If you have looked into flow batteries for any length of time, you will have found that the ferrocyanide/ferricyanide redox couple ( [Fe(CN)6] 4−, Fe(CN)6] 3−) is one of the most widely used in the field. This is because this redox couple has very high redox stability, great kinetics, significant solubility (0.7-1.2M depending on the exact salts used) and a redox potential that is lower to that of the Fe2+/Fe3+ redox couple (+0.22V and +0.5V respectively Vs saturated Ag/AgCl), with high stability under high pH conditions.
However you might have also noticed that there are no published examples of flow battery systems where ferrocyanide salts are used in a common electrolyte, symmetric system. That is, a system where the battery starts with the same exact electrolyte on both the catholyte and anolyte and the redox reactions happen from this mixed state. Common examples are ZnBr2, ZnI2, Vanadium and Fe systems using simple FeCl2 or FeSO4 salts.
Why is this the case? An initial reason is that ferrocyanide forms insoluble substances, Prussian blue or its analogues, with most heavy metal cations, so any battery that wants to do a metal reduction in the anolyte, such as the reduction of Zn2+ to Zn metal or the reduction of Fe2+ to Fe metal, would not work because you would precipitate these solids.
While the above reason makes things more difficult, it is solvable. We have known from the late 1940s that solids of this type can be easily dissolved by using pyrophosphates (see here) and other strategies with strong chelating agents also work. This might tempt you to make a symmetric battery with something like potassium pyrophosphate, zinc chloride and ferrocyanide, where you reduce zinc pyrophosphate at the anode to zinc metal and oxidize ferrocyanide to ferricyanide at the cathode. However this is a bad idea.
The reason is that anolytes using ferrocyanide, especially when the potential will subject the ferrocyanide side to low potential values, cause ferrocyanide to decompose at the anode to form some Fe metal but also free cyanide. The cyanide will then be free as Zn is already complexed by pyrophosphate but the Ferrocyanide inventory is hard to recover. We know this happens from studies of ferrocyanide solutions under reducing potentials (see here). There are some evident cathodic peaks here that clearly show ferrocyanide is NOT stable at reducing potentials below around -0.5V. When cycled to negative potentials for 100 times, it is clear that it is decomposing. Any free Fe will quickly form Prussian blue or hydroxides at basic pH which will then start destroying the battery. Add to that the fact that we are generating free cyanide and the inventory becomes much harder to deal with.
The above might make you think we can then go the other way and couple ferrocyanide with a higher redox potential couple instead, like these Fe phenanthroline complexes, use a ferricyanide anolyte instead of a ferrocyanide catholyte to create a flow battery in this way. The problem then is that we also have significant anodic current generated at high potentials that correspond to cyanide oxidation and ferrocyanide destruction as well.
The reality is, ferrocyanide is NOT well suited to work in a symmetric system. While it is very stable when working on its own in an asymmetric system, it is not well behaved when exposed to potentials either +1V or -1V away from its redox potential. This basically precludes the creation of a viable flow battery, as batteries with potentials below 1V are not going to be economically viable, especially given the additional costs involved in creating a symmetric system using ferrocyanide.
When using ferrocyanide, also consider that while this salt is relatively safe in its unaltered state, subjecting it to electrochemical abuse WILL generate free cyanide and it’s likely to pose a significant danger to you and others. For this reason, I would recommend to stay away from testing ferrocyanides in symmetric systems entirely, unless you are a trained professional and professionally well equipped to handle both the potential operational hazards and wastes generated from its decomposition products.