ssdd
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Chromate Oxidation Colors
Yesterday I ran this experiment:
50/50 ratio of salt and Ammonium Dichromate in a flask.
I then added a few ml of conc. sulfuric acid to produce a red gas.
I poured this red gas into a beaker containing hydrogen peroxide, it turned a nice shade of blue with this.
When the ammonium dichromate was added directly to the H2O2 it turned a deep blue, then to green, and eventually brown. When a small sum of acid was
added and then neutralized a yellow color was obtained.
I know that all these colors are due to the different oxidation states of chromium, but can anyone tell me what colors are associated with what
states?
Just a note: All of this work was done in a fume hood with two other people in case of a problem with the chromium gas or acids. 
Thanks
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sparkgap
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I seem to remember that such was tucked somewhere in woelen's site. You might want to check out his nice pics. 
The blue one is the nice peroxychromate ion that, unfortunately, is not long-lasting. Such a shame for a nice hue. After giving off the oxygen, all we
have left are chromium ions in, IIRC, the 3+ oxidation state.
sparky (~_~)
"What's UTFSE? I keep hearing about it, but I can't be arsed to search for the answer..."
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woelen
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I indeed did this experiment some time ago:
http://woelen.scheikunde.net/science/chem/exps/cro2cl2/index...
I also wrote pages about all kinds of chromium colors. The name of the element already tells you about its colorful properties. Chromium is called
after "chromos", meaning "color".
http://woelen.scheikunde.net/science/chem/solutions/cr.html
http://woelen.scheikunde.net/science/chem/exps/dichrom/index...
Finally, this page may also be interesting for you. It contains a fairly large section on chromium peroxo compounds.
http://woelen.scheikunde.net/science/chem/exps/peroxo/index....
Also a word of warning. Be careful with hexavalent chromium compounds, and especially with the chromyl chloride vapor/gas. Hexavalent
chromium is a carcinogen and exposure to the chromyl chloride is very easy, because it is a gas. With the solid dichromates the chance of exposure is
much less, provided you work cleanly and carefully.
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phj
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I've done this experiment just a little while ago.
Remarkable is the fact that the deep blue CrO5 yields a yellow/brown solution of tetraperoxochromate when some strong base is added. (NaOH)
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ssdd
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Thanks for the links, Ill have to take a look at them when I get home my school thought it would be amusing to block woelen's site. 
phj: the tetraperoxochromate explains why I got the yellow when the solution was nutralized. I was given a breif burst of yellow which quickly turned
back to blue, it seems this was because the tetraperoxochromate was breaking down (or according to what I just read) into hydrogen peroxide and some
other materials. When the H2O2 is introduced it changes the solution back to blue.
intresting... why is the tetraperoxochromate so unstabe it seems?
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woelen
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No, the change, back to blue is not because of breakdown of peroxochromate. It probably was because total amount of acid was larger than the amount of
base. If you add much more base, then you'll see that the color does not revert to blue.
[Edited on 24-5-07 by woelen]
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ssdd
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Ahhhhh... I did as you said and added a larger sum of acid to the leftovers from my experiment. And yes it stayed a green/yellow.
Sorry I had misinterpreted what I had read.
Now I have gotten a chance to check out the links provided, they make things much clearer... thanks.
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woelen
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If you like colorful reactions, with permanganate you can also do funny things:
http://woelen.scheikunde.net/science/chem/exps/chameleon/ind...
In fact, with almost every transition metal you can do remarkable and colorful reactions. That is one of the reasons why I am so interested in the
chemistry of the transition metals .
[Edited on 24-5-07 by woelen]
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phj
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| Quote: | Originally posted by woelen
In fact, with almost every transition metal you can do remarkable and colorful reactions. That is one of the reasons why I am so interested in the
chemistry of the transition metals .
[Edited on 24-5-07 by woelen] |
Amen.
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not_important
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Chromium and vanadium are the most colourful with simple reactions and compounds, when you move to complexes cobalt gives a wide range. Manganese does
well, but it's a little more trouble in reaching the various oxidation states.
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ssdd
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I just ran the Potassium Permaganate experiment you linked me to woelen.
It was by far a bit more impressive than the chromium reaction, at least in my eyes, because it was faster and over a wider range of colors. 
Wonder if I can find any vanadium compouds to work with in my schools chem room. Any ideas of what ones may work best?
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woelen
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Vanadium also is wonderful. Start off with vanadium pentoxide, which is easy to obtain in almost every potteries supply store for giving yellowish
colors to glass and ceramics.
If you don't have vanadium pentoxide, then ammonium metavanadate or sodium metavanadate can be used.
In any case, dissolve the solid in a warm solution of NaOH (5% concentration or so, this is not critical) and then strongly acidify this solution.
Have a look at the very special properties of the vanadium in the +5 oxidation state:
http://woelen.scheikunde.net/science/chem/solutions/cr.html
The other oxidation states also are quite interesting. This experiment describes how to make them:
http://woelen.scheikunde.net/science/chem/exps/vanadium/inde...
Have fun with the nice colors, but be careful with solutions of vanadium salts. They are quite toxic, much more so than solutions of manganese salts.
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phj
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| Quote: | Originally posted by not_important
Chromium and vanadium are the most colourful with simple reactions and compounds, when you move to complexes cobalt gives a wide range. Manganese does
well, but it's a little more trouble in reaching the various oxidation states. |
Cobalt's redox chemistry is a bit limited, but has some remarkable complexes.
Chromium has some colorfull complexes too, but has a richer redox chemistry.
Experimenting with manganese is trickier.
The only complex of magnesium I heard of is the faint pink tetrachloromanganese(II) complex.
Further chemistry of manganese is entirely restricted to redox.
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12AX7
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As far as chromium, right now I'm fascinated by its bright orange -- almost red -- potassium dichromate.
I'm also getting a lot of calcium chromate (or possibly dichromate). Bastard stuff sticks to glass badly! Damn our hard water.
Tim
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not_important
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Vanadium isn't real toxic, the main problem is with cronic exposure rather than acute poisoning. People have taken 25 to 50 mg/day of soluble
vanadium salts, for testing as a diabetes treatment, and suffered no ill effects; high doses lead to digestive upsets, green tongue, and fatigue. The
toxicity by ingestion is much lower than by inhalation, which is the main source of cases of vanadium poisoning. People who work making or using the
pentoxide, ceramics workers, working with V catalysts, servicing equipment were oil or coal is burned (fossil fuels carry traces of vanadium). The
body clears itself of vanadium fairly quickly, possibly because vanadium is an essential element so there are mechanisms to handle it.
Vanadium is more toxic that Mn(II) or Cr(III), but less toxic than the higher oxidation states of those elements; chromium(V) and (VI) are
carcinogenic. Vanadium is also less toxic than copper or nickel. So take normal care and you'll be alright. Do not use V2O5 as a dessert topping.
Yes, cobalt is limited to Co(II) and (III) for simple salts, in complexes you get +2, +3, and "+2.5". Its complexes come in just about every colour.
A solution of MnCO3 in aqueous KCN will give a yellow solution that slowly gives crystalline K4[Mn(CN)6]. Blow air or oxygen through the mixture it
will turn red then red-orange; addition of alcohol crystallises out K3[Mn(CN)6]. These are equivalent to ferro- and ferri- cyanides, but are less
stable.
Mn(III) also has several light-sensitive complexes with oxalic acid. Mn(II) forms few complexes, but does form a number of double salts such as
K2Mn(SO4)2.2H2O.
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woelen
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| Quote: | Originally posted by phj
The only complex of magnesium I heard of is the faint pink tetrachloromanganese(II) complex.
Further chemistry of manganese is entirely restricted to redox. |
All transition metals have a very rich coordination chemistry, including manganese. But much of this rich chemistry is at low oxidation states.
Manganese has a rich chemistry in its +1, 0 and -1 oxidation states, like most metals. Unfortunately, experiment with this is very hard for the home
chemist. Ligands, used with such low oxidation states are CO, PH3 and organic derivatives, C5H5(-) and other exotic (often bulky) organic ligands.
Manganese does have some interesting coordination chemistry, also for the home chemist, in its +3 oxidation state. A beautiful very brightly colored
purple complex is the pyrophosphato complex of manganese (III). This can be made by heating a small amount of any manganese (II) salt, added to
concentrated phosphoric acid, to which a tiny pinch of an oxidizer is added (e.g. KNO3, KClO3). On dilution with water, this complex
disproportionates, giving brown hydrous MnO2 and a manganese (II) salt.
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