# Water self ionization and PH

The pH scale

We have just seen how the acids and bases do to change the mutual concentrations of H 3 O + and OH - in solution and how they are placed in relation to the balance autoprotolysis.

In fact, already in this way we would be able to figure out whether a solution is acidic or basic obtaining [H 3 O +] and [OH -].

For reasons of practicality, however, it is chosen to define the acidity / basicity of a solution via the pH.

Definition of pH

The pH is no more than $pH&space;=&space;-log[H^+]$

Or the base-10 logarithm of the hydrogen ion concentration. Taking the same solution that we used as example above (a 0.1M HCl solution):

$pH&space;=&space;-log[0,1]&space;=&space;1$

You can also define the pOH, ie -log [OH -].

We had found that our solution 0.1 M of HCl had [OH -] equal to 1 · 10 -13.

$pOH&space;=&space;-log[1\cdot&space;10^{-13}]&space;=&space;13$

As well as [OH -] and [H 3 O +] are related, so are the pH and the pOH. The report is very interesting:

$K_w&space;=&space;[H^+][OH^-]$

$logK_w&space;=&space;log[H^+]&space;+&space;log[OH^-]$

$-logK_w&space;=&space;pH&space;+&space;pOH$

$pH&space;+&space;pOH&space;=&space;14$ (at 25 ° C)

For example, once you find pH = 1 for the previous year, the pOH would be simply given by:

$14&space;-&space;pH&space;=&space;pOH$ and then $pOH&space;=&space;14-1&space;=&space;13$

The pH scale is used because we almost always have to deal with very dilute solutions, in which the hydrogen ion concentration is very low, so that having to express as a negative power, of the type 10 -x, between 0 and 1. For this Therefore, the logarithmic scale is much more extensive and convenient to use than the linear one of the concentrations.

Here's the proof:

we took as example, three measures of concentration: 10 -7 M, 10 -3 M and 10 -2 M. The linear scale reading is almost impossible, because you have a large clump of all values ​​near zero. The same values ​​in the pH instead are three discrete points and easily distinguishable.

Neutral solutions, acids and bases

We have previously defined a neutral solution when:

$[H_3O^+]&space;=&space;[OH^-]&space;=&space;1\cdot10^{-7}&space;M$

• In light of what we learned on the pH we can say that a solution will be neutral when:

$pH&space;=&space;pOH$

then, given that $pH&space;+&space;pOH&space;=&space;14$

you will have the neutrality for $pH&space;=&space;7$

• acid for pH between 0 and 7 (7 excluded)
• basic to pH comprised between 7 and 14 (7 excluded)

Then the pH scale, for common use solutions, ranging from 0 to 14.

In fact, for very concentrated solutions of base or acid, it can be obtained respectively the pH values ​​greater than 14 and less than 0.

This comes from the pH math: if [H 3 O +] is greater than 1, the -log [H 3 O +] becomes negative. Similarly, if the concentration of [OH -] is greater than 1 it is the pOH to become negative and consequently the pH greater than 14.