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CHRIS25
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Generate oxygen for oxidation by pump
As I was thinking of alternative and more efficient ways to oxidize ferrous chloride without the inefficient and water laden use of 6% peroxide, but
wishing to stay with the air pump, I was hit with a 'brainstorm' moment.....Mmm...why not enrich/concentrate the air surrounding the pump? Now the
more sensible of you here may find an obvious fault with this: However is there any legitimate reason why a method could not be found?
1. Place the air pump inside a large/very large upside down container; Inside this container put either:
a) Electrolysis of water reaction, re-direct hydrogen outside/away
b) controlled Yeast and Peroxide reaction
c) huge amount of potted plants
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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Oscilllator
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Uh, you might want to think about option b a bit more carefully...
Option C will never work, no matter how many plants you use. The plants will only remove the CO2 from the air (~400ppm) and not the nitrogen, which of
course makes up 70% of the atmosphere.
Apart from that I would suggest the generation of chlorine gas with HCl and trichloroisocyanuric acid (get it in the pool section). Both of those
chemicals are quite cheap, and I imagine the oxidation will proceed smoothly.
Of course there are the dangers of working with chlorine gas, but it can be done easily and safely with the appropriate setup.
Of course the lazy chemist would just hook up the air pump and come back in a couple of days. No muss, no fuss.
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blogfast25
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Quote: Originally posted by Oscilllator  | [...] and I imagine the oxidation will proceed smoothly.
Of course there are the dangers of working with chlorine gas, but it can be done easily and safely with the appropriate setup.
Of course the lazy chemist would just hook up the air pump and come back in a couple of days. No muss, no fuss. |
I doubt that using chlorine gas will be much faster than air oxygen, apart from the obvious drawbacks of chlorine. Both oxygen and chlorine are poorly
soluble in water and sluggish oxidisers at RT.
As regards the latter, he's already tried that for weeks on end, it seems. Doesn't seem to be able to complete the reaction.
@Chris:
Using pure oxygen as opposed to air oxygen should be a little faster but I doubt that the difference is worth the fuss of constantly making the
oxygen. Compared to the methods already proposed, it's making things seriously complicated...
[Edited on 10-6-2014 by blogfast25]
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hyfalcon
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How do oxygen concentrators work? Could this tech be adapted to home use?
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DraconicAcid
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Why not use bleach? That should be rapid (especially since it will make the solution basic). You just have to be careful reacidifying it.
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.
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blogfast25
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| Quote: Originally posted by DraconicAcid | | Why not use bleach? That should be rapid (especially since it will make the solution basic). You just have to be careful reacidifying it.
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Bleach is very weak: about 4 - 5 % NaClO, so you need a lot of volume. That's why he objects to 3 % peroxide.
Bleach also introduces sodium which would then have to be separated from the Fe. And hypochlorite appears strong enough to oxidise Fe all the way to
ferrate (VI) (but that's unstable and then oxidises water and back to Fe (III))
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aga
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as Time is clearly not a factor, concentrate the peroxdide by slow heating over ages and ages and ages ...
That kinda works.
What exactly is the objection to more water in the ferric chloride anyways ?
Can't it be driven oof, at least partially by heating ?
[Edited on 10-6-2014 by aga]
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DraconicAcid
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| Quote: Originally posted by blogfast25 | | Bleach also introduces sodium which would then have to be separated from the Fe. And hypochlorite appears strong enough to oxidise Fe all the way to
ferrate (VI) (but that's unstable and then oxidises water and back to Fe (III)) |
But if the hypochlorite is not so concentrated, you get a nice precipitate of iron(III) hydroxide, which can be separated from the supernatant (which
is where the sodium ions are) and redissolved with hydrochloric acid.
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.
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CHRIS25
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Firstly, I want nothing to do with bleach or chlorine, both because of toxicity and because introducing anything extra within a confined compartment
where the pump is would be risky I think; (by the way, the plant idea was a joke oscillator).
After reading the following: Amount of Oxygen in air is about 0.28g/L Average human breathes in about 1.8 to 2.4g?L (this gives me something to
relate to since weights and densities of gases a bit too abstract at the moment); I did a calculation that if I produced 1 mole of Oxygen in a 5 litre
container that would be 6.4g/L ; that would be about 4 minutes of delivery. But to get just one mole would require 2 moles peroxide. This in turn
would require from my 6%: (100%/6% x 1 mole/34 g/mol peroxide ) x 2 moles peroxide = 978 mLs peroxide just for 4 minutes. If I have 0.2 moles of
Ferrous to oxidize to ferric then I only need 978/5 = 195 mLs 6% peroxide to mix with ???? grams of dried yeast.
Hey I am just doing the best I can to make this a bit more controlled, stabbing in the dark probably, stupid idea maybe, but is this feasible? or just
nonsense?
@Aga Hi, more water means Increased chances of iron hydrolyzing as a result of HCl being driven off, as a result you then have to add more HCl to
counter balance what one is losing and this becomes too finniglety and subjective I believe in my past experience.
@Draconic I am wanting to avoid all these extra complications of introducing foreign elements into a system where absolute accuracy is important and
extra elements means more risk of messing up.
[Edited on 10-6-2014 by CHRIS25]
[Edited on 10-6-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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aga
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as i understand it, the ferous solution is required, and some of it needs to be ferric.
so, separate the ferrous, then that part is done.
The rest, add 3% H2O2 then boil off the excess water.
Ferrous, known conc, ferric, known conc.
Now add dancing magnets !
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CHRIS25
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Hi, yes....you could, you need to split exactly in half so that if your, (for example) 100 mLs contains 0.5 moles, then 2 x 50 mLs will have 0.25
moles each. Oxidize one half of the solution and reduce back down to 50 mLs. Then for every 2 mLs of the ferric you use 1 mL of the ferrous for
example if you use 25 mLs of ferric this will be 0.125 moles and then use 12.5 mLs of ferrous will be 0.0625 moles
2FeCl3 + FeCl2 + 8NH3 + 4H2O = Fe3O4(s) + 8NH4Cl(aq)
[Edited on 10-6-2014 by CHRIS25]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Chris:
0.2 mol ferrous iron only needs 0.1 mol H2O2 to be oxidised to ferric iron (theoretically).
That’s 0.1 mol x 34 g/mol = 3.4 g pure H2O2.
6 % contains 6 g H2O2 / 100 g solution, so you need (3.4 / 6) x 100 = 57 ml of 6 % H2O2.
Ferrous solution and H2O2 both need to be ice cold, H2O2 solution added slowly, ml by ml, while stirring vigorously.
[Edited on 11-6-2014 by blogfast25]
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CHRIS25
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| Quote: Originally posted by blogfast25 | Chris:
0.2 mol ferrous iron only needs 0.1 mol H2O2 to be oxidised to ferric iron (theoretically).
That’s 0.1 mol x 34 g/mol = 3.4 g pure H2O2.
6 % contains 6 g H2O2 / 100 g solution, so you need (3.4 / 6) x 100 = 57 ml of 6 % H2O2.
Ferrous solution and H2O2 both need to be ice cold, H2O2 solution added slowly, ml by ml, while stirring vigorously.
[Edited on 11-6-2014 by blogfast25] |
Yes this is fine Gert in small quantities, but 0.2 was an example. In 0.5 moles this is 141 mLs with 132 being water. I don't think it is good idea
to add so much then have to simmer away so much water increasing the risk of hydrolysis, so to avoid this you then add more HCl - it's like a viscous
circle continually adding solution just to get rid of it safely
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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Try the alkaline route. Get some cheapo garden grade FeSO4.7H2O, dissolve it in water, precipitate with NH3 as Fe(OH)2, filter off and wash and air
oxidise that to Fe(OH)3 with some mild heating. The dissolve the Fe(OH)3 in excess HCl and concentrate to desired concentration.
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aga
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| Quote: Originally posted by blogfast25 | | Ferrous solution and H2O2 both need to be ice cold, H2O2 solution added slowly, ml by ml, while stirring vigorously. |
Ooops.
So just dunking the whole lot in at ST* wasn't a good idea then ?
* Shed Temperature, currently 30+ C
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CHRIS25
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| Quote: Originally posted by blogfast25 | | Try the alkaline route. Get some cheapo garden grade FeSO4.7H2O, dissolve it in water, precipitate with NH3 as Fe(OH)2, filter off and wash and air
oxidise that to Fe(OH)3 with some mild heating. The dissolve the Fe(OH)3 in excess HCl and concentrate to desired concentration.
|
Yep, I think I will try that, I have made plenty of Iron 2 Sulphate and co-incidently have a batch on now. Will give that a shot.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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| Quote: Originally posted by aga | So just dunking the whole lot in at ST* wasn't a good idea then ?
* Shed Temperature, currently 30+ C |
It depends on a lot of factors but it could potentially be dangerous. Oxidation reactions are generally strongly exothermic and oxidation by peroxide
is by no means an exception. Adding a lot of peroxide at once heats the solution strongly and can cause serious overboiling and thus splattering and
worse. For that reason it is recommended to start with cold reagents, add peroxide slowly and monitor temperature as you go. If temperature rises too
much, cool intermittently on an ice bath, then continue adding peroxide.
Too high temperature also favours the side reaction:
H2O2 ===> O2 + 2 H+ + 2e
... and that's just lost peroxide: it doesn't help doing what you're doing, which is to oxidise the other reagent. The side reaction attempts to do
the opposite!
[Edited on 11-6-2014 by blogfast25]
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aga
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Doh !
Im using OTC 3w% peroxide, and the heating wasn't significant, most likely due to the 97w% water accompanying it.
Seems like Ferric Chloride to me, and FeCl3 is an old, old friend.
It's already hugged me and left stains on my jeans.
Amazing how fast it boils down with and old PC fan on top of the beaker.
I worked it out at 1.6ml water leaving every minute !
[Edited on 11-6-2014 by aga]
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blogfast25
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| Quote: Originally posted by aga | Seems like Ferric Chloride to me, and FeCl3 is an old, old friend.
It's already hugged me and left stains on my jeans.
[Edited on 11-6-2014 by aga] |
Chris' also 'seemed' like ferric chloride, yet was mostly ferrous chloride. Check with K3Fe(CN)6 for ferrous iron. Concentrated Fe3+ masks things
because it's dark.
Stains? It'll do more than that: 20 % HCl eats through cotton fabric easily. Lab coat, anyone?
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aga
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The dark molasses colour, the yellow stains it leaves on everything, includng glass, the way it smells ...
It is entirely possible, even though i have bought and used about half a metric tonne of ferric chloride, that each time i was sold a dud that still
etched my pcbs.
Since the 'smells chlorine-y' versus actual Cl2 gas experience, i am completely open to suggestions regarding my assumptions.
Edit: boiling finished. SO fast with a fan !
only test i can do : etch a bit of Copper Clad PCB with it ...
colour change of the copper : normal - goes black when exposed to air after dipping in the solution ...
4 mins 36 seconds and all the copper has gone.
Looks, smells, stains and etched PCBs like ferric chloride to me.
i think i'll repeat that with the ferrous solution to see what happens.
[Edited on 11-6-2014 by aga]
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aga
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Amazing.
5 mins of the copper-clad board in the ferrous solution and nothing happened at all.
If anything it just seems cleaner. Shinier.
No going rose-coloured, no etching, no blackening when exposed to air.
I'll leave it overnight and see if that changes.
Certainily it is a different compound to the other one, which i strongly believe is pukka ferric chloride.
| Quote: | | 20 % HCl eats through cotton fabric easily. Lab coat, anyone? |
the stoichimetry of my particular reactions left a small excess of HCl to prevent CHRIS25's dreaded hydrolisation.
Jeans are cheaper than lab coats.
They also show the stains less !
[Edited on 11-6-2014 by aga]
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blogfast25
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| Quote: Originally posted by CHRIS25 | | Yep, I think I will try that, I have made plenty of Iron 2 Sulphate and co-incidently have a batch on now. Will give that a shot.
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Once filtered and washed, heat mildly on a low heat hot plate. Keep stirring and always keep it moist. Too dry will speed up the oxidation but may
render the ferric oxide less soluble in the HCl later on.
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WGTR
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Perhaps oxidation of a neutral ferrous salt with ozone?
http://www.eolss.net/sample-chapters/c07/e6-192-06-00.pdf
The precipitate could be filtered and washed, then acidified with HCl.
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blogfast25
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Sooner or later someone was going to come up with ozone, of course (but did it have to be you? )
The difficulties of preparing a steady stream of O<sub>3</sub> don't warrant this relatively simple problem.
Moreover ozone is also poorly soluble in water and may well be accordingly slow in these conditions, even though it's a powerful oxidiser...
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Fantasma4500
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electrolysis using iron iron, weak HCl solution anyone?
about the computer fan.. its a really underrated ghetto setup
infact the one i have has these 4 attachment points, the thing fits perfectly on my 1000 mL beaker, also upside down
the most effective is if you havent figured it out yet to have the fan drawing the water up from the beaker, or well the water vapour, instead of
blowing air into it, this would otherwise cool it down alot -- which could also be useful
i got my fan running almost 24/7 trying to get 400 mL CuCl2 into crystals at room temperature, so far getting to 200 mL
very useful for crystallization
the problem with using iron for electrolysis if you care for HHO is that the oxygen usually goes to oxidizing the iron creating Fe3O4, but i can
imagine it gets different if HCl is present
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