12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Chlorate as an all-purpose oxidizer?
As many of you know, I have a lot of chlorate on hand. As such, I look for ways to use it.
One I've proven is acidic molten salt oxidation: 2KClO3 + KCl + Cr2O3 --> K2Cr2O7 + Cl2 [unbalanced].
Chlorate is an excellent oxidizer in acidic solution as well. For example, I was dissolving granular copper wastes in HCl the other day. This
produces a viscous, deep brown solution, due to Cu(I) complexation. If left in air, oxygen will slowly oxidize Cu(I) to Cu(II), resulting in
precipitation of Cu(OH,Cl)2 and further corrosion of the metal. But this process is very slow. In lieu of the traditional hydrogen peroxide, I added
some sodium chlorate to the solution. This has no immediate effect, but when acid is added, the brown color immediately turns green, in proportion to
the amount of acid added. The green solution is then able to corrode the metal rapidly, heating the solution. The reactions are:
ClO3- + 6H+ + 6 Cu+ --> 6 Cu(2+) + 3 H2O + Cl-
Cu(2+) + Cu(0) = 2 Cu+
The decomposition of chlorate is highly dependent on the pH. Since dissolving metal causes the pH to rise, this is a very stable sort of oxidation,
not prone to runaway, regardless of the temperature! With excess chlorate in solution, as excess Cu(I) accumulates, just enough acid can be added to
oxidize the Cu(I) to Cu(II). An excess of acid will cause unnecessary decomposition, resulting in Cl2 or ClO2, so it's essential to this process that
the solution contains substantial amounts of Cu(I) and Cu(II). As long as there is Cu(I) present, the decomposition of chlorate couples extremely
well to it, and the reaction produces no corrosive fumes at all.
Such redox reactions are possible with any metal ion that has two or more soluble oxidation states. Iron is a popular one (think Fenton's Reagent),
and vanadium (between states 2+, 3+ and 4+) is being researched as a liquid battery. The above reaction, with copper chloride or iron chloride or
sulfate as the redox carrier, can be used to oxidize anything in, or into, solution. The main problem is, say you want to dissolve nickel in an iron
solution: then you have a mixture of salts. So it would still be useful for dissolving mixtures that are to be seperated anyway.
Tim
|
|
|
chloric1
International Hazard
Posts: 1071
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
Given that some of us can produce significant amounts with a minimal investment, use as a general oxidant can be favorable as long as sodium or
potassium can be tolerated int the mixtures. You know, I was thinking of using some of my sodium chlorate or potassium chorate to oxidize some
stannous chloride so I can produce stannic acid by precipitation. I may go for potassium as sodium may be too easily absorbed by the precipitate.
I especially like the idea of adding acid only when needed as Cl/ClO2 loss is only indicative of as oxidation potential loss. This might be a problem
with stannous salts as excess acid will be present for a clear solution. So more chlorate will need to be added compared to the pH controled method.
Fellow molecular manipulator
|
|
|
BromicAcid
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
I once had a chlorate mixture explode upon acidification, though there were some other factors at work it's something that I now refuse to do.
|
|
|
chloric1
International Hazard
Posts: 1071
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
WOW! Bromic I hope there was no harm done to you or any financial losses suffered via property damage.
Was there organics involved?
Fellow molecular manipulator
|
|
|
BromicAcid
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
Possibly. I covered this elsewhere, and the situation is somewhat complicated chemically so it's not applicable 1 to 1 but I was dissolving match
heads in hydrochloric acid. I had the mixture in a beaker covered with a watch glass and it was bubbling away, at the time I didn't know what was in
a match head really, so I thought the green gas in the beaker was Cl<sub>2</sub> however I started seeing sparks in the head space and
finally **Bang** and the beaker was gone.
Now, not to side track the topic any further, it is the molten salt oxidations with chlorate/hydroxide that have always caught my attention,
specifically the preparation of ferrates and high oxidation nickel compounds that I haven't seen described in recent literature.
|
|
|
chloric1
International Hazard
Posts: 1071
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
Aha well that will do it! Sounds like you had red P interacting with ClO2.
Where did you hear about the high oxidation nickel compoundsif they are not in the literature?
Fellow molecular manipulator
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
Bromic, could it just have been the ClO2 (deep yellow gas) exploding? It can explode on the slightest provocation, such as contact with organic
matter.
I have found chlorate to be very useful for the synthesis of bromine from a bromide salt and acid.
I once got a 98% yield of bromine after distillation that way.
|
|
|
bio2
Hazard to Others
Posts: 447
Registered: 15-1-2005
Member Is Offline
Mood: No Mood
|
|
.........As many of you know, I have a lot of chlorate on hand. As such, I look for ways to use it............
Works good as weed killer and is sold as such. Spray on the leaves a, fairly strong, solution.
|
|
|
BromicAcid
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
Chloric, modern match head compositions do not contain red phosphorus, they do however contain phosphorous sulfides (which have some interesting
properties themselves that are under discussed here. I myself am in agreement with Garage Chemist that it was just some ClO<sub>2</sub>
deciding to end its own cruel existence. As Wiki states (not the best of reference but more readily available than my bookshelf): "At
concentrations greater than 15% volume in air at STP, ClO2 explosively decomposes into chlorine and oxygen. The decomposition is initiated by light.
Thus, it is never handled in concentrated form...."
As for the nickel compounds I remember reading about them while looking for information on ferrates. Specifically on Ni<sup>+3</sup>,
Ni<sup>+4</sup>, and Ni<sup>+5</sup> compounds that people have thought to have created. I have seen literature on some of
these such as the fluorine compounds but not much on some of the nickel salts created or believed to be created in the oxidizing brew.
|
|
|
chloric1
International Hazard
Posts: 1071
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
OK it is phosphorus sulfide I forgot. For all intents and purposes though, especially when mixed with chlorate, it mind as well be red phosphorus
mixed with sulfur. I have added NaClO3 crystals to muriatic acid concentrate and got a bright yellow chlorine/chlorine dioxide gas evolving solution. The
chlorine was more volatile and evaporated within a half hour while the ClO2 remained alot longer possibly several hours to a day. Never dreamed of
heating these or adding easily oxidizables. In situations involving reducing metal ions, I think it should be relatively easy to avoid ClO2
generation by controlling either acid addition or chlorate addition. When ionic reducer is exhausted, H202 would probably destroy any ClO2 generated
at the close of a reaction.
Pentavalent nickel?? NICE! Love to oxidize the hell out of something with
that! I guess there is only one way to find out.
[Edited on 1/26/2008 by chloric1]
Fellow molecular manipulator
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I once did an experiment with a piece of nickel in molten NaOH/KClO3, and I obtained some black material. On dissolving in dilute hydrochloric acid,
this gave green Ni(2+) and besides that, the smell of chlorine appeared. So, the chlorate certainly oxidizes nickel, but whether this is to a higher
than +2 state I do not know. The smell of chlorine could be because of excess chlorate, left over from the mix.
Personally, I don't find chlorate the oxidizer of choice in many situations. It indeed is suitable for oxidizing metal salts mixes in concentrated
hydrochloric acid (e.g. copper, cobalt, nickel), but it is not very clean in other oxidations. With organics, it usually gives a lot of crap, due to
chlorine side production, resulting in chlorinated byproducts.
Even for the metal salts, I do not prefer chlorate, because the resulting solution contains both the metal salt and potassium/sodium ions. Simply
evaporating such a solution does not give the pure metal salt. With H2O2/HCl, simply heating the solution and then letting dry the material gives the
pure metal salt (provided it does not hydrolyze and form a basic chloride).
-------------------------------------------------------------------
Making high oxidation state nickel compounds is not difficult. You can even do that in aqueous solution.
- Dissolve a soluble nickel salt in same water and add a solution of NaOH. A slimy light green precipitate of Ni(OH)2 is formed.
- add a solutoin of Na2S2O8 (or any other peroxodisulfate salt). The precipitate quickly turns black.
When the black precipitate is rinsed, and this cleaned (still wet) precipitate is added to hydrochloric acid, then it dissolves, giving a green
solution of nickel(II) and a strong smell of chlorine appears. With hydrogen peroxide it also reacts. A lot of oxygen is produced, and light green
nickel hydroxide is formed again.
I have read somewhere that this black compound is NiO2.xH2O, but this is not really reliable data. It might also be Ni2O3.xH2O, or even some
indeterminate hydrous Ni(III)/Ni(IV) oxide.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
It's usually given the formula NiOOH, I think.
Segue-ing into nickel, can it be oxidized in neutral or mildly acidic solution? That would be a great way to seperate nickel from copper (e.g.,
plating bath) without dropping everything out of solution.
Tim
|
|
|
chloric1
International Hazard
Posts: 1071
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
Tim-Don't know for sure but I think nickel(II) is oxidized in alkaline solution. Two modes would be hypochlorite oxidation or alkaline persulfate
with silver ion.
Fellow molecular manipulator
|
|
|
JohnWW
International Hazard
Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by BromicAcid
As for the nickel compounds I remember reading about them while looking for information on ferrates. Specifically on Ni<sup>+3</sup>,
Ni<sup>+4</sup>, and Ni<sup>+5</sup> compounds that people have thought to have created. I have seen literature on some of
these such as the fluorine compounds but not much on some of the nickel salts created or believed to be created in the oxidizing brew.
|
Ni(V) has a special stability conferred by the fact that it has 5 unpaired "3d" electrons, one in each "3d" orbital, like Fe(III) and Co(IV) and
Mn(II). This is an especially stable configuration. However, the potential required for obtaining Cu(VI) is simply too great, although Cu(III) and
(IV) are known.
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Yet another oxidation
Today I observed another high temperature reaction of chlorate.
Some time ago, I was given some "carbide powder", which is the waste from grinding "carbide" cutting tools, which are most likely WC or W2C bonded
with a 10-20% Co matrix. This material is very fine, mixed with detritus (e.g., chips of waste carbide), and oily.
Without washing, this stuff seems to react with HCl, giving a deep green or red/brown solution (like a strong Cu(II) chloride or Fe(III) sulfate
solution), apparently depending very tightly on oxidation state. Much material is left unreacted, as expected.
So I decided to burn the material, first of all to get rid of the oil and other combustible residues, and second to oxidize the metal, if possible. I
don't know if I succeeded in oxidizing anything (I have a black powder, greenish in places), but I did toss on some sodium chlorate to see if I can
accelerate the process.
It turns out it reacts quite nicely with sodium chlorate, getting yellow hot on contact (where the chlorate melts into the mass) and sometimes giving
off blue jets of fire, suprising as this is sodium chlorate, no potassium to be found! The maximum temperature seems to be quite high for a chlorate,
up to perhaps 1500C.
The oxidized material seems to give off some HCl, so besides straight combustion, there must be something forming sodium salts. Hmm, tungstate could
be in there, then. I don't know about cobaltate.
Tim
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
