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Home
Science
Chemistry
Acids and Bases
pH and pOH
Introduction
In pure water, [H3O+] and [OH–] are equal to each other.
When an acid or base is dissolved in water, [H3O+] and [OH–] are
not
equal to each other.
Kw = [H3O+] [OH–] = 1.0 × 10–14
[H3O+] and [OH–] are inversely related to each other.
Acidic solution
: if [H3O+] > 1.0 × 10–7, then [OH–] < 1.0 × 10–7
At room temperature
Basic solution
: if [OH–] > 1.0 × 10-7, then [H3O+] < 1.0 × 10–7
pH
The acidity and basicity of a solution can be expressed in terms of its [H3O+] or [OH–].
These concentrations can span many orders of magnitude, therefore logarithmic values are commonly used.
pH = -log [H+]
pOH = – log [OH-]
[H+][OH-] = 1.0 x 10-14
Now taking the log of both sides gives us
Log[H+] + Log [OH-] = -14. Multiply by -1, and
-log[H+] - Log [OH-] = 14. But – log{H+] is pH!
So, pH + p OH MUST = 14 (at room temp.)
pH and pOH in a Neutral Solution
Neutral solution
: if [H3O+] = [OH–] = 1.0 × 10–7
At room temperature.
Like ANY equilibrium constant, the Kw is temperature dependent. Since ionization is ENDOTHERMIC, the Kw increases as temperature increases.
If the Kw is NOT 1.0x10-14 , then the pH of a neutral solution will NOT be 7.
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