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DerAlte
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Permanganates
I don’t know if there is any interest still in making KMnO4. There was a flurry of activity some time ago. I have developed a method that far excels
the traditional KOH plus MnO2 plus (oxidizer +) heat method. Any interest? It is a wet method using easily obtained chemicals. The gist of it is to
oxidize an Mn compound (several can be used) to the MnO4+ ion under somewhat carefully controlled conditions. It will take a fair time to explain, so
I wanted to see first whether any interest was out there ,,, Der Alte
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The_Davster
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Sure, why not. We have a big thread on this somewhere, I'll merge this with it as soon as I find the other one.
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DerAlte
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I saw the previous thread, but some time back. The industrial method of heating MnO2 with KOH and using air as an oxidizer is doomed to failure. The
industrial methods are all secret - but they carefully control temperature and blow air thru the melt. It is a fact that KMnO4 decomposes around 220 C
or so. I've tried this, years ago (~50!) and you get a bit of manganate, enough to color the solution green. People do not realize that even a
relative low solubility permanganate like KMnO4 when saturated produces a BLACK solution, only the meniscus being purple.
That being said, the amateur must seek other methods to get a reasonable yield. It ain't easy! But it can be done...
DerAlte
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DerAlte
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To make a permanganate, one needs a source of Mn, an oxidant, and a method. The process I am about to discuss is a wet method in H2O solution, at
modest temperatures. The Mn may be either an Mn(II) salt, MnO2 (or any other oxide, except Mn2O7 of course! Mustn’t make it too easy). It is always
convenient if the reagents are commonly available. So, you can use even MnO2 + crud from old batteries, IF you purify it with care. More anon.
For an oxidant, we have to look at the standard redox potentials of the conversion from Mn(II) or Mn(IV) to the oxidation state MnO4-. These are
conveniently shown on WebElements (WE) for Mn and possible oxidants. Using these data, we determine that the oxidation is best carried out in alkaline
solution from Mn++ (Mn(OH)2 in the WE figure – Mn will precipitate as hydroxide) at 0.34 volts. From MnO2 we need 0.60 volt. So we need an oxidant
with a redox potential of > 0.34 Volt. (or 0.60v for MnO2). The oxidant chosen has to work in alkaline solution – pH TBD.
H2O2, which works like a charm for making chromates, is decomposed catalytically by MnO2 and manganous salts. Chlorates (see WE redox diagram)
aren’t quite strong enough in alkaline solution. Forget K peroxydisulphate unless you have it on the shelf. Even then, it's marginal due to H2O2
formation.
Examination shows that hypochlorites might work, redox potential to Cl- of 0.89 in OH-. A bit close for MnO2, perhaps., so expect a slow reaction.
Let’s give it try…
To be continued…
DerAlte
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MadHatter
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KMnO4
Given the recent CPSC(assholes) victory against FireFox, I'm sure that any convenient
method will be appreciated. KMnO4 sales by pyro suppliers are now limited to 1 LB a
year. There's still some OTC sources for now but we don't know how long that'll last.
From opening of NCIS New Orleans - It goes a BOOM ! BOOM ! BOOM ! MUHAHAHAHAHAHAHA !
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guy
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are you suggesting using hypochlorite? Thats not really a good idea.
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Zinc
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Why?
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Pyrovus
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I've tried hypochlorites with limited success. I've found it to be very messy and annoying - no matter how much of an excess of hypochlorite I use,
the reaction never seems to go to completion. In fact, despite using large amount of hypochlorite I believe that the permanganate solutions prepared
this way are probably pretty dilute - after all, the colour of permanganate is pretty intense, yet the solutions weren't all that dark. Then of
course, is the problem of purifying it - the biggest problem being the presence of chloride ions in the solution, as in the presence of permanaganate
these are no mere 'spectator' ions, and I'm yet to find a viable way of getting rid of them easily.
However, if using hypochlorite as oxidant, use lots of base. The oxidation to permanganate for either Mn(II) or Mn(IV) involves the production of
acid, and in acidic solution permanganate will oxidise chloride (which is present from the hypochlorite being reduced) to Cl2. Cl2, of course, is not
fun, and in addition your (already probably low) yield will be decreased.
Never accept that which can be changed.
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JohnWW
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Exactly the same occurs, with lack of completion and difficulty in isolating the product as a solid, when one uses NaOCl to oxidize alkaline Fe(OH)3
or NaFeO2 to Na2FeO4, containing Fe(VI). The FeO4-- anion is intensely magenta in color, very similar to permanganate, and is used similarly in water
disinfection. Theoretically, FeO4- and FeO4, containing Fe(VII) and (VIII), could also exist, probably being intensely blue or blue-green in color due
to the charge-transfer absorption band being at an even longer wavelength, but I have not heard of them being prepared.
I understand that the usual modern industrial method of preparing Na and K permanganates(VII) and ferrates(VI), (and in similar ways chromates(VI),
plumbates(IV), bismuthates(V), etc.) is by electrolysis of a cold strongly alkaline solution containing dissolved Mn2O3 or MnO2 and Fe(OH)3, as
manganite(III) or manganite(VI) and ferrite(III), at suitable voltages. This avoids liberation of Cl- anions through hypochlorite decomposition, which
would be oxidized by the desired products.
[Edited on by JohnWW]
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DerAlte
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Pyrovus, I agree that hypochlorites are messy! See continuation.
But it happens to be a very strong oxidant that is commonly available. The only other wet methods I have tried are electrolysis: ferromanganese (70%
Mn) as anode in 30-40% KOH- works quite well, but where do you get the ferromanganese? I have also tried an anode of MnO2/graphite and you can get a
pink color.
DerAlte
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DerAlte
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…try the following to show that it can be done. I’ll assume you have the reagents.
(1) Heat MnO2 with 50% NaOH or KOH in a test tube. No reaction is noted (some authorities claim insoluble manganites, xK2O.yMnO2 are produced). On
Standing for some time a slight pink coloration may be noticed due to atmospheric O2.
(2) Instead of dissolving the NaOH in water, use household bleach as the solvent. This gives us a weak hypochlorite in about pH 14 solution. Heating
produces the green magnate color, due to MnO4— ion.
(3) Repeat (2) using a concentrated solution of sodium carbonate in bleach. On heating the pink permanganate color appear. pH is of the order of 11.
Manganate is only stable in very alkaline solution. If a permanganate is heated with strong KOH, it turns into manganate plus MnO2
The essence of the method, then, is to oxidize MnO2 or Mn++ with hypochlorite at pH 11.
To be continued…
DerAlte
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DerAlte
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Lest you think I’m stalling, let me say that my intent is to give some interesting chemistry, not to show how to produce a ton of KMnO4. This
process is NOT 100% efficient. Why any one would need a ton or even a kg escapes me, unless to indulge in dangerous pyrotechnomania or get the DEA
very interested. Permanganates are rather unstable, especially with metal powders. Use perchlorates instead. I grew out of pyrotechnomania by the
age of 18, managing to escape with no more than singed eyelashes and a few burns.
The main chemical usage of permanganates is in titration of reducing agents and in organic oxidations. It is possible to produce a working quantity
– 20 to 50 gm – by using the potassium process to be described.
BUT FIRST, a caution. Strong hypochlorites are nasty. Always use rubber gloves. Read an MSDS on the web for all reagents or products mentioned if you
are not sure about them.
If you indulge in chlorine production, remember that the stuff is noticeable at 10-100 ppb so don’t do it in an apartment, condo, row house or flat.
If you do and let it leak you may be invaded by Big Brother’s Hazmat Boys (clothe them all in yellow, ho!). Chlorine production demands a shed or
porch and good ventilation, an isolated plot and a brisk wind. Always absorb it rather than let it escape. You need good chemical apparatus for this.
Conc. HCl is used as a pool chemical and driveway cleaner and easily available. Be careful with it. Keep it a mile away from hypochlorites,
permangantes and other oxidizers, and all metals. Read the MSDS for everything.
Finally, to end this boring preamble, some means of weighing to about 0.1 gm is necessary for success. You can’t just throw quantities of the
reagents together and hope. Knowledge of volumetric analysis helps, too. If you’ve got a burette, use it! It helps to know the strength of the
hypochlorite solution (iodometry) and also the permanganate (acidic ferrous sulphate or a sulphite). I need a nap…
DerAlte
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DerAlte
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The Process
Manganese carbonate is reacted with an alkali hypochlorite and an alkali metal carbonate in aqueous solution. The putative reaction is
2MnCO3 + 5KClO + K2CO3 --> 2KMnO4 + 5KCl + 3CO2
(An alternative is to use Mn(OH)2 in place of the carbonate. Haven’t tried it but should be effective too. MnO2 can be used but the reaction is much
slower. For reasons to be stated, speed of reaction is vital.)
The KClO needs to be as strong as possible, circa 15%w/w. Now you’re not likely to have KClO but you can make it. More on making the components from
easy available stuff later. Same for MnCO3, but it can be found on sites catering to pottery enthusiasts. Potassium carbonate can be obtained from
sites providing material for amateur soap manufacture. Making it is not easy except from CO2 and KOH.
You can substitute Na for K but don’t expect to be able to crystallize out the NaMnO4 – it’s very soluble. It’s enough of a problem with the
potassium salt, due to the large amounts of chloride produced. And it’s no better if you try to precipitate KMnO4 by adding KCl.
This is Important if you want any worthwhile yield.: Do the reaction in a glass beaker on a water bath and keep the temperature below 65 Deg C but
above 50 deg. The reaction is slow and will take one to two hours at this temperature. DO NOT expect all the MnCO3 to react. Some will remain as
indeterminate oxides and hydrated oxides of manganese, a brownish black crud. .
The solution should quickly turn red and then BLACK. Even saturated KMnO4 at 0 deg, containing 2,84% W/W is black and opaque, which complicates the
separation by fractional crystallization.
When ready, dilute the solution with about 3X volume water and filter warm using a glass plug in a glass filter funnel. Do NOT Use Paper. KMnO4
attacks most organics, especially double bonds and –OH bonds.
That’s the bare bones. Enough for now! Next time, how to separate and get the permanganate and some notes on technique
DerAlte
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DerAlte
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EXTRACTING THE KMnO4:-
It is very helpful to determine the amount of product you actually have. The solution should be black. If it’s pink or red, you failed. If it
still stinks of hypochlorite, the reaction time was not long enough. Other products are KCl, unreacted K2CO3, and whatnot. We can ignore the carbonate
because of its high solubility.
If you are competent at volumetric analysis, titrate against ferrous (better, ferrous ammonium) sulphate. Hypochlorite may still be present so you may
overestimate the KMnO4. If you don’t have an accurate burette or can’t fabricate one, use a micropipette and count drips the way I do, using
rather dilute solutions. Crude but effective! An accurate balance is essential for titration – I use an antique chemical balance that deflects on
1mg.
One thing you do know is the approximate amount of K+ ion you have in solution (assuming you know the strength of the KClO solution used – see
later.). This helps to reduce the solubility of the KMnO4 considerably, by the common ion effect.
At 0 C the solubility of KMnO4 is about 2,8% w/w aq. That of KCl, 28%. Without doing a common ion calculation, it can be safely assumed that in
saturated KCl solution, KMnO4 solubility will be reduced well below 2%. So, cool in an ice bath Preferably in a refrigerator to delay the melting. You
will expect to get a blackish mess containing permanganate and chloride in all probability. Filter this off, keeping the filter cold, using a glass
wool plug in a filter, a glass mat filter or a glass frit filter but NOT Paper. Wash with a small amount of ice cold water.
Now is the time to reveal the real problem. I’m sure many have already seen it – THE HYPOCHLORITE DISPROPORTIONATES:-
3ClO- --> ClO3- + 2Cl-
So we also have to deal with an undetermined amount of chlorate. And here is a HUGE snag; it happens that KClO3 has almost an identical solubility
curve as the permanganate and it’s next to impossible to separate the two. It would take the patience of Job.
I can’t think of any chemical means. Permanganate is the stronger oxidizer. I cheated – I didn’t promise you’d make Solid or pure potassium
permanganate, merely permanganate. You got it, in solution. Now think of a solution. We could use the sodium salts instead of potassium …
2B continued…
DerAlte
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Filemon
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I have synthesized NaMnO4 with 2MnCl2 + 5NaClO (aprox.15%) + 6NaOH. The reaction is very quick. Surprisingly the NaMnO4 had precipitated. For the ions
Na+ of the NaCl?
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woelen
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Are you sure that it is not NaCl with some NaMnO4 mixed in? Permanganate is so dark, even a 5% mix with a colorless salt still is very dark, near
black.
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DerAlte
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woelen, Agree. Saturated soln of KMn04 @ O C is still very nearly black at 2.8%. MnO4- is about the most intensely colored ion in existence.
Filemon, I doubt the reaction is "quick". Sure, a pink color appears but this only means very, very dilute. Also, NaMno4 is exremely soluble like
sugar in a rainstorm, at about 200gm/100 gm aq. What you probably saw was a mixture of oxides and hyroxides of manganese. Mn is so avid for O that
even chlorides turn brown on being kept in air. Anhydrous MnCl2 can only be prepared in an atmosphere such as gaseous HCL
And, under pH 14 or so in strong alkali, green manganate shoulf be produced.
MnCl2 crystals, as hydrate, (can't remember which - ), turns acidic when kept. Strong hypochlorite acting directly on such will produce chlorine, so
take care. Addition of NaOH stops this by neuralizing the acid. Same is true off all Mn(II) salts, as regars oxidation. They all tun brown in time.
As for using Na instead of K salts in the method I outlined, guess what? The solubility of NaClO3 is of the same order as NaMnO4. Instead of being
fairly insoluble it's very soluble.
The only possible way to separate the permangante from the chlorate that I can think of is via the silver salts. Anyone got an ingot of silver and a
gallon of nitric acid?
Regards, To Be Contd,.
DerAlte
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Filemon
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| Quote: | Originally posted by woelen
Are you sure that it is not NaCl with some NaMnO4 mixed in? Permanganate is so dark, even a 5% mix with a colorless salt still is very dark, near
black. |
Probably it's blended with NaCl.
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G.i.B.
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Can you distill the KMnO4 from the mix at reduced pressure ?
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UnintentionalChaos
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That is like asking if you can distill salt out of salt water.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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G.i.B.
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I read somewhere, that you can distill solids from solids, I am not sure how.
I off course meant, first boil away all the water.
[Edited on 17-5-2007 by G.i.B.]
[Edited on 17-5-2007 by G.i.B.]
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not_important
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| Quote: | Originally posted by G.i.B.
I read somewhere, that you can distill solids from solids, I am not sure how. |
True, for example you can distill solid CO2 away from sand. So long as the vapour pressure is high enough at a temperature below the decomposition
point, you can distill a substance be it solid or liquid. I doubt that those requirements hold true for permanganates and chlorates.
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DerAlte
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Permanganates.
I am sure you all think I’m a lunatic by now! Back to disproportionation. Unless I’ve made a gross mistake in calculation , the reaction
3OCl- --> 2Cl- + ClO3-
is only a bit less energetic (SEP ~0.4v) as the oxidation of Mn++ we are trying to achieve (~0.5V). It is said that this disproportionation becomes
“significant” at 40 C. As we are using a fairly concentrated hypochlorite solution and a higher temp (55C) – on the premise of speeding up the
wanted reaction rate – this is even more likely to be true. I estimate that about 30-40% of the last product I got was chlorate after several
fractional crystallizations.
As for using silver salts, mad as it may seem, it should be feasible. A rather expensive idea, perhaps! (Recycle!) I don’t have any Ag salts on
hand so haven’t tried it. If you happen to have a bucket of AgNO3 and an adventurous spirit, first do yourself a favor by ensuring all the chloride
is gone from the product so you only have chlorate and permanganate, by repeated recrystallizations. Otherwise all you’ll get is AgCl pptd.. AgMnO4
has solubility 0.9g/100g at room temp; AgClO3 something like 18g/100g. AgNO3 is over 200g/100g. Easy separation.
Finally, react the pptd silver manganate with KCl solution to get inslouble AgCl and the desired KMnO4 in solution
If anyone is still interested, I shall expound my CRUD (Chemical Reagent from Utter Dross) method of getting a pure manganese product from old
alkaline cells, and also how to make 15% plus NaOCl or KOCl from crap used for pools.
I used, in another place at another time, to always use reagent grade chemicals; but Big Brother and his minions, the Hazmat police, and those who
believe that owning an Erlenmeyer flask is the prelude to being a drug czar, have made an experimenter’s life difficult (albeit from sound motives
combined with political correctness) to say nothing of the affront to my libertarian principles. So I am forced to turn crap into at least technical
grade reagent like the rest of you. And that can be fun, too…
DerAlte
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DerAlte
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How to prepare a strong solution of Hypochlorite.
After writing this offline I discovered that Garage Chemist had already posted an excellent dissertation on this under a new NaOCl thread. Please read
it. I agree with all he says.
This may be a bit redundant, in consequence, but having written it, I might as well post.....
You can get up to 15% NaOCl from pool supply companies (HTH). It will not keep well except in cold weather when you don’t want it for pools. If you
want any other hypochlorite, such as KOCl you’ll have to make it. You can absorb Cl2 gas in KOH or K2CO3 or use the CRUD method:
Calcium hypochlorite is a relatively stable solid form. Get some ‘pool shock treatment HTH’ type. This contains about 60-65% by weight Ca(OCl)2.
The rest is calcium hydroxide, chloride, chlorate, etc., all Ca compounds. Read the MSDS for this stuff, it’s nasty. React with sodium carbonate in
the correct proportion, assuming that the product is 100% Ca(OCl)2- this is close enough for government work. First mix the two as solids, then add
water sufficient to produce a 15% solution of NaOCl. (Remember the sodium carbonate may have up to 10H2O as hydrate – Dehydrate first at 150C in an
oven to Na2CO3.) Potassium carbonate is very soluble but sodium carbonate is most soluble near room temp.
You’ll now have a loathsome white mess! Calcium carbonate is precipitated as a very fine powder. Stir it up well and let it stand for a couple of
hours in the refrigerator at, say 10C. When settled, first add a few drip of carbonate solution to make sure all Ca++ ions are gone. Then filter over
a two stage filter of several fine weave pieces of glass cloth backed up by a glass wool plug in a glass filter funnel. Paper will be attacked by the
hypochlorite. Keep well stopped at a low temp. Color should be yellow to greenish.
This strength should be used rapidly. Even a week causes marked deterioration in OCl- content. Adding a little NaOH helps to keep it by buffering the
pH.
DerAlte
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DerAlte
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MANGANESE SALTS FROM used batteries: an old favorite. Some are under the delusion that the black mess is manganese dioxide. It does contain some,
true. But it also contains zinc oxide, lower oxides/hydroxides of Mn, (NH4Cl and Zn(NH3)2Cl2 and similar Zn complexes in the case of the old Zn/C
type), or KOH in alkaline plus potassium zincate and grapitic carbon dust. My CRUD process:
First clean away the solubles by boiling with water for ½ hr. This also breaks up lumps. Next wash in DILUTE COLD HCL or dilute acetic acid (5%)to
remove Zinc compounds like ZnO.
You could then try floating off the graphite with a stream of water but that doesn’t work very well. Or you can heat in an iron can to bright red
heat (500C+) to burn off the C and reduce the oxides to Mn2O3. Of course you destroy any dioxide present.
IN my CRUD process I just leave it. The next step should only be undertaken if you know and have the apparatus to deal with Chlorine properly. Absorb
it in NaOH, KOH, Na2CO3 or K2CO3 solution.
The mixture of oxides will not contain MnO, but contain MnO2, Mn2O3 and possibly Mn3O4 (And C). Dry off at 200C in an oven. React with strong HCl
(>30%) in a flask (Round is best, only 1/5 full to allow frothing room). Warning: You now have a chlorine factory. Heat gently on a sand bath or
gauze until reaction ceases. The stuff will not clear due to the presence of graphite but has to be judged by cessation of Cl2 emission. I use a
second tube in the gas generator flask to blow out the remaining chlorine (Not by mouth!)
Contents of the flask is now MnCl2, excess HCL (and C). Filter off when cool (paper is OK) to remove the C. A fine pore paper is essential: coffee
filters let a lot of fine graphite through. You can crystallize the MnCl2.4H20 as nice pink crystals from a solution acidified by HCl. Or, as I do,
convert to carbonate, an insoluble pink powder that keeps well, is not deliquescent, and can be used to make any Mn(II) salt.
Qualitative analysis shows that the salt from alkaline cells, the most common these days, is quite pure. That from Zn/C cells seems to have a slight
iron impurity.
Finally, if what you want is MnO2, you can regenerate it from the chloride (or carbonate) by precipitating the Mn(II) hydroxide with NaOH, washing it,
and adding strong NaOCl and warming. The hydroxide is creamy colored, carbonate whitish pink, and MnO2 black as the ace of spades. Mn compounds have
many colors, red, purple, green, blue, brown, black – all except yellow. At least, can’t think of one
DerAlte
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