pH and materials selction
The pH scale is merely a way of keeping track of the concentration a species (the hydrogen ion $\ce{H^+}$, $pH = -log\  a_{H^+} \simeq -log\  [H^+]$) in a medium (usually liquid, not necessarily aqueous). You can define arbitrary scales for other species, too. Analytical chemists are more familiar with scales such as pCl, with $pCl = -log\  a_{Cl^-} \simeq -log\  [Cl^-]$, and that's fine, but they're much rarer. The reason chemistry pays so much attention to hydrogen ions specifically is because the pH scale keeps track of what is arguably the most reactive species in all of chemistry (one could make a case for alpha particles, i.e. bare helium-4 nuclei $\ce{^4_2 He^{2+}}$ instead, but hydrogen ions are nevertheless comparable in strength and far more common). 

If you stop to think about it, a (non-solvated) hydrogen ion, $\ce{H^+_{(g)}}$ is a bare proton. Let me explain why this is an important realization. In chemistry, we talk a lot about how electrons get shoved around when different molecules meet (also known as chemical reactions), and the exact way in the electrons shuffle about is dependent on complicated electromagnetism at the quantum level. However, to a simple but enlightening approximation, one can imagine how substances behave based on how much electrical charge they concentrate in a region; simply, concentrated charges will display stronger electromagnetic effects.

In most of chemistry, we deal with charges of order unity relative to the fundamental electrical charge ($e=1.6022 \times 10^{-19}\ C$) spread in the volume of a few atoms (of order $10^{-3}\  nm^3$). The resulting charge density can be found by dividing the charge by the volume. Now here's the thing - the nucleus of an atom is tiny compared to the size of the entire atom (with its electrons). A lone hydrogen atom has a radius of approximately $0.1\ nm$, but a lone hydrogen nucleus has a radius of about $1\ fm = 0.000001\ nm$, a difference of five orders of magnitude. The volume difference is thus fifteen orders of magnitude. Therefore, the bare nucleus which comprises a $\ce{H^+}$ ion has a charge density 15 orders of magnitude greater than a similarly-charged atomic ion, for example $\ce{Li^+}$. In effect, this means that a bare $\ce{H^+}$ has such a huge positive charge density that it will tear opposite charges (electrons) from absolutely anything. Helium and neon are notoriously unreactive, but both react quite happily (and rather exothermically) with gaseous $\ce{H^+}$ forming hydrohelium(+) and hydroneon(+) as the respective products.

So why doesn't water (or any ready source of $\ce{H^+}$ ions, i.e. protic substance) immediately dissolve absolutely everything it touches? It so happens that, in condensed phases, whatever molecules close to a hydrogen ion are so attracted that they effectively encase the hydrogen ions into solvated hydrogen ions, forming a shell which protects everything else from the aggressive reactivity of the hydrogen ion. For example, in water, $\ce{H^+}$ is actually not a very correct description of reality. Rather, one finds $\ce{H(H_2O)_{n}^+}$ (often alluded to by formulae such as $\ce{H3O^+}$, $\ce{H5O2^+}$, $\ce{H9O4^+}$ and others), or more concisely (but less distinguishably) $\ce{H^+_{(aq)}}$. After being surrounded by water molecules, the hydrogen ion loses almost all of its reactivity, as the original charge is now effectively dispersed over several molecules, greatly dimishing the charge density of the solvated hydrogen ion compared to the non-solvated hydrogen ion. Even after the massive handicap, solvated hydrogen ions are still quite reactive. Thus, it is important to keep track of their amount, to figure out if certain materials are capable of withstanding the onslaught.

After all I said, it seems like it would be a good idea to minimize the amount of $\ce{H^+}$ in water, which can be done by going to high values in the pH scale. However, by doing so you begin to significantly increase the amount of hydroxide ions $\ce{OH^-}$ present, in order to obey the autodissociation reaction equilibrium of water. In some sense, hydrated $\ce{OH^-}$ ions in solution are about as reactive as solvated $\ce{H^+}$ (even though non-solvated $\ce{H^+}$ is much, much more reactive than non-solvated $\ce{OH^-}$). Thus, going to either end of the pH scale can be a bad idea. One could say that water solutions are least reactive* when their pH is close to 7 (including pure water), as one minimizes the sum of $\ce{H^+}$ and $\ce{OH^-}$ present.

* Least reactive in the specific sense that if a material resists acidic aqueous solutions or basic aqueous solutions, then it must resist pure water and neutral aqueous solutions, assuming that whatever else anything else in the water incapable of altering the pH value is also unreactive; a pH 7 sodium chloride solution for example can be quite corrosive to metals, but not because of the water itself.