I found a few journals on scirus yesterday which describe the use of ammoniumchlorate as cheap ClO2 source in the production of chlorites. One journal
even had the synthesis for ammoniumchlorate. Sadly, I can't find them anymore today (can't remember search string I used).
Has anyone reliable information about ammoniumchlorate?
I found a few articles
Polverone - 3-9-2002 at 10:59
And read a couple of them. One was on spontaneous decomposition of NH4ClO3 and the other on thermal decomposition. The only synthesis I found was
metathesis reaction of barium chlorate and ammonium sulfate; I don't know if this is helpful or not.
According to the spontaneous-decomposition article, NH4ClO3 is stable in cool aqueous solution but undergoes self-accelerating decomposition in the
solid state (unless kept at -5 C or colder, according to the thermal decomposition article), which is why it's so unwelcome in pyrotechnics.
Ammoniumchlorate is then isolated by drying in a dessicator with H2SO4.
Ammoniumchlorate consist of small needle-like crystals. Instabile, do not store, might explode without apparent stimilus, certain explosion above
100C. When spread out in a thin layer, ammoniumchlorate is safe to manipulation. Very water soluble.
The reactions with HClO3 use 40% HClO3 produced according to the following procedure:
322g of bariumchlorate are dissolved in 500ml boiling water. While stirring, a mixture of 98g conc sulfuric acid (53,3ml D 1,84) and 53,3ml H2O is
added. A slight excess of bariumchlorate is favourable over an excess of sulfuric acid. The bariumsulfate ppt is allowed to sit for one hour. Two
third of the liquid is decanted and the rest is filtered through a Buchner funnel. The filtrate is joined with the decanted solution and yields approx
660ml of approx 22% chloric acid (D 1,11).
When this is dried in a vac exsiccator over conc sulfuric acid, a 40% solution of chloric acid can be achieved, which complies with the formal
composition of HClO3.7H2O D 1,28
All taken from anorganische chemie by F. Huber and M. Schmeisser KABOOOM(pyrojustforfun) - 15-7-2003 at 18:19
<a href="http://swi.1av10.nu/dist/pyrok2k.zip">from Swedish Infomania:</a><br>
Ammoniumklorat (NH4ClO3)
Som kuriosa nنmns hنr ocksه ammoniumklorat, ett نmne som orsakat stora skador
dه det av misstag bildats nنr t.ex. kaliumklorat blandats med ett ammoniumsalt.
Pه grund av sin hِga kنnslighet och instabilitet نr det helt uteslutet som
sprنngنmne. Det exploderar om en 2 kg tung vikt faller 15-20 cm, och upphettat
till 40 °C sه exploderade det efter 11 timmar. Inneslutet in en glasflaska under
48 timmar sِnderfaller det explosivt och sprنcker flaskan. Ammoniumklorat
sِnderfaller sakta i rumstemperatur under bildandet av klorluktande gas. Det
exploderar enligt reaktionsformeln:
2NH4ClO3 ---> N2 + 3H2O + 2HCl + 3/2O2
Detonationshastigheten vid densiteten 0,9 g/cm³ fanns vara 3300 m/s och
blyblocksexpansionen 250cm³. Det tar eld frهn en stubin, och inneslutet sه
exploderar det. Skall man som ett experiment tillverka ammoniumklorat gِr man
som fِljer: 3,5 g kaliumklorat lِses 10 ml varmt destillerat vatten. Tillsنtt
till lِsningen 3,6 g ammoniumperklorat och rِr om. Kristaller av kaliumperklorat
bildas. Lهt lِsningen svalna och filtrera bort kristallerna. Koncentrera
lِsningen till 3 ml pه ett vattenbad. Kyl lِsningen och filtrera ut bildad
ammoniumklorat och torka det i rumstemperatur. Utbytet blir 1.9 g, dvs 55-63%
utbyte.
urea chlorate sounds very intersting but probably quite unstable (more than ammonium chlorate)
NH4ClO4 from NaClO4 ??
chemoleo - 26-7-2003 at 21:36
I always wondered this- how can I make NH4ClO4 from sodium perchlorate, if this the only starting material I got? I mean, NH4ClO4 without any other
salts in it of course (like, as vulture said 2KClO3 + (NH4)2SO4 -> 2NH4ClO3 + K2SO4, but how do I get rid of the potassium sulphate?) . I wouldnt
want to risk distilling perchloric acid (generated by adding H2So4 to NaClO4), which I then would add to ammoniumbicarbonate etc. The point is, it
seems almost impossible to make NH4ClO4 without acess to the free acid. I checked, NH4ClO4 has unfortunately a rel. high solubility, so preciptiation
is (unless under highly controlled conditions) no option. Any ideas??vulture - 27-7-2003 at 02:19
sure, but the point is, how do I get rid of the Na2SO4? that's what I really meant. I can't seem to find a way to make it pure, wthout any
other salts. (unless distillation, but thats a bit risky..)vulture - 28-7-2003 at 03:58
Freezing it out? IIRC, 70% perchloric acid freezes below -20C. At that temperature the solubility of Na2SO4 in 70% acid should be very low.
Don't use a paper filter! Glasswool is recommended.
at 0 c per 100 gr H2O solubility
NaClO4 167 gr
NaNO3 73 gr
NH4ClO4 12 gr
NH4NO3 118 gr
shure, you wont get 100 % yeild since some NH4ClO4 will remain in solution, but then it is just to add more NaClO4 or NH4NO3 to reduce the amount.
No tinkering about with perchloric acid
That is nasty stuff.
/rickard
/rickardchemoleo - 28-7-2003 at 16:05
cool. the precipitation method at 0deg C may be a good one. will try soon. by the way, where did u get the solubility data from? online? if so, whats
the URL?
thxrikkitikkitavi - 29-7-2003 at 00:27
I got the data from Langes handbook och chemistry. But from the ftp (when it is up agian) you can download Perrys Handbook of chemical engineering.
Chapter 2 in that book contains lots of physical data about various compounds, including solubilites for lots of organic and inorganic compounds.
/r
to chemoleo
the timeless - 4-9-2003 at 01:17
In the name of Jesus! Don´t distill explosive materials! It is suicidal, even if you filled the column with nitrogen!chemoleo - 4-9-2003 at 05:53
lol... who said I would? I am not totally suicidal, insane or retarded No, thats
not open to discussion vulture - 4-9-2003 at 06:38
I hate it when people read the first lines of a thread and then go into a jerky daisy chain of negative thoughts: perchloric acid ---> reaction
---> BOOM!Marvin - 14-9-2003 at 13:21
There is a complete synthesis of ammonium chlorate and barium chlorate (via ammonium) in Inorganic Synthesis from potassium chlorate and either
ammonium sulphate or chloride (I forget). This method is the better part of a century old though and this would be regarded as highly unsafe in
modern terms.
A similar method should be possible for perchlorate, but this is much less soluable, so this needs to be modified.
2NaClO4 + H2SO4 ---> 2HClO4 + Na2SO4
This sounds badly thought out. Perchloric acid is a strong acid, hydrogen sulphate isnt and is probably much more soluable. The general idea of
making an acid by precipitating a salt seems a little unsubtle.
From a neutral solution one way to cause sulphates to crash out is by adding ethanol. Though it doesnt make it any easier to control which sulphate
is the one that precipitates in a mixture and yes, this is the complete reverse of 'salting out'.
rickard, ammonium perchlorate is probably the least soluable, but taking the lowest value of a list isnt reliable for predicting which one will
precipitate first.vulture - 14-9-2003 at 22:56
Industrial method for NH4ClO4 is the following:
NaClO4 + HCl + NH3
Don't ask me how this actually works, but it is being done this way.sylla - 6-8-2004 at 12:36
I've done it and it's really easy. HCl + NH3 gives NH4Cl but at a lab scale we can use already done NH4Cl.
So, first prepare a saturated solution of aquous NH4Cl (about 50g/l at 20°C if I'm not wrong) and then add the correct amount of NaClO4 (which
is very soluble, 2Kg/l).
a part of the NH4ClO4 precipitate quickly due to its low sollubility (20g/l). You can be sure it is not NaCl because of its higher solubility.
Then just filter the ammonium perchlorate. It is a good idea to rince it with cold ethanol to remove any other salts.chemoleo - 6-8-2004 at 13:07
Doesnt this separation base on the common ion effect? I.e. surplus NH4Cl aids the precipitation of the nascent NH4ClO4.sylla - 6-8-2004 at 14:44
even if NaClO4 is more difficult to get, I would say it is a better idea to use an excess of perchlorate ion knowing it's very high solubility.
Actually, we are pretty limited by the solubility of ammonium chloride. One should try with more soluble salt such as ammonium nitrate.
Unfortunately it can't be done with ammonia because of beeing essentialy NH3 and NOT NH4+ (I mean, at 20°C of course...)hodges - 14-3-2005 at 16:07
I am in the process of making a small amount of ammonium chlorate as an experiment.
Everything is homemade, so I'm not positive about the quality of the chemicals. I still have a few grams KClO3 I made by electrolysis over a
year ago. I'm pretty sure that this is fairly good quality - it burns in a flash when mixed with sugar and that mixture also ignites in contact
with concentrated H2SO4. I made ammonium sulfate by adding sulfuric acid drain cleaner to clear household ammonia. I used an excess of ammonium
hydroxide, titrating until phenylthalian indicator was a very faint pink. The idea was to avoid forming the bisulfate. According to the CRC
handbook, ammonium sulfate does not decompose until 250C so I figured I was safe drying in an oven at 200C. However, later I read that ammonium
sulfate decomposes somewhat to the bisulfate in boiling water. So I'm not sure if I ended up with the sulfate or the bisulfate.
The replacement reaction I used was
(NH4)2SO4 + 2KClO3 ---> 2NH4ClO3 + K2SO4
132 2 * 123 2 * 101 176
I dissolved 0.3g of (NH4)2SO4 in the minimum amount of water required to dissolve at room temperature. Likewise I dissolved 0.6g of KClO3. I mixed
the solutions together. I added about an equal volume of denatured ethyl alcohol. The K2SO4 gradually separated out, and I moved the remaining
liquid to a small plastic cup using a pipette.
The liquid is presently drying. A drop of it dried on the side of the cup, so I scraped off the white crystal and ignited it with a barbecue lighter.
It burned with a slight fizz (may not have been completely dry) and produced a small puff of white smoke.neutrino - 14-3-2005 at 16:47
Quote:
Originally posted by hodges
I used an excess of ammonium hydroxide, titrating until phenylthalian indicator was a very faint pink. The idea was to avoid forming the bisulfate.
The problem here is that ammonium bisulfate is acidic.
ammonium chlorate
chloric1 - 15-3-2005 at 05:46
Well I happy to hear that you are having some success. I believe this was a method listed in "Handbook of Preparative Inorganic Synthesis"
actually this was a step in the barium chlorate proceedure if I understand correctly. Somewhere in my infinite stack of photocopies and articles I
have a proceedure that basically states that you can take a chlorate and add 70% nitric acid and evaporate then repeat a couple of times and you have
a 30% conversion to perchlorate.hodges - 15-3-2005 at 15:25
I ended up with 0.6g of completely transparent crystals. I understood they were supposed to be "small needles", and I imagined them looking
like pine needles. Instead, they are small and thin but fat prisms with a point on the end. They are about 1/4 inch long and about as wide as they
are fat. Here is the shape they look like:
=>
I tried to take a picture but they do not show up very well.
I placed a few of these on the end of a popsicle stick and lit with a barbecue lighter. They did not explode, but smouldered for several seconds,
occasionally flaring up and releasing a relatively large puff of white smoke. The smoke had a very foul smell - not that of chlorine but closer to
chloramines.
I don't presently detect any odor from the crystals, and no visible vapor is being given off. I measured their temperature with an infrared
thermometer and they are at the same temperature as their surroundings. Will leave them in a safe place near the sink and see if they start to heat
up or otherwise change.hodges - 19-3-2005 at 10:01
I'm not real convinced of the quality of my ammonium chlorate. For one thing, the "needles" are not like I had imagined them based on
their description. And I didn't see any explosive qualities. I did detect chlorine oxides by smell when I burned some, though.
I made the KClO3 myself. Its possible that it contains KClO4. I know there is some KClO3 since a mixture with sugar ignites with concentrated
sulfuric acid.
The ammonium sulfate is another unknown. One source lists its decomposition temperature as 250C, but another says it turns to the bisulfate when a
water solution of it is boiled. Since I dried it in the oven that may be what happened. Would there be an easy way to test for ammonium bisulfate
vs. ammonium sulfate? One way I can think of, which would not be particulaly quick, would be to weigh some , add concentrated NH4OH and evaporate,
then weigh again. The weight should increase slightly if it changed from bisulfate to sulfate.
NH4SO4 Decomp
kryss - 27-3-2005 at 07:43
We dry amm sulphate in an oven at 105 degs C , its a primary standard it doesnt decompose until much higher - it melts and forms bisulphateMr. Wizard - 27-3-2005 at 12:10
Ammonium chlorate is a precursor in the electrolytic production of the Per-chlorate. It is very unstable in the sense it tends to ignite anything
combustible. Workers have told me stories about splashing the solution on their pants, boots or shirt and then the clothes would catch fire later
from the friction of rubbing them. Wooden pallets would catch fire when they were touched. There WAS an Ammonium Perchlorate plant where I
lived, and as you can see by the attached URL, even the per-chlorate can be rather unpredictable. http://www.apechild.com
then go to the archives
then go to Feb 2004
look at the pepcon video
This was a big ass explosion, watch for the second, even bigger explosion. Sorry about the convoluted links, but a direct link doesn't seem to
work.
[Edited on 27-3-2005 by Mr. Wizard]hodges - 6-4-2005 at 17:06
I tried this again using an alternate synthesis (KClO3 and NH4ClO4, the KClO4 is much less soluable than any other product). I again found that the
resulting salt was not explosive when heated. Instead, it decomposed somewhat vigorously - about like ammonium dichromate or silver oxalate.
Certainly no bang or anything though. I'm guessing that this salt is not all that energetic unless mixed with something else (which would of
course be the case when it was produced acidentally in fireworks compositions). I think the real danger is probably its instability - i.e. it tends
to decompose by itself after a period of time, releasing chlorine dioxide which would ignite any organic matter nearby. I burned all the ammonium
chlorate I made but later will try making a small amount and keeping it around in a small disposable plastic cup in the sink to see if it ever
decomposes and if so how violently.
No, thats
not open to discussion 