Catalytic chlorination

The NCl3 generated will collect at the bottom of the reaction vessel. As this is a touch sensitive high explosive , I can see a slight problem...

Marvin...

"Read. More preferably "Chemistry of the elements" by Greenwood and Earnshaw than March."

Why, you... "Chemistry of the elements" is becoming my Bible...

"Theoretic,
"Can't oxidizers react to form more potent oxidizers?"
The reaction does not go forward, you have to force it in the same way water doesnt run uphill."

It's true that at the pH of the reaction the products are weaker oxidizers than the reactants, but if you change pH a lot, spectacular things happen. The ferrate ion is made using hypochlorite, whose oxidizing potential is WAY below ferrate's in acid solution, but in alkali ClO- is used to prepare FeO4--.

""A very apparent example is hypochlorites, which are extremely potent oxidizers - stronger than chlorine"

Wrong. "

You're right, I'm wrong.

"confusing HClO with ClO-"

Not really, since in acid solution the latter is transformed into the former (which has higher potential though).

""Even chlorates are stronger oxidizers than chlorine"

Wrong."
Why, you... Chlorine's potential is 1.35 V, chlorate's potential is 1.47 V.

""Another thing is that chlorine in a 1+ oxidation state (yes, that's chlorine's oxidation state in NCl3!) "

Wrong."

I don't think! How come then NCl3 hydrolyses into NH3 + 3HClO and not into HNO2 + 3HCl? During hydrolysis of a compound the element's formal oxidation number is unchanged. And nitrogen's electronegativity is ever so slightly bigger than chlorine's...

"Additionally you seem to be confusing... rate of reaction with oxidising potential"
Maybe, because untill now I somehow assumed that ClO- is a more potent oxidizer than ClO3- (which it isn't) ... or in theory anyway, since HClO (produced by hydrolysis of ClO- in aqueous solutions) HAS got a higher oxidizing potential (1.50 V) than ClO3-.

Theoretic,

I find "chemistry of the elements" too similar to other uni inorganic books for me to consider it a bible. It is a lot better for werner complex chemistry, and discussion of redox.

"It's true that at the pH of the reaction the products are weaker oxidizers than the reactants, but if you change pH a lot, spectacular things happen."
Yes, you send the reaction backwards to chlorine. Unless you isolate, and then acidfy in which case you force creation of a stronger oxidising agent. The important thing here though is the reaction is running downhill unless you force it the other way by making that the new downhill.

"The ferrate ion is made using hypochlorite, whose oxidizing potential is WAY below ferrate's in acid solution, but in alkali ClO- is used to prepare FeO4--. "
I have a better example. If you make MnO2 alkaline enough, it will absorb oxygen from the air and turn into manganate(VI). When you make this acid, it disproportionates into Manganate(VII) and MnO2. Now because you are aparently 'making' it from oxygen, would it be fair to say that oxygen therefore is the stronger oxidising agent, or that permanganate is and you are forcing the reaction externally?

"...in acid solution the latter is transformed into the former (which has higher potential though). "
In alkaline solution HClO essentially doesnt exist, and in acid solution ClO- essentially doesnt exist. This is a big difference, and the difference in oxidising potential between acid and alkali solutions is huge.

"Chlorine's potential is 1.35 V, chlorate's potential is 1.47 V. "
This is correct for 1M acid solution, but its not true for neutral solutions and its not true for basic solutions. Since its not generaly true the statement is not true. More importantly, where you were going with your argument was essentially 'Chlorine makes chlorates, and they are stronger oxidising agents than chlorine', except under all conditions chlorine does make chlorates they are a less strong oxidising agent. This is why the statement is completely wrong in your context.

"During hydrolysis of a compound the element's formal oxidation number is unchanged."
Perhaps youd like to explain that to polysulphide ions, the halogen gasses, tha alkali metals and anything else capable of finding a back door in high school chemistry.

"How come then NCl3 hydrolyses into NH3 + 3HClO and not into HNO2 + 3HCl? "
You are placing a great deal of faith in a model for compounds at the other extreme end of bonding to that which NCl3 uses. Formal oxidation state is of limited use at the very best of times in my experience. To me it seems more of a preuniversity teaching tool and method of naming compounds than a means for working out how something unknown should react. I have to admit at this point, I was under the impression that chloramine would hydrolyse with alkali to hydroxylamine, and this does not seem to happen.

"And nitrogen's electronegativity is ever so slightly bigger than chlorine's."
Depending on who you read, and most people use Pauling's scale, chlorine is higher than nitrogen. Since you are under the impression this is not the case, I'm guessing you are using Alled and Roschow.

If you ask the question, if a nitrogen atom next to a chlorine atom had to ionise, which way would it go?
Youd find the most stable form would be N+ Cl- becuase the ionisation energies are similar, but chlorine has a high electron affinity, whereas nitrogens electron affinity is virtually nothing.

Lastly the 'accepted' state of affairs for naming the compound is 'nitrogen (III) chloride', with a formal charge of +3 on the nitrogen and -1 on each chlorine. A trip to webelements will confirm this, though its not clear from which of their references they got it from.

My conclusion about the oxidising strength of NCl3 is that it has to be less strong than chlorine becuase its formed in high yeild from chlorine. If we swing the reaction towards high pH, is the oxidation potential of the NCl3 really going up, or is it just the potential of the hypochlorite ion is dropping so rapidly this becomes the prefered form in basic solution? NCl3 does not seem to be a strong oxidising agent, and strong oxidising agents are not what we need to speed up reactions anyway.

If this thread was your intended 'chlorination without substitution thread' I have more bad news btw

It would be interesting to know what dilute NCl3 would do to double bonds, and if this could be useful. This at anyrate would really not be substitution.


unionised,

"Chlorate can oxidise chloride to chlorine...
And the reverse can also happen, you are making the same mistake. You are changing the conditions, telling yourself that is only a very minor alteration and that the resulting change of oxidising potential is therefore 'free'. Catalysis does not work like this, it cannot work like this.

What I think you are thinking is that the AlCl3 is making the NCl3 a 'stronger oxidising agent'. This is possible in this case, but if true it wouldnt be catalysis, the AlCl3 would be used up, as adding just NCl3 would not be able to regenerate the adduct from the waste products of the reaction. What is also possible, and how catalysts work is they lower the activiation energy instead, usually by stabilising the intermediate 'high energy' state that limits the rate of the reaction.

The point that lewis acids catalyse halogenation of arenes is a good one, but what is required for this isnt a better oxidising agent, its electrophilic bromine. If it was about oxidising strength it wouldnt go any faster.

"Which way is uphill? towards the better oxidant or towards the separate halide and halogen?"
Uphill is to higher energy, in this case towards the production of a small population of electrophilic bromine.

I agree that NCl3 is not suitable for these sorts of home experiments.

Marvin:
""During hydrolysis of a compound the element's formal oxidation number is unchanged."
Perhaps youd like to explain that to polysulphide ions, the halogen gasses, tha alkali metals and anything else capable of finding a back door in high school chemistry."

Alkali metals? Halogen gases? I SAID A COMPOUND! And more, you seem to be confusing redox reactions (alkali metals or fluorine + water) and disproportionation reactions (halogen gases + water) with hydrolysis. DO read. Polysulfides? Sorry for not being specific, I actually meant hydrolysis of bonds between different elements.

"In alkaline solution HClO essentially doesnt exist"
Well, solutions of hypochlorites do hydrolyse, don't they? And they create an alkaline pH, right? And what is the product of hydrolysis if not HClO? The fact that hypochlorite solutions are such strong oxidants is because of the self-created HClO they contain, ClO- ions are themselves weak oxidants. Another thing if you mean ALKALINE alkaline solutions (the standart pH of 14 created by NaOH). Under such conditions yes, hydrolysis is hindered and all HClO that was in solution reacts. The resulting solution is a much weaker oxidant than simple hypochlorite.

"It would be interesting to know what dilute NCl3 would do to double bonds, and if this could be useful. This at anyrate would really not be substitution."

I say: add across them! Then either the resulting intermediate hydrolyses to dichloroamine and chlorohydrin, or the reation takes on an altogether different path. Sulfur dichloride (SCl2) is known as a reagent in preparation of mustard gas (S(CH2CH2Cl)2) by its addition to ethylene, so I suspect NCl3 would behave the same way - by forming tri-2-chloroethylamine (N(CH2CH2Cl)3). A third pathway might be addition of NCl3's chlorine to the double bond and evolution of free nitrogen.
I've been talking about C=C bonds. I leave C=N, C=O, C=S and others to someone else.
And I don't see why NCl3 would pose a significant ristk if it would be generated a little at a time and (I hope) reacted as soon as generated? It can react slower than we want, thus accumulating, though.

What I mean is bubbling chlorine through a solution of the chemical to be chlorinated, together with some ammonium salt. The idea of extracting NCl3 in a solvent then using it so as to carry out the chlorination with much less bother is very good.
But why only tertiary alkanes?
For NCl3 preparation I recommend Ca(ClO)2 + ammonium carbonate... Oh sorry, my science teacher is looking over my shoulder. How rude...
Anyway, then add a sourse of CO2 to precipitate out alkaline Ca(OH)2 as neutral carbonate. You can use excess AC (ammonium carbonate), but that makes unpleasant ammonia solution. Ca(HCO3)2 is ideal, NaHCO3 less so, although acceptable, if enough is used, Na2CO3 is unacceptable since the aim is to neutralize the solution. This greatly improves the yield, (three times, I assume, by loking at the equation)esspecially looking at the fact that the Ca(OH)2 produced will sink to the bottom of the NCl3 layer, performing evil things...
And I don't think you need to dry it, since no water would be extracted into the inert organic solvent.
Then, if you want an all-out chlorination, add the organic compound to the NCl3 solution slowly, dropwise. If you are afraid for the integrity of your organic molecule or don't want an all-out chlorination, then add NCl3 solution dropwise, slowly.
Have fun!

OK, I DIDN'T prepare any NCl3 (or even tried it). But...
Quote:
"Thankyou for giving us the benefit of your inexperience."
Yes, I don't experiment, I just talk. The "have fun" bit was inconsiderate, and I know a (maybe essential) minimum about making NCl3, the reaction is unpredictable and dangerous (both the making and reaction of NCl3), but someone on E&W did succesfully prepare NCl3 by AC + bleaching powder, that's why I assumed it would work again... ...the person obviously didn't read any literature about making NCl3, and it was a bad chice, but it worked.
And OK, it was an out-of-the-deep-blue-sea assumption of mine that NCl3 reacts with organics in a desirable way.
Do compounds with a N-Cl bond generally work well as chlorinators? Trichloroisocyanuric acid can be made by Cyanuric acid + chlorine, then, if the reaction goes the way we want, CUA is regenerated... Or are N-Cl compounds generally useless and should never be let to touch anything organic but tertiary alkanes?

Thanks for pointing out the fact that reactions are generally faster with hot reactants; I'm sure that, like me,all the other folks on this board, didn't know that.
Now look at what I said. The AlCl3 catalyses the acylation reaction. That makes it a catalyst. The fact that it subsequently reacts with the product does not detract from this. It does this by reacting with one of the reactants to form an intermediate that reacts more quickly with the other reactant than the original material would. In this case it forms a complex with the acid chloride that acylates the arene faster than than acid chloride would do.
If you accept that NCl3 is formed from NH3 and Cl2 and that it reacts with the alkane quicker than the chlorine does (which is doubtful) and that it is converted back to NH3 in this process then it is doing the same job as a castalyst and I don't see what is wrong with calling it one.
The problem is that you need to add a lewis acid to get the NCl3 to react and that will be nobbled by the product. Its just possible that some zeolite or molecular sieve has an acidic site that would catalyse the NCl3 + alkane reaction but be unaffected by the ammonia (I doubt it, but that's not the point) then the reaction would work with the NCl3 acting as a catalyst. It would be diffiult to see what else you could call this system (NCl3 and "magic zeolite" apart from a catalyst or a catalyst and an activator.

"Unfortunately, it doesn't dissolve in water very well."
True, and it's a godsend for those mad enough to make it. Surely it would be nice having any limbs left after distilling a NCl3 solution?

What I said applies for chlorinations in an organic solvent (which tend to dissolve the organic compound to be nitrated, which is very useful) - as you say:
"It might be possible to chlorinate some tertiary alkanes by forming NCl3 by chlorination of ammonia, isolating it in solution in some innert solvent (perhaps CH2Cl2) drying it , adding a lewis acid catalyst and using this to chlorinate the alkane."


[Edited on 30-11-2003 by Theoretic]