|The pH Scale|
|Every aqueous solution is either acidic, basic or neutral.
is a quantitative relationship between the concentration of hydronium
hydroxide ions in the solution.
Neutral solution [H3O+] = [OH-]
Acid solution [H3O+] > [OH-]
Basic solution [H3O+] < [OH-]
|The brackets as usual denote molar concentrations.
|The pH scale is a numerical scale which, for most
from 0 through to 14. The numbers on the scale represent the relative
of solutions and can be converted into actual hydronium ion
|The pH scale is based on the self-ionization of pure water.
|Two water molecules will sometimes combine into hydronium
H2O + H2O <------> H3O+
|Pure water is considered to neutral and the hydronium ion
is 1.0 x 10-7 mol/L which is equal to the hydroxide ion
[H3O+] = [OH-] = 1.0 x 10-7
|The equilibrium law for this reaction should be
|Keq = [H3O+][OH-]
(1.0 x 10-7)(1.0 x 10-7) = 1.0 x 10-14
|You will please note that at neutrality the molarity of
ion is 10-7. The 7 plays a part in the pH scale by
neutrality. The scale reaches a maximum at 14. Please note again that
hydronium and hydroxide concentrations multiply out to 10-14
M. The pH scale was derived around this relationship:
|ie. [H3O+] =
|So the pH is the -log of the [hydronium ion].
|What is the pH of an HCl solution which has a [H3O+]
= 1.0 x 10-3?
pH = -log[H3O+] = -log[1.0 x 10-3] =
|What is the pH of an acetic acid solution whose [H3O+]=2.5
|What is the hydronium concentration of nitric acid if the
= 10-pH = 10-(4) = 1.0 x 10-4 mol/L
|What is the [H3O+] of HCl if the pH
|What is the pH of 0.010 mol/L hydrochloric acid?
|The pH of a solution may be determined by the use of an
instrument known as a pH meter, or through the use of chemical
Acid-base indicators are dyes which undergo slight changes in molecular
structure and colour when the pH value of the solution changes.
|Specific colours correspond to specific pH values. Some examples are: litmus, phenolphthalein, bromothymol blue, etc. There is a list of acid-base indicators in the databook.|