The pH Scale
We have discussed that a substance is considered an acid when it dissolves in water to release H+ ions; whereas a base dissolves in water to release 0H- ions.
This is only partially true.
In truth, water or H20, always dissociates to create the presence of both types of ions even if the concentrations of these ions may be extemely low and the life time of the individual ion in its free state, extremely short.
In chemistry we express concentrations is by the use of square brackets. [H+] represents the concentration of H+ ions and [OH-] represents the concentration of OH- ions.
For acid solutions [H+]>[OH-]; for basic solutions [H+]<[OH-] and for neutral solutions or for pure water [H+]=[OH-]
In fact using this notation we can express a relationship between the relative concentrations of the two types of ions in a way that is ALWAYS true, whenever we are considering aqueous solution of any kind.
[H+ ][OH− ] = Kw = 10−14
Therefore in the case of a neutral solution where [H+]=[OH-]:
[H+ ][OH− ] = 10−710−7 so that [H+ ]= 10−7
In the case of concentrated Sulphuric Acid [H+] could equal 10−1 and in the case of concentrated NaOH solution, the [OH-] might be 10−1 thus forcing [H+] to be 10−13.
So whether the OH- ion dominates, or the H+ ion dominates, the value of [H+] always tells us what is the acid/base state of the solution.
In order to express this value in a more convenient form we normally state only the exponent or LOG value of [H+] as a positive quanitity. We call this quantity the pH of the solution. The pH will alway indicate the relative acidity, basisity or neutrality of a solution.
The illustration above shows the pH values for some familiar solutions.