[wilson]

Hi,
I used phenolphthalein in the lab for acid-base titration before and I realized that after the end point was reached (persistent faint-pink colour), the faint-pink colour faded after a while, i.e. the solution became colourless again. Thus, it is not really persistent, so to speak.

What is the reason for quirky phenolphthalein to exhibit such behaviour?
Perhaps it could have been due to carbon dioxide in the air reacting with my base like NaOH... But are there more plausible reasons relevant to the behaviour of phenolphthalein?

[nj_bartel]

how long did it take for the color reversion?

[LQ43]

What is the reason for quirky phenolphthalein to exhibit such behaviour?
Perhaps it could have been due to carbon dioxide in the air reacting with my base like NaOH

That is good thinking. What do you get with CO₂ dissolving in an aqueous solution?

Or it could be dependent on the acid that you were titrating.

[wilson]

how long did it take for the color reversion?

It was quite short, about under half a minute. One instant it was faint-pink, I turned around for a moment, turned back, and the next thing it was colourless again.

[wilson]

What is the reason for quirky phenolphthalein to exhibit such behaviour?
Perhaps it could have been due to carbon dioxide in the air reacting with my base like NaOH

That is good thinking. What do you get with CO₂ dissolving in an aqueous solution?

Or it could be dependent on the acid that you were titrating.

Carbonic acid--the acidic impurity that reacts reacts with the sodium hydroxide.

For a simple example, the acid I used was HCl.

I really hope the colour reversion wasn't because of poor technique. Assuming that it is not due to operator's error, what could be the reason for the colour reversion? I heard something about phenolphthalein's characteristics that makes it not quite a good indicator.

And since we are on the general topic of titration--how can I reduce the error caused by carbon dioxide reacting with sodium hydroxide? It is considered a systemic error right? Will titrating the NaOH to a standard (like KHP) help at all?

[LQ43]

Yes, carbonic acid can be produced in significant amount to react with excess NaOH.  With such weak acids, titrations can give difficult endpoints with indicators. 

Using any indicator means that you are doing a second titration as the indicator itself is a weak acid and some of the NaOH is being used to convert the indicator to its base form. Using as small amount of indicator as reasonable is recommended (that is only a few drops instead of shooting in a dropperful)

CO2 error can be reduced by using only just previously boiled deionized water for sample preps and any extra used for washing down. 

Standardization of NaOH with KHP means that you are doing another titration, and usually with phenolthalein/suitable indicator. So same careful techniques apply (to reduce error)

If there are very accurate balances available, simply diluting a charged amount of NaOH(s) in a volumetric flask to the proper concentration can be done instead of KHP standardization. Consider the error on the balance and the volumetric flask

[wilson]

Yes, carbonic acid can be produced in significant amount to react with excess NaOH.  With such weak acids, titrations can give difficult endpoints with indicators. 

Using any indicator means that you are doing a second titration as the indicator itself is a weak acid and some of the NaOH is being used to convert the indicator to its base form. Using as small amount of indicator as reasonable is recommended (that is only a few drops instead of shooting in a dropperful)

CO2 error can be reduced by using only just previously boiled deionized water for sample preps and any extra used for washing down. 

Standardization of NaOH with KHP means that you are doing another titration, and usually with phenolthalein/suitable indicator. So same careful techniques apply (to reduce error)

If there are very accurate balances available, simply diluting a charged amount of NaOH(s) in a volumetric flask to the proper concentration can be done instead of KHP standardization. Consider the error on the balance and the volumetric flask

Yes, I understand.

Just two clarifications:
1. If I dilute a known amount of NaOH, is it still possible for CO2 in the environment to pose a great error to my titration due to reaction with NaOH? Let's say the time from the dilution to titration is <1hr, and the molarity of NaOH is about 0.5M.

2. Besides being a weak acid itself, are there other reasons for phenolphthalein to revert in colour/give false end-points? I keep on asking for more reasons because I heard something special about phenolphthalein not observed in many other indicators (not the pKa/pH range). It has something got to do with side reactions if I remember correctly.

Thank you.

[Borek]

If there are very accurate balances available, simply diluting a charged amount of NaOH(s) in a volumetric flask to the proper concentration can be done instead of KHP standardization. Consider the error on the balance and the volumetric flask

No, that's a very bad idea. In practice NaOH is always contaminated with Na₂CO₃ - it very easily adsorbs atmospheric CO₂. Not to mention fact that it is hygroscopic. That's why you have to normalize it against some primary substance like KHP and you can't prepare solution of exactly known concentration by weighting.

[Borek]

1. If I dilute a known amount of NaOH, is it still possible for CO2 in the environment to pose a great error to my titration due to reaction with NaOH? Let's say the time from the dilution to titration is <1hr, and the molarity of NaOH is about 0.5M.

CO₂ is always a problem, but if you are reasonably quick error should be neglectable.

Besides being a weak acid itself, are there other reasons for phenolphthalein to revert in colour/give false end-points? I keep on asking for more reasons because I heard something special about phenolphthalein not observed in many other indicators (not the pKa/pH range). It has something got to do with side reactions if I remember correctly.

No side reactions that I am aware off. In high pH phenolphthalein will loose its color (reversibly) again, but that's another story, your pH is not high enough for that.

[wilson]

1. If I dilute a known amount of NaOH, is it still possible for CO2 in the environment to pose a great error to my titration due to reaction with NaOH? Let's say the time from the dilution to titration is <1hr, and the molarity of NaOH is about 0.5M.

CO₂ is always a problem, but if you are reasonably quick error should be neglectable.

Besides being a weak acid itself, are there other reasons for phenolphthalein to revert in colour/give false end-points? I keep on asking for more reasons because I heard something special about phenolphthalein not observed in many other indicators (not the pKa/pH range). It has something got to do with side reactions if I remember correctly.

No side reactions that I am aware off. In high pH phenolphthalein will loose its color (reversibly) again, but that's another story, your pH is not high enough for that.

Yea. That's the one I have been looking for. About the different forms of phenolphthalein... Thanks.
However, I realized that the fading of pink colour to form a colourless solution at ~pH>12 has been observed at different concentrations from different sources. Some sources say that the effect is only visible at 4M NaOH, while some sources say that the effect is visible even from low concentrations i.e. <1M NaOH. I think it is still possible for pink colour to fade away even at low concentrations of NaOH. I guess it would just take longer since the step to change from P2- to POH3- is slow and its rate is dependent on the concentration of OH-.

[Borek]

Note that change of the color in the high pH is reversible, just like the change used for the endpoint detection when titrating, so as long as the pH is stable color of the solution should be constant.

[wilson]

Note that change of the color in the high pH is reversible, just like the change used for the endpoint detection when titrating, so as long as the pH is stable color of the solution should be constant.

But at the end point of a titration, there is bound to be an excess of NaOH. Because the concentration of this excess NaOH is not high, it takes a while for the pink colour to fade away. Besides the colourless (pH0-8.2) to faint pink is a faster step than the slower pink to colourless (pH >12).

[LQ43]

If there are very accurate balances available, simply diluting a charged amount of NaOH(s) in a volumetric flask to the proper concentration can be done instead of KHP standardization. Consider the error on the balance and the volumetric flask

No, that's a very bad idea. In practice NaOH is always contaminated with Na₂CO₃ - it very easily adsorbs atmospheric CO₂. Not to mention fact that it is hygroscopic. That's why you have to normalize it against some primary substance like KHP and you can't prepare solution of exactly known concentration by weighting.

I was not aware of the contamination with Na2CO3 but in practice in our labs we keep NaOH under tightly sealed conditions although that doesn't discount the error in mass. Thanks for the correction