Sciencemadness Discussion Board - Chlorate as an all-purpose oxidizer?

Chlorate as an all-purpose oxidizer?

12AX7 - 25-1-2008 at 10:43

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 - 25-1-2008 at 16:10

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.

BromicAcid - 25-1-2008 at 16:21

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 - 25-1-2008 at 17:07

WOW! Bromic I hope there was no harm done to you or any financial losses suffered via property damage.

Was there organics involved?

BromicAcid - 25-1-2008 at 17:33

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 Cl2 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 - 25-1-2008 at 18:26

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?

garage chemist - 25-1-2008 at 18:35

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 - 25-1-2008 at 18:51

.........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 - 25-1-2008 at 20:02

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 ClO2 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+3, Ni+4, and Ni+5 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 - 26-1-2008 at 17:25

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! :cool:

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]

woelen - 28-1-2008 at 00:06

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).

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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 - 28-1-2008 at 09:53

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 - 28-1-2008 at 14:33

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.

JohnWW - 28-1-2008 at 18:50

Quote:

Originally posted by BromicAcid
As for the nickel compounds I remember reading about them while looking for information on ferrates. Specifically on Ni+3, Ni+4, and Ni+5 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.

Yet another oxidation

12AX7 - 8-2-2008 at 14:37

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