Sciencemadness Discussion Board - Making I2 with NO2, success

Making I2 with NO2, success

Pyro - 7-11-2012 at 08:16

Hello everyone,

A few days ago I had success making I2 by passing NO2 through a solution of KI, this has a few advantages over the HCl, H2O2 and other methods. It doesn't use expensive H2O2, just nitric acid, sugar and KI. and it produces KNO2, a useful reagent.

ORIGINS OF THE IDEA
I got the idea for this when Plante1999 and I were talking and I mentioned I needed NaNO2, So he suggested I react KI and NO2. he told me he got this idea from reading about nitrous acid being used to react with iodide making NO and iodine. NO reacted with NO2 to mke nitrite. So he thought about using NO2 as the oxidant to get almost all nitrite so that the formed nitric and nitrous acid make NO as well when oxidizing iodine and the NO would react with NO2 to make nitrite.
for NO2 production sugar is best as it wastes none of the available nitrogen
thanks to DJF90 for providing plante1999 with this process

2NO2+2KI->2KNO2+I2.

The KNO2 can be recuperated by boiling down the filtrate after filtering out the I2.

MY SETUP
I set up an NO2 generator with a flask and vacuum takeoff adapter with a hose and pipette leading into a cold solution of KI on a stirplate.
The first time i made a little mistake and put too much sugar into my nitric acid and had masses of NO2

. my entire lab was full of it but I had my I2 really quick.
The second time I got my NO2 production under control and that is where these pics are from

1)the setup I used, to the right of my hotplate you can see my yield of I2 from 100g KI, it is about 60-70g It isn't dry yet.

2)I used sugar and nitric acid to make NO2. but since it needed heating and my hotplate was occupated I got to use my new bunsen burner

may it never run out of gas

3)this is how I lead NO2 through my solution, it would be a good idea to wrap something protective over that clamp. mine is quite discoulored from all the I2 and a little NO2

4)this is my KI solution after approx. 220s

5)this is the KI solution after approx. 600s

I havn't been able to finish this run, I will finish it and post more pics this week end.

EXPERIMENTAL:

required equipment:
-Gas generator (flask, gas takeoff adapter or vacuum takeoff adapter and a stopper)
-hose (NO2 resistant, I used a gas hose made for propane)
-beaker (400ml)
-pipette or even better a fritted bubbler
-clamps
-A way to filter out the I2

required reagents:
-166g Potassium iodide (KI)
-150ml distilled water
-a means of generating 1mol NO2 (36 HNO3+C12H22O11=6 H2C2O4+23 H2O+36 NO2) so you need about 65g 100% HNO3, you need to figure out how much more diluted HNO3 you will need. And about 10g C12H22O11(sugar)

Start by dissolving the KI in the cold water,
then put all your HNO3 and a few grams of sugar into your gas generator. (DO NOT put in more sugar! otherwise you will have a disaster on your hands like I did)

put the hose onto a pipette or a fritted bubbler and into the solution. put the other side onto the gas takeoff adapter. then heat your flask of sugar and nitric acid until it has reached reaction temperature, then quickly put the gas takeoff adapter unto it and clamp it in place. now whenever the NO2 production slows add a few more grams of sugar until it all has reacted away.

now you should have a pulpy solution of iodine and aqueous KNO2. filter out the iodine and wash it with a tiny amount of water. set this out to dry and recrystallize if you want.
now boil down the solution of KNO2 and I2 until you start seeing crystals forming at the top. then cool it down in the freezer to get as much KNO2 out as possible.

Thanks to Plante1999 for this process!
and thanks to DJF90 for providing the source of the reaction

[Edited on 8-11-2012 by Pyro]

elementcollector1 - 7-11-2012 at 08:46

Very nice, but 3% H2O2 isn't exactly expensive compared to nitric acid, and I use 3% pharmacy stock for all my iodine syntheses.

Hexavalent - 7-11-2012 at 08:50

Nice work, Pyro!

The only disadvantage I guess is that nitric acid, a somewhat difficult to acquire reagent, is used, as well as the toxicity of the NO_x.

Pyro - 7-11-2012 at 09:00

true. but my HNO3 costs 2 eur/l for 60% lab grade. you are right about the toxicity of the NO2, It ruined my keck clip and left me breathing badly for the entire next day. though if you are careful and don't put in too much sugar into your HNO3 like I did

you shouldn't have too much trouble as the NO2 reacts almost completely with the KI.
the upside is that you make oxalic acid, KNO2 and I2. all useful reagents.

Nicodem - 7-11-2012 at 09:31

Why don't you post the experimental? What is the point in posting about an experiment, with pictures even, but no experimental? If you prepared iodine and potassium nitrite with a single preparation, then I'm sure some will be interested about the procedure.

Pyro - 7-11-2012 at 09:33

ok. ill edit it

AJKOER - 7-11-2012 at 10:40

Some reactions might be helpful if your interest is to have some KNO2 formed. In particular:

2 NO2 + H2O <--> HNO3 + HNO2

KI + HNO2 = HI + KNO2

KI + HNO3 = HI + KNO3

2 HI + NO2 = I2 + H2O + NO
----------------

Or, on net:

3 NO2 + 2 KI + H2O = I2 + KNO2 + KNO3 + NO + H2O

or on eliminating the common H2O:

3 NO2 + 2 KI (aq) = I2 + KNO2 (aq) + KNO3 (aq) + NO

However, as per Wikipedia (http://en.wikipedia.org/wiki/HNO3 ):

NO2 + HNO2 --> HNO3 + NO

it is bested to avoid an excess of NO2 as more nitrate than nitrite could be formed.

[Edited on 7-11-2012 by AJKOER]

Pyro - 7-11-2012 at 10:45

good point, but the NO2 should react with the KI first, and some NO2 will bubble through the solution without reacting.

Magpie - 7-11-2012 at 14:03

I like this synthesis because I have never seen NO2 per se used as an oxidant. I think there are many such synthetic methods that use common chemicals but are not well known.

If you write this up as a procedure in the PrePublication section you will have a month to edit it instead of just 24hr.

You can place your picture titles under the pictures. Your readers will appreciate this. Here's how:

1. Write your titles first, in order, with 1 or two spaces before and after each one. I've been doing this in bold to help keep them distinct from adjacent regular text.
2. Upload your pictures, in order.
3. Cut and paste each picture (a file found at the bottom of the page) to its place above your titles.

[Edited on 7-11-2012 by Magpie]

[Edited on 7-11-2012 by Magpie]

bfesser - 7-11-2012 at 14:30

I find that I2 is much easier to purchase than HNO3, and typically cheaper (though nowhere near as cheap as 3% H2O2!), in my area. Even cheap 'technical grade' I2 is a breeze to purify via sublimation. Still, nice use of <a href="http://en.wikipedia.org/wiki/Nitrogen_dioxide" target="_blank">NO2</a> <img src="../scipics/_wiki.png" /> as an oxidizer. Thanks for sharing.

[Edited on 7/9/13 by bfesser]

Pyro - 7-11-2012 at 15:37

thanks magpie. is that better?

Magpie - 7-11-2012 at 15:55

Quote: Originally posted by Pyro

thanks magpie. is that better?

Yes! - much better.

And I do recognize that Plante1999 invented (or owns) the process.

Pyro - 7-11-2012 at 16:02

that has been there the whole time.
plante and I will write a prepub. later.

DJF90 - 7-11-2012 at 16:46

I mentioned this method to Plante1999 in a private communication last month. I had seen the reaction mentioned in a JChemEd paper, and Plante tested the reaction as he found it interesting. The original source of the information used in this work can be found attached to this post.

Plante1999 claims to have had the idea from iodine recovery in Brauer.

[Edited on 8-11-2012 by DJF90]

Attachment: microscale gas generator.pdf (197kB)
This file has been downloaded 996 times

[Edited on 8-11-2012 by DJF90]

plante1999 - 7-11-2012 at 17:22

I apologize DJF90 because he was not mentioned in the process.

I had tought of the process when I read the iodine recovery from brauer, but lacking of money I didn't tried my idea, latter on, DJF90, when we where talking about something irevelent for this thread gave me the reference, armed with the reference I tried it on very small scale since I didn't have enough money for normal one. Then Pyro asked me and I told him my idea. But I didn't said that DJF90 was involved. Then when he made is post I told him about DFJ90 but it was found not being important so it was removed.

Magpie - 7-11-2012 at 18:48

Quote: Originally posted by Pyro

...then put all your HNO3 and a few grams of sugar into your gas generator. (DO NOT put in more sugar! otherwise you will have a disaster on your hands like I did)....

I think that if you dissolved the sugar in some water in the flask, then attached a pressure equalized addition funnel with the HNO3, you could keep the NO2 generation under perfect control. Just add HNO3 as needed with the funnel stopcock.

Sedit - 7-11-2012 at 20:18

Quote: Originally posted by Pyro

true. but my HNO3 costs 2 eur/l for 60% lab grade. you are right about the toxicity of the NO2, It ruined my keck clip and left me breathing badly for the entire next day.....

Shit dude you got a dose like that? I did also one time and have not breathed right since.... It destroys the fine blood vessels inside of your lungs IIRC and scars up.

I have found it much harder to breath after back to back accidents sometime ago, one involving concentration of H2SO4 on a very damp day that left my lab filled with a thick smog of H2SO4 followed within the month by NOx exposure.

AJKOER - 8-11-2012 at 05:07

Note, if one is solely interested in a nitrite salt, note the following reaction given in the paper "A Novel Microscale Gas Generator" cited by DJF90 above which, I suspect, occurs in a non-aqueous environment:

2 NO2+ 2 KI → 2 KNO2 + I2

It is interesting to contrast this reaction with my expectation in an aqueous environment which I presented above:

3 NO2 + 2 KI (aq) = I2 + KNO2 (aq) + KNO3 (aq) + NO

Now, it does appear that the dry salt approach is superior in yield of KNO2 and separation issues. But by heating, this is not such an issue as per the reaction (see Wikipedia http://en.wikipedia.org/wiki/KNO2 ) occuring at 400 C (note, subsequent decomposition of the KNO2 at 440 C

and detonation

upon melting at 537 C).

2 KNO3 → 2 KNO2 + O2

and the Iodine would sublime

starting at 184 C. I also suspect again that excess NO2 may result in less KNO2 and more KNO3. Also, even without an excess of NO2, uneven exposure of the salt to the NO2 gas may result in some KNO3.

[Edited on 8-11-2012 by AJKOER]

woelen - 8-11-2012 at 07:14

I also tried this reaction and I think that AJKOER is right. I passed NO2 over a concentrated solution of KI and this results in immediate formation of I2, which forms a solid layer over the surface of the liquid and under the layer, bubbles of a colorless gas are formed. The material starts foaming somewhat and the gas inside the foam is colorless.

When the gas bubbles break up, then brown gas appears again. This is due to oxidation of the NO to NO2 by oxygen from air.

So, I think that the reaction yields KNO2 only at most at 50% yield, the remaining material being KNO3. The reaction goes according to AJKOER's equation.

An interesting experiment would be to bubble pure NO2 (which is not easy to obtain!) in a test tube, filled with a solution of KI. If the reaction of Pyro occurs, then no gas remains, all gas dissolves in the liquid. If nitrate is formed instead, then the colorless gas NO will collect in the test tube. One way to make pure NO2 might be to make a volume of NO under water in an inverted testtube (which is easy to obtain) and then fill this test tube with solution of KI and then bubble O2 in the test tube, half the volume of NO. If you do this, and Pyro's reaction occurs, then no gas should remain, if AJKOER's reaction occurs, then 1/3 of the NO gas remains after adding the O2.

vmelkon - 8-11-2012 at 10:11

What fuel do you use for your bunsen burner and where does it come from?

AJKOER - 8-11-2012 at 11:10

Quote: Originally posted by woelen

....
So, I think that the reaction yields KNO2 only at most at 50% yield, the remaining material being KNO3. The reaction goes according to AJKOER's equation.
....

Expecting side reactions, I would suspect that the yield would be "only at most 50%" as well. Some other possible reaction paths:

2 NO + O2 --> 2 NO2

that is, as observed by Woelen the presence of any air would add to NO2 formation. This could be followed by:

2 NO2 + H2O --> HNO3 + HNO2

These acids could now react with the newly formed Iodine from the postulated reaction:

3 NO2 + 2 KI (aq) = I2 + KNO2 (aq) + KNO3 (aq) + NO

by either of these two frequency quoted reactions:

10 HNO3 + 3 I2 = 10 NO + 2 H2O + 6 HIO3

10 HNO3 + I2 = 10 NO2 + 4 H2O + 2 HIO3

so some of the Iodine formed could (if there was not an excess of KI employed to absorb additional HNO3) be converted to Iodate via Nitric acid. One source, however,(http://www.webelements.com/iodine/chemistry.html ) cites the need for hot concentrated HNO3. Another source in the case of very concentrated Nitric acid, suggests Iodine nitrate (in agreement with Mellor, page 291 to 292, http://books.google.com/books?id=AnnVAAAAMAAJ&pg=PA287&a... ) and the formation of HNO2 (see http://www.sciencedirect.com/science/article/pii/00221902758...) and as Nitrous acid can decompose as follows when dilute and cold (see http://en.wikipedia.org/wiki/Nitrous_acid ):

2 HNO2 → NO2 + NO + H2O

or, when warm and concentrated as follows:

3 HNO2 → HNO3 + 2 NO + H2O

the frequently quoted reactions may be based on the actual formation of HNO2 decomposing under different conditions.

Note, any liberated NO2 from this reaction could attack HNO2:

NO2 + HNO2 --> HNO3 + NO

as previously noted forming more NO and HNO3 that could reduce Nitrite formation.

In addition, in a non-neutral solution, the newly formed Iodine may dissolve more readily with time:

I2 + H2O <--> HI + HOI

apparently as a function of initial Iodine concentration, temperature and pH (see https://docs.google.com/viewer?a=v&q=cache:F8ZxjpvpO30J:... ). The formation of Hypoiodous acid is favored by high initial I2 concentration, higher temperatures and high pHs leads to its decomposition (to iodate). The HOI is unstable and rapidly undergoes disproportionation:

3 HOI --> 2 HI + HIO3

so more Iodine can also be converted to Iodate.

Interestingly, one cannot actually employ an excess of KI due to the formation of the soluble tri-iodide ion:

I2 + I- = I3-

Bottom line, I would not be surprised for a lower yield of Iodine (when working with conc KI solution and too little NO2 or excess KI, or no stirring forming local conc HNO3 in a hot solution) and less Nitrite (if excess NO2, or vigorous stirring in air allowing NO+ O2 to react and dissolve) when conditions permit the creations of tri-iodides, iodates and nitrates.
----------------------------------
[EDIT] There may be interesting parallel reaction between NO2 and ClO2 with respect to their reaction with KI.
Per Mellor, page 289, http://books.google.com/books?id=AnnVAAAAMAAJ&pg=PA289&a... to quote:

"Iodine separates from an acidified soln. of potassium iodide: 2Cl02+10HI=2HCl+4H20+5I2; in neutral soln.: 6Cl02+10KI=4KI03 +6KCl+3I2; and in the bicarbonate soln.: 2Cl02+2KI=2KCl02+I2, whereby 80 per cent, of the chlorine dioxide is converted into the chlorite."

[Edited on 9-11-2012 by AJKOER]

woelen - 9-11-2012 at 02:38

I think that with respect to formation of iodine this reaction is good, very good and that nearly 100% yield is obtained. Oxidation to iodate does not occur under the conditions, present in this system. I have tried oxidation of iodine to iodate with quite a few oxidizers and dilute nitric acid, nor NO/NO2 in the presence of water are capable of oxidizing iodine. Much higher concentrations and/or elevated temperatures are necessary.

So, when it comes to making I2, this reaction is good (provided you can get HNO3 cheaply and easily), when it comes to making nitrites I think that this reaction is not that good and that very impure nitrites are obtained.

Pyro - 9-11-2012 at 10:05

Wow, this has recieved a lot of attention

@Magpie:
you are correct, but i see no need to make more dishes than needed
@vmelkon:
I use propane from a 10,5kg bottle it's quite expensive but in my case this was better as we use the same type of bottle for cooking. a bottle that big will last about 10 years at the rate i use it. how about yours?
@Sedit:
i get it from: De wagenaere NV in 5l pails, a place in my area. while i hate NO2 and the smell of it makes me want to vomit, I find SO2 worse

White Yeti - 9-11-2012 at 13:41

It's a nice method, but I'm just curious. Does anyone use the bleach and hydrochloric acid method?

elementcollector1 - 9-11-2012 at 13:45

Quote: Originally posted by White Yeti

It's a nice method, but I'm just curious. Does anyone use the bleach and hydrochloric acid method?

I do; it's my main source of iodine. In fact, I have some ampouled somewhere around here...

AndersHoveland - 9-11-2012 at 15:03

I think the reaction, in the presence of water, is actually initially:

2 KI + 3 NO2 + H2O --> 2 KNO3 + 2 HI(aq) + NO

However, when the concentration of the hydroiodic acid in solution starts to get more concentrated, the equilibrium will shift, and iodine will begin to be formed.

2 KI + 2 NO2 --> 2 KNO2 + I2

Normally, I would tend to think, nitrogen dioxide would preferentially oxidize nitrite to nitrate before it would displace iodide from its salts, but in this situation things are different. All that concentrated hydroiodic acid in solution shifts the equilibrium, by making the environment acidic and reducing.

While iodine will oxidize nitrite to nitrate under alkaline conditions, under acidic conditions nitric oxide can reduce nitrate to nitrogen dioxide.

Quote:

“Chlorine, bromine, and iodine also oxidize nitrous acid solutions to nitrate. The reaction between aqueous iodine and nitrite ion is measurably slow, and the rate in buffered solutions (pH = 6 to 7) has been studied by Durrant, Griffith, and McKeown [Trans. Faraday Soc., 32, 999 (1936)]. The net reaction is
NO2- + I2 + H2O = NO3- + 2I- + 2H+

Hydroiodic acid acts as a reducing agent towards nitrous acid solutions, the reaction products being iodine and nitric oxide.

RATE CONSTANTS FOR THE OXIDATION OF NITRITE ION BY IODINE for Systematic Inorganic Chemistry (1946)

KNO2 + NO2 --> KNO3 + NO

KNO2 + I2 + H2O --> 2 HI(aq) + KNO3

NO + I2 + H2O --> 2 HI(aq) + NO2

3 KNO2 + 2 HI(aq) --> KNO3 + KI + H2O + 2 NO

4 KNO2 + I2 --> 2 KNO3 + 2 KI + 2 NO

I think the best equation to describe the reaction, after hydriodic acid has already accumulated and made the solution acidic, may be:
4 KI + 4 NO2 --> 2 KNO3 + 2 KI + I2 + 2 NO

Some of you may be wondering why the iodine does not just oxidize the nitric oxide to nitric acid. Here are my thoughts on that. Nitric acid is a strong oxidizing agent, but generally only at high concentrations (>60%). However, even 20% conc. nitric acid is still an oxidizing agent. It is just that its reaction rate as an oxidizer is very very slow. Nevertheless, this slow reaction rate would still tend to make the equilibrium favor nitric acid, since the rate of reaction forming nitric acid would be much faster than the reverse action. The big factor here is probably all those nitric oxides floating around, which act as catalysts to reduce back the nitric acid. While dilute nitric acid is essentially not an oxidizer, nitric oxide which is a reactive free radical can reduce it to lower oxides of nitrogen which can act as much more reactive oxidizing agents. Again, although dilute nitric acid is a more "powerful" oxidizing agent than nitrous acid, it is just not a reactive one.

[Edited on 9-11-2012 by AndersHoveland]

AJKOER - 19-11-2012 at 14:13

Quote: Originally posted by White Yeti

It's a nice method, but I'm just curious. Does anyone use the bleach and hydrochloric acid method?

The only concern I have with this Chlorine basis method is the possible formation of ICl and ICl3 with reduced yield if it happens that excess Cl2 is employed or Iodine crystals are exposed locally to Cl2 gas. Basis, see Wiki (http://en.wikipedia.org/wiki/Iodine_monochloride ), to quote:

"Preparation of iodine monochloride entails simply combining the halogens in a 1:1 molar ratio, according to the equation

I2 + Cl2 → 2 ICl

When chlorine gas is passed through iodine crystals, one observes the brown vapor of iodine monochloride. Dark brown iodine monochloride liquid is collected. Excess chlorine converts iodine monochloride into iodine trichloride in a reversible reaction:

ICl + Cl2 <--> ICl3 "

White Yeti - 20-11-2012 at 13:25

That's why iodide is used in slight excess. It's easier to deal with triiodide than with assorted halogen halides. The reaction is conducted in aqueous solution; you may have overlooked the fact that ICl decomposes in water.

AJKOER - 21-11-2012 at 13:36

Quote: Originally posted by White Yeti

..... The reaction is conducted in aqueous solution; you may have overlooked the fact that ICl decomposes in water.

Actually, ICl rapidly reacts as follows:

ICI + H2O —> Cl– + OI– + 2H+

And, under acidic conditions, the final products (through a series of intermediaries) are:

Acid:
5 ICl (aq) + 3 H2O <--> HIO3 + 2 I2 + 5 HCl

Base:
3 ICl(aq) + 6 NaOH <--> NaIO3 + 2 NaI + 3 NaCl + 3 H2O

Source: "Kinetics of hydrolysis of iodine monochloride measured by the pulsed-accelerated-flow method" by Yi Lai Wang, Julius C. Nagy, Dale W. Margerum, published in
J. Am. Chem. Soc., 1989, 111 (20), pp 7838–7844. Online link
http://pubs.acs.org/doi/abs/10.1021/ja00202a026

So the formation of ICl in an aqueous solution can lead to some Iodine loss via the formation of Iodate. Hence, my prior comment "formation of ICl and ICl3 with reduced yield".

[Edited on 21-11-2012 by AJKOER]

White Yeti - 21-11-2012 at 15:15

If you're going to start splitting hairs again, I'm gonna piss off. Before I do, I'd like to say that the disproportionation reaction in base is unfavourable unless there is a large excess of hydroxide and a huge amount of iodine chloride is present. This would not be the case the reagents are used in stoichiometric amounts. The loss of iodine as iodate is as insignificant as considering the formation of chlorate when dissolving chlorine in water.

woelen - 21-11-2012 at 23:35

The loss of iodine to iodate only occurs when there is excess chlorine. As long as there is excess iodide no iodate/iodic acid is formed at all (if any were formed, then in the acidic solution, usually used to make iodine from iodide and bleach it would be converted to iodine at once).

So, in practice the chlorine/bleach method works fine as long as no excess amount of bleach is used and one assures that the solution remains acidic. If the solution becomes alkaline, then the reaction does not work anymore. In that case you get iodate as final product.

AJKOER - 22-11-2012 at 06:49

OK, I apology for the seemingly hair splitting, but my concern lies with my original comment "reduced yield if it happens that excess Cl2 is employed or Iodine crystals are exposed locally to Cl2 gas." As no one has challenged the local concentration issue, I have some theoretical concerns on how it may impact Iodine yield (and sometimes my theory are how things can go wrong are all too accurate).

Assume no stirring permitting local Cl2 concentration issues and then, with even perfect stoichiometric amounts, some ICl and then Iodate is formed (or formed by another route). Then via the reaction (see http://paws.wcu.edu/bacon/vitamin%20c.pdf ):

IO3- + 6 H+ + 8 I- = 3 I3- + 3 H2O

we have a significant conversion of the Iodide to the tri-iodide and not Iodine.

The comments presented do not appear, however, to suggest that this is a significant issue observed in practice.

[EDIT] Per wikipedia (http://en.wikipedia.org/wiki/Iodine_clock_reaction ):

IO3− + 6 H+ + 5 I− → 3 I2 + 3 H2O

so as long as we avoid a stoichiometric excess of Iodide, loss due to tri-iodide (from iodate and excess iodide) may not be an issue after all.

[Edited on 22-11-2012 by AJKOER]

woelen - 22-11-2012 at 06:58

AJKOER, what you describe in your last post certainly will occur, also in practice, but it is not an issue. Suppose, due to some locally concentrated spot with lots of Cl2 that ICl and/or HIO3 are formed, then this is not a true loss. Somewhere else there will be insufficient chlorine to convert all iodide to iodine. So at one spot, iodide remains present (either as free iodide, or bound to iodine as I3(-)) and at other spots excess Cl2 leads to formation of HIO3. As soon as the mix is stirred the iodine will be formed anyway. HIO3 and I(-) immediately react with each other to iodine. What I write is true under the following conditions:

- total amount of chlorine is not in excess, any local excess is compensated with other spots where there is too little chlorine.
- solution is acidic, otherwise iodate and iodide do not react to form iodine.

A good way to make iodine from iodide and bleach is first to titrate some of the bleach such that its concentration is known and then add a stoichiometric amount of potassium iodide (heated in an oven to 100 C for some time to drive off any water from the crystals). When all iodide is dissolved, then acid is added in sufficiently large amounts.

vmelkon - 24-1-2013 at 09:51

This reaction got me interested since it produces KNO2. I'm interested in having a nitrite salt since some organic chemistry experiments call for it.

Is the only way to separate the nitrite and nitrate is by taking advantage of solubility differences?

Another question : when you melt KNO3, it decomposes to KNO2 and O2, right?