Sciencemadness Discussion Board - Bromide to bromine???

Gentlemen.

It does all depend a lot on the kind quantities one wants to produce. I prepared about 0.5 ml of wet bromine when I was about 15. I oxidised KBr and collected the wet, undistilled Br2 with a small separating funnel. Enough to verify some its properties and to marvel at this dark liquid.

At such low quantities the risk to oneself is very small, especially if you still take the necessary care.

i just want a little sample for my element collection

Ramium - 12-2-2015 at 11:12

I will buy some condensing gear soon
I assume that is what i need

diddi - 12-2-2015 at 13:03

re glassware. there are a number of ebay sellers who offer organic chem sets and distillation sets. look on youtube for a demonstration of the Br extraction (the ones that use gloves and fume hoods are prolly SM members). there is plenty of discussion about ebay sellers on SM

Ramium - 12-2-2015 at 16:00

Thanks heaps well have a look

Texium - 12-2-2015 at 16:36

The first time that I isolated bromine, I used 31% HCl with solid potassium permanganate and sodium bromide in a ground glass simple distillation apparatus. It worked, but the yield was not great, and was contaminated with water. I used HCl because it was the only strong acid I had at the time. Sulfuric acid is preferred because it absorbs whatever water is left in the mix and allows the pure bromine to distill over.

I was definitely not very experienced at the time that I did this, and it still turned out alright, so I'd say go for it if you're only trying to get a few mL. It's not as dangerous as some make it out to be, and is fairly straightforward even if you don't have much experience. As long as you understand the risks, listen to the advice of others (which it looks like you are), and take reasonable precautions, you should be fine. My own advice to add is to just make sure that you're outside and well away from other people when you do it. Also, if it manages to escape the apparatus, don't try to contain it, just stop what you're doing and get away from there until the fumes clear. If you go into panic mode in that scenario, that's when you're most likely to get yourself hurt.

blogfast25 - 12-2-2015 at 17:42

Quote: Originally posted by zts16

Sulfuric acid is preferred because it absorbs whatever water is left in the mix and allows the pure bromine to distill over.

Care to explain that? Dilute H2SO4 doesn't absorb water.

Conc. H2SO4 does but cannot be used here because of the KMnO4 (possible formation of Mn2O7).

With a BP of 59 C the bromine can be distilled off easily.

Wet bromine can be dried further with conc. H2SO4, then distilled again if need be.

HCl isn't recommended here because KMnO4 can oxidise it to chlorine.

KMnO4 + 8 H+ + 5 e ==== > Mn2+ + K+ + 4 H2O

[Edited on 13-2-2015 by blogfast25]

Molecular Manipulations - 13-2-2015 at 12:46

You could probably get away with using ~85% sulfuric acid, which is still hygroscopic but most likely won't make much or any manganese heptoxide.
But potassium permanganate is not the best option, especially in combination with HCl (aq) as blogfast25 said.
I personally use 98% sulfuric acid and 35% hydrogen peroxide, I've found nothing that works better so far.
I've found potassium chlorate can oxidize bromide with dilute sulfuric or conc. hydrochloric acid, but it's very slow and not really that practical.
No matter how I isolate it I always dry it with 98% sulfuric acid, re-distill and ampule it.

Ramium - 13-2-2015 at 15:26

http://orgchem.colorado.edu/Technique/Procedures/Distillatio...

would this distillation setup work for distilling the bromine???

if so where would I get all the parts (I live in new Zealand)

thanks

[Edited on 13-2-2015 by Ramium]

Metacelsus - 13-2-2015 at 15:32

The only thing I would be worried about is the rubber seal on the thermometer. If you can get a ground glass seal, that would be better. You could try ordering from Laboy glass; I think they ship to New Zealand.

diddi - 13-2-2015 at 16:13

agree. get a GG thermo adapter.

Ramium - 13-2-2015 at 16:20

thanks guys!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!! !!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!! !!

Ramium - 13-2-2015 at 16:24

I found this at laboy glass
http://www.laboyglass.com/chemistry-distilling-set-24-40.htm...

is it suitable???

Texium - 13-2-2015 at 18:58

HCl isn't recommended here because KMnO4 can oxidise it to chlorine.

Yep, I'm aware of that, but it worked fine for me, and I was simply saying that although it's not ideal, if that's all that he had access too, he could do it. Plus, chlorine produced should react further with the bromide to liberate more bromine.

Also, I apologize about the sulfuric acid blunder, I did not really think about that before I posted it.

[Edited on 2-14-2015 by zts16]

blogfast25 - 14-2-2015 at 08:31

Quote: Originally posted by zts16

Plus, chlorine produced should react further with the bromide to liberate more bromine.

Also, I apologize about the sulfuric acid blunder, I did not really think about that before I posted it.

The time I made that bit of Br2, I actually used Cl2 gas. But Cl and Br form inter-halogens like BrCl, another reason not to use HCl or Cl2.

[Edited on 14-2-2015 by blogfast25]

Ramium - 14-2-2015 at 11:56

Quote: Originally posted by zts16

HCl isn't recommended here because KMnO4 can oxidise it to chlorine.

why cant I just use KMNO4 without any other oxidisers???

[Edited on 14-2-2015 by Ramium]

blogfast25 - 14-2-2015 at 12:26

why cant I just use KMNO4 without any other oxidisers???

Nobody said you can't. You do need some acid too.

[Edited on 14-2-2015 by blogfast25]

Ramium - 14-2-2015 at 12:57

why cant I just use KMNO4 without any other oxidisers???

Nobody said you can't. You do need some acid too.

[Edited on 14-2-2015 by blogfast25]

how about sodium bisulphate???

blogfast25 - 14-2-2015 at 13:30

how about sodium bisulphate???

That would work, but you'll need 8 mol of it per mol of KMnO4.

Please edit your post upstairs, the one with all the !!!!!, it's buggering up the view of all posts on this page!

KMnO4 + 8 H+ + 5 e ==== > Mn2+ + K+ + 4 H2O

or:

KMnO4 + 8 NaHSO4 + 5 e ==== > MnSO4 + 1/2 K2SO4 + 4 H2O + 4 Na2SO4 + 5/2 SO42-

And to complete it:

5 NaBr === > 5/2 Br2 + 5 Na+ + 5 e

[Edited on 14-2-2015 by blogfast25]

Ramium - 14-2-2015 at 17:13

how about sodium bisulphate???

sorry about the !!! Post i was very exited it wont let me edit it i posted it on another computer .So could you tell me how much Potassium bromide sodium bisulphate and potassium permanagnate in grams???

[Edited on 15-2-2015 by Ramium]

[Edited on 15-2-2015 by Ramium]

gdflp - 14-2-2015 at 17:27

It depends on how much you want to synthesize. Blogfast gave you all of the information you need, you need a molar ratio of 1:5:8 KMnO₄:KBr:NaHSO₄ respectively.

blogfast25 - 14-2-2015 at 18:15

Quote: Originally posted by gdflp

It depends on how much you want to synthesize. Blogfast gave you all of the information you need, you need a molar ratio of 1:5:8 KMnO₄:KBr:NaHSO₄ respectively.

Ramium isn't well versed in moles. He needs to look that up and acquaint himself with it. Age, young or old, is no excuse here.

Ramium - 14-2-2015 at 19:25

lets say I want to make 20 ml of bromine
how many grams of potassium bromide potassium permanganate and sodium bisulphate would I need for that??????

[Edited on 15-2-2015 by Ramium]

[Edited on 15-2-2015 by Ramium]

Magpie - 14-2-2015 at 22:19

We get nervous when someone who doesn't know how to use moles asks how to make a very hazardous chemical.

Learn some chemistry first. You will also have more fun that way.

Ramium - 14-2-2015 at 23:05

Quote: Originally posted by Magpie

We get nervous when someone who doesn't know how to use moles asks how to make a very hazardous chemical.

Learn some chemistry first. You will also have more fun that way.

I researched moles and work this.

If i started with 1 mole of KMNO4 i would need
KMNO4 =158.03 g
KBr = 5*119.00g = 595g
NaHSO4 = 8*120.06g = 960.48g

That would be too much, so 8 could divide by 10
So
KMnO4 = 15.80g
KBr = 59.5g
NaHSO4 = 96.05g

How much bromine would that make?

[Edited on 15-2-2015 by Ramium]

Texium - 14-2-2015 at 23:52

You need to know how to convert between moles and grams because that's the only way you can interpret chemical equations properly and find out how many grams you need. If you hope to someday design your own procedures rather than following step-by-step instructions, it's the first basic skill that you'll need.

Sure, we could spend a few minutes to do the calculations for you, but we call that "spoonfeeding" here, and that's not what we do, because you don't learn anything, and will just keep asking these simple questions. Instead, how about if you go and find out how to do it yourself? I'll even get you started: try googling "converting between moles and grams"

Once you've got that down, go check out oxidation-reduction and half-reactions, and that will help you with interpreting blogfast's equations.

Ramium - 15-2-2015 at 00:02

Sorry I researched moles and work this.

If i started with 1 mole of KMNO4 i would need
KMNO4 =158.03 g
KBr = 5*119.00g = 595g
NaHSO4 = 8*120.06g = 960.48g

That would be too much, so 8 could divide by 10
So
KMnO4 = 15.80g
KBr = 59.5g
NaHSO4 = 96.05g

Does this sound right so far?

[Edited on 15-2-2015 by Ramium]

[Edited on 15-2-2015 by Ramium]

blogfast25 - 15-2-2015 at 05:41

Yes. Use a slight excess of bromide, like 20 %.

Ramium - 15-2-2015 at 11:09

I dont think i am quite ready for this experiment. maybe i will practise moles on five or so other experiments before i try this .do u think that is a good idea?

[Edited on 15-2-2015 by Ramium]

gdflp - 15-2-2015 at 11:17

Yes, I think that it would be a good idea for you to get some more lab experience before you attempt this.

blogfast25 - 15-2-2015 at 11:59

Quote: Originally posted by gdflp

Yes, I think that it would be a good idea for you to get some more lab experience before you attempt this.

Yup. I second that. There's plenty of interesting stuff to do w/o significant danger.

Ramium - 15-2-2015 at 13:35

Will be back after i learn a few things

MrHomeScientist - 17-2-2015 at 11:22

Might I recommend copper chemistry as a great introduction to the subject. There's tons of things you can do with readily available chemicals. Start with blue copper sulfate, which you can buy as a root killer in the plumbing section of your local hardware store (probably - I'm in the US and it might be different elsewhere). With your HCl, you can then make green copper chloride solution, and adding baking soda to that will yield greenish basic copper carbonate as a solid. If you can't find any of that, buy some malachite online and crush that up - it's mostly copper carbonate. Surely you have access to rocks where you live!

Figuring out the chemical equations involved and calculating the proper amounts of each to add together in a reaction is a great exercise. Try googling things like "copper sulfate hydrochloric acid reaction" to find the formulas. Then look up the molecular weight (g/mol) of each compound, and use that in your stoichiometry (the math of chemistry) to figure out how much of compound B to add to your chosen amount of compound A. So if you have
A + B --> C + D
you can choose how much A you want to start with, and calculate how much B you need to add to produce a certain amount of C or D.

Also for completeness, allow me to shamlessly self-promote my video on bromine extraction from spa-grade NaBr! https://www.youtube.com/watch?v=NKjyM2AkZSY
It's probably a bit more complex than you are interested in, but it has the advantage of completely avoiding distillation if you don't have the apparatus for it. As I mention in the video, be sure to work with all glass equipment! Bromine and plastic don't get along too well.

Your enthusiasm is great; don't lose it! It's a very good sign that you are able to recognize that something might be beyond your capabilities, and are willing to try other things to improve your skills first. Making bromine was one of my favorite experiments, but you definitely need to be prepared before you attempt it. Keep it up!

Ramium - 18-2-2015 at 21:33

Thanks man!! I have made copper hydroxide, copper carbonate, copper phosphate, copper oxide, copper chloride and copper acetate. I plan to make ethyl acetate, caffeine(pure), glacial acetic acid ,Gold hydroxide and some esters . Maybe, if u dont mind, i could show you what i learned after that and you could tell me if u think i need to know more to do bromine or not?

[Edited on 19-2-2015 by Ramium]

Theoretical bromide oxidation without acid

Molecular Manipulations - 21-2-2015 at 21:12

I haven't tried this out yet, but I think it will work.
A lot of people don't have access to strong mineral acids, and while I do, I wanted to come up with a way to isolate bromine without them or electrolysis or chlorine.
Here's the equation: 2 NaBr + H₂O₂ + CuCl₂ --> Br₂ + 2 NaCl + CuO + H₂O -203 kJ/mol.
Since this reaction is exothermic by 203 kJ/mol there's only two reasons I can think of for it not to work.
1: Kinetics make it too slow. Because it's an aqueous reaction with mostly ions this seems unlikely, but without experimental data I can't verify this.
2: In my experience hot copper chloride tends to catalyze the decomposition of hydrogen peroxide, so yields may suffer from this.
If time permits, I'd like to try this out soon.
In the meantime, can anyone else think of a reason this wont work?

[Edited on 22-2-2015 by Molecular Manipulations]

Caustic Window - 22-2-2015 at 04:33

Hello everyone.

This may be interesting in theory, but even if it does work, the bromine is going to be full of CuO. Assuming no reaction between the Br and CuO, which I don't think will happen, perhaps you could acidify the bromine later and wash out the copper salt but that's a lot of effort. I can't see the oxide settling out nicely. And the CuO is going to make the bromine look ugly, and the prettiness of bromine is a main reason a lot of people seek to make it for element collections and such.

j_sum1 - 22-2-2015 at 05:08

Quote: Originally posted by Molecular Manipulations

I haven't tried this out yet, but I think it will work.
A lot of people don't have access to strong mineral acids, and while I do, I wanted to come up with a way to isolate bromine without them or electrolysis or chlorine.
Here's the equation: 2 NaBr + H₂O₂ + CuCl₂ --> Br₂ + 2 NaCl + CuO + H₂O -203 kJ/mol.
Since this reaction is exothermic by 203 kJ/mol there's only two reasons I can think of for it not to work.
1: Kinetics make it too slow. Because it's an aqueous reaction with mostly ions this seems unlikely, but without experimental data I can't verify this.
2: In my experience hot copper chloride tends to catalyze the decomposition of hydrogen peroxide, so yields may suffer from this.
If time permits, I'd like to try this out soon.
In the meantime, can anyone else think of a reason this wont work?

[Edited on 22-2-2015 by Molecular Manipulations]

My initial thought is no, this won't work.
Because you will precipitate solid CuO and (hopefully) liquid Br2, both of these and certainly the first involve a decrease in entropy. Now, I haven't done a calc of Gibb's free energy, but I think it will be unfavourable.

deltaH - 22-2-2015 at 05:31

Quote: Originally posted by Molecular Manipulations

I think this could work with some modification. I think the tetrachlorocoprate ions here can act as a catalyst. I think the stoichiometric equation should be:

2Br[-] + H2O2 + 2HCl =>CuCl2(cat.)=> 2Cl[-] + Br2 + 2H2O

So you would need to add stoichiometric amounts of hydrochloric acid. The copper chloride would act as the catalyst and the peroxide as the oxidant. Note, copper won't change its oxidation state overall and remain in solution as copper chloride.

[Edited on 22-2-2015 by deltaH]

j_sum1 - 22-2-2015 at 05:36

I am sure that does work. However the stated aim was to do it without mineral acids.

deltaH - 22-2-2015 at 05:38

Oops... my bad sorry :mad:

j_sum1 - 22-2-2015 at 05:46

It does raise an interesting question though. Are there syntheses of Br2 from Br- that do not require acidic environments. I am not aware of any. But my experience on such things is quite limited.

Molecular Manipulations - 22-2-2015 at 07:44

Caustic WindowSorry, I forgot to mention I had the intention of distilling it. Thought this goes without saying.
How would copper oxide react with bromine, there's no hydrogen to complete the acid-base reaction. CuO + Br₂ --> CuBr₂ + 1/2 O₂ has got to be unfavorable, but I didn't check either.

blogfast25 - 22-2-2015 at 08:14

Quote: Originally posted by Molecular Manipulations

CuO + Br₂ --> CuBr₂ + 1/2 O₂ has got to be unfavorable, but I didn't check either.

The Standard Heat of Formation of CuO = - 156.06 kJ/mol (NIST), for CuBr2 I get - 141.8 kJ/mol (Wolfram Alpha).

Entropic effects aside, that suggests that:

CuO + Br₂ --> CuBr₂ + 1/2 O₂

... at STP would have an equilibrium constant fairly close to 1. Without saying anything about rates, of course...

[Edited on 22-2-2015 by blogfast25]

Molecular Manipulations - 22-2-2015 at 09:51

Well you perhaps didn't but I did, that 203 kJ/mol was Gibbs Free energy, calculated at 298 K, not just the enthalpy. You do realize that a decrease entropy at or below room temperature doesn't make a very big difference. It's measured in J/mol not kJ.
Thanks for that Blogfast, I didn't expect that to be that close, and since oxygen is much more volatile I'm guessing entropy will drive the equilibrium even more to the right?

[Edited on 22-2-2015 by Molecular Manipulations]

blogfast25 - 22-2-2015 at 10:09

Quote: Originally posted by Molecular Manipulations

[Thanks for that Blogfast, I didn't expect that to be that close, and since oxygen is much more volatile I'm guessing entropy will drive the equilibrium even more to the right?

Not in a practical sense of the word. It's more than likely that to achieve equilibrium in a reasonable amount of time you need to run this at at least 200 C or so, volatility is then equalised of course.

To estimate the equilibrium constant K, one needs to correct the Gibbs Free Energies of Formation of all species involved, for temperature (not hard to do, if the thermochemical data are available) and then calculate ΔG at that temperature T (e.g 200 C).

Since as this is hardly a reaction of 'practical' value, I won't waste my time on that calculation. :cool:

[Edited on 22-2-2015 by blogfast25]

Molecular Manipulations - 22-2-2015 at 10:28

Agreed.
I just tried the experiment anyway, and it was a failure I must admit.
I added about a gram of both sodium bromide and anhydrous coper chloride (didn't actually weigh them).
Then 30% hydrogen peroxide was added drop-by-drop, a very violent reaction occurred releasing a colorless gas, oxygen no doubt. The solution got very dark brown, which I at first mistakenly thought was a copper oxide precipitate. No bromine was observed. Then I added 5 mLs of 6% hydrogen peroxide, the solution slowly fizzed (releasing more oxygen) and became very light green, implying that copper and chloride ions where more or less unchanged.
I'm guessing that the dark color was CuBr₄^- ions, which are very dark in concentrated solution.

blogfast25 - 22-2-2015 at 10:44

MM:

That's hardly a CuO + Br2 testing experiment, though? To stand a chance you need to run this in anhydrous conditions. Water and peroxide seriously skew things.

Still, not really worth doing so, as the outcome is likely what we predicted and it has little practical value.

Molecular Manipulations - 22-2-2015 at 10:56

Yeah, that experiment was just to see if bromide could be oxidized I'm the way I suggested several posts above.
I don't think I have an copper (II) oxide right now, it's easy to make sure, but like you said, hardly worth it.

Oscilllator - 22-2-2015 at 22:15

It does raise an interesting question though. Are there syntheses of Br2 from Br- that do not require acidic environments. I am not aware of any. But my experience on such things is quite limited.

Chlorine gas will displace bromide to form bromine in a classic displacement reaction. The question is then how to obtain Cl2 without the use of an acid - electrolysis of NaCl comes to mind.

j_sum1 - 22-2-2015 at 22:23

Good point.
Amazing how easy it is to overlook the obvious.

(But then if you are doing that, you could always do electrolysis of NaBr!!) (Overlooked by both of us

And everyone who read what I wrote without catching the glitch.

)

blogfast25 - 23-2-2015 at 05:18

Quote: Originally posted by Oscilllator

Chlorine gas will displace bromide to form bromine in a classic displacement reaction.

Cl and Br form the interhalogen BrCl though. I think it's for that reason that method is rarely used. But it's a nice demonstration.

Molecular Manipulations - 23-2-2015 at 07:35

Good point.
Amazing how easy it is to overlook the obvious.

(But then if you are doing that, you could always do electrolysis of NaBr!!) (Overlooked by both of us

And everyone who read what I wrote without catching the glitch.

)

Literally overlooked by you. MrHomescientist linked a video by him on it earlier in this thread. I personally dislike most electrolysis isolations, but that's just me. They take a long time, are usually open to air and thus can get contaminated, plus my MMO gets degraded by bromine, and I hate using carbon rods.

deltaH - 23-2-2015 at 08:51

One might be able to form bromine by melting and then distilling a mixture of ferric nitrate nonahydrate (melting point 47.2 °C) with a bromide salt. The hypothetical reaction for the sodium salt, for example, that I'm thinking of would correspond to:

10NaBr(s) + 4Fe(NO3)3.9H2O(l) + heat => 4FeOOH(s) + 32H2O(l) + 10NaNO3(aq) + 5Br2(g) + 2NO(g)

There could be some contamination with nitrosyl bromide as there is an equilibrium between bromine and nitric oxide.

Again, heed the warning of the EXTREME toxicity and corrosiveness of some of the hypothetical products! Chemical pneumonitis is easily induced from these gases even in very small amounts.

[Edited on 23-2-2015 by deltaH]

blogfast25 - 23-2-2015 at 10:00

10NaBr(s) + 4Fe(NO3)3.9H2O(l) => 4FeOOH(s) + 32H2O(l) + 10NaNO3(aq) + 5Br2(g) + 2NO(g)

I doubt if that would work. The low temperature works against it, I think.

When the ferric nitrate nonahydrate 'melts' it essentially dissolves into its own crystal water. Nitrate ions alone can't oxidise bromide (better check that against a reduction potentials table!).

Edit:

NO3- + 4 H+ + 3 e ===> NO + 2 H2O

Ered = + 0.95 V

Br- ===> 1/2 Br2 + e

Eox = - 1.07 V

So E = Ered + Eox = - 0.12 V < 0

Which suggests this won't work. Oxidation with nitrates at higher temperatures might work by removing volatile Br vapours.

[Edited on 23-2-2015 by blogfast25]

deltaH - 23-2-2015 at 10:12

10NaBr(s) + 4Fe(NO3)3.9H2O(l) => 4FeOOH(s) + 32H2O(l) + 10NaNO3(aq) + 5Br2(g) + 2NO(g)

Surprisingly, I had done a reaction where a chloride salt was heated with molten Fe(NO3)3.9H2O. This produced a very noxious gas with a strong chlorine odour, so I suspected formation of at least some chlorine or nitrosyl chloride (which can decompose to chlorine anyhow in an equilibrium with NO) and or NOx.

If that was the case, then the bromide version should proceed more readily.

But again, the product might be contaminated by the other gases [or not].

Quote:

Edit:

NO3- + 4 H+ + 3 e ===> NO + 2 H2O

Ered = + 0.95 V

Br- ===> 1/2 Br2 + e

Eox = - 1.07 V

So E = Ered + Eox = - 0.12 V < 0

Which suggests this won't work. Oxidation with nitrates at higher temperatures might work by removing volatile Br vapours.

Those reduction potentials are very close blogfast.

I don't know what would happen if one used nitric acid and a bromide salt as is, the reason for stating the iron nitrate version is because of that experimental observation I had (chlorine case).

With aqua regia, there is also a complex equilibrium between nitric acid, water, hydrochloric acid, nitrosyl chloride, NOx and chlorine that is suggestive (chlorine case).

Finally, this route of using ferric nitrate sticks to the spirit of the thread, i.e. avoiding the use of liquid acids. However, nitric acid can be considered to be generated in situ by the easy decomposition of ferric nitrate.

[Edited on 23-2-2015 by deltaH]

blogfast25 - 23-2-2015 at 12:06

E = 0 V does mean K approx. 1. Maybe a bit of sodium bisulphate would help too.

deltaH - 23-2-2015 at 12:17

E = 0 V does mean K approx. 1. Maybe a bit of sodium bisulphate would help too.

Why?

Thinking of the fused sodium bromide + sodium nitrate + sodium bisulfate variant?

Could work, at high temperatures, nitrosyl bromide is probably not favoured thermodynamically, so generated NOx should pass straight through the condenser in part (aside from what is soluble in bromine).

However, generating HBr instead is a competing parallel reaction that could be a big problem with the addition of "dry acid".

Would you be in a position to safely try out the ferric nitrate version I have described by any chance?

[Edited on 23-2-2015 by deltaH]

Molecular Manipulations - 23-2-2015 at 16:30

How about a dry method? Again theoretical:
2 NaBr + Ca(OCl)₂ + Mg --> Br₂ + 2 NaCl + CaO + MgO.
I haven't yet checked out the Gibbs Free Energy for this, everything is forming a very stable ionic compound, except bromine. Could this work?

blogfast25 - 23-2-2015 at 17:39

E = 0 V does mean K approx. 1. Maybe a bit of sodium bisulphate would help too.

Why?

Thinking of the fused sodium bromide + sodium nitrate + sodium bisulfate variant?

No, I was still thinking of your ferric nitrate nonahydrate idea.

blogfast25 - 23-2-2015 at 17:42

Quote: Originally posted by Molecular Manipulations

How about a dry method? Again theoretical:
2 NaBr + Ca(OCl)₂ + Mg --> Br₂ + 2 NaCl + CaO + MgO.
I haven't yet checked out the Gibbs Free Energy for this, everything is forming a very stable ionic compound, except bromine. Could this work?

ΔG almost certainly quite negative. But possibly a can of worms in terms of byproducts, I think, not sure...

Interesting thought though. :cool:

Might want to add a low melting flux to promote contact.

Or why not leave out the Mg altogether? Ca(ClO)2 is a strong oxidiser as such.

NaBr + Ca(ClO)2 === > NaCl + CaO + 1/2 Br2

('I may be a dreamer but I'm not the only one...'

)

[Edited on 24-2-2015 by blogfast25]

Molecular Manipulations - 23-2-2015 at 20:20

Because if you look closely you'll notice that can't be balanced, there's two oxygens and two chlorines in calcium hypochlorite, no matter how you try it, either oxygen or chlorine must evolve. Chlorine is in the +1 oxidation state (I know, oxidation states don't really exist), hence the necessary reducing agent.

[Edited on 24-2-2015 by Molecular Manipulations]

blogfast25 - 23-2-2015 at 20:31

Quote: Originally posted by Molecular Manipulations

Because if you look closely you'll notice that can't be balanced, there's two oxygens and two chlorines in calcium hypochlorite, no matter how you try it, either oxygen or chlorine must evolve. Chlorine is in the +1 oxidation state (I know, oxidation states don't really exist), hence the necessary reducing agent.

[Edited on 24-2-2015 by Molecular Manipulations]

Ooopsie. That Brandy and Coke is taking its toll.

deltaH - 23-2-2015 at 22:22

I don't think it's a good idea to add magnesium to an oxidant if you desire to make bromine.

blogfast25 - 24-2-2015 at 06:21

I don't think it's a good idea to add magnesium to an oxidant if you desire to make bromine.

Flash powder alert!

Molecular Manipulations - 24-2-2015 at 07:17

I don't think it's a good idea to add magnesium to an oxidant if you desire to make bromine.

I was thinking the same because of magnesium oxide contamination. Is this why, or some other reason? Glass wool could be use at the neck of the flask to stop some of it, but it would probably need to be distilled twice.
Of course it could just not work, or go way too fast.
Or the magnesium oxidation could be fast, but it would take a while for chlorine to oxidize bromide, I think the sodium bromide would need to be fused, and even still it will take a while.

[Edited on 24-2-2015 by Molecular Manipulations]

blogfast25 - 24-2-2015 at 07:49

Quote: Originally posted by Molecular Manipulations

How about a dry method? Again theoretical:
2 NaBr + Ca(OCl)₂ + Mg --> Br₂ + 2 NaCl + CaO + MgO.

How about replacing that 'nasty' Mg with something a bit tamer like Zn?

Mg with chlorates is explosive, not sure about with molten hypochlorites.

I don't get your point about MgO: magnesium powder is pretty stable and forgiving.

[Edited on 24-2-2015 by blogfast25]

Metacelsus - 24-2-2015 at 07:50

Instead of adding magnesium, what about adding an acid?

blogfast25 - 24-2-2015 at 07:52

Quote: Originally posted by Cheddite Cheese

Instead of adding magnesium, what about adding an acid?

Equation?

Molecular Manipulations - 24-2-2015 at 08:19

It's just a matter of avoiding contamination, magnesia is very stable, I just don't want it in my bromine.
Zinc seems like a good idea, it would melt easily and not react to fast. Also whether zinc or magnesium it won't be powdered, probably chips, so flash powder isn't going to happen.
As for Cheese's acid, first, this entire point is not to use acid, if acid was to be considered I'd just use the regular acid + bromide + oxidizer.
Besides, calcium oxide will react with acids...

[EDIT] Shouldn't this be is R and A Acquisition I've already waisted so many posts here that I'm now most active in 'beginnings'... [shudder]

[Edited on 24-2-2015 by Molecular Manipulations]

blogfast25 - 24-2-2015 at 08:23

Quote: Originally posted by Molecular Manipulations

It's just a matter of avoiding contamination, magnesia is very stable, I just don't want it in my bromine.

But how would it get in there? You distil off the bromine, so the worst that can happen is mechanical entrainment and that can happen with the other reaction products too.

Sodium bisulphate is a salt.

[Edited on 24-2-2015 by blogfast25]

deltaH - 24-2-2015 at 08:33

What's wrong with ferric nitrate and bromide salt route?

Molecular Manipulations - 24-2-2015 at 08:40

Yeah, I was referring to mechanical 'entrainment'. I had the image of flash powder in the back of my mind when I wrote that, it makes a lot of fine magnesium oxide dust (among other things). I guess if the reaction is slow this might not be a problem.
You're quite right, sodium bisulfate is indeed a salt, and an acidic one at that. Is there a point to that, or are just saying words now?

Tin tetrachloride is a covalent liquid...

blogfast25 - 24-2-2015 at 08:44

Quote: Originally posted by Molecular Manipulations

Is there a point to that, or are just saying words now?

The emoticon should have been a bit of a giveaway (it was a joke, i.o.w.)

Post with 'you can't use XYZ' never have much sway with me. NaHSO4 is a case in point: so OTC no one can reasonably object to recommending it in an amateur synthesis.

blogfast25 - 24-2-2015 at 08:45

What's wrong with ferric nitrate and bromide salt route?

We've already discussed that, no?

10NaBr(s) + 4Fe(NO3)3.9H2O(l) => 4FeOOH(s) + 32H2O(l) + 10NaNO3(aq) + 5Br2(g) + 2NO(g)

Quote:

Without some source of H3O+ that won't work too well, if you rely on the 'classic' oxidising power of nitrates.

Also, why ferric nitrate hydrate? Why not Al nitrate hydrate?

[Edited on 24-2-2015 by blogfast25]

deltaH - 24-2-2015 at 09:10

Quote:

Without some source of H3O+ that won't work too well, if you rely on the 'classic' oxidising power of nitrates.

Also, why ferric nitrate hydrate? Why not Al nitrate hydrate?

Ferric nitrate solutions are highly acidic. Also ferric nitrate is not very stable, tending to decompose to ferric oxide/hydroxide and so generating nitric acid, even on gentle heating. Plus, you're overlooking that I've had encouraging results with the chloride version.

As for the aluminium nitrate, in principle that could also work, but ferric nitrate is easier to obtain IMHO and not as aggressive as aluminium nitrate.

[Edited on 24-2-2015 by deltaH]

blogfast25 - 24-2-2015 at 09:20