A misunderstanding about the energy profile of reactions with a catalyst involved
The upper graph is just the most simple way to visualize the effect of a catalyst on a reaction $\ce{S -&gt; P}$:

The activation energy is lowered.

The activation energy for the reaction in that direction is the difference of the energies of the starting material $S$ and a transition state $TS^\#$. 

Since it is the same starting marterial in the presence or absence of the catalyst, the energy of the transition state is different.

Can the same transition state have two different energies - just through the remote distance magical action of a catalyst located somewhere? Probably not!

It is much more plausible that - in the absence and presence of a catalyst - two completely different (different in structure and different in energy) transition states are involved.

Exactly this situation is described in the second graph!

The catalyst "reacts" with the starting material, either by forming a covalent bond, by hydrogen bonding, etc. and thus opens a different reaction channel with different intermediates and a transition state not found in in the non-catalyzed reaction.

In the end, the same product $P$ is formed and the catalyst is regenerated, but this doesn't mean that the catalyst wasn't heavily involved in the reaction going on.