How does one classify aqueous solutions of ions as acidic, basic and neutral? The following is a homework question:    


  Classify the aqueous solutions of the following ions as acidic, basic or neutral:
          $\ce{C6H5NH3+, Cl- , Fe^3+ , Mg^2+ , HSO4- , NO3- , NH4+ , Al^3+ , C5H5NH+, H2PO4-, CH3CO2-, Cu^2+, Na+, Ag+, ClO4-, PO4^3-, CN-, SH-, Ba^2+}$


I am looking for an approach to solve it. 

My thought process so far:


If any of the radical is highly reactive with $\ce{H+}$, then the solution will be basic as $\ce{OH-}$ is left. Similarly, if it reacts with $\ce{OH-}$, the solution will be acidic.
However, this is where I am confused. Take $\ce{Fe^3+}$ for instance. It is known that it reacts to form $\ce{Fe(OH)3}$, which is a strong base. But, according to my logic above, the solution will be acidic.


Can someone help me with this and possibly give me hints at solving the rest?
Your approach thus far is correct. $\ce{Fe^3+}$ is acidic; $\ce{Fe(OH)3}$ is not a strong base. Indeed, if $\ce{Fe(OH)3}$ were a strong base, it would dissociate and form $\ce{Fe^3+}$, but, as you've mentioned, the opposite process happens.

Your confusion may arise from the fact that $\ce{Fe^3+}$ is a Lewis acid, and produces protons in a "different" way than Arrhenius acids:
\begin{align*} \ce{HCl &amp;&lt;=&gt; H+ + Cl-}\\ \ce{Fe^3+ + H2O &amp;&lt;=&gt; Fe(OH)^2+ + H+ }\end{align*}

Either way, your thought process (1.) will work for both cases.