Velzee
Hazard to Others
Posts: 379
Registered: 19-8-2015
Location: New York
Member Is Offline
Mood: Taking it easy
|
|
Ferricyanide reduction to Ferrocyanide?
Using the SE, I haven't come across a viable method (besides reduction using H2O2 and KOH) of reducing potassium ferricyanide to potassium
ferrocyanide without H2O2. But I am known to miss things right in front of me, so I ask; does anyone know of an effective method to
reduce the salt to ferrocyanide? I have a decent amount of it, and I find I often have more of a need of the ferrocyanide salt.
Thank you in advance!
Check out the ScienceMadness Wiki: http://www.sciencemadness.org/smwiki/index.php/Main_Page
"All truth passes through three stages. First, it is ridiculed. Second, it is violently opposed. Third, it is accepted as being self-evident."
—Arthur Schopenhauer
"¡Vivá Cristo Rey!"
—Saint José Sánchez del Río |
|
|
clearly_not_atara
International Hazard
Posts: 2691
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
According to Wiki the standard electrode potential of ferricyanide -> ferrocyanide is +0.36V, so many things should be able to reduce it. Ascorbate
would be my first guess. According to the SEP (which is quite high actually), zinc metal, Sn2+, Cr2+, bisulfite, and thiosulfate can all reduce it, as
well as many reactive metals, but whether they work in practice I don't know. I think hydrogen peroxide is preferred because it is very cheap and
there are no byproducts.
But if even H2O2 is unavailable then the right answer to this question must depend on what you have at hand. Ferricyanide is an oxidizer, so should be
easy to reduce.
[Edited on 04-20-1969 by clearly_not_atara]
|
|
|
Velzee
Hazard to Others
Posts: 379
Registered: 19-8-2015
Location: New York
Member Is Offline
Mood: Taking it easy
|
|
Reduction by sulfite (Na2SO3 + KOH) seemed to work  Now to scale it up and
isolate the salt...
Check out the ScienceMadness Wiki: http://www.sciencemadness.org/smwiki/index.php/Main_Page
"All truth passes through three stages. First, it is ridiculed. Second, it is violently opposed. Third, it is accepted as being self-evident."
—Arthur Schopenhauer
"¡Vivá Cristo Rey!"
—Saint José Sánchez del Río |
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
How do you know whether all fericyanide is reduced to ferrocyanide or not? A mix with a little ferricyanide in it may be hard to distinguish from pure
ferrocyanide.
Did you use precise stoichiometric amounts? Tell us a little more about the precise conditions/procedure of your experiment, it may be quite
interesting, also for other members.
|
|
|
pantone159
National Hazard
Posts: 586
Registered: 27-6-2006
Location: Austin, TX, USA
Member Is Offline
Mood: desperate for shade
|
|
I have used, for kinetics experiments calibration, reduction with ascorbic acid, in a pH 4.4 KH2PO4 buffer. I noted the rate at this pH is supposedly
1.2/sec. Reduction of the ferricyanide was measured by absorbance at 320 nm and 418 nm, the reaction seemed to go to completion in about 8 sec with
the concentration that I used. I made no attempt to isolate the ferrocyanide.
Martins, L, and de Costa, J. Journal of Chemical Education 65, 176-178 (1988). The Oxidation of Ascorbic Acid by Hexacyanoferrate(III) Ion in Acidic
Aqueous Media.
|
|
|
Velzee
Hazard to Others
Posts: 379
Registered: 19-8-2015
Location: New York
Member Is Offline
Mood: Taking it easy
|
|
Quote: Originally posted by woelen  | How do you know whether all fericyanide is reduced to ferrocyanide or not? A mix with a little ferricyanide in it may be hard to distinguish from pure
ferrocyanide.
Did you use precise stoichiometric amounts? Tell us a little more about the precise conditions/procedure of your experiment, it may be quite
interesting, also for other members. |
I first made a solution of 5 grams of K3[Fe(CN)6] in ~50mL (too much in my opinion) and 1 gram of KOH , a little more than the 0.85g required to
account for any water, which forms an orangish-red colored solution with a slight yellow touch; I then added 1 gram of Na2SO3 (again, slightly more
than the stoichiometric amount of 0.96g) to just enough water to dissolve it. I then mixed these two solutions, and within seconds, the color had
changed to the familiar pale/dull yellow solution of potassium ferrocyanide. Upon sitting overnight, though, orange specks of what I presume to be
Iron hydroxides were found at the bottom of the beaker.
Check out the ScienceMadness Wiki: http://www.sciencemadness.org/smwiki/index.php/Main_Page
"All truth passes through three stages. First, it is ridiculed. Second, it is violently opposed. Third, it is accepted as being self-evident."
—Arthur Schopenhauer
"¡Vivá Cristo Rey!"
—Saint José Sánchez del Río |
|
|
The Volatile Chemist
International Hazard
Posts: 1981
Registered: 22-3-2014
Location: 'Stil' in the lab...
Member Is Offline
Mood: Copious
|
|
Ability to form Prussian blue is how I usually distinguish between the salts, especially because its detectable at a low level. If I recall, you can
add Fe(II) to an Ferrocyanide solution and it shouldn't form Prussian blue unless there's ferricyanide present. Could be out my rear though, sorry if
I'm wrong. I'll check some notes or just try it.
I also brought up on the Skype group that I recall isopropanol being slowly oxidized by ferricyanide. So if you need a gentle method and can prove
that goes to completion, that might work.
|
|
|
nimgoldman
Hazard to Others
Posts: 303
Registered: 11-6-2018
Member Is Offline
|
|
Ferricyanide can be reduced to ferrocyanide with e.g. ascorbic acid [1].
Other suitable reducing agents include e.g. Mg, SO2, H2S, Na2SO3, SnCl2 [2].
Ferrocyanide can be worked up by the addition of soluble calcium and potassium salts (e.g. CaCl2, KCl). If both are present, ferrocyanide will
precipitate as nearly insoluble double salt:
Na4Fe(CN)6 + Ca2Fe(CN)6 + 4 KCl → 2 K2CaFe(CN)6(s) + 4 NaCl
This is then dissolved by heating potassium carbonate solution, leaving behind a precipitate of calcium carbonate:
K2CaFe(CN)6 + K2CO3 → K4Fe(CN)6 + CaCO3(s)
Here is much more info about ferrocyanide/ferricyanide chemistry [3].
It is also possible to just displace the Fe(3+) with Fe(2+) in the complex by adding a solution of some Fe(II) salt (e.g. FeSO4) to the
ferricyanide solution.
This precipitates so called so called Turnbull's Blue which is chemically identical to Prussian Blue, except darker:
3 FeCl2 + FeCl3 + 3 K3[Fe(CN)6] → Fe4[Fe(CN)6]3(s) + 9 KCl
Note that original Prussian Blue is produced by adding Fe(III) solution to ferrocyanide solution.
The pigment dissolves in sodium or potassium hydroxide. Iron(III) precipitates as hydroxide while iron(II) stays in solution as ferrocyanide
[5]:
Fe4[Fe(CN)6]3 + 12NaOH = 3Na4[Fe(CN)6] + 4Fe(OH)3(s)
Analysis
Presence of ferrocyanide in sample can be determined with iron(III). Similarly,presence of ferricyanide can be determined with iron(II). Both
reactions produce deep blue precipitate. Works in neutral/acidic solution to prevent dissolution of the pigment - perhaps use a pH buffer.
Presence of iron(III) can be determined with salicylic acid at pH 2-3, forming a dark violet complex.
Iron(III) also reacts with thiocyanate, forming a blood-red coloured complex.
The exact amounts of iron(II) and iron(III) can be determined with COD assay. A known amount of the sample is completely oxidized in an acidic
solution (use fume hood!) with a known excess amount of oxidizer (e.g. H2O2, KMnO4). The unreacted oxidizer is then
determined by titration against standard ferric ammonium sulfate (FAS) with ferroin (o-phenanthroline) as an indicator.
[1] https://www.researchgate.net/post/Is-there-any-way-and-react...
[2] https://www.aplustopper.com/changing-iron-ii-ions-iron-iii-i...
[3] https://www.911metallurgist.com/blog/ferrocyanides
[4] https://en.wikipedia.org/wiki/Prussian_blue
[5] https://en.crystalls.info/Sodium_hexacyanoferrate(II)
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
