Compounds that change color at a certain pH level are called acid-base indicators. Usually, they are weak acids or bases, their color changes in color correspond to the protonation or deprotonation of the indicator itself.
An indicator is a Bronsted-Lowry conjugate acid-base pair in which the acid is a different color than the base. During titration, the endpoint is when the indicator changes color.
In titration, the equivalence point is when the amount of acid and base is just enough to cause complete consumption of the acid and the base.
What happens if you use the wrong indicator?
Using the wrong indicator for any type of titration can result in relatively large errors. As we know indicator works within the specific pH range, if we use an incorrect indicator can result in titrant consumption that is either extremely low or very high, making it impossible to determine the exact molarity or normality of the titrand (Sample solution).
A suitable indicator for a titration will change color over a narrow pH range and have a distinct color at low pH and a different, distinct color at high pH.
For example,
We cannot use phenolphthalein as an indicator for titrating strong acids and weak bases, since phenolphthalein changes color between pH 8.1 - 10, which is slightly below 7. Phenolphthalein is a basic indicator that can only distinguish changes in the range of basic pH.
People also ask:
Does any indicator affect titration?
Why is it important to use the same indicator in a titration?
Would a different indicator be better in this titration Why?
What is indicator error in titration?
What would happen if you forgot to add phenolphthalein to your reaction mixture?
Why doesn't the indicator affect the titration results?
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