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Current time:0:00Total duration:7:05

AP Chem: SAP‑10 (EU), SAP‑10.A (LO), SAP‑10.A.1 (EK), SAP‑10.A.2 (EK)

- [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.

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