Why is phenolphthalein an appropriate indicator for titration of a strong acid with a strong base?
If you view the titration curve for $V_
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_
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?
Was your solution made up precisely to the graduation in the volumetric flask?
Did you pipette the exact volume of the aliquot or were you off by a drop or two?
Are you able to state the volume added from the burette to arbitrary precision or is there some uncertainty beyond the two decimal places given by the graduated lines?