Sciencemadness Discussion Board - Generate oxygen for oxidation by pump

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.

CHRIS25 - 10-6-2014 at 13:21

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]

aga - 10-6-2014 at 13:42

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 !

CHRIS25 - 10-6-2014 at 14:24

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]

blogfast25 - 11-6-2014 at 05:05

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]

CHRIS25 - 11-6-2014 at 06:40

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

blogfast25 - 11-6-2014 at 08:11

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.

aga - 11-6-2014 at 08:25

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

CHRIS25 - 11-6-2014 at 09:01

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.

blogfast25 - 11-6-2014 at 09:22

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]

aga - 11-6-2014 at 09:26

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]

blogfast25 - 11-6-2014 at 09:31

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?

aga - 11-6-2014 at 09:46

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]

aga - 11-6-2014 at 10:06

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]

blogfast25 - 11-6-2014 at 10:53

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.

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.

WGTR - 11-6-2014 at 11:03

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.

blogfast25 - 11-6-2014 at 12:16

Quote: Originally posted by WGTR

Perhaps oxidation of a neutral ferrous salt with ozone?

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...

Fantasma4500 - 12-6-2014 at 02:55

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

blogfast25 - 12-6-2014 at 04:29

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]

aga - 12-6-2014 at 07:52

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]

Zyklon-A - 12-6-2014 at 08:39

as Time is clearly not a factor, concentrate the peroxide by slow heating over ages and ages and ages ...

That kinda works.

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]

aga - 12-6-2014 at 12:00

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.

blogfast25 - 14-6-2014 at 05:07

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.

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]

Air pump tests confirmed:

CHRIS25 - 14-6-2014 at 08:44

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)?

aga - 14-6-2014 at 09:03

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]

CHRIS25 - 14-6-2014 at 09:10

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]

aga - 14-6-2014 at 09:25

Also if I doubled the output of the air pump

The idea of you hooking up multiple pumps made this spring to mind :-
http://simpsons.wikia.com/wiki/French_Chef?file=French_Chef3...

blogfast25 - 15-6-2014 at 04:33

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]

After 14 Hours - iodometric Titration

CHRIS25 - 16-6-2014 at 04:02

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]

blogfast25 - 16-6-2014 at 05:12

But alas after 10 minutes another blue solution?

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.

CHRIS25 - 16-6-2014 at 08:36

But alas after 10 minutes another blue solution?

well that is an eye opener, had no idea that oxidation of ferrous was that fast, obviously though the iodide is extremely sensitive.

blogfast25 - 16-6-2014 at 08:57

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]

CHRIS25 - 16-6-2014 at 10:04

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]

AJKOER - 16-6-2014 at 17:17

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]

blogfast25 - 17-6-2014 at 04:40