BUFFERS
Definition: A buffer solution is one in which there is
very little change in its pH when a relatively
large amount of strong acid or strong bases
is added to it.
How to make a buffer
A buffer contains a mixture of a weak acid and its
conjugate base OR a weak base and its conjugate acid.
e.g. (1) weak acid / conjugate base
CH3COOH / CH3COOIf strong acid is added to the buffer solution, excess
H3O+ ions are removed from the solution by the
conjugate base:
CH3COO- + H3O+  CH3COOH + H2O
If strong base is added to the buffer solution, excess
OH- ions are removed from the solution by the weak acid:
CH3COOH + OH-  CH3COO- + H2O
e.g. (2) weak base / conjugate acid
NH3 / NH4+
If we add acid: NH3 + H3O+  NH4+ + H2O
If we add base: NH4+ + OH-  NH3 + H2O
Buffer solutions work by changing OH- and H3O+ into
H2O and other ions/molecules that already form part of
the buffer.
Uses
 Human blood must maintain pH within the range
7.35 – 7.45; buffer H2CO3/HCO3 Maintain pH of soil PO43-/HPO4-
Determining the pH of a buffer solution
If the buffer solution contains HA (weak acid) and A(conjugate base), the following equilibrium is established:
HA + H2O
A- + H3O+
Ka = [A-] x [H3O+]
[HA]
taking the –log of the whole expression, we get:
pKa = pH - log [A-]
[HA]
Rearranging gives, the Henderson-Hasselbach equation:
pH = pKa + log [A-]
[HA]
Example
Calculate the pH of a buffer solution containing
0.025molL-1 ethanoic acid and 0.010molL-1 sodium
ethanoate. pKa (ethanoic acid) = 4.75
Note: pH < pKa because the buffer contains a higher
concentration of acid than base, so pH is more acidic
(lower) than the pKa.
Buffer Examples
1. Calculate the pH of a 0.1molL-1 nitrous acid solution
(HNO2) which has been combined with 0.1molL-1 sodium
nitrite (NaNO2).
Ka (HNO2) = 5 x 10-4
2. Calculate the pH of a solution that contains equal
volumes of 0.1molL-1 CH3COOH and 0.05molL-1
NaCH3COO.
Ka(ethanoic acid) = 1.74 x 10-5
3. Continuing Chemistry p 133, Q 1, 2, 3, 6, 7
The Henderson-Hasselbach equation can also be used to
prepare a buffer solution of known pH, because if we
make [conjugate base] = [weak acid] then log 1 = 0 and
pH = pKa of the acid used.
By carefully selecting an acid that has a pH close to that
required, this is a simple process.
Note: buffers are usually effective at controlling
pH ± 1 unit.
Example
1. Given samples of HX and NaX explain how you would
prepare a buffer solution of Ph=4 given Ka(HX) = 10-4
2. Calculate the ratio of ethanoate ion to ethanoic acid
in a buffer solution of pH 5.25.
pKa(ethanoic acid) = 4.75
Note: Adding water to a buffer solution does not change
the pH since the concentration of the weak acid and
weak base are both changed by the same factor.
TITRATION CURVES
If you plot pH vs volume of base (or acid) added during a
titration, a pH curve is produced.
e.g. strong acid – strong base
NaOH added to HCl
25mL aliquots of 0.1molL-1 HCl, were titrated with
0.1molL-1 NaOH
Volume of NaOH added/mL
There are characteristic points on the curve that can be
identified:
 Initial pH
 Equivalence point – pH and volume
 Final pH
Remember:
End point – is the point at which the indicator changes
colour
Equivalence point – the point in the titration where there
are equivalent proportions of each reactant
Usually (if you have carefully selected your indicator)
the end point = equivalence point.
There are characteristic shapes to the following
titration curves:
a) Strong acid – strong base
b)Weak base – strong acid
c) Weak acid – strong base
(Weak acid – weak base – doesn’t have a typical curve
because it is very difficult to detect the equivalence
point)
a) strong acid – strong base
- pH range at equivalence = 3 - 11
- pH of salt produced = 7
b) weak base – strong acid
- pH range at equivalence = 7 – 3
- salt produced is acidic
e.g. NH3 + HCl  NH4+ + ClSalt produced is acidic since NH4+ + H2O NH3 + H3O+
Volume of HCl added/mL
c) weak acid – strong base
- pH range at equivalence = 11 - 7
- salt produced is basic
e.g. HCOOH + NaOH  HCOO- + Na+ + H2O
Salt produced is basic since:
HCOO- + H2O  HCOOH + OH-
Volume of NaOH added/mL
Titration curves have many uses, including how to select
a suitable indicator for titration.
Selecting indicators for acid-base titrations
An effective indicator for a titration MUST change
colour over the pH range covered by the flat vertical
part of the titration curve.
Strong acid-strong base
Strong acid – weak base
Strong base – weak acid
pH range 3 - 11
pH range 3 – 7
pH range 7 - 11
Indicators change colour over a small pH range of
approximately 2 pH units.
e.g. methyl orange changes colour over the pH range 3-5
phenolphthalein over 8 – 10
Indicators are all weak acids themselves and are usually
extracted from plants.
(pH range of indicator = pKa of indicator ± 1)
Indicator
Phenolphthalein
Bromothymol
blue
Methyl red
Methyl orange
Methyl violet
HA colour pKa
Colourless 9.4
Yellow
7.0
A- colour pH range
Pink
8.6 – 10.2
Blue
6.2 – 7.8
Red
Red
Yellow
Yellow
Yellow
Blue
5.1
3.7
1.1
4.3 – 5.9
2.9 – 4.5
0.3 – 1.9
Looking back at the titration curves – appropriate
indicators can be chosen 
USING TITRATION CURVES
A titration curve can be used to obtain information:
1. Starting pH
2. Determine the pKa of the weak acid
3. pH of the salt formed
4. Identify the buffer region
5. Equivalence volume
6. Final pH