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Ramium
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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]
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blogfast25
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Yes. Use a slight excess of bromide, like 20 %.
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Ramium
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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]
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gdflp
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Yes, I think that it would be a good idea for you to get some more lab experience before you attempt this.
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blogfast25
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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.
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Ramium
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Will be back after i learn a few things
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MrHomeScientist
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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!
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Ramium
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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]
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Molecular Manipulations
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Theoretical bromide oxidation without acid
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]
-The manipulator
We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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Caustic Window
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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.
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j_sum1
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| 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.
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deltaH
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| Quote: Originally posted by j_sum1 | | 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. |
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]
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j_sum1
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I am sure that does work. However the stated aim was to do it without mineral acids.
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deltaH
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Oops... my bad sorry 
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j_sum1
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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.
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Molecular Manipulations
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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.
-The manipulator
We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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blogfast25
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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]
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Molecular Manipulations
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| Quote: Originally posted by j_sum1 |
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. |
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]
-The manipulator
We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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blogfast25
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| 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?
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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. 
[Edited on 22-2-2015 by blogfast25]
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Molecular Manipulations
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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.
-The manipulator
We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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blogfast25
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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.
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Molecular Manipulations
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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.
-The manipulator
We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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Oscilllator
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| Quote: Originally posted by j_sum1 | | 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.
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j_sum1
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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. )
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blogfast25
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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.
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