I used methyl orange as my indicator to titrate a weak acid:

At first, I added 3 drops of methyl orange to a solution of the weak acid, $\ce{CH3COOH}$, and the solution turned red. Then, after adding a few drops of the strong base $\ce{NaOH}$ solution, the solution became yellow. Why are only a few drops of $\ce{NaOH}$ needed to make the solution yellow?

In addition, I then added a few drops of the strong acid $\ce{HCl}$, and the solution became red again. Why are only a few drops needed here as well?

  • $\begingroup$ If you add an equimolar amount of NaOH and HCl, you have effectively dissolved a bit of table salt in your solution, and the pH is back where it was. Check the concentrations of your weak acid and your HCl and NaOH solution, and do the math. $\endgroup$
    – Karl
    May 14, 2019 at 22:12

1 Answer 1


Using methyl orange for titration of weak acids is wrong idea.

The indicator has range $\mathrm{pH}=3.1~\mathrm{(red)} - 4.4~\mathrm{(yellow)}$.

If acetic acid with $\mathrm{p}K_\mathrm{a}=4.75$ is to be determined, the indicator starts to turn toward yellow, when majority of acid is not titrated yet.

The concentration ratio $\frac {\ce{[CH3COONa]}}{\ce{[CH3COOH]}}$ is $0.022$, resp. $0.45$ for full red resp. full yellow. $99.9\%$ of the acid is titrated at $\mathrm{pH}=7.75$.

Therefore, indicator shows a change after a small addition of $\ce{NaOH}$. A small addition of $\ce{HCl}$ returns things back.

Methyl orange is suitable for titration of strong or weak bases by strong acids, turning yellow to red.

Much more suitable and frequently used indicator for weak acid titration is phenolphthalein with range $\mathrm{pH}=8.2~\mathrm{(clear)} - 10.0~\mathrm{(violet)}$

Aside of suitable indicator $\mathrm{pH}$ range, the transition from weak color(clear, yellow) to intensive colour (red, violet) is easier to catch than vice versa.


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