Titration Curves
Purpose:
In this lab you will experimentally determine the titration curves for several acids and
find the dissociation constants, Kas, for the acids.
Overview:
Acid–base titration is a useful technique for determining the concentration of an acidic or
basic solution. The titration of a weak acid or base is also useful for determining the Ka
or Kb for an acid or base, respectively.
In this experiment you will use a pH meter and buret to generate plots of titrations, called
titration curves. By plotting pH versus the volume of titrant the curve shows how the pH
of the solution changes as an acid, or base, is added. The shape of a titration curve is
influenced by both the concentration as well as the nature of the acid, or base. By
titrating both strong and weak acids you will see different titration curves and how they
change with the strength of the acid.
For this experiment you will perform three different titrations. First, you will titrate a
strong acid, HCl, with a strong base, NaOH. The concentration of the HCl will be
known, and the results of the titration will allow you to calculate the concentration of the
NaOH. The equation for the reaction is:
(1)
HCl + NaOH → NaCl + H2O
Secondly, you will titrate a weak acid, potassium hydrogen phthalate, KHC8H4O4,
(abbreviated KHP), with a strong base. Phthalic acid is a weak, diprotic acid; KHP is the
potassium salt of phthalic acid and is itself a weak acid. From the results of this titration
you will be able to calculate the Ka for KHP. The equation here is:
(2)
KHC8H4O4 + NaOH → NaKC8H4O4 + H2O
Lastly, you will titrate a weak polyprotic acid, maleic acid, H2C4H2O4, with a strong base.
The results of this titration will allow you to calculate Ka for the removal of each
successive H+ ion in the acid. It will also be interesting to compare the shape of the
titration curves in the three cases. The individual and net equations are:
(3)
(a) H2C4H2O4 + NaOH → NaHC4H2O4 + H2O
(b) NaHC4H2O4 + NaOH → Na2C4H2O4 + H2O
Net: H2C4H2O4 + 2NaOH → Na2C4H2O4 + 2H2O
In all cases you will need to identify the endpoint of the titration. The endpoint is the
point at which the number of moles of OH– added to the solution equals the moles of H+
originally present. The pH of the solution at the endpoint depends on what is left in the
flask once you have reached the endpoint. Before beginning the experiment think about
what ions will be present at the endpoint in each titration, and decide whether you think
the pH at the endpoint will be acidic, basic, or neutral.
Materials:
Chemicals:
0.1M HCl solution
sodium hydroxide
potassium hydrogen phthalate
maleic acid
phenolphthalein indicator solution
Equipment:
buret
250 mL volumetric flask
Balance
250 mL beaker
electronic pH meter
Erlenmeyer flask
Hazards:
You will be handling strong acids and bases in this lab. Goggles must be worn
throughout this experiment.
Procedure:
1. Prepare 250 mL of an approximately 0.1 M solution of sodium hydroxide.
2. Obtain 30-35 mL of the 0.1 M HCl.
3. Clean a 50 mL buret and rinse it with the NaOH solution. Fill the buret with the
NaOH and set up a ring stand and beaker for titration.
4. Accurately, using a pipette if available, measure 25.00 mL of 0.10 M HCl into the
beaker. Add a few drops of phenolphthalein.
5. Check your pH meter for calibration using buffer solutions.
6. Position the HCl solution below the buret. Place the pH electrode in the solution.
7. Gently swirl the solution and measure the pH. Record the buret volume and the
pH.
8. You are now ready to begin the titration. This process goes more smoothly if one
person manipulates and reads the buret while another person reads the pH meter.
Your goal will be to generate enough data points to give a smooth curve. In some
regions of the titration the pH changes very slowly and you can add several mL of
base at a time. Nearer the endpoint, the pH changes very quickly and you will
need to add base one drop at a time.
Add the first increment of NaOH, enough to raise the pH about 0.15 units.
Record the pH and buret volume. Continue adding NaOH solution in increments
that raise the pH by about 0.15 units and enter the buret reading after each
increment. When a pH of about 3.5 has been reached, change the increment of
NaOH to single drops. Be sure to continue to enter the buret readings after each
increment.
When a pH of about 10 is reached again add larger increments that raise the pH
by about 0.15 pH units, and input the buret reading after each increment.
Continue adding NaOH solution until the pH remains constant.
9. Graph the data. This is your titration curve for the hydrochloric acid. The
endpoint is the volume at the steepest part of the curve (called the inflection
point). Determine the volume of NaOH needed to reach the endpoint and use this
value to calculate the exact molarity of the NaOH.
10. Repeat the procedure using a 0.50g sample of KHP dissolved in approximately 50
mL of distilled water instead of the HCl solution. Determine the acid dissociation
constant for KHP (equal to Ka2 for phthalic acid).
11. Repeat the procedure using a 0.50g sample of maleic acid dissolved in
approximately 50 mL of distilled water instead of the HCl solution. Determine
the acid dissociation constants (Ka1 and Ka2) for maleic acid.
Hints for Finding Ka:
A theoretical titration curve for acetic acid is shown below.
Titration curve of acetic acid
For this titration, a 10 mL volume of 0.5 M acetic acid is titratated with 0.5 M NaOH.
Because the concentration of acid and base are the same, the equivalence point is reached
when 10 mL of base are added to the solution. The equivalence point, when enough base
is added to react with all the acid present, exhibits the sharp increase in pH, as discussed
above. This happens when the moles of base added equals the moles of acid in the
sample. It is this point that is used to determine the concentration.
The flat portion of the titration curve before the equivalence point is called the buffer
region. In this part of the pH scale, the acid and conjugate base are both present in
significant concentrations and the solution resists changes in pH. As base is added to a
solution in this buffer region, acetic acid reacts with it to form acetate ion, without a large
change in pH. If additional acid were added to a solution in the buffer region, it would
react with the conjugate base, acetate ion, and, again, the pH would not change
appreciably.
In the middle of the buffer region lies the half-equivalence point. Here the volume of
base added is half that required to reach the equivalence point and half the acetic acid has
been converted to the conjugate base, acetate ion. This means that the concentrations of
acetic acid and acetate ion are equal.
If we examine the equilibrium expression at the half-equivalence point, we find
something interesting:
At the half-equivalence point, [CH3COOH] = [CH3COO-], so
Ka = [H3O+]
And taking the log and multiplying both sides by -1 yields
pKa = pH
So at the half-equivalence point, half the acetic acid has been converted to the conjugate
base, acetate ion, and the pH will be equal to the pKa of the acid. This gives us an
experimental way to determine the pKa of a weak acid.
This technique will also work for polyprotic acids if the Kas are not too close. If they are
at least a factor of 100 apart, the interaction of the equilibria will be minimal and can be
ignored. Ka1 is determined as above. Subsequent Kas are determined using the pH at half
volume from the previous equivalence point.
Adapted from: http://www.ivygreen.ctc.edu/knutsen/chem160/titracrv.htm and
www.dartmouth.edu/~chemlab/chem3-5/acid1/full_text/chemistry.html