Weak base–strong acid titrations | Acids and bases | AP Chemistry | Khan Academy

- [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.