KoiosPhoebus
Hazard to Self
Posts: 53
Registered: 23-1-2023
Member Is Offline
|
|
Oxygen scavengers for working with Iron(II) and Manganese(II) salts?
Hi everyone!
I'm trying to prepare some complexes of Fe(II) and Mn(II) (I've got some discussion of my progress on the glycine complexes at the bottom of this
topic - https://www.sciencemadness.org/whisper/viewthread.php?tid=15...).
Unfortunately, Fe(II) salts oxidise very readily in solution to Fe(III). Mn(II) is more oxidation-resistant but will also follow suit in an organic
solvent and/or at alkaline pH (or if left exposed to air for a while). The usual solution of working at a low pH is not a viable option as the
complexes often form adducts with acids. e.g. the conjugate base of the amino acid glycine has sufficiently strong affinity for H+ that the
following reaction is able to proceed:
Mn(Gly)2 + 2 H+ + SO42- → Mn2+ + 2 GlyH + SO42-
With the product being ferrous glycine sulphate or manganous glycine sulphate instead of the desired ferrous bisglycinate/manganous bisglycinate.
Additionally, many acidic reducing agents like oxalic acid or citric acid form their own complexes with Fe(II) and Mn(II), so instead of synthesising
Fe(Gly)2.H2Oxalate, I might end up with Fe(Oxalate).(GlyH)2.
Another alternative I've read about is working at elevated temperatures to reduce the solubility of gases. This does seem to work somewhat; however in
my experience oxidation still occurs after some time, especially for the Fe(II) salts. Additionally, I have one methodology (https://doi.org/10.1016/S0020-1693(00)82257-4) where slow crystallisation from ethanol/water mixture is required - constant heating would make
that difficult.
Does anyone have any suggestions on how I might prepare a non-oxidising environment for me to work with Fe(II) or Mn(II) salts? I've read some stuff
on inert atmospheres but a lot of it seems to require specialised equipment which I don't currently have e.g. glove boxes. Is there an equivalent for
the home chemist or another reasonably inexpensive way to replicate an inert atmosphere or deoxygenate a solvent at home?
Thanks in advance for any advice!
|
|
|
DraconicAcid
International Hazard
Posts: 4363
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
If you put a bit of metallic iron in the solution, then the iron will stay at Fe(II),
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.
|
|
|
unionised
International Hazard
Posts: 5129
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Have you considered bisulphite.
I don't think it's completely immune from complex formation but it's not bad.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Huh? Example? 
|
|
|
KoiosPhoebus
Hazard to Self
Posts: 53
Registered: 23-1-2023
Member Is Offline
|
|
Thanks, I'll keep that in mind. I'm a little unsure as to how well that works
in an organic solvent as I've always seen it used in water; the low solubility of most Fe salts in organic solvents might reduce the rate at which
Fe0 is able to reduce any Fe3+ formed. Additionally, I'm not sure how this would work if I'm trying to crystallise out the
ferrous complex - how would I go about separating the iron from the desired salt?
Bisulphite (as in HSO3-) is acidic so there's a good
chance it'll react with the complex. Sulphite (SO32-) is an interesting option I'll consider, but as far as I can tell, sodium
sulphite and potassium sulphite are considered insoluble (or "sparingly soluble") in ethanol. They might be useful when I'm running aqueous reactions
(and I'll keep them in mind for those), but not so much when I'm working in organic (or mixed aqueous/organic) solvents.
A couple of other options I've been considering:
Reducing agents from photography, e.g. pyrogallol, pyrocatechol, metol, hydroquinone. However, these are quite pricey and I'm not
very familar with their use (e.g. which one has more reducing power) and whether they might complex the metal cations.
Working in a sealed bag with reducing gases. e.g. placing an evaporating dish in the bag with dissolved potassium metabisulphite, or preparing
my initial metal salt using a strong acid (so instead of adding manganese chloride, I'd add manganese metal to hydrochloric acid to evolve hydrogen
gas). This could work but I'm unsure as to how effective it would be - I've heard of people using similar methods a la the "balloon transfer" method
and it apparently wasn't very effective - https://www.researchgate.net/post/What-are-the-possible-ways...). Additionally I'm a little iffy on how safe it is to work with gases like
SO2.
Working in a sealed bag, but using an "oxygen trap" to remove the oxygen within. e.g. placing a dish of cold ascorbic acid solution to remove as
much of the oxygen as possible into the bag, then running my reaction in said bag. Again, I'm unsure how effective this is as well as the
practicalities of working in such an environment.
If anyone has used any of these methods before, or has any other suggestions I'd love to hear them!
|
|
|
Texium
Administrator
Posts: 4624
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Phenols aren’t a good option: they complex iron (I don’t know about manganese, but I wouldn’t be surprised if they did) and their iron(III)
complexes are an intense dark purple color.
|
|
|
KoiosPhoebus
Hazard to Self
Posts: 53
Registered: 23-1-2023
Member Is Offline
|
|
Quote: Originally posted by Texium  | | Phenols aren’t a good option: they complex iron (I don’t know about manganese, but I wouldn’t be surprised if they did) and their iron(III)
complexes are an intense dark purple color. |
Ah, okay. Thanks for letting me know! (and saving me quite a bit of money in
the process)
I do wonder if they could still be used in an "oxygen trap", but there's cheaper chemicals (like ascorbic acid) which could also be used
there.
|
|
|
woelen
Super Administrator
Posts: 8034
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Maybe you could use sodium dithionite. That's a strong reductor. It is important though that the pH of the solution does not become low, because
dithionite is unstable at low pH and decomposes to sulfur, sulfur dioxide, and other (acidic) oxo species of sulfur.
|
|
|
j_sum1
Administrator
Posts: 6338
Registered: 4-10-2014
Location: At home
Member Is Offline
Mood: Most of the ducks are in a row
|
|
There are ways of making a glovebox using plastic containers. You might also rig up an inert environment with an argon balloon attached to some normal
glassware.
I don't have a lot of experience but I have seen both done with reasonable success.
|
|
|
KoiosPhoebus
Hazard to Self
Posts: 53
Registered: 23-1-2023
Member Is Offline
|
|
| Quote: Originally posted by woelen | | Maybe you could use sodium dithionite. That's a strong reductor. It is important though that the pH of the solution does not become low, because
dithionite is unstable at low pH and decomposes to sulfur, sulfur dioxide, and other (acidic) oxo species of sulfur. |
Thanks for the input woelen. Is there a particular reason you suggest
dithionite instead of metabisulfite? The latter is more commonly available where I am, and I'm under the impression that it's also more
reduced?
| Quote: Originally posted by j_sum1 | There are ways of making a glovebox using plastic containers. You might also rig up an inert environment with an argon balloon attached to some normal
glassware.
I don't have a lot of experience but I have seen both done with reasonable success. |
Thanks for the input, I might look into the rigging up a glovebox stuff. The
ResearchGate thread I posted above did however suggest that the "balloon transfer" method (unsure if this is what you're referring to?) resulted in
carbon dioxide and water vapour levels returning to atmospheric levels after a few minutes. It's probably not the biggest deal with Mn(II) but my
experience is that Fe(II) compounds oxidise given relatively small exposure to oxygen when in basic solutions.
Does anyone know if a vacuum desiccator would also work for this purpose? I'm considering buying one anyway as some of the complexes simply will not
dry out even under strong heating or being placed into a bag with calcium chloride. I've never used one before though so I have no idea if it'll purge
the insides of oxygen, and they're rather pricey so I'd want to be sure it'll do the job before buying it.
|
|
|
teodor
National Hazard
Posts: 928
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
So, when, for example, you mix water solutions of MnSO4 and barium glycinate, do you experience any Mn(II) oxidation?
Also there is another published method, mixing alcoholic solution of Mn(NO3)2 and aqueous lithium glycinate in the molar ration 1:2. The precipitate
should be washied with alcohol and dried in vacuum at 70C
[Edited on 21-3-2023 by teodor]
|
|
|
woelen
Super Administrator
Posts: 8034
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Dithionite is a MUCH stronger reductor than metabisulfite. Dithionite must not be confused with the very tame dithionate! It contains sulfur in
(average) oxidation state +3, while metabisulfite (and also sulfite) contain sulfur in oxidation state +4. Dithionite is such a strong oxygen
scavenger, that it even can ignite in air, when it is made humid. If you take some solid Na2S2O4 and make the powder damp, while keeping it in contact
with air, then the powder becomes hot, due to the reaction with oxygen from the air.
I myself have used Na2S2O4 in reactions with Mn(OH)2, which I wanted to keep pale pink. Without the Na2S2O4 the suspension of Mn(OH)2 immediately
turns brown in contact with air, but with Na2S2O4, the mix remains white or pale pink.
|
|
|
KoiosPhoebus
Hazard to Self
Posts: 53
Registered: 23-1-2023
Member Is Offline
|
|
| Quote: Originally posted by teodor | So, when, for example, you mix water solutions of MnSO4 and barium glycinate, do you experience any Mn(II) oxidation?
[Edited on 21-3-2023 by teodor] |
I use calcium glycinate or histidinate in place of barium, but broadly
speaking as long as the temperature is kept close to boiling I have not seen much oxidation. The complexes formed are dark red (and Mn(III) is brown)
so I've tested for oxidation by pipetting a small amount into a solution of NaOH; typically a white/cream precipitate forms with at most traces of
brown.
For Mn(II), the problem with the double displacement is not that it oxidises (it's mostly fine as long as kept hot), but that I could not get the
complexes formed out of solution. So far, I've attempted:
Precipitation with organic solvents (acetone/ethanol). Introducing the organic solvent has thus far produced a creamy organic
solvent layer and a water layer, which then becomes this oily, gooey mess. Stirring does not help as the complex/water mixture just becomes these
sticky clumps; filtering off the organic solvent layer (or precipitating in organic solvent and filtering that off) has so far also failed as the
precipitate simply blocks filters. When in organic solvent, the complex also oxidises much more readily, probably due to the increased solubility of
oxygen in said solvents.
Very strong heating, practically boiling. The complex/water mix gets to a point with very little water, and then forms this viscous goo. Further
heating results in vigorous bubbling and the formation of a sticky, harder goo with the texture of melted rock candy. Further heating simply resulted
in the formation of more bubbles but not any change in the texture or structure of the hard goo.
Heating it to reduce the volume of liquid, and then placing it into a sealed bag with calcium chloride. After leaving it for a couple of hours,
the solution remains as a relatively flowable liquid, with no apparent drying-out. Additionally, while I never got around to testing it for oxidation,
I suspect more of the Mn(II) would have oxidised in this method as the dish containing the complex would cool down over time in the bag, increasing
the solubility of oxygen.
Note that the "minimal oxidation while kept hot" statement only applies to Mn(II). With Fe(II), the complex oxidises even when the solution is kept at
near-boiling temperatures, forming red ferric oxide on the side of the dish and a layer of Fe(II,III) oxide on the surface of the solution which can
stick to magnetic stir bars etc.
Fe(II) will also oxidise even when citric acid or L-ascorbic acid are added into the solution. I prefer not to add compounds which might complex with
the metal, but I saw this done in a patent and in a separate publication so I thought I'd try it out as well; unfortunately it did not prevent
oxidation.
Additionally with ascorbic acid, the solution rapidly turned purple, which is a sign of the formation of Fe(Ascorbate). I suspect citric acid does
something similar as the solution remained green, which is usually a sign of Fe(Citrate) formation as Fe(Gly)2 is yellow-brown in
solution. However I'm not 100% sure if Fe(Citrate) was formed as Fe2+ is also green in solution.
| Quote: Originally posted by teodor |
Also there is another published method, mixing alcoholic solution of Mn(NO3)2 and aqueous lithium glycinate in the molar ration 1:2. The precipitate
should be washied with alcohol and dried in vacuum at 70C
[Edited on 21-3-2023 by teodor] |
Yep, I've heard of this method as well though with FeCl2 instead.
The problem is that oxygen dissolves much more readily in alcohol solvents than in water, so the Mn(II) complex oxidises very readily in alcohol-water
mixtures. Within minutes, the orange/cream mixture (depending on which amino acid is used) turns brown, presumably oxidising to Mn(III).
| Quote: Originally posted by woelen | Dithionite is a MUCH stronger reductor than metabisulfite. Dithionite must not be confused with the very tame dithionate! It contains sulfur in
(average) oxidation state +3, while metabisulfite (and also sulfite) contain sulfur in oxidation state +4. Dithionite is such a strong oxygen
scavenger, that it even can ignite in air, when it is made humid. If you take some solid Na2S2O4 and make the powder damp, while keeping it in contact
with air, then the powder becomes hot, due to the reaction with oxygen from the air.
I myself have used Na2S2O4 in reactions with Mn(OH)2, which I wanted to keep pale pink. Without the Na2S2O4 the suspension of Mn(OH)2 immediately
turns brown in contact with air, but with Na2S2O4, the mix remains white or pale pink. |
Thanks for correcting me on that.I'm not sure I can find dithionite but I'll
keep it in mind.
Would anyone happen to know if a vacuum desiccator would prevent oxidation if I pulled a vacuum on the dish?
|
|
|
DraconicAcid
International Hazard
Posts: 4363
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by KoiosPhoebus | | Precipitation with organic solvents (acetone/ethanol). Introducing the organic solvent has thus far produced a creamy organic solvent layer and a
water layer, which then becomes this oily, gooey mess. |
Don't use acetone for that- it far too easily gets salted out. Methanol is your best bet.
That being said, I had similar trouble trying to isolate nickel glycinate complexes- I eventually just gave up.
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.
|
|
|
teodor
National Hazard
Posts: 928
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by KoiosPhoebus |
For Mn(II), the problem with the double displacement is not that it oxidises (it's mostly fine as long as kept hot), but that I could not get the
complexes formed out of solution. |
So, I think those are important points:
1. The molar ratio. Mn(C2H4NO2)2 * 2H2O could be isolated, but the solution is typically the mixture of Mn(HGly)2+ and MnGly(HGly)+ ions (HGly =
zwitterion)
2. The temperature colse to 70C. I think this is the zone where dihydrate crystallizes.
3. The method with lithium glycinate wich forms "white precipitate" probably is dependent on the initial concentration of solutions, try to vary it
(the original publication was in a soviet union magazine, I have no idea how to access its archives).
[Edited on 22-3-2023 by teodor]
|
|
|
KoiosPhoebus
Hazard to Self
Posts: 53
Registered: 23-1-2023
Member Is Offline
|
|
| Quote: Originally posted by DraconicAcid |
Don't use acetone for that- it far too easily gets salted out. Methanol is your best bet.
That being said, I had similar trouble trying to isolate nickel glycinate complexes- I eventually just gave up. |
Hmmm, I'll think about using methanol. I've also tried the same procedure, but
with histidine as the ligand and water/ethanol as the medium - water for the lithium histidinate and ethanol for the manganese chloride. This
procedure was used in another paper to prepare ferrous histidinate dihydrate (https://doi.org/10.1016/S0020-1693(00)82257-4). For the most part it
seemed to work - initial formation of a cream precipitate followed by dissolution of precipitate and the solution changing to orange. However the
relatively high solubility of oxygen in ethanol meant that the mixture oxidised very rapidly, going to brown and then black.
Interestingly someone else in the metal glycinates thread (https://www.sciencemadness.org/whisper/viewthread.php?tid=15...) managed to prepare and isolate nickel glycinate by heating on a steam bath. I'm
not sure why this worked and heating on a hotplate failed - maybe it has something to do with getting more even heating?
It's been incredibly frustrating to me as I've been able to crystallise compounds which are described as deliquescent, but somehow a couple of metal
amino acid complexes elude me.
| Quote: Originally posted by teodor |
So, I think those are important points:
1. The molar ratio. Mn(C2H4NO2)2 * 2H2O could be isolated, but the solution is typically the mixture of Mn(HGly)2+ and MnGly(HGly)+ ions (HGly =
zwitterion)
2. The temperature colse to 70C. I think this is the zone where dihydrate crystallizes.
3. The method with lithium glycinate wich forms "white precipitate" probably is dependent on the initial concentration of solutions, try to vary it
(the original publication was in a soviet union magazine, I have no idea how to access its archives).
[Edited on 22-3-2023 by teodor] |
Thanks for sharing that! I've tried both methods - the first is difficult to
crystallise (apparently Fe(Gly)2 is also challenging to isolate this way - see: https://www.researchgate.net/post/How_can_I_synthesize_pure_...), while the second oxidises readily in contact with air.
My hotplate doesn't have a temperature setting, but I've tried drying at quite a few variations of temperature, from near-boiling to relatively cool
temperatures (50 deg C). It's possible that I just need to keep the temperature at a steady 70 deg C, but I suspect the bigger factor here might be in
the use of a vacuum to dry the solution.
I'm starting to think that the only way out here might be to use a vacuum desiccator to purge the insides of oxygen while drying, and to properly
degas the solvents (so far I've just been boiling everything). The paper that I had on preparing iron(II) complexes of amino acids (https://doi.org/10.1016/S0020-1693(00)82257-4) also seems to suggest drying under vacuum or under nitrogen ("They were dried by continuous pumping
for several hours"). I have no idea how I'm going to heat it at 70 deg C under vacuum though - the vacuum cabinets large enough for something like
that are way out of my price range.
If anyone has any ideas on either drying very stubborn metal complexes or maintaining oxygen free environments when drying a precipitate, I'd love to
hear more! The full vacuum desiccator setup is rather pricey and I'd prefer not to have to go down that route if possible.
Also if anyone has any thoughts on whether a vacuum desiccator would actually do the job of keeping the mixture oxygen-free, I'd very much appreciate
your input. It'd be annoying to buy the full setup and find that it didn't purge oxygen fast enough to prevent oxidation or something similar.
|
|
|
teodor
National Hazard
Posts: 928
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by KoiosPhoebus |
It'd be annoying to buy the full setup and find that it didn't purge oxygen
fast enough to prevent oxidation or something similar.
|
KoiosPhoebus, for your experiment you need:
1. The kettle like this one.
It can keep the constant temperature of water inside. The price - 40EUR.
My can work perfectly with 250ml flasks if I detach the cap of the kettle.
2. A round-bottom flask.
3. A water aspirator. There are good cheep plastic models.
4. A rubber hose.
5. An adapter of hose to a ground glass joint.
6. A vacuum lubricant.
7. Something to fix the flask inside the kettle.
To test your setup put a bit warm water in a flask, apply vacuum and wait till it boils, then boil it until it stops boiling (because it will cool
down). Then measure the boiling temperature and find the pressure by a table. Typically, to get good vacuum you should pay attention to the quality of
rubber, lubrication of joints and tighting the hose connections as much as possible. The test with boiling water will proof you have the proper setup
(the preassure from an aspirator typically is in the range 10-25 mmHg, depending on the time of year and quality of your hose connections).
Also I think in the lithium method they use saturated solution.
I have not so big experience of Fe(II) salts preparation but the method of zink metal in acidic solution worked perfectly for me (to keep the small
current of "nascent" H2 in the mixture which prevents Fe(II) from oxidation.
Update:
it is known that H2 bubbling can reduce some compound (e.g. permanganate), I am wondering, can hydrogen reduce Fe(III) to Fe(II)
(solvent/temperature/catalyser) ?
[Edited on 23-3-2023 by teodor]
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
