pKa of indicators

[epsilonprodigy]

Anyone know the possible consequence of using an indicator whose pKa is much higher or much lower than the pH of the solution at the equivalence point?

This is not a question I've run into, but rather one that I can see being asked....


This is what I'm thinking. Let's assume you're using an acidic indicator and adding it to a base. If the pka of the indicator is far below the pH @ equivalence, it would keep dissociating even after you'd passed the equivalence point. Therefore, you WOULD NOT observe a color change at the equivalence point, but rather much later, after the pH had gone much lower.

If the pKa of the indicator is a lot higher than the pH at equivalence point, it will stop dissociating prematurely, as in, before you reach the equivalence point. You'll then see the color change, meaning the presence of un-dissociated indicator, before you've reached the equivalence point.


Does that sound right?

---

Members don't see this ad.

 

[ilovemcat]

An indicator is no different from any other buffer. If placed in an extremely acidic or extremely basic solution, the indicator will exist primarily in it's conjugate base form or conjugate acid form. The conjugate base might exhibit one color (say blue) and the conjugate acid might exhibit another color -- or even no color at all (for our purposes well say yellow). The ideal scenario for the indicator is an equivalence point around it's buffer region (where it resists changes in pH). This point is of signficant importance because we can observe a color change -- it's this color change that tells us the equivalence point has been reached.

Let's say you titrated HCl with NaOH -- we know for a strong acid/base titration equivalence point is at 7. However, you instead decide to use an indicator with a pKa of 5. Inititally, the solution of the beaker (with HCl) has a pH of 2. Here, the indicator exists primarily in it's acidic form (so the indicator appears yellow). As you add more and more NaOH, the pH will rise -- eventually it'll reach a pH of 5. At a pH of 5, the indicator behaves as a buffer (equal amounts of conjugate base and acid). If the conjugate base appeared blue and the conjugate acid appeared yellow then during this buffer region it'll appear as a mixture of the two -- say green. The problem with this is that (had you not known this was a strong acid/base titraton) you would of thought that the equivalence point was reached -- but yet we know that's not the case since the equivalence point for any (monoprotic) strong acid base titration is at a pH of 7. As you continue adding more NaOH, the conjugate base form of the indicator will predominate and the indicator will appear more blue - until eventually all the organic acid is deprotonated into conjugate base. At this point, you will observe no color change.That's why it's of crucial important to use an indicator near or close to the equivalence point.

 

[ilovemcat]

Anyone know the possible consequence of using an indicator whose pKa is much higher or much lower than the pH of the solution at the equivalence point?

This is not a question I've run into, but rather one that I can see being asked....


This is what I'm thinking. Let's assume you're using an acidic indicator and adding it to a base. If the pka of the indicator is far below the pH @ equivalence, it would keep dissociating even after you'd passed the equivalence point. Therefore, you WOULD NOT observe a color change at the equivalence point, but rather much later, after the pH had gone much lower.

If the pKa of the indicator is a lot higher than the pH at equivalence point, it will stop dissociating prematurely, as in, before you reach the equivalence point. You'll then see the color change, meaning the presence of un-dissociated indicator, before you've reached the equivalence point.


Does that sound right?

By the way, what you said was right. 🙂 Just thought I'd explain why that's the case in hopes that would understand it better - but it seems you already got it.

 

[indianjatt]

nvm