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blogfast25
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Quote: Originally posted by Antiswat  | electrolysis using iron iron, weak HCl solution anyone?
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 |
If by 'HHO' you mean the well known 'fuel saving' scam, any reference to that bit of pseudo-science on this forum belongs in 'Detritus'.
http://www.aardvark.co.nz/hho_scam.shtml
[Edited on 12-6-2014 by blogfast25]
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aga
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Fuel saving is simple, and doesn't need OHH HOH HHO or HoHoHo.
Simply travel less, stay in the lab and get more experimenting done.
[Edited on 12-6-2014 by aga]
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Zyklon-A
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| Quote: Originally posted by aga | as Time is clearly not a factor, concentrate the peroxide by slow heating over ages and ages and ages ...
That kinda works.
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It works quite well actually, with slow heating one can get nearly 70% H2O2 with about 60% yields.
With STP evaporation, one can get up to 85% H2O2 with about 70% yields IIRC.
But this is very time consuming of course.
Months ago, alexleyenda said that he could concentrate H2O2 by boiling, he got pretty great yields too.
I tested this, and he was right. I got 30% H2O2 by boiling a 3% solution with 75% yields. Now I doubt that you could get much
higher conc. without sacrificing yields, but this application doesn't require much higher yields.
[Edited on 12-6-2014 by Zyklonb]
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aga
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The 3% works.
Tried it.
Blogfast25's warning regarding oxidation reaction is true, however it seems the large volume of water sucks the heat from the 3% H2O2 doing the work.
Boiling out all the excess water is quick and easy with the fan as well.
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blogfast25
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| Quote: Originally posted by aga | Blogfast25's warning regarding oxidation reaction is true, however it seems the large volume of water sucks the heat from the 3% H2O2 doing the work.
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Yes, both the water in the ferrous chloride solution and in the peroxide acts as a heat sink. Simply put and assuming no heat losses or cooling, ΔH =
m Cp ΔT, with ΔH = the heat of reaction, m = the mass of water, Cp = the heat capacity of water and ΔT = the increase in temperature. Obviously for
small m, ΔT becomes larger: in concentrated solutions the temperature profile can become a serious concern.
Have you tested your solution for Fe<sup>2+</sup>?
[Edited on 14-6-2014 by blogfast25]
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CHRIS25
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Air pump tests confirmed:
0.5 mol Ferrous chloride soln fully prepared, Solution Level = 187 mLs
Delivered air with a 50 Hz 3.2W fish tank pump:
After 17 hours and again after 23 hours:
15 mL sample taken and diluted with 200 mLs water;
0.15M thiosulphate prepared (7.44 g)
1.2 KI
50 mL sample extracted from the 200 mLs to be titrated;
Final mole Calculations for sample 1 based on 187 mLs
Final mole calculations for sample 2 based on 172 mLs
@ 16 hours molarity is 0.24, 0.04489 mol ferrous oxidised
@ 23 hours molarity is 0.32, 0.055 mol ferrous oxidised
this is 0.00258 moles per hour after first 16 hours and then 0.00239 moles per hour after a further 7 hours totaling 23 hours. It would then take 200
hours of pumping air to fully oxidise the 0.5 moles of ferrous. (based upon average 0.0025 moles per hour)
.....Peroxide anyone? (though just for confirmation I will continue for another 10 hours or so and then see if the oxidation rate is really this
consistent)?
‘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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Nice one CHRIS25.
So now we know.
Peroxide! Yeah baby, even if it is weak.
@blogfast25 - no i have not tested it for Fe2+ as i have not looked up how to do that.
It is so much like ferric chloride that unless chemistry has the equivalent of cheap chinese copies, i would bet that it is FeCl3.
I did test it in making a ferrofluid, and that kinda worked out ok.
[Edited on 14-6-2014 by aga]
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CHRIS25
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Ahh, but I have just realised something, I think the rate of oxidation has slowed down because I may not have enough HCl in solution. The original
was 0.5 mol Fe and 1.27 mol HCl. I would need another 0.5 mol HCl to complete the oxidation process. If 0.055 is oxidised then that would have used
up 0.11 HCl of the 0.27 remaining in solution. Hope all this theorizing computes into something sensible!
Also if I doubled the output of the air pump, that 200 hours becomes 4 days, not long at all really especially if one is not in a hurry.
[Edited on 14-6-2014 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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The idea of you hooking up multiple pumps made this spring to mind :-
http://simpsons.wikia.com/wiki/French_Chef?file=French_Chef3...
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blogfast25
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Good work, Chris.
Trust me, the lack of HCl would not have played much part: the oxidation would still proceed but producing Fe(OH)3 rather than FeCl3. If anything,
less HCl probably speeds things up slightly.
aga: yours probably is mostly FeCl3 but you should not be surprised to find a small amount still as Fe (II): the K3Fe(CN)6 test for Berlin Blue is
very sensitive.
[Edited on 15-6-2014 by blogfast25]
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CHRIS25
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After 14 Hours - iodometric Titration
Another 14 hours and the results are:
0.544 molarity
0.106 moles of ferric in a 0.5 mole solution ferrous/ferric mixture
0.0076 moles per hour
I added another 0.25 moles (9g) 37% HCl before beginning this third round.
All results compared:
@ 17 hours 0.0026 moles per hour
@ 7 hours 0.0078 moles per hour
@ 14 hours 0.0076 moles per hour
Whatever the scenario, an average of 0.0076 moles of ferrous oxidized per hour is...well....a pretty daft scenario. What I have come to understand is
that all those people I heard say that you could oxidize this amount of ferrous with a month or so of patience did not do any real experimenting/or
simply judged their solutions by the appearance of a golden yellow colour just like I did.
The following happened yesterday but I ignored it, it just happened again: Did my write up, turned back after 10 minutes to clean up and the titrated
solution still sitting on the stirrer with the burette had turned back to the blue colour that you get when you first add corn starch, once again I
allowed drops of thiosulphate into the solution and after 0.8 mL we had a clear solution again. I DID wait after the titration had finished and left
the stirrer going for a few seconds before turning off, just to make sure that the clear solution remained clear. But alas after 10 minutes another
blue solution?
[Edited on 16-6-2014 by CHRIS25]
[Edited on 16-6-2014 by CHRIS25]
[Edited on 16-6-2014 by CHRIS25]
[Edited on 16-6-2014 by CHRIS25]
[Edited on 16-6-2014 by CHRIS25]
[Edited on 16-6-2014 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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A bit of your remaining ferrous oxidised to ferric in that time, the ferric then oxidised more iodide to iodine, et voila: blue colour.
Titrate quite fast and to the first disappearance of blue: that's your end point and don't look back!
In the presence of so much ferrous that titration will never be 100 % accurate but it's ok for your purpose.
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CHRIS25
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| Quote: Originally posted by blogfast25 |
A bit of your remaining ferrous oxidised to ferric in that time, the ferric then oxidised more iodide to iodine, et voila: blue colour.
Titrate quite fast and to the first disappearance of blue: that's your end point and don't look back!
In the presence of so much ferrous that titration will never be 100 % accurate but it's ok for your purpose. |
well that is an eye opener, had no idea that oxidation of ferrous was that fast, obviously though the iodide is extremely sensitive.
‘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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The difference between 'a bit of blue' and 'no blue at all' is essentially 1 drop (about 0.05 ml) of thiosulphate solution: that is 0.1 mol/L x 0.0005
L = 0.05 millimol of iodine. Not much at all. Only 0.1 millimol of ferrous ion needed to be oxidised to cause that bit of blue.
[Edited on 16-6-2014 by blogfast25]
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CHRIS25
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In deep thought the last half hour and Actually I believe I calculated wrongly. And these new figures highlight something interesting.
The first Run was for 17 hours: 0.045 moles were oxidized TOTAL: Therefore this is 0.045 moles actually oxidized. This is 0.0026 moles per hour.
The second run was for 7 hours: 0.055 moles was the TOTAL: Therefore 0.055 - 0.045 = 0.01 moles were actually Oxidized: This is 0.0014 moles per
hour.
Third run was for 14 hours: 0.106 moles were oxidized TOTAL: Therefore 0.106 - 0.055 = 0.05 moles were oxidized: This is 0.0036 moles per hour.
This is me learning maths getting it wrong then realizing my ignorance Sorry! Anyway what is interesting is the fact that after the 7 hour run I
added HCl, the surplus HCl was pretty low during the 17 hour run. The oxidation rate during the 14 hour run is higher than the other two. This can
either be the HCl, Or, the fact that I blocked the air outlet on the pipe and pricked about 30 tiny holes along the length of the tube to decrease
bubble size and this therefore increases the rate at which oxygen can dissolve.
[Edited on 16-6-2014 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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AJKOER
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Here is an interesting discussion (no claim on usefulness or accuracy) using colloidal Silver as a transport for captured active forms of oxygen
(link: http://www.thesilveredge.com/Silver%20carries%20oxygen.shtml... ) that you may find to be at least an interesting experiment. To quote:
"As stated in the report, "The Development and Functions of Silver in Water Purification and Disease Control," by Richard L. Davies and Samuel F.
Etris of The Silver Institute in Washington, DC:
"Atomic (nascent) oxygen adsorbed onto a bed of silver atoms or ions in solution readily reacts with the sulfhydryl (H) groups surrounding the surface
of bacteria or viruses to remove the hydrogen atoms (removed as water) causing the sulfur atoms to form an R-S-S-R bond; respiration is blocked and
the bacteria expire."
The researchers also state in the same report:
"It has long been known that oxygen is adsorbed on the surface of silver in its atomic state. Also that oxygen diffuses more freely within silver than
within any other metal.
Ronald Outlaw, working at NASA/Langley, undertook a fundamental study of oxygen diffusion with the objective of producing atomic oxygen for the
evaluation of the degradation of organic materials in space. He discovered the most prolific source of nascent oxygen to be metallic silver.
Atomic oxygen fits very well in the octahedral holes of gold, silver, and copper. In gold, the electron cloud of oxygen tends to be repelled by the
lattice electrons of the gold atoms stopping movement through the holes. With copper, the oxide is formed resulting in a barrier.
Silver, with an almost a perfect fit, offers so little repulsion that a little thermal energy will readily move it from hole to hole.
...Molecular oxygen is present and silver readily adsorbs it converting it to nascent oxygen which is available to oxidize bacterial enzymes and other
organics. Their reaction with the atomic oxygen is instantaneous."
Possible sources for "nascent oxygen" have been attributed (questionable) to H2O2, NaClO, O3 and the electrolysis of water. There is actually a good
reference for the last method (see "Epoxidation of cyclohexene with the nascent oxygen generated by electrolysis of water", by Kiyoshi Otsuka,
Masahiro Yoshinaka and Ichiro Yamanaka, in the J. Chem. Soc., Chem. Commun., 1993, 611-612, DOI: 10.1039/C39930000611, link: http://www.sciencemadness.org/talk/post.phpaction=edit&f... ).
[Edit] The action of sunlight on O2 may be a simple path to so called sinlet oxygen (see http://en.m.wikipedia.org/wiki/Singlet_oxygen ) in the presence of a photosensitizer pigment, likely distinct from nascent oxygen. Interestingly,
longer lasting singlet oxygen is chemically produced by the action of H2O2 on NaClO (same Wiki article).
Employing singlet Oxygen in your experiment (with or without colloidal Silver) may prove quite interesting!
[Edited on 17-6-2014 by AJKOER]
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blogfast25
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| Quote: Originally posted by CHRIS25 | The oxidation rate during the 14 hour run is higher than the other two. This can either be the HCl, Or, the fact that I blocked the air outlet on the
pipe and pricked about 30 tiny holes along the length of the tube to decrease bubble size and this therefore increases the rate at which oxygen can
dissolve.
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That seems plausible. But these reaction rates remain very small.
@AJ:
A catalyst could be very useful here and that thought has certainly occurred to me.
But finely divided silver is likely to become deactivated quickly in the presence of strong HCl and oxygen: it will turn into silver chloride I think.
It could be worth a try but colloidal silver (as homeopathic 'medicine') is really expensive...
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AJKOER
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Blogfast:
Note my edit suggesting the easily formed (NaClO + H2O2) so called singlet Oxygen.
Here is an education site describing the following demonstration experiment (link
http://www1.chem.leeds.ac.uk/delights/texts/expt_27.html )
"In this demonstration the singlet oxygen is prepared by oxidising hydrogen peroxide with chlorine gas. The initial product formed when chlorine is
passed through an alkaline hydrogen peroxide solution is hypochlorite:
Cl2 + 2 OH–==> OCl– + Cl– + H2O (27.1)
The hypochlorite ion probably reacts with hydrogen peroxide to give the chloroperoxide ion, which splits off chloride to form oxygen:
H2O2 + OCl–==> ClOO– + H2O (27.2)
ClOO–==> 1O2 + Cl– (27.3)
According to the law of conservation of spin oxygen is formed in the excited singlet state rather than in the triplet state, which is the ground state
of molecular oxygen. Triplet oxygen can be converted into electronically excited singlet oxygen by supplying energy. The excited oxygen molecules
emit red light on returning to the triplet ground state.
1O2==> 3O2 + hn (l = 634 nm) (27.4)
The chemical behaviour of the two oxygen species is also quite different.
Preparation. Prior to the experiment prepare in a 250 ml beaker the solution of 20 g of NaOH in 140 ml of water. In a 50 ml beaker pour
30 ml of 30% H2O2. Cool both beakers in the ice bath. Prepare a 500 ml wash bottle with a glass frit and connect it to a chlorine cylinder (a small
lecture bottle is perfect). Alternatively chlorine can be generated from KMnO4 and concentrated hydrochloric acid).
Demonstration. The two well-cooled solutions are poured into the wash bottle and the room is darkened. A rapid stream of chlorine gas is
now passed through the solution, which at once glows bright red."
This activated form of Oxygen may provide me with some interesting cool experiments!
[Edited on 17-6-2014 by AJKOER]
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Fantasma4500
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about H2O2 concentrating i recall plante1999 talked about he had 90% yield based on oxygen gas titration and reached +80%
i might be slightly off but i do recall there was minimal loss of H2O2... using a fan or something else to suck out the water will greatly increase
yield / lower loss
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