# Why is phenolphthalein an appropriate indicator for titration of a strong acid with a strong base?

Please explain based on the pH at the equivalence point and the transition range for phenolphthalein. Thanks for the help.

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If you view the titration curve for $V_{added}$ vs $pH$:

you can see that the equivalence point occurs at $pH$ = 7.

Phenolphthalein is fuchsia from $pH$s roughly between 8.2 and 12, and is colorless below 8.2. When the number of moles of added base is equal to the number of moles of added acid (or vice versa; example on valid for strong monoprotic acids/bases assuming 100% dissolution), the $pH$ is equal to 7.

You might say, if the pH needed is 7, and phenonphthalein changes at $pH$s around 8.2, how can you use this indicator?

Well, again looking at the curve, from $pH$ = 11 to about $pH$ = 4, pH changes very rapidly with from an infinitesimally small change in $V_{added}$. Since one drop of added titrant will cause this large change, even though the change in color of phenolphthalein does not occur right on the equivalence point, it is within approximately one drop. This kind of uncertainty is "acceptable uncertainty" in using titration to volumetrically determine concentrations.

To clarify what I mean by "acceptable uncertainty", you should realise that each of your measurements has some kind of uncertainty to them:

• When you weighed out the primary standard to titrate against, was the balance perfect?

In the scheme of things, $\pm$1-2 drops will not be a significant factor in getting an accurate result, but you should most definitely acknowledge that there is uncertainty in your answer.