| Pages:
1
2 |
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
chromium compounds
Hi every one.recently I searched for making chromium salts from dichromate from sceincemadness but nothing
I Know how to make chrom alum,but this is a double salt.the separation of chromium sulfate from pottasium sulfate is the question.there is a way to
adding ammonia to gain the Cr(OH)3.the trouble is this hydroxide contains the inert trioxide of chromium and would not dissolve in nitric or
hydrochloric acid,ofcourse the hydroxide would dissolve but it is not economical at all in my country Iran and if I want to buy the chromium salts,the
american and the germanian are just exist in iran and I should buy them by pay a high price!!
Recently I found the chromium sulfate with a low price,but it is fairly soluble in water.I think it is the anhydrous form of it according to the
wikipedia but my sulfate was green powder and the wiki states that the anhydrous is reddish brown!
I want you guys to tell how to separate the chromium sulfate and potassium sulfate from the solution of reducted dichromate by sulfuric acid and
ethanol and guess what is my mysterious green compound?
please post your suggestions.sorry for the bad english language
|
|
|
Metacelsus
International Hazard
Posts: 2531
Registered: 26-12-2012
Location: Boston, MA
Member Is Offline
Mood: Double, double, toil and trouble
|
|
Are you trying to make a chromium(iii) or chromium(vi) salt?
The green color is due to chromium(iii). Chromium(vi) is orange.
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
chromium 3 is the target.separate the chromium sulfate from the potassium sulfate solution and obtain the Cr2(SO4)3.n H2O
I know the green color is due to chromium 3 but it is not soluble in water just a little
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
I ask this question from an expert of chemistry lab in my college but he said I dont konw how to separate the chromium sulfate from pottasium sulfate
solution!! fool!!I dont know how he is a professor and working in the lab!!
|
|
|
elementcollector1
International Hazard
Posts: 2684
Registered: 28-12-2011
Location: The Known Universe
Member Is Offline
Mood: Molten
|
|
Seriously? Lighten up! Just because he's a professor doesn't mean he knows the answer to whatever question you might ask. This isn't a matter of
general knowledge of chemistry, it's more a required specific knowledge of chromium and its salts.
Your 'mysterious green compound' is likely chromium sulfate hydrate, hexahydrate if I remember correctly.
I'm not sure I'm remembering this right, but can't double salts be 'broken' via acidification?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
I don't quite understand why precipitating chromium hydroxide, filtering, washing, and redissolving in sulfuric acid would not work for you. This
should separate insoluble chromium hydroxide from soluble potassium hydroxide and return you to a solution of hydrated chromium sulfate alone. I
haven't tried this myself so I can't say from experience.
Wikipedia states "heating chromium(III) sulfate leads to partial dehydration to give a hydrated green salt," and based on other information it seems
to me that the full hydrate is reddish brown, partial hydrate is green, and anhydrous is purple.
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
last night I put some of that green powder in a glass,add a lot of water,didnt dissolve.I add some calcium chloride to the solution because I thought
that if that powder was chromium sulfate,then it would react with CaCl2 in the solution:Cr2(SO4)3(s)+3CaCl2(aq)=2CrCl3(aq)+3CaSO4(s)
today I saw that some white powder(CaSO4) and a lot of unreacted green powder was in the bottom of glass and the solution was deep green maybe due to
the CrCl3 solution.(probably the green powder is the chromium sulfate because of the creation of white powder at the bottom of glass,maybe calcium
sulfate that is not soluble in water)
But...but why its not soluble in water if it is sulfate??
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Sasan:
I think you should really start all over again because something isn't quite right here.
Reduce your potassium dichromate with sulphuric acid and ethanol or methanol. Allow to cool. Add strong ammonia SLOWLY until pH is about 8 - 9. The
precipitate formed is Cr(OH)3.nH2O. Filter it off and wash it with copious amounts of water.
This hydroxide, if you've done things properly, will dissolve in any strong, dilute acid to form a solution of the corresponding Cr (III) salt. A
little gentle heat may be required.
All this is straightforward and should pose you no particular problems whatsoever.
[Edited on 6-3-2014 by blogfast25]
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
Ok adding ammonia is the best path to Cr salts reacting the hydroxide with acid
I put some aluminium powder in the reducted solution of dichromate,because I thought that it can make chromium and the Al goes to solution instead of
Cr.but after a day all of Al was at the bottom of reducted solution.is it possible Al goes to the solution instead of Cr3+?
At the end I think adding ammonia is the best selection and obtain the hydroxide
Please post your suggestions.sorry for bad language
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Chromium is very peculaliar and unfortunately things are not as easy as blogfast25 writes. His suggestions can be followed, but great care must be
taken.
Heating a solution of a chromium(III) salt in many cases leads to problems. Chromium(III) in such cases forms complexes with anions which are very
hard to break and may take weeks or even months to break again, even in the presence of ammonia or hydroxide.
Reduction of the dichromate with ethanol or methanol must be done very carefully, with cooling and good stirring. If the solution becomes warmer than
appr. 40 C, then a green sulfato-complex is formed, which is amazingly hard to crystallize and amazingly hard to break down. If the reduction is done,
such that the liquid remains cold, then you end up with a dark purple/grey solution and from that you can prepare quite pure Cr(OH)3 with ammonia. If
the solution is green, then the sulfato-complex is formed.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Woelen, I don't dispute what you say and maybe I've just been 'lucky' but I've never encountered these problems, having extracted Cr(III) quite a few
times. Perhaps using acetic acid (instead of sulphuric) to effectuate the reduction of dichromate solves that problem for once and for all?
I do have a sample of Cr alum solution that I overheated and months later there's still no sign of it crystallising.
[Edited on 7-3-2014 by blogfast25]
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
seems that every body have problems with chromium,woelen is right,chromium make such a strong complexes that Al powder cant reduct the Cr(3+) to form
elemental chromium
If I use nitric acid instead of sulfuric,can Al goes to the solution instead of Cr(3+)?maybe there would no any strong complexes of chromium with the
sulfate anions
Another question is that the reaction of HF(aq) and the K2Cr2O7 can release flourine gas?like reacting with HCl(aq) that releases chlorine vapors.if
no then what would be happen?or no reaction occurs???
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
HF and K2Cr2O7 do not react. Fluorine is one of the strongest chemical oxidizers and no common chemical is capable of oxidizing fluoride ion to
fluorine gas. Not even the strongest available oxidizers like perxenates or peroxodisulfate or ozone can make fluorine from fluorides.
|
|
|
chornedsnorkack
National Hazard
Posts: 522
Registered: 16-2-2012
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by sasan  |
Another question is that the reaction of HF(aq) and the K2Cr2O7 can release flourine gas?like reacting with HCl(aq) that releases chlorine vapors.if
no then what would be happen?or no reaction occurs??? |
Something may happen depending on the HF/H2O ratio.
HF is notorious for being a weak acid when dilute yet very strong when concentrated. Whereas H2Cr2O7, while a strong acid is not very strong, and
nowhere close to H2SO4 or concentrated HF.
First step:
K2Cr2O7+2HF<->2KF+H2O+2CrO3
Second step:
CrO3+2HF<->CrO2F2+H2O
What are the respective HF/H2O ratios where the first and the second reaction go to right?
Chemical synthesis of fluorine is notoriously not straightforward. It can be done, but in two steps not one. And the oxidant is manganese, not
chromium.
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
Failed attempt at hexaamminechromium(III)
I attempted to produce the compound hexaamminechromium(III) sulfate using chromium sulfate and household ammonia. Instead, I got a bluish-grey
precipitate that doesn't redissolve on the addition of excess ammonia. From what I've read chromium complexes tend to be non-labile and substitution
of the coordinate sphere occurs only slowly.
What other interesting coordination compounds of chromium can I make in a school lab?
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by Brain&Force | I attempted to produce the compound hexaamminechromium(III) sulfate using chromium sulfate and household ammonia. Instead, I got a bluish-grey
precipitate that doesn't redissolve on the addition of excess ammonia. From what I've read chromium complexes tend to be non-labile and substitution
of the coordinate sphere occurs only slowly.
What other interesting coordination compounds of chromium can I make in a school lab? |
If you want it to work, you need to reflux the chromium(III) sulphate in ammonia with a catalytic amount of zinc. Chromium(III) complexes are very
non-labile; chromium(II) complexes are labile and will substitute quickly.
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.
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
I had never thought of making a chromium(II) complex - I'll try that instead. I'll also see if I can plate chromium metal using a chromium(III)
solution. I know most industrial processes use chromates instead, but I just want to figure out if it's possible.
And I'll also try to make a Cr(OH)63- complex with NaOH.
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
chornedsnorkack states that: First step:
K2Cr2O7+2HF<->2KF+H2O+2CrO3
Second step:
CrO3+2HF<->CrO2F2+H2O
the reaction of chromyl chloride with water at first makes the blue complex of chromium oxyperoxide then decomposes to Cr(3+) with bluish green tint
Chromyl flouride like chromyl chloride decomposes in the water to Cr(3+) maybe chromium flouride
Brain&force,see handbook of "Handbook of preparative inorganic chemistry"VOLUME 2-second edition edited by Gerog barauer,there you will see a lot
of chromium complexes in its all states specially in the (3+).It will be very helpful for you to make complexes of chromium,search it in google and
download it
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
I'm running a few experiments to check whether complexed Cr<sup>3+</sup> can be reduced with Al or not. Quite surprising results so far.
But more tests are needed.
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
I think Al (like Zn) can reduce Cr3+ into Cr2+.not to the elemental chromium,but Im not sure
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Check the Electrochemical Series: 'naked' Cr<sup>3+</sup> is reduced by Al to Cr (0) with volts to spare (so to speak). But so far I've
not seen any reduction. More tests to follow 'soon'.
[Edited on 11-3-2014 by blogfast25]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Here are some results from attempts to reduce Cr(III) with aluminium.
Firstly a bit of chromium alum solution was prepared at room temperature (RT) and to half of it was added some ammonium sulphate and some heat: the
solution turned dark green. On cooling it didn’t revert back to the original colour.
Some of that solution, as well as some of the cold Cr alum solution, was loaded into separate test tubes and a piece of aluminium sheet, lightly
sanded, added to each. This is what that looked like but bear in mind strong colour distortion due to flash (left is green solution, right is blue
solution):
After over a week (second photo), no change was noted in EITHER of the tubes, clearly no plating of Cr(0) had taken place (more about the third tube
below). I have to conclude tentatively that in BOTH cases the chromium is complexed too strongly to be reduced by Al.
I then went on to prepare a solution of Cr(NO3)3 as follows. The required amount of KCr(SO4)2.12H2O was dissolved in about 400 ml of water, at RT and
the Cr precipitated as Cr(OH)3.nH2O with the equivalent amount of 33 % NH3. The precipitate was filtered and the filter cake washed several times on a
Buchner and sucked dry.
Quantitatively transferred into a large beaker, it was dissolved in the required amount of 70 % HNO3, yielding a deep blue solution which was diluted
to about 250 ml. The molarity of the Cr(NO3)3 solution was about 0.35 mol/L. The deep blue is of course caused by the
[Cr(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup> complex ion.
The third tube in the second photo is that solution (slightly diluted), with a piece of Al added, to see if reduction take place. Nitrate ions
aren’t suitable ligands for Cr(III). So far (a few hours) no change has been observed.
A couple of simple tests were carried out too. Two tubes loaded with the solution, slightly diluted, were heated on steam bath. No change in colour
was observed. To the first (left in the picture below) was then added a good pinch of ammonium sulphate, to the second (centre in the picture below)
ammonium chloride. Third tube is Cr(NO3)3 + aluminium.
Left (sulphate) and centre (chloride) tubes changed colour to deep green quite quickly:
This suggests ligand exchange.
Another interesting observation was made by adding an excess 33 % NH3 to the two green tubes. The one with the suspected sulphate complex turned mauve
immediately (but not all precipitate has dissolved yet). The one with the suspected chloride complex first precipitated a green precipitate, somewhat
later this then turned mauve too. The mauve is due to the Cr(III) ammonia complex and these results seem to suggest the strength of complexation is
NH3 > Chloride > Sulphate > H<sub>2</sub>O.
[Edited on 15-3-2014 by blogfast25]
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
I told you that it would not reduce with Al,just to the Cr2+ that is blue,for certainty if you add sodium acetate then red chromous acetate will be
form but if no precipitate of acetate occurs,so Al cant reduce Cr3+ even to the 2+
blogfast try reducing with Mg,but I think maybe it would react with water to from hydroxide in precence of other ions.
I should purchase Chromium powder,treating with hydrochloric to form chromous chloride,add Al to see what will be form
I`m not professional in electrochemistry but I think electrolysis the solution of Cr3+ with Al/Cr(s) electrodes can be useful.sorry for bad english
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Sasan:
If complexation is the cause then reduction to Cr(II) won't work either. I'll check my tubes with acetate later on. Both green tubes didn't go blue
though.
Last night I also tried Mg (with an even greater oxidation potential compared to Al) but it didn't seem to work: after chewing up the acid reserve in
the solution, Cr(OH)3 started precipitating.
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
Chromium like his uncle(!!)manganese is somewhat complicated
I think Mg would not react with water only in the presence of other ions in water to form Mg(OH)2 and because of this Cr(OH)3 precipitated instead of
Mg(s) + Cr(3+) = Cr(s) + Mg(2+),otherwise magnesium is much more electropositive than chromium
wikipedia:An aqueous solution of a Cr(III) compound is first reduced to the chromous state using zinc.[5] The resulting blue solution is treated
with sodium acetate, which results in the rapid precipitation of chromous acetate as a bright red powder.
2 Cr3+ + Zn → 2 Cr2+ + Zn2+
2 Cr2+ + 4 OAc- + 2 H2O → Cr2(OAc)4(H2O)2
The synthesis of Cr2(OAc)4(H2O)2 has been traditionally used to test the synthetic skills and patience of inorganic laboratory students in
universities because the accidental introduction of a small amount of air into the apparatus is readily indicated by the discoloration of the
otherwise bright red product
Producing chromous acetate wont be easy because of air oxidation,watch here:www.tulane.edu/~inorg/.../Experiment%203.pdf
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
