This article is about volumetric titration. For other uses, see
Titration (disambiguation).
In
medicine, titration is the process of gradually adjusting the dose of a medication until the desired effect is achieved.
Titration setup: the titrant drops from the
burette into the
analyte solution in the
flask. An indicator present then changes color permanently at the
endpoint.
Titration is a common laboratory method of
quantitative chemical analysis that is used to determine the unknown
concentration of a known
reactant. Because volume measurements play a key role in titration, it is also known as volumetric analysis. A
reagent, called the titrant or titrator
[1], of a known concentration (a
standard solution) and
volume is used to react with a solution of the
analyte or titrand
[2], whose concentration is not known. Using a calibrated
burette to add the titrant, it is possible to determine the exact amount that has been consumed when the endpoint is reached. The endpoint is the point at which the titration is complete, as determined by an indicator (see below). This is ideally the same volume as the
equivalence point - the volume of added titrant at which the number of
moles of titrant is equal to the number of moles of analyte, or some multiple thereof (as in
polyprotic acids). In the classic strong acid-strong base titration, the endpoint of a titration is the point at which the pH of the reactant is just about equal to 7, and often when the solution permanently changes color due to an
indicator.
There are however many different types of titrations (see below).
Many methods can be used to indicate the endpoint of a reaction; titrations often use
visual indicators (the reactant mixture changes colour). In simple
acid-base titrations a pH indicator may be used, such as
phenolphthalein, which becomes pink when a certain pH (about 8.2) is reached or exceeded. Another example is
methyl orange, which is red in acids and yellow in alkali solutions.
Not every titration requires an indicator. In some cases, either the reactants or the products are strongly coloured and can serve as the "indicator". For example, an
oxidation-reduction titration using
potassium permanganate (pink/purple) as the titrant does not require an indicator. When the titrant is reduced, it turns colourless. After the equivalence point, there is excess titrant present. The equivalence point is identified from the first faint pink color that persists in the solution being titrated.
Due to the logarithmic nature of the pH curve, the transitions are, in general, extremely sharp; and, thus, a single drop of titrant just before the endpoint can change the pH significantly — leading to an immediate colour change in the indicator. There is a slight difference between the change in indicator color and the actual equivalence point of the titration. This error is referred to as an indicator error, and it is indeterminate.
