BarryScott
Harmless
Posts: 3
Registered: 29-8-2015
Member Is Offline
Mood: No Mood
|
|
Unknown Products: Electrolysis of MgSO4 + NaHCO3 and Iron Cathodes.
Hello, Barry Scott here (very new here so this may be in the wrong area!)
I'm relatively new to electrolysis and while i was messing around attempting to get the insoluble Magnesium carbonate, i decided to electrolyze an
aqueous solution of Magnesium Sulphate and Sodium Bicarbonate. both my anode and cathode are low grade steel.
i was expecting to get a reaction somewhat like this
2MGSO4(aq) + 2NaHCO3(aq) -DC-> 2MgCO3(s) + 2NaSO4(aq) + H2(g)
However i have had a reaction more like this
2MGSO4(aq) + 2NaHCO3(aq) + Fe(s) -DC-> ??? + ??? +??? + H2?
I have received 3 precipitates; one which is dark swampy green on the surface with some light brown foam. the second, a white chunky precipitate
distributed evenly in the solution. The third precipitate is between the aqueous solution and the white precipitate, it is dark, dull blue.
A flammable gas also is formed during the reaction. i believe it is hydrogen.
the reaction has also been exothermic as the solution is warm to the touch.
Another possible theory is that, i may be producing sulphuric acid and it is reacting with the iron anode and cathode (but only one of the cathodes
look degraded). But then what is the flammable gas?
if anyone could help me identify the precipitates formed in the reaction it would be of great assistance!
Many thanks ,
Barry.
UPDATE:
The swampy green substance on the surface dropped down into the dark blue precipitate, which suggests they were the same chemical complex. after
filtering the precipitates i believe a large amount of iron oxide dissolved into solution (classic brown color and odor), the origin of which i do not
know.
[Edited on 30-8-2015 by BarryScott]
|
|
|
BarryScott
Harmless
Posts: 3
Registered: 29-8-2015
Member Is Offline
Mood: No Mood
|
|
UPDATE 2:
after leaving the precipitate to dry, the once dull blue precipitate has turned a wet, rust brown. It does not have a distinguishable odor. Perhaps it
has reacted with the oxygen in the air?
|
|
|
kecskesajt
Hazard to Others
Posts: 299
Registered: 7-12-2014
Location: Hungary
Member Is Offline
Mood: No Mood
|
|
Light brown foam: Iron hydroxide,carbonate,oxide, etc...
Swampy green: water soluble iron compounds
White precipitate: Magnesium carbonate
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
To prepare MgCO3, which I have performed on numerous occasions, simply first dissolve NaHCO3 in a concentrated solution of MgSO4. Then, I usually heat
the solution in a microwave to decompose the Mg(HCO3)2 into a fine white suspension of MgCO3 (but likely some basic magnesium carbonate as well, see,
for example, http://www.researchgate.net/publication/237947862_Precipitat... ) . Reactions:
MgSO4 + 2 NaHCO3 = Mg(HCO3)2 + Na2SO4
Mg(HCO3)2 + H2O --Boiling--) MgCO3 (s) + 2 H2O + CO2 (g)
--------------------------------------------------------------------
The electrolysis approach is a bad choice for several reasons, in my opinion. Generally speaking, electrolysis is slow and the general chemistry can
get complex (due, for example, the formation of active radical species for one, see discussion and cited references at http://chemistry.stackexchange.com/questions/12315/electroly...
). As an illustration in the current case, I would expect hydroxyl radicals, .OH, which in the presence of carbonate, CO3(2-), or the bicarbonate
anion, HCO3-, can further create the reactive carbonate radical anion, .CO3-, and depending on pH, .HCO3. Some radical based chemistry (reference,
please see https://www.google.com/url?sa=t&source=web&rct=j&... ):
HCO3- + .OH = .CO3- + H2O
CO3(2-) + .OH = .CO3- + OH-
.CO3- + H+ = .HCO3
Also, I would expect oxygen to get into the mix sourced from the electrolysis of water and/or air contact. As Iron + O2 in an electrolyte (MgSO4)
forms an Iron-air battery (more electrochemistry, for details on this galvanic cell, see for example: http://m.jes.ecsdl.org/content/159/8/A1209.full ), so expect a ferrous salt from this experiment. From your iron electrodes, possibly also some
Fenton, Electro-Fenton and even Photo-Fenton chemistry in the presence of light could also enter the picture (as some background, see for example,
"Fundamental Mechanistic Studies of the Photo-Fenton Reaction for the Degradation of Organic Pollutants" at https://www.google.com/url?sa=t&source=web&rct=j&... ) forming ferric salts, more hydroxyl radicals, carbonate radical anions,...
Then, of course, there is a host of possible standard chemical side reactions involving more than one iron species.
Bottom line, I seriously doubt a simple (and accurate) explanation is possible surrounding your electrolysis.
[Edit] With respect to Kecskesagt response on the insoluble swamply green salt, my guess is possibly a mixed ferrous salt (that is, a double salt). I
agree on the white salt is MgCO3 if the reaction mix got to over 50 C, otherwise hydroxyl anions, -OH, could also be around (resulting in Mg(OH)2 or a
basic salt). The latter OH-, for example, steming from a possible galvanic (Iron-air battery) cell reaction, or via the creation of the carbonate
radical anion (see reaction above), or as an intermediate product from the electrolysis of water, which then reacts with MgSO4 forming Mg(OH)2.
[Edited on 1-9-2015 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Actually found a paper "The Oxidation of Iron(II) with Oxygen in NaCl Brines", published in 2007 by J. Michael Trapp and Frank J. Millero, fully
available at
https://www.google.com/url?sa=t&source=web&rct=j&url=http://www.rsmas.miami.edu/users/jtrapp/files/The%2520Oxidation%2520of%2520Iron(II)%2
520with%2520Oxygen%2520in%2520NaCl%2520Brines.pdf&ved=0CB0QFjAAahUKEwjKyaHixtjHAhXDmIAKHbZMA74&usg=AFQjCNF5zZw5rz0Dl07mUFHYHB2pPddhmg&sig2
=9ealO3hLAuOlZMh6Ggn8cQ
To quote from the journal page 1481:
"These studies show that at the levels of carbonate in natural sea water, the hydroxyl species of iron dominate the rates of the oxidation.This study
also demonstrates that the rates of oxidation of Fe(II) at nano-molar level were different than earlier studies at micro molar levels.These
differences have been attributed to a change in the reaction mechanism [18,19]. At low levels, iron was not able to compete for the intermediates
formed in the Harber-Weiss [20] mechanism
Fe(II) + O2 → Fe(III) + O−2· (15)
Fe(II) + O−2· + 2H+ → Fe(III) + H2O2 (16)
Fe(II) + H2O2 → Fe(III) + OH· + OH− (17)
Fe(II) + OH· → Fe(III) + OH− (18) "
Now, in the current context of an electrolysis with iron electrodes in the presence of O2 and an electrolyte, the ability of Iron to compete for
intermediates above is not an issue.
[Edited on 2-9-2015 by AJKOER]
|
|
|
BarryScott
Harmless
Posts: 3
Registered: 29-8-2015
Member Is Offline
Mood: No Mood
|
|
Wow!
Thanks to everyone for explaining this to me.
I feel somewhat stupid now. but thanks!
Barry
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
