For B.Pharm Sem I

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Theories of acid base
indicators
Sapan K. Shah
Assistant Professor
Priyadarshini J. L. College of Pharmacy,
Nagpur-440016

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Theories of neutralization (Acid base)
Indicators
1. Ostwald Theory
2. Resonance Theory (Quinonoid Theory)

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1) Ostwald Theory
• Indicators are weak acids or weak bases which have different
colors in their conjugate base and acid forms (two color
indicators);
• Others are one color indicators, and have one form colored with a
colorless conjugate form.
• Most indicators in common use are intensely colored, and can be
used in dilute solution in such small quantities that the acid-base
equilibrium which is under examination is not disturbed by the
addition of the indicator.

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• As weak acids or weak bases, they are able to
reach instantaneous equilibrium with the
system, and
• the color of the solution will range between the
extreme colors of the two forms as the
proportion of acidic and basic forms
automatically adjusts itself to the pH of the
solution.

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• The following equilibrium will apply for an
indicator functioning as weak acid:
HIn H+ + In–
• In acid solution, the excess of H+ ions will depress
the ionization of the indicator. The concentration
of In– will be small, and of HIn large, and the color
will be that of the unionized form.
• Alkali will promote removal of hydrogen ions from
the system with an increase in the concentration of
the ionized form (In–) so that the solution acquires
the ionized color.

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• Similarly for an indicator functioning as a
weak base, the following equilibrium will
apply:

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2) Resonance (Quinonoid) Theory
• Although the behavior of indicators can be
explained in terms of ionization of weak acids
and bases, as above, the equilibrium is actually
more complex,
• the color changes being brought about by
tautomeric changes in the structure of the
molecule.

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This is illustrated by the behavior of
phenolphthalein in solution-
The red color in alkaline solution is due to
the quinonoid structure, with the resulting
increased possibilities for resonance between
the various ionic forms.

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pH range of indicators(color change interval):
• The observed color of a two color indicator is
determined by the ratio of the concentrations of
ionized and unionized forms.
• Observable color changes are, however, limited
by the ability of the human eye to detect
changes of color in mixtures.
• This is particularly difficult where one color
predominates and in practice is almost impossible
when the ration of the two forms exceeds 10 to
1.

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• Thus the limit of visible color change will be represented
by the introduction of the term log10 for log[In-]/[HIn] in
the above expression so that
• The average color change interval of an indicator is,
therefore, about two pH units.
• The observed color changes within the indicator range
are seen as gradual change of tint or shade which ranges
from one extreme color to the other.
• The shade of color is independent of the amount of
indicator present, but the use of too much indicator
should be avoided as slight changes are then more
difficult to detect.

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• With a single color indicator, such as
phenolphthalein, the intensity of color is
important and not shades difference.
• The actual concentration of indicator is therefore
significant, and should be carefully controlled.
• Since the useful range of an indicator only
extends over approximately two pH units, it is
essential to have a series of indicators available
to cover the complete pH scale.

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A list of such indicators in common
use, together with their color
changes is given in table.
Indicator pH range Color changes
Phenolphthalein 8.3-10 colorless pink Red
Methyl orange 2.9-4.6 Red Orange yellow
Methyl red 4.2-6.3 Red Orange Yellow
Phenol red 6.8-8.4 Yellow Orange Red
Thymolphthalein 9.3-10.5 Colorless Pink Red

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