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I-2
When both CH3COOH and CH3COO
–
are added to pure water the following equilibrium is established.
CH3COOH
(aq)
+ H2O
(l)
1H3O
+
(aq)
+ CH3COO
-
(aq)
(3)
If H3O
+
is added to this solution, the position of the equilibrium shifts to form more reactant,
consuming most of the added H3O
+
. Therefore, no dramatic change in the pH is observed. Similarly,
addition of OH
-
ion to the buffer solution shifts the position of equilibrium to form more product, thus,
consuming most of the added OH
-
. Again, no dramatic change in the pH is observed. This is the
chemical principle that makes buffer solutions effective.
If the initial concentrations of CH3COOH and CH3COO
–
used to prepare the solution are 0.010 M or
greater, their equilibrium concentrations, will not change appreciably. Also, the concentrations of
CH3COOH and CH3COO
–
are large when compared to [H3O
+
]. In fact, [H3O
+
] and the pH are
determined by the concentrations of CH3COOH and CH3COO
–
, through the acid dissociation constant:
K
a
=
[H3O
+
][CH3COO
-
]
[CH3COOH]
(4)
Rearrangement gives:
[H3O
+
] =
K
a
[CH3COOH]
[CH3COO
-
]
(5)
and taking the negative logarithm of both sides gives:
pH = pK
a
- log
[CH3COOH]
[CH3COO
-
]
or
pH = pK
a
+ log
[CH3COO
-
]
[CH3COOH]
(6)
The Henderson-Hasselbalch equation (7) is the general form of this relationship that can be applied
to many weak acid - conjugate base pairs:
pH = pK
a
+ log
[A
-
]
[HA]
(7)
The Henderson-Hasselbalch equation is commonly used to calculate the pH of buffer solutions. All
that is required is the K
a
of the weak acid and the concentrations of both the weak acid and its
conjugate base. By selecting the appropriate weak acid and its conjugate base and manipulating the
concentrations, the pH of a solution can be buffered to almost any desired value.
Generally, buffers will maintain a relatively constant pH as long as both HA and A
-
are present in the
solution in significant concentrations. The capacity of the buffer, therefore, is determined by the
