Thanks for this interesting question, a question that is critical in
understanding acid-base titrations using visual indicators in
Essentially because an indicator is an acid/base substance itself a
small quantity of acid (or base depending on which is the titrant)
must be added past the equivalence point in order to cause the
indicator to change colour (assuming the appropriate indicator has
been chosen for the type of acid-base titration being performed)
The following quote summarizes the intent of a titration -
"A titration is the controlled addition and measurement of the amount
of a solution of known concentration that is required to react
completely with a measured amount of a solution of unknown
concentration. In a neutralization reaction, the point at which there
are equivalent quantities of hydronium and hydroxide ions is called
the equivalence point. The point in a titration where an indicator
changes color is called the end point of the indicator."
"Determining the Equivalence Point during an Aicd Base Titration"
Itasca Learning Center
In the titration of a strong acid and a strong base the equivalence
point (when stoichiometrically equivalent amounts of the reactants are
present) occurs at about a pH of 7. To indicate when this occurs an
indicator that changes colour at about a pH of 7 such as Bromothymol
Blue should be used. (Indicators change colour over a range of about
1.5 pH units)
If a weak acid and strong base titration is being carried out the pH
at equivalence point will be higher than 7 (due to the stronger base
component) and an indicator like Phenolphthalein should be used.
Conversely a strong acid - weak base titration will reach equivalence
at a pH below 7 and an indicator such as Methyl Orange should be used.
For example HCl(aq) + NaOH(aq) ->NaCl(aq) + H2O
Equivalence point is reached when exactly equimolar quantities of HCl
and NaOH are present in the reaction flask. But at that point an
indicator such as Bromothymol Blue is still mainly in its original
form and it takes a very small amount of the titrant (a slight excess)
to change the indicator to its conjugate and thus cause the colour
For example (representing the acidic form of Bromothymol Blue as BBH+
and the basic form as BB)
BB + HCl -> BBH+ + Cl- Colour change (assuming HCl is the titrant)
Usually only a drop or two of indicator is used in acid-base
titrations. The more indicator added the more excess titrant is
reqwuired to initiate the colour change and thus the further away the
end point (when the indicator changes colour) will be from the
Some links worth looking at for further information and clarification
of any of the terms I have used in this answer.
A very informative Think Quest on the topic
A lecture titled "Acid - Base Titration" with clear instructive
An "Acid - Base Titration" page that contains some basic theory along
with practical prodecures.
"An unknown acid" (Practical sheet with good background theory on
selecting the right indicator
Search strategy using Google key words titration, equivalence point,
end point, acid, base, indicator.
Hope this answers your question clearly enough
All the best with your Chemistry!