Acid-Base Calculations
Acid-Base Calculations
The Ion-Product Constant for Water, Kw Water undergoes ionization to a small extent:
H20(l) H+(aq) + OH–(aq)
The equilibrium constant for the reaction is the ion-product constant for water Kw:
(1)
This is a key equation in acid-base chemistry. Note that the product of [H+] and [OH–] is a constant at a given temperature (Eq(1) value is for 25oC). Thus as the hydrogen ion concentration of a solution increases, the hydroxide ion concentration decreases (and vice versa).
The pH scale is widely used to report the molar concentration of hydrogen ion H+(aq) in aqueous solution. The pH of a solution is defined as
(2)
Similarly, pOH and pKw are defined as
(3) (4)
If you take the log10 of both sides of Eq(1), multiply the resulting equation by (-1), and use the definitions of pH, pOH and pKw above, the result is the very useful equation
pH + pOH = pKw = 14.00 (5)
Equations (2) and (3) above may be solved for [H+] and [OH–] respectively to give
(6) (7)
(Here we use the well known rule that if , then .) In practice, the pH scale is only used when [H+(aq)] is less than 1.0 M.
Acidic, basic, and neutral solutions can be distinguished as shown below:
Type of Solution pH [H+]
Color of litmus
Acidic
< 7.00
>
pink
Neutral
= 7.00
=
in between
Basic
> 7.00
<
blue
pH and [H+] Calculations for Strong Acids and Bases
By definition, strong acids and bases are 100% ionized in water solution. Ionization of a strong acid gives rise to H+ ions, and ionization of a strong base produces OH– ions. The equilibrium constant for a strong acid or strong base is undefined, since the reaction the ionization is complete. There is no equilibrium!
In nearly all cases of practical interest the [H+] for a strong acid (or the [OH–] for a strong base) is determined completely by the stoichiometry of the reaction. Once the [OH–] or pOH is known for a base, the [H+] or the pH
The Ion-Product Constant for Water, Kw Water undergoes ionization to a small extent:
H20(l) H+(aq) + OH–(aq)
The equilibrium constant for the reaction is the ion-product constant for water Kw:
(1)
This is a key equation in acid-base chemistry. Note that the product of [H+] and [OH–] is a constant at a given temperature (Eq(1) value is for 25oC). Thus as the hydrogen ion concentration of a solution increases, the hydroxide ion concentration decreases (and vice versa).
The pH scale is widely used to report the molar concentration of hydrogen ion H+(aq) in aqueous solution. The pH of a solution is defined as
(2)
Similarly, pOH and pKw are defined as
(3) (4)
If you take the log10 of both sides of Eq(1), multiply the resulting equation by (-1), and use the definitions of pH, pOH and pKw above, the result is the very useful equation
pH + pOH = pKw = 14.00 (5)
Equations (2) and (3) above may be solved for [H+] and [OH–] respectively to give
(6) (7)
(Here we use the well known rule that if , then .) In practice, the pH scale is only used when [H+(aq)] is less than 1.0 M.
Acidic, basic, and neutral solutions can be distinguished as shown below:
Type of Solution pH [H+]
Color of litmus
Acidic
< 7.00
>
pink
Neutral
= 7.00
=
in between
Basic
> 7.00
<
blue
pH and [H+] Calculations for Strong Acids and Bases
By definition, strong acids and bases are 100% ionized in water solution. Ionization of a strong acid gives rise to H+ ions, and ionization of a strong base produces OH– ions. The equilibrium constant for a strong acid or strong base is undefined, since the reaction the ionization is complete. There is no equilibrium!
In nearly all cases of practical interest the [H+] for a strong acid (or the [OH–] for a strong base) is determined completely by the stoichiometry of the reaction. Once the [OH–] or pOH is known for a base, the [H+] or the pH