Acid–base indicators are compounds that change color when they become protonated or deprotonated. Because this color change occurs over a specific pH range, indicators can be used to approximate the equivalence point of an acid–base titration. Created by Jay.
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- Isn't the pKa = pH at the midpoint? Why are we saying, for some of them, that the pKa = pH at the equivalence point? (e.g.5:40)
Am I misunderstanding something?(3 votes)
- what is the meaning of prononated(1 vote)
- Being protonated means that something has gained a hydrogen nucleus. Protonating something means that we’ve donated a hydrogen nucleus to another thing. In essence something acting as a base (by accepting a proton) or acting as an acid (by donating the proton). The vast majority of hydrogen atoms are composed of a single proton and an electron. If we remove that electron somehow, then we’re left with the hydrogen’s nucleus, or simply a proton.
Hope that helps.(2 votes)
- [Instructor] 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 HIn. 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. So the deprotonated form is represented by In with a negative charge. And 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 three at 25 degrees Celsius. And let's say the pH of the solution is equal to two. In that case, the pH of the solution is less than the pKa value for our acidic proton. So if you 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 two, 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 that a pH of two would appear to be red. Next, let's think about the color of the solution if the pH of the solution is equal to four. Since the pKa value is equal to three, if the pH is equal to four, the pH is greater than the pKa value. And 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 four, 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 three. Since the pKa value is also equal to three, the pH is equal to the pKa value. And 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. And 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 three, the solution would be orange. Next let's think 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 three, therefore three plus one is equal to four and three minus one is equal to two, and two to four is the approximate pH range over which our indicator changes color. And we already know our solution with the acid-base indicator in it at a pH of two is red and at a pH of three the solution is orange and at a pH of four the solution is yellow. Therefore, if we were to change the pH from two to four, 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 based 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 four to six. So a little bit over four methyl red starts to change from red to orange, and then eventually to yellow by the time you hit a pH of about six. However, the pH at the equivalence point for this titration looks to be somewhere between eight 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 five. And the goal is to match the pKa value as closely as possible to the pH at the equivalence point. But since the pH at the equivalence point is between eight and 10, that's too far away from five. 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 it has a different pH range. At a pH of about eight, phenolphthalein is colorless. However, as the pH changes from eight 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. And thinking about the pKa value for phenolphthalein which is approximately nine, that falls in the range of eight 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 based strong acid titration. For a weak based strong acid titration, the equivalence point occurs at a pH less than seven. If we try to use phenolphthalein for this titration, remember phenolphthalein changes from eight to 10, or in this case, it'd be changing from 10 to eight. So we'd start at this relatively high pH here, and if we tried to use phenolphthalein and we stopped at 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 nine, which is not a good fit for the pH at the equivalence point which appears to be between four and six. Methyl red has a pH range of about four to six over which it changes color and a pKa value of approximately five. And since the equivalence point, the pH of the equivalence point is between four and six, methyl red would be a good choice as an acid-based 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.