Preparation of Buffer Solutions and
Colorimetric Determination of the pH
pH is defined as the negative of the logarithm of [H3O]+. (Strictly speaking we
should use the activity of H3O+, a dimensionless quantity, but we substitute the numerical
value of the molarity of H3O+ for its activity.)
pH = −log[H3O+]
The term to buffer means to prevent changes or to lessen the shock of changes. In
chemistry, a buffer is a solution that prevents a drastic pH change. In other words, it
tends to maintain a steady acidity level even when a strong acid or strong base is added.
A chemical buffer is prepared by
a) dissolving both a weak acid and its salt in water (e.g. CH3COOH, CH3COONa),
b) dissolving both a weak base and its salt in water (e.g. NH4OH, NH4Cl),
c) dissolving acidic salts of weak polyprotic acid in water (e.g. NaH2PO4, or
Na2HPO4).
The number of moles of strong acid or strong base needed to change the pH of 1
liter of buffer by one unit is called the buffer capacity. The larger the buffer capacity, the
more resistant the buffer is to changes in pH. The capacity of a given buffer is determined
by the concentrations of acid and salt and also by the ratio of acid/salt or base/salt. E.g.
for a given acid/salt ratio, the capacity of a buffer system increases with increasing
concentrations of acid and salt.
E.g. for a buffer of a weak acid (HA) and its salt with strong base the following
equilibrium must be taken into account:
HA ↔ H+ + A−
[HA]≅ cacid; [A−]≅ csalt;
K a=
[ H  ][ A− ]
[ HA ]
[ H  ]=K a
(Ka= acid ionization constant)
c acid
[ HA ]
=K
a
c salt
[ A− ]
The most effective pH range for any buffer (the pH is most resistant to change) is at or
near the pH where the acid and salt concentrations are equal (that is pKa).
Most efficient buffer: pKa= pH
Acceptable buffer: pKa= pH ± 1.
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Useful rules of thumb for selecting buffer mixtures are:
1.
A good buffer mixture should have about equal concentrations of the weak acid
and its salt or of the weak base and its salt.
A buffer solution has generally lost its usefulness when one component of the buffer pair
is less than about 10% of the other.
2.
Weak acids and their salts are better as buffers for pH<7; weak bases and their
salts are better as buffers for pH>7.
3.
The capacity of a buffer system increases with increasing concentrations of acid
and salt, or base and salt.
Acid-Base Indicators
Certain organic substances change colour in dilute solution when the hydrogen
ion concentration reaches a particular value. For example, phenolphthalein is a colourless
substance in any aqueous solution with a hydrogen ion concentration greater than 1.0x108
M (pH less than 8.0). In solutions with a hydrogen ion concentration less than 1.0x10-8
M (pH greater than 8.0), phenolphthalein is red or pink. Substances like phenolphthalein,
which can be used to determine the pH of a solution, are called acid-base indicators.
Acid-base indicators are either weak organic acids, HA, or weak organic bases, BOH,
where the letters A or B stand for complex organic group.
The equilibrium in a solution of the acid-base indicator methyl orange, a weak
acid, can be represented by the equation
HA
↔ H+ + A−
red
yellow
The anion of methyl orange is yellow, and the nonionized form is red. If acid is added to
the solution, the increase in the hydrogen ion concentration shifts the equilibrium toward
the red form in accordance with the law of mass action.
K a=
[ H  ][ A− ]
[ HA ]
The indicator's colour is the visible result of the ratio of the concentrations of the two
species A− and HA. For methyl orange:
K
[ A− ] [ subs tan ce with yellow color ]
=
= a
[ HA ]
[ subs tan ce with red color ]
[H ]
2
When [H+] has the same numerical value as Ka, the ratio of [A−] to [HA] is equal to 1,
meaning that 50% of the indicator is present in the red acid form and 50% in the yellow
ionic form, and the solution appears orange in colour. When the hydrogen ion
concentration increases to a pH of 3.1, about 90% of the indicator is present in the red
form and 10% in the yellow form, and the solution turns red. No change in colour is
visible for any further increase in the hydrogen ion concentration.
Addition of a base to the system reduces the hydrogen ion concentration and shifts the
equilibrium toward the yellow form. At a pH of 4.4 about 90% of the indicator is in the
yellow ionic form, and a further decrease in the hydrogen ion concentration does not
produce a visible colour change. The pH range between 3.1 (red) and 4.4 (yellow) is the
colour-change interval of methyl orange; the pronounced colour change takes place
between these pH values. In general, the colour-change interval of an indicator is the pH
range, where pronounced colour change takes place; the borders of this interval can be
estimated by pKa−1 and pKa+1.
Some acid-base indicators
Indicator
Colour in the more
acid range
Methyl violet
Yellow
Thymol blue
Pinkish red
Benzyl orange
Orange
Bromophenol blue
Yellow
Congo red
Blue
Methyl orange
Red
Methyl red
Red
Bromocresol purple
Yellow
Bromothimol blue
Yellow
Phenol red
Orange
Crezol red
Reddish yellow
Thymol blue
Yellow
Phenolphthalein
Colourless
Tymolphthalein
Colourless
Alizarin yellow R
Yellow
pH range
(colour-change interval)
0-2
1.2 - 2.8
1.9 - 3.3
3.0 - 4.6
3.0 - 5.2
3.1 - 4.4
4.2 - 6.3
5.2 - 6.8
6.0 - 7.6
6.8 - 8.4
7.2 - 8.8
8.0 - 9.6
8.0 - 9.8
9.3 - 10.5
10.0 - 12.1
Colour in the more
basic range
Violet
Yellow
Yellow
Violet
Red
Yellow
Yellow
Purple
Blue
Red
Blue
Blue
Pink
Blue
Red
Colorimetric Determination of the pH
The basis for what the chemist calls colorimetric analysis is the variation in the
intensity of the colour of a solution with changes in concentration (or pH). The colour
may be due to an inherent property of the constituent itself (e.g. MnO4− is purple) or it
may be due to the formation of a coloured compound as the result of the addition of a
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suitable reagent (e.g. indicator). By comparing the intensity of the colour of a solution of
unknown concentration (or pH) with the intensities of solutions of known concentrations
(or pH), the concentration of an unknown solution may be determined.
Follow the following procedure:
1.
Estimate the pH of the unknown solution using an universal indicator paper.
2.
Select an indicator.
3.
Select a buffer system.
4.
Prepare a series of solutions with various pH values (10 ml from each, pH steps in
the series are 0.2 - 0.2)
5.
Add 1 or 2 drops (strictly the same amount) of indicator to each of the solutions
and finally to the unknown solution. Compare the colour of the unknown solution to the
solutions with known pH.
Solutions indicated in the diagram:
„Glikokoll”
7.505 g glycocol and 5.85 g NaCl in 1000 cm3 solution
„Sósav”
0.1 M HCl solution
„Citrát”
21.008 g citric acid monohydrate and 200 cm3 1 M NaOH in 1000 cm3
solution
„NaOH”
0.1 M NaOH solution
„Prim. foszfát”
9.078 g KH2PO4 in 1000 cm3 solution
„Szek. foszfát”
11.876 g Na2HPO4·2H2O in 1000 cm3 solution
„Borát”
12.404 g boric acid and 100 cm3 1 M NaOH in 1000 cm3 soln.
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