Acid–base indicators | Acids and bases | AP Chemistry | Khan Academy

Acid-base indicators are used in titrations to determine when the equivalence point is reached. Let's look at a hypothetical indicator. In the protonated form, the indicator has the formula H-I-N. So this would be the acidic proton on this protonated form. When base is added, the protonated form is converted into the deprotonated form.

So we lose an H and we lose a positive charge. The deprotonated form is represented by I-N with a negative charge. If we add an H plus to the deprotonated form, we would make the protonated form of the acid-base indicator. For this hypothetical acid-base indicator, the protonated form is red and the deprotonated form is yellow.

Next, let's think about the color of a solution containing this acid-base indicator at different pH values. Let's say that the pKa value for the acidic proton on the protonated form is equal to 3 at 25 degrees Celsius. And let's say the pH of the solution is equal to 2. In that case, the pH of the solution is less than the pKa value for our acidic proton.

If we think about the protonated form being a weak acid and the deprotonated form being the conjugate base, when the pH is less than the pKa value for a conjugate acid-base pair, the concentration of the weak acid is greater than the concentration of the conjugate base. At a pH of 2, we have a lot more of the protonated form, which is red, than the deprotonated form, which is yellow. Therefore, the color of the solution at a pH of 2 would appear to be red.

Next, let's think about the color of the solution if the pH of the solution is equal to 4. Since the pKa value is equal to 3, if the pH is equal to 4, the pH is greater than the pKa value. For a conjugate acid-base pair, when the pH of the solution is greater than the pKa value, the concentration of the conjugate base is greater than the concentration of the weak acid. Or, as I have written here, the concentration of the weak acid is less than the concentration of the conjugate base.

So at a pH of 4, we have a lot more of the deprotonated form, which is yellow, than the protonated form, which is red. And since we have more yellow than red, the solution would be yellow. Finally, let's think about the situation where the pH is equal to 3. Since the pKa value is also equal to 3, the pH is equal to the pKa value.

For a conjugate acid-base pair, when the pH is equal to the pKa, the concentration of the weak acid is equal to the concentration of the conjugate base. If we have equal amounts of the protonated form and the deprotonated form, we have equal amounts of red and yellow. Therefore, at a pH of 3, the solution would be orange.

Next, let's think about the pH range of the color change for this hypothetical indicator. The approximate pH range over which an indicator changes color is equal to the pKa value plus or minus one. The pKa value for our hypothetical indicator was equal to 3; therefore, 3 plus one is equal to 4, and 3 minus one is equal to 2.

And 2 to 4 is the approximate pH range over which our indicator changes color. We already know our solution with the acid-base indicator in it at a pH of 2 is red, at a pH of 3 the solution is orange, and at a pH of 4 the solution is yellow. Therefore, if we were to change the pH from 2 to 4, we would see the color of the solution go from red to orange to yellow.

Now let's see how to choose the right acid-base indicator for a titration. Our goal is to choose an indicator whose color change occurs as close as possible to the pH of the equivalence point. For a weak acid-strong base titration, the equivalence point occurs at a pH greater than seven.

First, let's look at the acid-base indicator methyl red. Methyl red has a pH range of approximately 4 to 6. So a little bit over 4, methyl red starts to change from red to orange and then eventually to yellow by the time you hit a pH of about 6. However, the pH of the equivalence point for this titration looks to be somewhere between 8 and 10.

Therefore, if we used methyl red and we stopped the titration when the color changed, we'd be stopping the titration too early. So we might be stopping it somewhere in here before we reach the equivalence point. So methyl red would not be a good choice as an acid-base indicator for this titration.

Another way to think about this is with pKa values. For methyl red, the pKa value is approximately 5, and the goal is to match the pKa value as closely as possible to the pH of the equivalence point. But since the pH at the equivalence point is between 8 and 10, that's too far away from 5. So by looking at the pKa value, we know that methyl red is not a good fit for this titration.

Phenolphthalein is an example of another acid-base indicator and has a different pH range. At a pH of about 8, phenolphthalein is colorless. However, as the pH changes from 8 to 10, phenolphthalein goes from colorless to pink. Because the color of the indicator changes in the same range where we would find the equivalence point, phenolphthalein is a good choice as an acid-base indicator for this titration.

Thinking about the pKa value for phenolphthalein, which is approximately 9, that falls in the range of 8 to 10 where we find our equivalence point. So we could think about it either in terms of the pH range or the pKa value.

Next, let's choose an indicator for a weak base-strong acid titration. For a weak base-strong acid titration, the equivalence point occurs at a pH less than 7. If we try to use phenolphthalein for this titration, remember phenolphthalein changes from 8 to 10 or, in this case, it'd be changing from 10 to 8.

So we'd start at this relatively high pH here, and if we try to use phenolphthalein and we stopped it when the color change occurred, we'd be stopping the titration too early. So we might be stopping the titration somewhere in here. So phenolphthalein is not a good choice as an acid-base indicator for this particular titration.

Thinking about using pKa values, the pKa value for phenolphthalein is approximately 9, which is not a good fit for the pH at the equivalence point, which appears to be between 4 and 6. Methyl red has a pH range of about 4 to 6 over which it changes color and a pKa value of approximately 5.

Since the pH of the equivalence point is between 4 and 6, methyl red would be a good choice as an acid-base indicator for this titration. So to summarize, when choosing an acid-base indicator for a titration, choose an indicator whose color change occurs as closely as possible to the pH at the equivalence point.