Why does the Anode and not the Cathode corode during eletrolysis of water into $H_2$ and $O_2$? After experimenting with the production of Hydrogen from salt water, I kept on noticing that the Anode (-) gets coroded and not the Cathode (+). Is it perhaps that water in itself is a naturaly positively charged medium and that the Anode gets striped of its own natural electrons to feed back to the power source and keep the circuit flowing?
 A: Metals tend to be oxidized, not reduced. Oxidation happens at the anode (positive charged one during electrolysis, it wants electrons), thus the cathode is protected look into cathodic protection.
The long answer:
Metals tend to become positive ions thus they are easily oxidized not reduced. Oxidation happens at the anode, reduction at the cathode. During electrolysis the anode is connected to the positive terminal thus it wants electrons, though a lot of the electrons will come from the electrolyte Cl$^-$ and H$_2$O some electrons are also stripped from the metal electrode causing it to rust/corrode. We know that thermodynamically a reaction is spontaneous if its Gibbs free energy is negative. In electrochemistry the Gibbs free energy is given by the Nernst equation
$$
\Delta G=-nfE
 $$
so in electrolysis we need to supply a voltage (E) just as atleast as high as those calculated from the Nernst equation theoretically to cause electrolysis/decomposition. We can tell which reactions are occurring using the standard reduction potential table and comparing them to the supplied voltage. Though we must consider kinetics as well not just thermodynamics for the real solution.
