desman
Harmless
Posts: 5
Registered: 9-10-2018
Member Is Offline
|
|
How to make anhydrous NaClO3 from hydrated NaClO3 ?
Hi guys,
Is there a technique that can make anhydrous NaClO3 from hydrated NaClO3 ?
Thanks.
|
|
|
fusso
International Hazard
Posts: 1922
Registered: 23-6-2017
Location: 4 ∥ universes ahead of you
Member Is Offline
|
|
Heat it above 100C but below 300C until all water evaporated. Done.
|
|
|
desman
Harmless
Posts: 5
Registered: 9-10-2018
Member Is Offline
|
|
The H2O molecule is trapped with NaClO3.
You can heat wet KClO3 to get anhydrous KClO3 because KClO3 is not hygroscopic. But in this case NaClO3 is very hygroscopic (attracted to water).
|
|
|
fusso
International Hazard
Posts: 1922
Registered: 23-6-2017
Location: 4 ∥ universes ahead of you
Member Is Offline
|
|
Quote: Originally posted by desman  |
The H2O molecule is trapped with NaClO3.
You can heat wet KClO3 to get anhydrous KClO3 because KClO3 is not hygroscopic. But in this case NaClO3 is very hygroscopic (attracted to water).
|
Therefore you need to heat it to boil away the water. This is a good way to make many substances
dry/anhydrous.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by fusso | | Quote: Originally posted by desman |
The H2O molecule is trapped with NaClO3.
You can heat wet KClO3 to get anhydrous KClO3 because KClO3 is not hygroscopic. But in this case NaClO3 is very hygroscopic (attracted to water).
|
Therefore you need to heat it to boil away the water. This is a good way to make many substances
dry/anhydrous. |
There are compounds which will decompose if you try to dehydrate them by heating, but those are mostly compounds in which the metal has a high
oxidation state, so the oxide is very, very stable (eg, aluminum salts, chromium(III) compounds, etc). Sodium, however, doesn't have a high oxidation
state, so you really don't have to worry about that. Just don't get it so hot that the chlorate ion decomposes (which happens around 300 C).
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I would not recommend heating a large quantity of possibly impure hydrated NaClO3, could detonate.
Reason, with larger quantities heating is unequal.
Also, with say a copper impurity, the chlorate can becomes a more sensitized energetic compound!
---------------------------------
Try mixing with alcohol and passing dried air into the solution.
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
Desication
Using sodium hydroxide
Or a stronger desacant?
Passing warm air through the solution
Plastic welders or tiny hair dryer might work just need to attach a hose and bubble stone.
Distill off the water with reduced pressure
Vacuum distill
[Edited on 10-10-2018 by symboom]
[Edited on 10-10-2018 by symboom]
[Edited on 10-10-2018 by symboom]
|
|
|
Tsjerk
International Hazard
Posts: 3022
Registered: 20-4-2005
Location: Netherlands
Member Is Offline
Mood: Mood
|
|
| Quote: Originally posted by AJKOER | I would not recommend heating a large quantity of possibly impure hydrated NaClO3, could detonate.
Reason, with larger quantities heating is unequal.
Also, with say a copper impurity, the chlorate can becomes a more sensitized energetic compound!
---------------------------------
Try mixing with alcohol and passing dried air into the solution. |
NaClO3 will not detonate. Also not when impure. Unequal heating definitely is not the reason for this. It just decomposes to NaCl and oxygen.
[Edited on 10-10-2018 by Tsjerk]
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
| Quote: Originally posted by AJKOER | I would not recommend heating a large quantity of possibly impure hydrated NaClO3, could detonate.
Reason, with larger quantities heating is unequal.
Also, with say a copper impurity, the chlorate can becomes a more sensitized energetic compound!
---------------------------------
Try mixing with alcohol and passing dried air into the solution. |
This is not an issue. NaClO3 will nog detonate nor self-ignite when heated. It can decompose (even violently so), but only at high temperatures (250+)
in the presence of certain catalysts. When you heat the solid to 130 C or so in an oven and occassionally stir the solid mass and crumble it to expose
other parts of the solid, then the material will become nicely dry.
NaClO3 is somewhat hygroscopic, but not insanely so. NaClO4 is another matter. It nearly is impossible to store that in a normal container. I once had
some NaClO4 but it had changed into a wet mud. A few years ago I used up all of it in an experiment to make transition metal complexes, which could be
well crystallized as perchlorate salt.
|
|
|
desman
Harmless
Posts: 5
Registered: 9-10-2018
Member Is Offline
|
|
| Quote: Originally posted by woelen | | Quote: Originally posted by AJKOER | I would not recommend heating a large quantity of possibly impure hydrated NaClO3, could detonate.
Reason, with larger quantities heating is unequal.
Also, with say a copper impurity, the chlorate can becomes a more sensitized energetic compound!
---------------------------------
Try mixing with alcohol and passing dried air into the solution. |
This is not an issue. NaClO3 will nog detonate nor self-ignite when heated. It can decompose (even violently so), but only at high temperatures (250+)
in the presence of certain catalysts. When you heat the solid to 130 C or so in an oven and occassionally stir the solid mass and crumble it to expose
other parts of the solid, then the material will become nicely dry.
NaClO3 is somewhat hygroscopic, but not insanely so. NaClO4 is another matter. It nearly is impossible to store that in a normal container. I once had
some NaClO4 but it had changed into a wet mud. A few years ago I used up all of it in an experiment to make transition metal complexes, which could be
well crystallized as perchlorate salt. |
I'm curious, is it possible to make anhydrous NaClO4 from hydrated NaClO4 by heating or vacuum ?
|
|
|
fusso
International Hazard
Posts: 1922
Registered: 23-6-2017
Location: 4 ∥ universes ahead of you
Member Is Offline
|
|
| Quote: Originally posted by desman | | Quote: Originally posted by woelen | | Quote: Originally posted by AJKOER | I would not recommend heating a large quantity of possibly impure hydrated NaClO3, could detonate.
Reason, with larger quantities heating is unequal.
Also, with say a copper impurity, the chlorate can becomes a more sensitized energetic compound!
---------------------------------
Try mixing with alcohol and passing dried air into the solution. |
This is not an issue. NaClO3 will nog detonate nor self-ignite when heated. It can decompose (even violently so), but only at high temperatures (250+)
in the presence of certain catalysts. When you heat the solid to 130 C or so in an oven and occassionally stir the solid mass and crumble it to expose
other parts of the solid, then the material will become nicely dry.
NaClO3 is somewhat hygroscopic, but not insanely so. NaClO4 is another matter. It nearly is impossible to store that in a normal container. I once had
some NaClO4 but it had changed into a wet mud. A few years ago I used up all of it in an experiment to make transition metal complexes, which could be
well crystallized as perchlorate salt. |
I'm curious, is it possible to make anhydrous NaClO4 from hydrated NaClO4 by heating or vacuum ?
|
yes
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Per Bretherick Vol 1 (free to download online) to quote page 1371:
"Although most explosive incidents have involved mixtures of the chlorate with combustible materials, the exothermic decomposition of the chlorate to
chloride and oxygen can accelerate to explosion if a sufficient quantity and powerful enough heating are involved. A case history of a fire-heated
explosion of a store of 80 t of chlorate is given. The more stable sodium chlorate will also explode under similar conditions [1]. "
So yes, the oven heating is a very good idea for larger quantities, avoids 'powerful heating' scenarios.
However, electrolysis prepared chlorate in the presence of residual metal electrode corrosion products, I would have concerns. It is reported that
perchlorate salts acidified (with a stronger acid than HClO4) in the presence of transition metals (or organic material) presents an explosion hazard
(see page 145). Perhaps true also with chlorates.
----------------------------------------
Thanks for this thread as my examination of the action of say hydroxyl radicals on aqueous ClO3- looks very interesting:
.OH + ClO3- = OH- + .ClO3
.ClO3 + .ClO3 = Cl2O6 = [ClO2]+[ClO4]−
and reactions with dichlorine hexoxide apparently proceeds with the creation of perchlorates and products (see https://en.wikipedia.org/wiki/Dichlorine_hexoxide).
-----------------------------------------------------------
A side comment on an interesting scenario which may have involved more chemistry than just heating aqueous NaOCl by itself to form chlorate, which
normally produces very little NaClO3. To quote from p, 1388:
"A saucepan of vegetable stew had been heated too long and had formed a thick carbonised adherent cake. In an attempt to clean the pan, 1 l of
domestic bleach was added and the pan was left to heat on an electric hotplate. Again it was left too long and after all the water had evaporated, the
residue exploded violently."
So a possible more complex chlorate formation path, per my recent comments on the ability of activated carbon to reduce ferric (formed by the attack
of the steel or more likely cast iron by food acids) to ferrous. Also, residual food acids could make some hypochlorous acid from the bleach. Next,
the interaction of Fe(ll) and a mix of HOCl/NaOCl in a fenton-type reaction:
Fe(ll) + HOCl = Fe(lll) + .OH + Cl- (since pH >5 in presence of excess NaOCl)
.OH + ClO- = OH- + .ClO
.OH + .ClO = H+ + ClO2-
ClO2- + HOCl/NaOCl = ClO3- + HCl/NaCl
HCl + NaOCl = HOCl + NaCl
Fe(lll) + C* = Fe(ll) + C**
and a possible reaction cycles continues increasing chlorate concentration.
[Edited on 10-10-2018 by AJKOER]
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Thanks, I will bear it in mind the next time I plan to heat a few truck loads of chlorate.
|
|
|
Tsjerk
International Hazard
Posts: 3022
Registered: 20-4-2005
Location: Netherlands
Member Is Offline
Mood: Mood
|
|
Fe(lll) + C* = Fe(ll) + C**
You are gaining two electrons there
[Edited on 10-10-2018 by Tsjerk]
[Edited on 10-10-2018 by Tsjerk]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Tsjerk | Fe(lll) + C* = Fe(ll) + C**
You are gaining two electrons there
[Edited on 10-10-2018 by Tsjerk]
[Edited on 10-10-2018 by Tsjerk] |
I try to describe the action of AC to resemble a Fe/Cu redox couple:
Fe(lll) + Cu(l) = Fe(ll) + Cu(ll)
The original source (see 'Enhanced production of reactive oxidants by Fenton-like reactions in the presence of carbon materials' by Jiwon Seo, et al,
at https://www.researchgate.net/publication/276120744_Enhanced_... , scroll down) Eq (3) uses the notation:
Fe(lll) + C red = Fe(ll) + C oxid
The authors do note that "The fundamental mechanism of the Fe(III) reduction by carbon materials has not yet been fully established, but literature
has reported that carbon materials have reducing power."
|
|
|
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
Do you NEED it as solid Sodium Chlorate ?. If you need the solid it is much more convenient to convert it to KClO3. KClO3 is sparingly soluble and can
be precipitated from a solution of NaClO3 by adding any soluble potassium salt. It is also easy to dry and obtain anhydrous.
If you're not part of the solution, you're part of the precipitate.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
