Sciencemadness Discussion Board

Quick Question-Thermal Reaction of NaOCl

The Volatile Chemist - 12-5-2014 at 08:23

In regards to the reaction that goes on when Sodium Hypochlorite is heated, are the only products NaCl and NaClO3? The Wikipedia article says that -OCl converts to "Chlorine, Oxygen, and various Chlorates". So will bleach turn to "various Chlorates", like perchlorate, or just NaClO3?
Thanks!
Nathan

Zyklon-A - 12-5-2014 at 08:53

It will not turn into perchlorate.
The reaction is as follows: 3 NaClO + heat → 2 NaCl + NaClO3.
This is equation doesn't occur: 4 NaClO + heat → 3 NaCl + NaClO4.
BTW, I've seen sodium hypochlorite written as NaClO and NaOCl. I assume either is correct, but which is more scientific?
I use NaClO because generally I write the more electronegative atoms last.

HgDinis25 - 12-5-2014 at 10:18

Quote: Originally posted by Zyklonb  
It will not turn into perchlorate.
The reaction is as follows: 3 NaClO + heat → 2 NaCl + NaClO3.
This is equation doesn't occur: 4 NaClO + heat → 3 NaCl + NaClO4.
BTW, I've seen sodium hypochlorite written as NaClO and NaOCl. I assume either is correct, but which is more scientific?
I use NaClO because generally I write the more electronegative atoms last.


Oxygen is more electronegative than Chlorine, so the negative Zone of the Ion is in the Oxygem Atmom. Therefore, the ionic bond happens between Sodium+ and -OCl.
Following this logic, NaOCl should be used. However, I think there are specific rules in this cases.

Zyklon-A - 12-5-2014 at 10:27

Ok, thanks. I also thought the opposite because NaCl is a more stable compound then Na2O or NaOH.

blogfast25 - 12-5-2014 at 11:58

Quote: Originally posted by HgDinis25  
Therefore, the ionic bond happens between Sodium+ and -OCl.
Following this logic, NaOCl should be used. However, I think there are specific rules in this cases.


It's now customary to write all oxoanions of the halogens as XO<sub>n</sub><sup>-</sup>, including ClO<sup>-</sup>.

In solid NaClO (which is highly unstable) Na<sup>+</sup> and ClO<sup>-</sup> ions alternate in a three dimensional lattice held together by the electrostatic attraction between the ions. That is what's called ionic bonding.

By your reasoning chlorite should be written as [O<sub>2</sub>Cl]<sup>-</sup> but we don't do that either.

The Volatile Chemist - 12-5-2014 at 14:02

I use both interchangably, but I guess it should be NaClO. I think Oxychloride when I see OCl- and Hypochlorite when I see ClO-, but I don't think it really matters.

HgDinis25 - 12-5-2014 at 14:41

Quote: Originally posted by blogfast25  
Quote: Originally posted by HgDinis25  
Therefore, the ionic bond happens between Sodium+ and -OCl.
Following this logic, NaOCl should be used. However, I think there are specific rules in this cases.


It's now customary to write all oxoanions of the halogens as XO<sub>n</sub><sup>-</sup>, including ClO<sup>-</sup>.

In solid NaClO (which is highly unstable) Na<sup>+</sup> and ClO<sup>-</sup> ions alternate in a three dimensional lattice held together by the electrostatic attraction between the ions. That is what's called ionic bonding.

By your reasoning chlorite should be written as [O<sub>2</sub>Cl]<sup>-</sup> but we don't do that either.


In an Ionic Bond, the main interaction happens between the Positive Pole of the Cation and the Negative Pole of the Anion (yes Anions can have Positive Poles).
Quote:

By your reasoning chlorite should be written as [O2Cl]- but we don't do that either.


True, that's why I said there had to be special rules for this sort of cases.

blogfast25 - 13-5-2014 at 04:53

Quote: Originally posted by HgDinis25  
In an Ionic Bond, the main interaction happens between the Positive Pole of the Cation and the Negative Pole of the Anion (yes Anions can have Positive Poles).


True, that's why I said there had to be special rules for this sort of cases.


You're over simplifying: have a look at the calculation of the Madelung constant for ionic lattices.


Special rules? Writing hypochlorite as OCl<sup>-</sup> would be an exception to a simple rule: the more electronegative element follows the less electronegative one. See also chloro and fluoro anions. Even oxo cations (TiO<sup>2+</sup> e.g.) [Edit: 'cations' instead of 'anions', of course]

No, 'OCl<sup>-</sup>' is merely a relic from chemistry's rich and convoluted past.

[Edited on 13-5-2014 by blogfast25]

The Volatile Chemist - 13-5-2014 at 05:26

Quote: Originally posted by blogfast25  
Quote: Originally posted by HgDinis25  
In an Ionic Bond, the main interaction happens between the Positive Pole of the Cation and the Negative Pole of the Anion (yes Anions can have Positive Poles).


True, that's why I said there had to be special rules for this sort of cases.


You're over simplifying: have a look at the calculation of the Madelung constant for ionic lattices.


Special rules? Writing hypochlorite as OCl<sup>-</sup> would be an exception to a simple rule: the more electronegative element follows the more electronegative one. See also chloro and fluoro anions. Even oxo anions (TiO<sup>2+</sup> e.g.).

No, 'OCl<sup>-</sup>' is merely a relic from chemistry's rich and convoluted past.


What's wrong with remembering our past? :)

Zyklon-A - 13-5-2014 at 05:58

Nothing is wrong. It's just completely useless to use outdated formulas. And it makes you look like you don't know what you're talking about - hence the reason I wanted to know the right way. Thanks Blogfast.

The Volatile Chemist - 13-5-2014 at 08:00

Quote: Originally posted by Zyklonb  
Nothing is wrong. It's just completely useless to use outdated formulas. And it makes you look like you don't know what you're talking about - hence the reason I wanted to know the right way. Thanks Blogfast.

I see. I'll use -ClO from now on (Or at least try, I forget a lot.)

blogfast25 - 13-5-2014 at 10:39

I've noticed that for sodium chlorate (NaClO<sub>3</sub>;), some have started to use 'sodium chlorate(V)' and for sodium perchlorate (NaClO<sub>4</sub>;), 'sodium chlorate(VII)'. I don't know if this is a IUPAC initiative or what. If so, perhaps one day the recommended nomenclature will be: 'chlorate(I)', 'chlorate(III)' etc.

Who knows?

Zyklon-A - 13-5-2014 at 11:18

It seems logical to use "potassium chlorate(VII)" rather than potassium perchlorate", because it isn't even really a peroxide at all. It contains no O-O single bonds.

Potassium_perchlorate.png - 5kB

woelen - 14-5-2014 at 00:45

If you want to write formally correct formulas, then you would have to write

NaOCl for sodium hypochlorite
NaOClO for sodium chlorite
NaOClO2 for sodium chlorate
NaOClO3 for sodium perchlorate

In practice, however, one usually writes NaClOx (1 <= x <= 4) for the different compounds.

I never have seen names like sodium chlorate(III) for the chlorite or sodium chlorate(VII) for the perchlorate in any paper. Not in practical texts (e.g. pyrotechnics stuff), nor in scientific papers. But indeed, formally, these are better names, but old traditions are very hard to change. Especially the name "perchlorate" is a bad one, the ion does not have any peroxo-group, the prefix "per" only tells that it has extra oxygen, compared to chlorate. If these ions were discovered nowadays, they most likely would have received different, more formal, names.

AJKOER - 15-5-2014 at 10:45

Quote: Originally posted by The Volatile Chemist  
In regards to the reaction that goes on when Sodium Hypochlorite is heated, are the only products NaCl and NaClO3? The Wikipedia article says that -OCl converts to "Chlorine, Oxygen, and various Chlorates". So will bleach turn to "various Chlorates", like perchlorate, or just NaClO3?
Thanks!
Nathan


In reality of the lab, one cannot heat solid Sodium hypochlorite. It just explodes, so the question of chlorate formation without specifying a solid or an aqueous reaction is one, possibly fatal, mistake. As an aqueous disproportionation in the lab for this Sodium salt is only likely in a hot, highly concentrated and ionic solution, alas a few more oops. In sunlight, a wholly different path to chlorate and even some perchlorate creation takes place.

There are also compounds that can accelerate Chlorate formation (like acetates and chlorides, but generally neutral chlorides are needed to avoid decomposition of any formed chlorate as pH is another factor). The optimal temperature range for aqueous reactions is also important (generally, below boiling).

But, do not take my word for it, seach SM and the web for actual quality references.

Please be wary of arrogance educators with good teaching skills but, at times, dangerously deficient in actual knowledge of chemistry. I speak as a US native who have to endure a high school chemistry teacher, in an otherwise good suburban school, who also taught gym.

[Edited on 15-5-2014 by AJKOER]

The Volatile Chemist - 15-5-2014 at 12:26

Quote: Originally posted by AJKOER  
Quote: Originally posted by The Volatile Chemist  
In regards to the reaction that goes on when Sodium Hypochlorite is heated, are the only products NaCl and NaClO3? The Wikipedia article says that -OCl converts to "Chlorine, Oxygen, and various Chlorates". So will bleach turn to "various Chlorates", like perchlorate, or just NaClO3?
Thanks!
Nathan


In reality of the lab, one cannot heat solid Sodium hypochlorite. It just explodes, so the question of chlorate formation without specifying a solid or an aqueous reaction is one, possibly fatal, mistake. As an aqueous disproportionation in the lab for this Sodium salt is only likely in a hot, highly concentrated and ionic solution, alas a few more oops. In sunlight, a wholly different path to chlorate and even some perchlorate creation takes place.

There are also compounds that can accelerate Chlorate formation (like acetates and chlorides, but generally neutral chlorides are needed to avoid decomposition of any formed chlorate as pH is another factor). The optimal temperature range for aqueous reactions is also important (generally, below boiling).

But, do not take my word for it, seach SM and the web for actual quality references.

Please be wary of arrogance educators with good teaching skills but, at times, dangerously deficient in actual knowledge of chemistry. I speak as a US native who have to endure a high school chemistry teacher, in an otherwise good suburban school, who also taught gym.

[Edited on 15-5-2014 by AJKOER]


I was actually planning on using off-brand bleach, I don't have powdered NaClO. I was going off of an experiment Mr. Home Scientist posted on.

Zyklon-A - 15-5-2014 at 12:43

Yeah... Have any of you guys tried (maybe even succeeded:o) in making pure NaOCl/NaClO?

The Volatile Chemist - 15-5-2014 at 13:08

No. Could you get small amounts by chilling it so some crystallizes? On the subject, I saw it sold at a hardware store, something like a chemical for a water purifier... or something....

blogfast25 - 16-5-2014 at 10:33

Quote: Originally posted by The Volatile Chemist  
No. Could you get small amounts by chilling it so some crystallizes? On the subject, I saw it sold at a hardware store, something like a chemical for a water purifier... or something....


What you saw was probably calcium hypochlorite, not sodium hypochlorite.

Zyklon-A - 16-5-2014 at 11:27

Quote: Originally posted by The Volatile Chemist  
No. Could you get small amounts by chilling it so some crystallizes?

No, not with a household bleach solution. I have heard that a 12% solution is possible to get OTC, but I've never found it. (And even with 12% it won't work)
The highest I can find is 8.25%. At 0°C ~29 grams dissolve in 100 mL of water, so you'll need to concentrate the solution first. I haven't been able to find any solubility curve over different temperatures, all I could find is Wikipedia's info which convenitly is at 0°C. Does anyone have any data on this? I don't know how to do this, as I've heard that it will form the chloride and chlorate ( 3 NaClO + heat → 2 NaCl + NaClO3), even at low temps if you try to evaporate it. I don't have a ref. for this, so I might be wrong.

One Idea I have is dissolve as much sodium chloride as possible in household bleach solution, and preform electrolysis at < 30°C. Continue to add NaCl as it's used up to maintain saturation. As soon as you see crystals, (Which are not NaCl) stop electrolysis, and cool to 0°C and collect pure NaOCl.
This should absolutely not be done inside or in a glass vessel, and only on a very small scale. Not only does NaOCl explode upon heating, but it's quite shock sensitive.
Does anyone know what the equation is when it decomposes? (Explodes, not when you heat it in solution).
My guess is: 2 NaOCl → 2 NaCl + O2. Or 2 NaOCl → Na2O + Cl2 + 1/2 O2. I doubt that any chlorine oxides are formed...


[Edited on 16-5-2014 by Zyklonb]

The Volatile Chemist - 16-5-2014 at 12:00

Quote: Originally posted by Zyklonb  
Quote: Originally posted by The Volatile Chemist  
No. Could you get small amounts by chilling it so some crystallizes?

No, not with a household bleach solution. I have heard that a 12% solution is possible to get OTC, but I've never found it. (And even with 12% it won't work)
The highest I can find is 8.25%. At 0°C ~29 grams dissolve in 100 mL of water, so you'll need to concentrate the solution first.

[Edited on 16-5-2014 by Zyklonb]

I've heard of very high concentrations used in non-chemistry jobs like cleaning and as a pre-siding (house siding) wash. I don't know how concentrated, though.

Zyklon-A - 16-5-2014 at 12:05

Oh yeah, they certainly do exist. Not really OTC, but still legal as long as they're below 40% - Which is then considered a very strong oxidizing solution, and cannot be sold anywhere but a professional laboratory.