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AP®︎/College Chemistry
Weak base–strong acid titrations
For the titration of a weak base with a strong acid, the pH curve is initially basic and has an acidic equivalence point (pH < 7). The section of curve between the initial point and the equivalence point is known as the buffer region. At the half-equivalence point, the concentrations of the buffer components are equal, resulting in pH = pKₐ (where pKₐ refers to the conjugate acid of the weak base). Created by Jay.
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I might be a little paranoid, but at1:27I'm really confused why the initial solution has a more steep titration curve than the "buffer region". I feel like it would be much more effective in terms of neutralizing acid, since looking at the Henderson Hasselbalch equation, we have pH=pKa+log([NH4+]/[NH3]). The larger NH3, the less drastic the change in [NH4+]/[NH3] once we add acid. Thus, I'm really confused. What am I missing?(3 votes)
Before any titrant is added, the analyte solution is mostly the weak base, ammonia, with a very minor amount of weak conjugate acid, ammonium, present. When we add the acid titrant, ammonia reacts with the acid vigorously and shifts the ammonia neutralization reaction to the right. If you think of it as an equilibrium problem, adding acid disturbs equilibrium and Le Chatelier’s principal speeds up the forward reaction to relieve the stress and restore equilibrium.
After this initial burst of neutralization, we begin to produce appreciable amounts of ammonium. And now a solution with appreciable amounts of a weak base and its conjugate acid is a buffer solution which obstructs the drastic change in pH. So we would expect the change in pH to slow as we create a buffer with increased titrant volume.
The steep slope of the titration curve in the beginning corresponds to a large change in pH per mL of titrant added, while the gentler slope of the buffer region corresponds to a small change in pH per mL of titrant added.
Hope that helps.(3 votes)
1:27Video transcript
- [Instructor] Ammonia is
an example of a weak base and hydrochloric acid is an
example of a strong acid. And if we're doing a weak
base-strong acid titration, that means that ammonia is the analyte, the substance we're analyzing, and we're titrating ammonia
with hydrochloric acid and therefore hydrochloric
acid is the titrant. And when ammonia reacts
with hydrochloric acid, the product is an aqueous
solution of ammonium chloride. For our complete or
overall ionic equation, since ammonia is a weak base, we show it as NH3 in our
complete ionic equation. However, since hydrochloric
acid is a strong acid that ionizes 100%, we show it as breaking up into its ion, so H+ and CL-. Ammonium chloride is a soluble salt, therefore we would show ammonium chloride in aqueous solution as ammonium cations and chloride anions. To write the net ionic equation, we leave out spectator ions. And since we have chloride
anions on the left side and on the right side, chloride anions are the spectator ions. And leaving those out, we
get the net ionic equation, which is ammonia NH3 plus H+ goes to NH4+. So this is one way to write
to the net ionic equation for this weak base-strong acid titration. Next, let's look at the titration curve for our weak base-strong acid titration. pH is on the y-axis and milliliters
of acid is on the x-axis because we're adding our strong
acid to our aqueous solution of our weak base. Looking at the first point
on our titration curve, the pH is relatively basic. So this is before any
strong acid has been added. The reason why the pH is
basic is because we have an aqueous solution of
our weak base, ammonia, which reacts with water to
produce ammonium cations and hydroxide anions. And it's these hydroxide
anions that cause the pH to be relatively high. However, the equilibrium
favors the reactants for this reaction. So we have mostly ammonia
and very little ammonium at this point in the titration curve. Next, we think about adding some acid to our aqueous solution of ammonia. And from our net ionic equation, when ammonia reacts with H+, that forms the ammonium cation, NH4+. Looking at the titration curve, as we add more and more acid, the pH starts to decrease. However, in this range, there's
a slow decrease in the pH. As more acid is added, more ammonia is turned
into the ammonium cation. Eventually, we reach a point where all of the initial
ammonia has been neutralized by the addition of the acid. This point is called
the equivalence point. And the way to find the equivalence point on our titration curve is to first look for this
sharp decrease in the pH. And then we can draw a straight line here. And approximately halfway
down that straight line is a good estimate of
the equivalence point for this titration. To find the pH of the solution
at the equivalence point, we simply go over to where
the equivalence point is on the y-axis. And so for this pH, we can see the pH at the equivalence point is less than seven. So let me go ahead and
write that down here. The pH is less than seven for a weak base-strong acid titration. The reason why the pH is less than seven at the equivalence point is because all the ammonia
that we started with has been completely neutralized and turned into the ammonium cation, NH4+. The ammonium cation is a weak acid and reacts with water to
form hydronium ions, H3O+, and ammonia, an aqueous solution. At 25 degrees Celsius,
water has a pH of seven. However, since the ammonium
cation is a weak acid and we're increasing the
concentration of hydronium ions in solution, that decreases the pH, therefore the pH is less than seven at the equivalence point. In addition to ammonium ions, there are also chloride
anions in solution. However, chloride anions
do not react with water and therefore do not affect the pH. Going back to the equivalence
point on our titration curve, if we dropped down here to the x-axis, we can see the equivalence point
occurs after 50 milliliters of acid has been added. Therefore, if it took
50 milliliters of acid to neutralize all of the ammonia
that was initially present, it would take half that volume
or 25 milliliters of acid to neutralize half of the ammonia. So if we go back up here
and we draw a dashed line to our titration curve, this point on our
titration curve represents the half equivalence point. So this point represents
the half equivalence point on our titration curve. And since we've neutralized
half of the ammonia that was initially present, that means there are equal
concentrations of ammonia and the ammonium cation at this point. Let's go back to our equivalence points where all the ammonia that we started with has been neutralized. Therefore, if we add some
more acid to the solution, there's no more ammonia
for it to react with. And therefore we see the pH drop. So this portion of the titration curve is the region of excess acid. Let's go back to the
half equivalence point on our titration curve. Remember at that point, the concentration of ammonium cation is equal to the concentration of ammonia. The ammonium cation and ammonia are a conjugate acid-base pair. And when there are significant amounts of a weak conjugate acid-base pair, there's a buffer solution. Therefore, at the half equivalence point, we have a buffer solution, and we can calculate the pH at that point by using the
Henderson-Hasselbalch equation. So pH is equal to the
pKa of the weak acid, plus the log of the concentration
of the conjugate base, divided by the concentration
of the weak acid. For this example, the
base is ammonia, NH3, and the conjugate acid is
the ammonium cation. NH4+. Therefore, this pKa value in the Henderson-Hasselbalch equation is referring to the pKa value of ammonium. And because the concentrations
of ammonium and ammonia are equal at the half equivalence point, the ratio of their
concentrations is equal to one and the log of one is equal to zero. Therefore, at the half equivalence point, the pH is equal to the pKa
value of the weak acid. So if we wanted to find the pKa value for the ammonium cation, we would find the half equivalence point and we'd draw our dotted line over to where the intersects on our y-axis and whatever pH that is, that's the pKa value of ammonium. So in this case, it looks
to be a little bit over nine as a good estimate for the pKa value of the ammonium cation. Next, let's think about how
our titration curve can tell us about the relative concentrations of our weak conjugate acid-base pair. We know that at the
half equivalence points where the pH is equal to the pKa value, the concentration of ammonium cations is equal to the concentration of ammonia. So let's think about a
point just to the left of our half equivalence point, which I'm gonna call point P. At point P, the pH is greater than the pKa value. And we know the initial
point on our titration curve was almost all weak base, almost all NH3. Because point P is in
between the initial point where there was almost all NH3, and the half equivalence point
where there was equal amounts of NH3 and NH4+, at point P, there must be more NH3 than NH4+. Therefore, when the pH of
the solution is greater than the pKa value, we know the concentration
of ammonia is greater than the concentration of the ammonium cation. Or you could say the
concentration of ammonium is less than the concentration of ammonia. We could have also figured this out using the Henderson-Hasselbalch equation. However, it's often
easier just to think about the shape of the titration curve and where the point in question is in relation to important points. For example, in this case, the initial point and the
half equivalence point. Next, let's think about
a point just to the right of the half equivalence point. And I'm gonna call this point Q. At point Q, the pH of the
solution is less than the pKa. Point Q is in between the
half equivalence point and the equivalence point, which is approximately here
on the titration curve. Remember at the equivalence point, all the ammonia that we started with has been converted into ammonium, NH4+. And because point Q is in between
the half equivalence point where the amount of NH3 is
equal to the amount of NH4+, and the equivalence
point where all the NH3 has been converted into NH4+, all of the initial NH3. That means that at Q, there
must be more NH4+ than NH3. Therefore, when the pH is
less than the pKa value, we can say the concentration
of ammonium, NH4+, is greater than the
concentration of ammonia, NH3. The half equivalence
point, point P and point Q are all a part of the buffer
region on the titration curve. Remember that buffers
resist large changes in pH, and that's why we see
a slow decrease in pH as acid is added at this
part of the titration curve. At the very beginning of the titration, we had almost all ammonia
and therefore we did not have a buffer solution. However, as acid was added
and the ammonia was converted into the ammonium cation, NH4+, when significant amounts
of both are present, we do have a buffer solution. And that represents the buffer region on our titration curve, so in here. However, as more and more acid is added, we can see a sharp change in pH start to occur right about here, so we're no longer in the buffer region as we approach the equivalence point. So when we think about the titration curve of a weak base-strong acid titration, and we think about the
half equivalence point where the pH is equal to the
pKa value of the weak acid, it's important to remember
that there's a buffer region or a buffer zone around
that half equivalence point.