (as created in our bodies) and HOCl (from chlorine treated water). Also, related arguments related to the creation of excited
states.
Apparently, even a low exposure level alone is not likely a place where some authorities feel they can support (see full discussion at http://www.sciencemadness.org/talk/viewthread.php?tid=89653#... and cited references).
An extract of some of my prior comments:
| Quote: Originally posted by AJKOER |
.....
I especially find interesting the comment, to quote:
"Following the conclusion of the 2013 report, ECHA commissioned an assessment of the mode of action of the cobalt salts, which has concluded that the
cobalt salts are genotoxic carcinogens by inhalation with a non-threshold mode of action; RAC4 supported this conclusion. This work answered one of
the key uncertainties raised in the previous report. "
with the key word being 'non-threshold', which I am interpreting, per my extracted comments above, as relating to dosing level assessment.
[Edited on 13-9-2018 by AJKOER] |
And, also:
| Quote: Originally posted by AJKOER |
......
Per this article, "Advanced Oxidation Process Based on the Cr(III)/Cr(VI) Redox Cycle", by Alok D. Bokare and Wonyong Choi, published in Environ. Sci.
Technol., 2011, 45 (21), pp 9332–9338, DOI: 10.1021/es2 . To quote from the abstract:
"Oxidative degradation of aqueous organic pollutants, using 4-chlorophenol (4-CP) as a main model substrate, was achieved with the concurrent
H2O2-mediated transformation of Cr(III) to Cr(VI). The Fenton-like oxidation of 4-CP is initiated by the reaction between the aquo-complex of Cr(III)
and H2O2, which generates HO• along with the stepwise oxidation of Cr(III) to Cr(VI). The Cr(III)/H2O2 system is inactive in acidic condition, but
exhibits maximum oxidative capacity at neutral and near-alkaline pH. Since we previously reported that Cr(VI) can also activate H2O2 to efficiently
generate HO•, the dual role of H2O2 as an oxidant of Cr(III) and a reductant of Cr(VI) can be utilized to establish a redox cycle of
Cr(III)–Cr(VI)–Cr(III). As a result, HO• can be generated using both Cr(III)/H2O2 and Cr(VI)/H2O2 reactions, either concurrently or
sequentially. The formation of HO• was confirmed by monitoring the production of p-hydroxybenzoic acid from [benzoic acid + HO•] as a probe
reaction and by quenching the degradation of 4-CP in the presence of methanol as a HO• scavenger. The oxidation rate of 4-CP in the Cr(III)/H2O2
solution was highly influenced by pH, which is ascribed to the hydrolysis of CrIII(H2O)n into CrIII(H2O)n-m(OH)m and the subsequent condensation to
oligomers. The present study proposes that the Cr(III)/H2O2 combined with Cr(VI)/H2O2 process is a viable advanced oxidation process that operates
over a wide pH range using the reusable redox cycle of Cr(III) and Cr(VI)."
Link: http://pubs.acs.org/doi/abs/10.1021/es2021704
.....
Per Wiki on Cr(VI), to quote:
"Hexavalent chromium compounds are genotoxic carcinogens. Due to its structural similarity to sulfate, chromate (a typical form of chromium(VI) at
neutral pH) is transported into cells via sulfate channels.[5] Inside the cell, hexavalent chromium (chromium(VI)) is reduced first to pentavalent
chromium (chromium(V)) then to trivalent chromium (chromium(III)) without the aid of any enzymes.[5][6] The reduction occurs via direct electron
transfer primarily from ascorbate and some nonprotein thiols.[5] Vitamin C and other reducing agents combine with chromate to give chromium(III)
products inside the cell.[5] The resultant chromium(III) forms stable complexes with nucleic acids and proteins.[5] This causes strand breaks and
Cr-DNA adducts which are responsible for mutagenic damage.[5] According to Shi et al., the DNA can also be damaged by hydroxyl radicals produced
during reoxidation of pentavalent chromium by hydrogen peroxide molecules present in the cell, which can cause double-strand breakage.[6]"
"For drinking water the United States Environmental Protection Agency (EPA) does not have a Maximum Contaminant Level (MCL) for hexavalent chromium.
California has finalized a Public Health Goal of 0.02 parts per billion (ppb or micrograms per liter)[12] and established a MCL of 10 ppb.[13]"
"Workers who are exposed to hexavalent chromium are at increased risk of developing lung cancer, asthma, or damage to the nasal epithelia and skin.[2]
Within the European Union, the use of hexavalent chromium in electronic equipment is largely prohibited by the Restriction of Hazardous Substances
Directive."
Link: https://en.wikipedia.org/wiki/Hexavalent_chromium
[Edited on 30-9-2017 by AJKOER] |
[Edited on 6-11-2018 by AJKOER]
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teodor
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Cr(OH)3 * nH2O aerial oxidation?
Do you know what is the product of aerial oxidation of the wet Cr(OH)3 * nH2O? On the photo is just Cr dissolved in conc. HCl precipitated with 6M KOH
solution. It is not (well) washed. It takes few minutes to change the colour from green to brown olive on contact with the air.
[Edited on 7-11-2020 by teodor]
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