In the standard brown ring test for the nitrate ion, the brown ring complex is: $$\ce{[Fe(H2O)5(NO)]^{2+}}$$
In this compound, the nitrosyl ligand is positively charged, and iron is in a $+1$ oxidation state.
Now, iron has stable oxidation states +2 and +3. Nitrosyl, as a ligand, comes in many flavours, of which a negatively charged nitrosyl is one.
I see no reason why the iron doesn't spontaneously oxidise to +3 and reduce the $\ce{NO}$ to −1 to gain stability. But I don't know how to analyse this situation anyway. I think that there may be some nifty backbonding increasing the stability, but I'm not sure.
So, why is iron in +1 here when we can have a seemingly stable situation with iron in +3?