chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
Sodium Chlorate and Potassium Chlorate
hi there,
thought I miight contribute this for some needy souls out there.
First off, NaClO3 is often found in garden shops, normally 75%, as the main weed killer ingredient. We even get it in the UK (which is tight on such
things), so they certainly have it in the US. In France, you get 99% for a few francs ahem euros.
Anyway, it is actually not hard to make it from normal cooking salt yourself.
Just dissolve cooking salt NaCL in water, put it in a jar with a lid and a pressure release hole. Make two holes into the lid and put two graphite
electrodes into it (of course you can use steel etc but it gets dissolved eventually). Seal the sides on the electrodes, so that no gases can escape
(i used silicon). Then electrolyse this at 40 Volts, more, or less. Solution will hopefully become hot, and acquire a green hue. even graphite gets
corroded by the nascent chlorine, so you will have black particles floating about. dont worry about this for now (you can filter them once the
electrolysis is complete).
Continue electrolysing this until .. well u dont have masses of chlorine coming off the anode anymore... I myself did it for several hours. make sure
taht the chlorine remains in the jar, you obvioulsy need it for forming hypochlorite (i.e. Cl2 + 2NaOH --> NaOCL +NaCl ). This, during the second
round, forms chlorite (i.e. NaClO2) and finally chlorate (NaClO3). So, the only gases that should escape your electrolysis cell is hydrogen and
ogygen. Once electrolysis is complete (as I said, I never tested how long this has to be done... rather do it as long as u can), dry it in an oven at
50 deg C, or whatver you chose. It is extremely hygroscopic, so store in an airtight jar (so drying wont really work at room temp).
Testing it - just mix with icing sugar (1:1 by volume) and ignite... and enjoy the yellow light of sucess.
By the way, when the NaClO3 is melted until O2 bubbles form, sodium perchlorate is produced which may be useful for the below.
Potassium Chlorate/ Perchlorate
This is where the neat stuff begins. Dissolve stoichiometric (i.e. equimolar) amounts of potassium carbonate K2CO3 and NaClO3 in solution. The
resulting sodium carbonate Na2CO3 is extremely soluble, while the KClO3 is not. The same holds for the potassium perchlorate KClO4 (if you do the melt
as mentioned above, the resulting KCl is soluble, too, so no problem). Thus, filter the solution and wash with ice cold water.
Purity is tested easily  - FLAME TEST!
just mix with sugar, or sulfur or whatever, and see whether the flame is still yellow. It shouldnt be of course, purple/violet white is the colour you
should get.
I managed to have it at high yields (>500g), and made beautiful little fireworks with Calcium (red) or barium (green). Of course you can do much
more than this.
have fun 
|
|
|
blip
Hazard to Others
Posts: 133
Registered: 16-3-2003
Member Is Offline
Mood: absorbed
|
|
You can also use KCl and most other potassium salts, not just K<sub>2</sub>CO<sub>3</sub> for those beginners out there.
I have a question kind of related to this on making pure chlorites. Could ClO<sub>2</sub> simply be reacted with a metal to form a metal
chlorite? What about 2NH<sub>3</sub> + H<sub>2</sub> + 2ClO<sub>2</sub>
<sup><u> </u></sup>> 2NH<sub>4</sub>ClO<sub>2</sub> for ammonium chlorite?
2NH<sub>4</sub>ClO<sub>2</sub> <sup><u> <font
face="symbol">D</font> </u></sup>> N<sub>2</sub>Cl<sub>2</sub> +
4H<sub>2</sub>O
N<sub>2</sub>Cl<sub>2</sub>'s structure:
Cl-N=N-Cl
Sometimes polymers of that may be created, although N<sub>2</sub>Cl<sub>2</sub>'s formation gives an increase in entropy.
Maybe it'd just produce nitrogen and chlorine gases, but that'd be a bummer.
E.g.:<tt>
Cl
|
Cl-N-N-N-Cl
</tt>
Also, chlorites likely have other interesting uses.
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
no... as far as I remember I once played with ClO2 and I don't recommend it. ClO2 is so reactive there is no way it will simply react with a
metal to form chlorite. Instead, it will simply decompose and oxidise (in the form of radicals), possibly chlorinate.
N2Cl2 is definitely very unlikely to exist, simply because the N=N bond is unstable by itself, and the chlorine is bound to rip electrons out of the
double bond system, making it even more unstable. even NCl3 by itself is extremely unstable. Making polychlorinated N-compounds are therefore
probably impossible to make  sorry
I wouldnt want to play with such compounds anyway! I think the maximum repeat of nitrogens is 3 (i.e. single bonded, N-N-N) and this is in organic
compounds that stabilise it. hope i am not talking bollocks 
|
|
|
blip
Hazard to Others
Posts: 133
Registered: 16-3-2003
Member Is Offline
Mood: absorbed
|
|
Umm, doesn't hydrazine have a nitrogen double bond? And why would chlorine take electrons from that double bond when it's already sharing
one with nitrogen to complete its octet? I know NCl<sub>3</sub> is unstable, but I heard that NHCl<sub>2</sub> and
NH<sub>2</sub>Cl aren't. Btw, I was also thinking that NH<sub>4</sub>OCl's decomposition could produce
NH<sub>2</sub>Cl.
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
hydrazine is NH2NH2 as far as i remember.
the -N=N- bond does exist of course, i.e. in heteroaromatics and in C6H6-N=N-C6H6 - this is used for many syntheses. But Cl-N=N-Cl??? Never heard of
this compound and whould think its unlikely to be even remotely stable. the octett isnt complete then, btw, chlorine will be Cl- and nitrogen will be
[N=N]2+ - as the chlorine is so electronegative it is unlikely to share it well with the nitrogen to form an ordinary covalent bond. So thats why the
chlorines are at least disturbing the double bond system by favouring the formation of N2 (normal nitrogen gas triple bonded) and Cl2. Also, stability
does not necessarily mean an energetically well-balanced molecule, as it still may be highly reactive. just like HNCl2 and H2NCl
also seem to remember that ammoniumhypochlorite is unstable, think it will disproportionate to nH4Cl and chlorate/perchlorate. could be completely
wrong on that one tho
[Edited on 28-7-2003 by chemoleo]
[Edited on 28-7-2003 by chemoleo]
|
|
|
blip
Hazard to Others
Posts: 133
Registered: 16-3-2003
Member Is Offline
Mood: absorbed
|
|
Yeah, sorry about that, I was thinking hydrazine was HN=NH for some weird reason. Hypochlorite disproportionates into chlorates/perchlorates and
chlorides when the solution is warm. Ok, nitrogen's electronegativity is <a
href="http://www.webelements.com/webelements/scholar/elements/nitrogen/compounds.html" target="_blank">3.04</a> whereas
chlorine's is <a href="http://www.webelements.com/webelements/scholar/elements/chlorine/compounds.html"
target="_blank">3.16</a> (I could have sworn my chem book said 3.5, but I'm not near it right now) with the difference being
0.12. That would be a non-polar covalent bond meaning there would be no ionization according to:<tt>
0.0-0.4 Non-polar covalent
0.5-2.0 Polar covalent
2.1-4.0 Ionic</tt>
It might not be stable, but on the other hand it may be. I haven't been able to find any info on it. Is there some reason it would be unstable, like something with orbitals (I haven't had time to read about that
kind of stuff)?
I'm pretty sure the atoms share electrons like this:<tt>
·· ·· ·· ··
:Cl··N::N··Cl:
·· ··</tt>
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
ok, sorry I didnt mean to phrase it as to there being an actual [N=N]2+ ion. all I meant was that the electron density will shift to the chlorines,
simply because it is more e-negative and because its orbital can be filled. That leaves you with the partially positive nitrogens, making it
partially +I - and does this exist? the most common ones are -III, +II, +III, +IV, +V. This maybe part of the reason making this unstable.
In terms of orbitals, this would be an sp2 orbital, and planar. an empty electron pair (from the nitrogen) would be situated next to the chlorine, at
a 120 deg angle, with the p orbital sticking out of the plane. Sounds to me that maybe Cl, by sheer spacial electron density, will mess with the free
e- pair next to it and make the system unstable. This is speculation though, I don't really know
[Edited on 29-7-2003 by chemoleo]
|
|
|
tryptamine
Harmless
Posts: 16
Registered: 20-3-2003
Member Is Offline
Mood: elated in love
|
|
my suggestions
A few suggestions for chlorate cells:
- Use a very small interelectrode distance and maximize the surface area of your electrodes.
- Run it at ~3 volts
- As high an amperage as you can muster
- Try and find some titanium and stainless steel for the electrodes(I've got a pile of ti welding rod and some bolts)
- If you use graphite you must filter the final solution through celite and or a micron filter. You don't want finely divided carbon in your
chlorate!
-If you just electrolyse the hell out of it without monitoring it you will over do it and end up with alot of perchlorate.
-In fact the only way to get pure chlorate would be to crystalize out the chlorate and wash to remove the excess salt.
|
|
|
alwynj48
Harmless
Posts: 3
Registered: 2-8-2003
Location: orlando
Member Is Offline
Mood: No Mood
|
|
This is a side subject and I am sorry if it’s a dumb question. I understand that bonding is all about reducing entropy i.e. exchanging electrons and
shearing is lower energy hence the bonding energy. But can you orbital savvy guys explain why you get chlorates and perchlorates. Now aluminates,
borates, plumbates silicates, ect. that seems ok. The core atom is not very electronegative they have spare electrons. But why chlorine and why so
many oxygen ates or is there an equivalent set of ates for other atoms different than oxygen? I am guessing but I assume the answer is similar to the
explanation for inert gas compounds.
One more question are there the usually ates for fluorine?
Bob
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
oxidation of chlorine...
you get chlorates because oxygen is in fact more e-negative(only slightly I think) than chlorine. Thus, chlorine gets oxidised. Due to similar
e-negativities, there are therefore also chlorine oxides, i.e. where the chlorine accepts electrons from the oxygen. THese compounds are generally
very unstable, tho. Chlorates are more stable than chlorine oxides , but still they they decompose at high temps, as you probably know.
There are Bromates and iodates (and the per-ates too), too, which are more stable than chlorates, as far as I remember. Most likely this is due to the
greater e-negativity difference AND the atom size, i.e. that of bromine and iodine.
I once checked whether fluorates exist, and they do, but are superbly unstable, and exist only at extremely low temperatures or something. so nothing
useful for most people
Dont know whether there are "ates" for atoms other than oxygen - like sulphur etc. I doubt they are common though.
dan
|
|
|
blip
Hazard to Others
Posts: 133
Registered: 16-3-2003
Member Is Offline
Mood: absorbed
|
|
I always thought oxygen would bond to the chlorine using coordinate covalent bonds from chlorite to perchlorate, but also have a single bond to a
negatively charged oxygen atom in all four examples (sorry about the sucky ASCII drawings, and resonance structures weren't helpful to me when I
started out):<tt>
hypochlorite
+- -+
|·· ·· | -1
|:O··Cl:|
|·· ·· |
+- -+
chlorite
+- -+
|·· ·· ··| -1
|:O··Cl:O:|
|·· ·· ··|
+- -+
chlorate
+- -+
|·· ·· ··| -1
|:O··Cl:O:|
|·· ·· ··|
| :O: |
| ·· |
+- -+
perchlorate
+- -+
| ·· |
| :O: |
|·· ·· ··| -1
|:O··Cl:O:|
|·· ·· ··|
| :O: |
| ·· |
+- -+</tt>
Sulfur can also do coordinate covalent bonding along with many other elements:<tt>
sulfur dioxide
· · ··
:O::S:O:
· · ··
sulfur trioxide
· ··
:O::S:O:
· ·· ··
:O:
··
sulfate
+- -+
| ·· | -2
| :O: |
| · |
|·· · ··|
|:O··S:O:|
|·· ·· ··|
| :O: |
| ·· |
+- -+
sulfite
+- -+
| ·· | -2
| :O: |
| · |
|·· · ··|
|:O··S:O:|
|·· ·· ··|
+- -+</tt>
| Quote: | | I am guessing but I assume the answer is similar to the explanation for inert gas compounds. |
Yes, e.g. XeO<sub>4</sub> (I'm pretty sure it exists):<tt>
··
:O:
·· ·· ··
:O:Xe:O:
·· ·· ··
:O:
··</tt>
| Quote: | | One more question are there the usually ates for fluorine? |
Chemoleo answered that already, but hypofluorites are known to be stable with the structure of hypochlorites.
I'm pretty sure I caught all of the errors in the diagrams....
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
NaClS4 ???
blip, I think what alwyn meant was whether there are compounds like NaClS4 (i.e. analogous to NaClO4), or NaClSe4 etc. there I thought they are not
particularly common, if they exist.
I know there are things such as FeCuS4, but they are polysulphides rahter than Cu VII +.
We have this thing about covalent/ionic bonds etc lolz - sure they are covalent in the case of the ClO4- ion, but the chlorine atom has a formal
charge of VII+. this is because the chlorine gets oxidised. nothing to do with an ionic bond though.
pz
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
