Which reactions occur during the oxidation and dissolution of bromide/bromine? I'm trying to understand the fate of bromide anions during and after they undergo anodic oxidation. I have a solution of 94% DMF/6% water (by volume) with 0.1 M tetrabutylammonium  bromide (TBAB) as the supporting electrolyte. I've used this electrolyte to reduce a cathode at -2.5 V or less v. the Ag/AgCl reference electrode. During the reaction, the solution near the platinum counter electrode (the anode) turned yellow. I would assume this is due to the reaction: 
$$\ce{2Br–(aq)-&gt; Br2(l) + 2e– }$$

The next day, I noticed that the yellow liquid in the solution had completely disappeared. I'm assuming that the bromine had dissolved into the solution into $\ce{Br-}$ ions again. However, I'm not 100% sure of the chemical reaction associated with this. Also, would there be any side products associated with the dissolution? Or, do we end up with the species we initially started with in terms of bromine? 

I ask because I would like to be able to reuse the electrolyte multiple times and so want to figure out exactly how the electrolyte changes after the reaction. At the cathode, there will be reactions such as the hydrogen evolution reaction which will make the solution more basic. However, I'm not sure what happens with the bromine.
A possibilty is that elementary bromine formed -- causing the the yellow color -- then solubilized in water yielding bromine water.



(source)

It does not dissolve well (2.8% ref -- but do not forget the high density (about 3.10 g/mL) of liquid bromine here). It then may disproportionate into then colorless hypobromit and bromide

$\ce{Br2 + H2O -&gt; Br- + BrO- + 2H+}$ (ref.)

especially under light, forming an acidic solution.  The hypobromit may act as an oxidizer.

To draw an analogy, chlorine gas does dissolve into water, too (ref, ref).

Proceed with caution since bromine is corrosive, though.