Chromium Revisited
Chromium Revisited
Chromium Revisited
(3a) Heat a small amount of the violet (purple by transmitted light) solution of Cr[III]
sulphate to boiling for a few minutes. The solution turns from violet to deep green.
On cooling it preserves the green color. It will stay green for several days to weeks
but finally return to the original violet. The green form is chemically different from
the violet form as can be shown by precipitation of the sulphate ion with a barium
salt, and also by electric conductivity or freezing point measurement. The green salt
does not form a precipitate. The mystery was solved in the early 1900s; the violet
ion is the complex Cr(H2O)6+++; the dark green one maybe [Cr(H2O)4(SO4)]+; or
perhaps [Cr2(H2O)l0(SO4)]++++ or the covalent [Cr2(H2O)6(SO4)3], or a mixture
of all three.
Only the violet form will crystallize. If the green form is heated to dryness, only
green scales result. Further cautious heating produces an orange-yellow product
which is the anhydrous form with perhaps some decomposition. Strong heating
(bright red heat) causes decomposition and SO3 evolution: Cr2(SO4)3 Cr2O3 +
3SO3
For the chloride of Cr[iii], Wiki has:
Chromium(III) chloride (also called chromic chloride) is a violet coloured solid
with the formula CrCl3. The most common form of CrCl3 sold commercially is a dark
green hexahydrate with the formula [CrCl 2(H2O)4]Cl.2H2O. Two other hydrates are
known, pale green [CrCl(H2O)5]Cl2.H2O and violet [Cr(H2O)6]Cl3. This unusual feature
of chromium(III) chlorides, having a series of [CrCl 3n(H2O)n]z+, each of which is
isolable, is also found with other chromium(III) compounds.
This indicates the facility with which Cr forms complexes, in this case with water. It
is also interesting to note that both Cr[III] sulphate and chloride, if anhydrous, are
virtually insoluble in water. The crystalline form of sulphate is Cr2(SO4)3.12H2O or,
according to the above, [Cr(H2O)6]2(SO4)3.6H2O, and is very soluble.
(4) Cr[III] Compounds
{The temptation to insert a Pourbaix or Eo/pH diagram here was great if you
understand them they tell a lot at a glance. But many do not and it might merely
confuse.}
The salts of Cr+++ with strong acids all give an acid reaction in solution in other
words Cr(OH)3 is a weak base; being a bit more technical, the ion Cr+++ has a pKa
of about 4 whereas a strong base ion like Na+ has pKa~14.
(4a) Take a test-tube of the violet sulphate solution and drip in some fairly
concentrated solution of NaOH (or KOH). Until the hydroxide neutralizes some of
the acidity nothing happens, then a flocculent precipitate of a light gray-green color
settles out. If too much hydroxide is added, this precipitate re-dissolves. The
precipitate is hydrated Cr[III] hydroxide, Cr(OH)3.xH2O. If instead the green form of
sulphate is used, the formation is slower but the same precipitate obtained. If done
carefully, the liquid becomes colorless and depleted of Cr[III] ions. If the precipitate
is left long enough (or centrifuged) the color becomes a very distinct green color
and compacts, probably due to loss of water.
Reaction is simply Cr+++ + 3(OH)- + xH2O Cr(OH)3. xH2O (s); x is approx. 3
when dried carefully. In alkaline solutions Cr[III] always exists as the hydroxide,
except when the pH is high.
(4b)Addition of excess hydroxide causes the precipitate to redissolve giving a light
green solution. This is usually stated to be due to chromite ion Cr(O2)-, or a
hydrated form of this. This is still a Cr[III] ion, probably O=CrO-. Reducing the pH
with acid re-precipitates the hydroxide.
(4c) Using NH4OH a similar precipitate with a grayish color is first formed. As further
hydroxide is added carefully, again the liquid will clear. If excess ammonia is added,
and the solution left for an hour or so, this gelatinous ppt. will turn a light purple,
quite different in color from the sodium hydroxide case. On standing for a few days,
or on prolonged boiling, the purple color disappears as ammonia is lost, leaving a
gray-blue hydroxide. My guess is that the purple is a pentamminechromium
complex; [Cr(NH3)5.(OH)x](3-x)(OH). yH2O.
{The colors of the Cr[III] ammine chloride complexes are; tri+++, yellow; tetraCl2+,
green;
pentaCl++, purple; non ionic tri Cr(NH3)3Cl3, violet and the ion
Cr(NH3)2Cl4-, orange-red (Pauling, General Chem.)}.
See Brauer, p1345ff, for more on hydroxides and the organic complexes of Cr[III].
(4d) The hydroxide has a very low solubility product in water, about 6x10^-31. It
can be purified by washing with large amounts of water and can be dried as a dull
green powder, hydrated, and is a useful source of Cr[III] salts with common mineral
acids. On heating to red heat it loses water to become a bright green oxide,
2Cr(OH)3 - -> Cr2O3 + 3 H2O. This oxide, the most stable form of Cr[III] under mild
conditions, is usually insoluble in acids but can be reacted with fused alkali
hydroxides to form chromates see later on. Structure is O=Cr-O-Cr=O.
Most of the Cr[III] salts exist as blue-violet or green types when hydrated. As large
crystals the violet forms look black; as powder, reddish. Chromium and potassium
sulphate are isomorphous and crystallize together in all proportions, giving crystals
that vary from lightest pink to almost black. Chrome alum is the equimolecular
mixture. It can form massive octahedral crystals with patience I have grown them
up to 7cms. long looking as black as coal.
Insoluble Cr[III] salts are the normal phosphate {blue} CrPO4, can be hydrated;
sodium carbonate solution produces a similar precipitate to hydroxide that may be a
basic carbonate or Cr2(CO3)2.xH2O. The color is purple Cr[III] also produces CO2
complexes.
Correction to Part 1 Above.
sodium carbonate solution produces a similar precipitate to hydroxide that may
be a basic carbonate or Cr2(CO3)2.xH2O. The color is purple Cr[III] also produces
CO2 complexes..
I re-ran this one today and now find that the color quoted is nearer lavender or light
blue gray. Nor does it dissolve in excess Na2CO3. A short boiling does not convert
the color, except to maybe reduce the bluish tint to slightly more gray. Try it and
see!
But it is not green as several texts have it. (But I have not yet tried the green form
of the sulphate) All texts also say this is not a carbonate but a hydroxide and on
further effort I tend to agree. If the precipitate is washed many times with water, to
eliminate carbonate contamination, and treated with acid, very little CO2 is
evolved. It dissolves to give the violet form of Cr[III]+++ aqua complex.
I think the difference in color is probably due to crystal size, a frequent cause of
such phenomena. As for CO2 complexes, this too is ruled out. I think this is a Der
Alte aberration; I must have been thinking of the carbonyl, Cr(CO)6.
Remembering that the (violet) ion is [Cr(H2O)6]+++ and is acidic, we have in water
[Cr(H2O)6]+++(aq) + H2O(liq) < -- > [Cr(H2O)5(OH)]++(aq) + H3O+(aq);
So one would expect Na2CO3 + Cr+++ salt to produce CO2 provided the pH is low
enough. It does (best seen in conc. Cr salt solution, else the CO2 remains dissolved).
As the salt is gradually neutralized by OH- ions from the dissolved Na2CO3 or NaOH
a sequence of exchanges between the H2O ligands and the OH- ligands can be
assumed to take place:
[Cr(H2O)6]+++(aq) => [Cr(OH)(H2O)5]++(aq) => [Cr(OH)2(H2O)4]+(aq)
=> [Cr(OH)3 (H2O)3] (Solid precipitate);
And if the pH exceeds ~11 afforded by Na2CO3 in solution, as in the case of NaOH
solution, this precipitate dissolves to form negative aqua complex ions:
=> [Cr(OH)4(H2O)2] - => [Cr(OH)5(H2O)] - - =>
[Cr(OH)6] - - -
102V, primary sodium chromite is formed, whilst above 102V the soln. contains also
tertiary sodium chromite. Potassium chromite is similarly produced ; below 82Valkali only the primary chromite is formed, whilst above 82V the soln. contains also
secondary chromite. From soln. of potassium chromite which have stood for a long
time, needle-shaped crystals of the formulaCr2O3.3K2O.8H2O have been obtained.
(V stands for the inverse of concentration, liters/mol, IIRC). Best of luck if you want
to try it!
The precipitated hydroxide aqua complex does dissolve in strong NH3 solution, in
spite of its weakness as an alkali. It switches water for NH3 in the hexa coordinated
complex,
[Cr(H2O)6]+++
For such beauties as Hexaamminechromium (III) Chloride, etc., see Brauer you are
on your own.
The overall message is that Cr[III] preferentially forms hexa-coordinated
compounds, ionic and covalent, unless anhydrous. And anhydrous CrCl3, eg, is
insoluble if pure and covalent.
I performed the original experiments over a few months last year. These notes come
from that record.
(see
www.argentumsolutions.com/publications/CorrVol39p4881983.pdf
for
a
discussion of issues re Eo/pH diags for Cr/water: it has several but they cannot be
copied secured pdf)
In other words, CR(OH)3 is metastable at RT and the equilibrium lies well to the
right in the above equation. You can show this experimentally as follows. {There, I
knew I ought never to have mentioned Pourbaix diagrams.}
(4e) This takes a bit of time to carry out. One must first produce a working quantity
of Cr2(OH)3.xH2O from a Cr+++ compound as outlined previously aim to get 5-10
g reckoned as unhydrated. Wash copiously with water several times. The let settle
and decant, finally filtering through a fine mesh filter paper. Dry carefully in an oven
at ~ 120-140C for several hours, until the mud becomes a powder of a dull green
color.
Try reacting some of this with dilute acid it should react fairly rapidly and give the
violet form of the Cr[III] salt. Scrape the powder from the filter paper and place in an
evaporating dish and heat in the oven to around 250C. This still does not totally
dehydrate the hydroxide but also does not seriously convert it to oxide. This can be
proven by reaction with acid or strong alkali. Now take the powder and heat slowly
on an iron spoon over a propane flame. Somewhere around a dull red heat (500600C, I would guess) if you are lucky you may see it caloresce suddenly brighten
and then dim. {This did not work the first time I tried it it needs a fair quantity (at
least 5g) to work, IMHO}. Now heat strongly. On cooling the power still has a dull
green color but it almost totally non-reactive to strong acid or alkali solutions,
showing that Cr2O3 is very unreactive.
The experiment demonstrates that Cr(OH)3 is changed to the oxide by heating
strongly enough, but is not exactly unstable. In fact it can be kept indefinitely at RT
in the dried state without losing reactivity. The calorescence is probably due to a
change of state (entropy change) as the amorphous structure left on ignition of the
hydroxide changes to crystalline Cr2O3 at some transition temperature..
Other ways to make chromic oxide are to heat ammonium dichromate (NH4)2Cr2O7
or K2Cr2O7 with sulphur. Both ways tend to produce a bright green product,
especially the second (which is fun to perform!).
(NH4)2Cr2O7 -- > N2 + 2H2O + Cr2O3
K2Cr2O7 + S - -> K2SO4 + Cr2O3
(4d)Heat some solid potassium hydroxide (or NaOH) with about the same quantity
of potassium nitrate (or NaNO3; chlorates can be used) in a nickel (or SS) crucible
or dish until melted together. Add some Cr2O3 made by one of the above methods
or otherwise available. It reacts to produce Chromate, K2Cr2O4, a yellow substance.
The Cr[III] state is oxidized to the Cr[VI] state easily under the extreme alkaline
conditions even air will do this, slowly. More of this later.
It is worth noting that Chromium is quite a reactive metal, slightly less so than Zinc
but more so than iron, yet chromium is plated on to iron (over Ni or Cu) and is very
resistant to corrosion. This is believed to be because of Cr2O3 formation. The metal
does dissolve in fairly dilute HCl but nitric acid passivates it, and sulphuric also has
little effect, and even when concentrated has to be heated to react with any speed.
Chromium Revisited Part IIIa
IUPAC Gold Book defines a transition element as one whose atom has an incomplete
d sub-shell, or which can give rise to cations with an incomplete d sub-shell]. The
electronic structure of Cr in the ground (unionized gaseous) state is [Ar].3d5.4s1; it
has six valence electrons outside the argon electron structure.
For technical reasons (as the physicists are wont to say whenever the explanation is
too involved in a cloud of Quantum Mechanics) the 4s level has a lower energy level
than 3d in the atoms K, atomic number (Z) =19 and Ca Z=20. That is, 4s is bound
more strongly to the nucleus. And hence the d levels are not occupied first. The
simplified explanation is that electrons in the s orbitals penetrate the Ar shells and
hence see a higher part of the nuclear charge, while the d orbitals do not. As Z
increases the Ar core decreases in size due to the increased nuclear charge and the
4s levels begin to approach the 3d levels.
The first appearance of the 3d orbitals is at Sc, [Ar].3d1.4s2; past Zn, [Ar].3d10.4s2,
the 3d and 4s shells are filled and the 4p subshell starts at Ga. At vanadium Z=23
the 4s and 3d levels approach equal energy, and at Cr Z=24 the 4s electron actually
has a higher energy by 0.96 eV (Pauling). This means the 4s electron is lost first in
compounds of Cr.
The 3d block elements have no chemically similar elements of lower atomic weight
and represent a new series with new chemical properties. Neither zinc nor scandium
ions have any colored compounds AFAIK (unless with transition metal anions, of
course!) although they are part of the d block elements. They can form complexes,
of course.
I dont consider them true transition elements, only those from Ti to Cu. These
elements form at least one chemically stable ion with a partly filled d sub-shell. The
d orbitals can, due to the Pauli exclusion principle, each have two electrons with
opposed spin, but the d shell first fills up each with only one in each, due to the fact
this is the lowest energy state and these are parallel rather than opposed.
Ti has configuration [Ar].3d2.4s2, Cu has [Ar].3d10.4s1. But Cr does not follow the
usual Building Rule (in German Aufbau) for the electronic configuration. Vanadium
-- >
Cr++(aq)
Eo = -0.42 V
Zn++(aq) + 2 e-
So 2Cr+++(aq) + Zn(s)
reaction
< -->
Eo = 0.782 V
2Cr++(aq) + Zn++(aq)
dE 0.386 V overall
However, if we look at the overall reaction as (sulphate case with H2SO4, Zn)
Cr2(SO4)3 +Zn -- > 2CrSO4 + ZnSO4, it looks as if adding acid is not necessary.
But it is; so also
2H+(aq) + 2e- - - >
H2(g)
Eo=0.000V
Eo=0.782V
dE = + 1.64 V
oxidation of Cr++
Eo = 0.000 V
Cr2+ E0 = - 0.410 V
d E = + 0.410 V
reduction of H+ in
These are energetic, especially the oxidation, so air must be somehow excluded.
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 off 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].
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.
Chromium Revisited Part IV
The current version of the Periodic table is in some respects less genuinely chemical
in flavor than Mendeleefs 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 or get mnOXY.doc from scipics).
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 (Ill leave it to you to work out the equation!).
(6d) Neutralize the yellow solution produced in (6c) to slightly acidic 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) - (havent 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.
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. (Theres 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)
-- >
Yellow liquid
Temp C = 0
30
60 100
Na2CrO4
32
88
115 126
Na2SO4
41
45
43
36
36
37
39
NaCl
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; its 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 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 (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 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.
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 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 havent 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
More on Chromates II
[b](7d)Cr2O7- - to Cr+++ Reduction using alcohols[/b]
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 dont 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).
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 heating gently if necessary until dissolved and
digest 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:
+ H2CrO4
-- > K2SO4
H2O
CrO3
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 I 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[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. But no Chromium - Epilogue
Dont 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 (~2030%) (but not H2SO4 or HNO3, they will probably just passivate the SS), keeping
metal in excess. Pour off liquid, evaporate to dryness but dont overheat. Just
dissolve the chlorides in water and add slowly a solution of 20% NaOH in bleach (56% 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? Dont 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
Der Alte