Sodium iodate from sodium chlorate

Looking at the older threads it is clear that most nerds on SM are trying to turn iodate salts into iodide salts or free iodine. So my desire to turn iodine into sodium iodate is going to be little short of heresy .

There is an excellent video by Chemplayer on the preparation of potassium iodate from potassium chlorate and iodine. So I decided to use this as my basic reaction model. But then I spotted in Brauer (p323) under the preparation of sodium periodate a similar preparation, this time of sodium iodate. Great, I thought, precisely what I wanted.

But when I began to look at the two procedures a major difference became apparent. In Chemplayer’s video they use a molar ration of iodine to chlorate of 1.105 while Brauer uses a ratio of 0.671. In other words Brauer is using roughly twice as much chlorate per mole of iodine. I can't imagine the sodium and potassium change this ratio as they are really only spectator ions.

In my experiment I decided to use an initial ratio of 1. The reaction occurred nicely, exactly as Chemplayer described. When the early vigorous reaction began to subside I warmed it at 80-90 C for about an hour but some iodine remained and the reaction appeared to have stopped so I added a bit more sodium chlorate and the remaining iodine largely dissolved but the solution became orange. I presumed this was due to unreacted iodine in solution as the triiodide ion so I added a pinch more sodium chlorate but this made no difference at all, by this point the iodine to chlorate ratio had been reduced to 0.838; about halfway between Chemplayer and Brauer.

Then a dense white ppt began to form. I presumed this was due to the slow evaporation of water so I topped up the solution to its original volume, more precipitate formed that would not dissolve even on boiling so I added more water raising the volume of solution to about 33% above starting volume.

I decided at this point to neutralise the suspension since all of the iodine had now gone and I suspected that a sodium hydrogen iodate was starting to precipitate and I wanted the meta-iodate (IO3-). I took 56ml of 50% NaOH solution, dilute 45ml of it to 100ml and commenced to pour the pale orange suspension into the alkali solution. Initially I got a clear solution almost immediately and then a white ppt began to form; then it happened! The iodate solution began to turn black and copious clouds of iodine vapour formed. After a few seconds at first but requiring longer and longer following each addition the colour discharged.

When I had added all of the iodate suspension the suspension was full of iodine and somewhat acid so I slowly added the extra NaOH solution and after about 9ml the colour had almost discharged. The suspension of sodium iodate was hot, almost boiling, so was left stirring to cool. The final suspension was pure white and the dense solid settled rapidly. I am happy that at least some sodium iodate has formed but what is the orange colour and why does it liberate iodine only briefly under partial neutralisation. Once the solution is neutral or slightly alkaline the iodine disappears.

Clearly the reaction is complex. If the reaction occurred through displacement chlorine from chlorate ie

Chlorate + iodine = iodate + chlorine

This would involve a simple 5 electron transfer and the ratio of iodine to chlorate would be 1.

However, if the chlorine ends up a chloride ions or hydrochloric acid it would accept 6 electrons and the ratio of iodine to chlorate would need to be 1.2.

Chemplayer’s ratio is in between these two values and they state that only a little chlorine is evolved. This is my observation too. So why does Brauer recommend such a large excess?

Having looked into this a bit more I see that in some iodate titrations in very acid conditions iodine monochloride is an important intermediate but this is unstable in less acid conditions. This may explain the sudden appearance and then disappearance of iodine during neutralisation and also the lack of iodine fumes in spite of the orange solution being at almost boiling point during the preparation.

Has anyone else tried this preparation? Or have any other ideas on the varied chemistry of this reaction?

@Chemplayer, thanks for the response. It appears though from your video and response that you did not see the orange colour at the end (when all of the iodine had dissolved) that I found and caused me to believe that iodine was still present in solution because there was insufficient chlorate to oxidized all of the iodine to iodate. I now believe that the colour is due to iodine monochloride in solution as ICl2- ions. The formation of this species is well known from iodate redox titrations. It is not the concentration of iodate that determines the electron transfer number but the acidity of the solution. It is possible that as the liberated acid is partly neutralised the iodine monochloride disporportionate:

5ICl + 3H₂O -> HIO₃ + 5HCl + 4I

Because the solution is still acid, hot and there is still excess chlorate present it drives the reaction forward towards iodate again. So if the neutralisation is carried out on the hot solution it drives the reaction to practical completion. This is supported by the fact that I got a crude yield of 97.7%!!! When I tried to pecipitate the remaining iodate with strontium chloride I got a little ppt but it represents a further 4%! I like these sorts of reactions; >100% yield . These recoveries are based on iodine since that was the limiting reagent.

I am about to recrystallises my sodium iodate product which, like you say is a snow white powder but must be a little impure given the impossibly good yield. Watch this space.

Ah, just found my sodium iodate is roughly a monohydrate so the yield is more like 90% and then 4% as strontium iodate monohydrate. I have just acidified the clear filtrate and copious iodine has precipitated, this means that some iodide must also have formed.

[Edited on 8-5-2020 by Boffis]

Quote: Originally posted by chemplayer...

One of those reactions where the overall stoichiometry can be summarised into a neat 'swapsy' equation but this appears to be nature just trying to foil us into thinking it's simple and con us into thinking we know what's going on.

Yes I agree totally! I have been investigating the literature on the reaction of iodate in the presence of chloride/HCl and iodide/ other oxidizable species and it is insane. There is actually a lot of published data, sometimes in really obscure places because they are looking at some really weird angle on this reaction. It is clearly complex.

Interestingly I have repeated the reaction using your ratios exactly except that I used a stoichiometric equivalence of sodium chlorate in place of K chlorate. This time I used sodium chlorate from a different source (Naclo, a Turkish company that used to sell sodium chlorite and chlorate on Ebay) and I got a near quantitative yield based on iodine!! From 20g of iodine I got 31g of crude sodium iodate (ie >100%!) and managed to recover 1.28g of strontium iodate from the filtrate (about 5% of the iodine imput) !!. I suspect there are purity issues here but certainly overal recoveries are good and using your "recipe" I didn't get the orange colour when the iodine had dissolved and no transient iodine ppt during the neutralisation step either.