1. EXPERIMENT V
POTENTIOMETRIC TITRATION
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INTRODUCTION
Many Acid-Base titrations are difficult to accomplish using a visual indicator for one of several
reasons. Perhaps the analyst is color-blind to a particular indicator color change; there may not
be a suitable color change available for a particular type of titration or the solutions themselves
may be colored, opaque or turbid. It may be desired to automate a series of replicate
determinations. In such situations, potentiometric titration, using a glass hydronium ion selective
electrode, a suitable reference electrode and a sensitive potentiometer (a pH meter) may be
advantageous.
THEORY
Any acid-base titration may be conducted potentiometrically. Two electrodes, after calibration
[to relate potential in millivolts (mV) to a pH value] are immersed in a solution of the analyte.
One is an indicator electrode, selective for H3O+
and the other a stable reference electrode. The
potential difference, which after calibration is pH, is measured after the successive addition of
known increments of acid or base titrant.
When a potentiometric titration is being performed, interest is focused upon changes in the emf
of an electrolytic cell as a titrant of known concentration is added to a solution of unknown. The
method can be applied to all titrimetric reactions provided that the concentration of at least one
of the substances involved can be followed by means of a suitable indicator electrode. The
critical problem in a titration is to recognize the point at which the quantities of reacting species
are present in equivalent amounts. The titration curve can be followed point by point, plotting as
ordinant, successive values of the cell emf (pH) vs the corresponding volume of titrant added. A
typical titration curve is presented in Figure 2. Figure 3 represent another method for
determining the equivalence point from the titration curve data. Table I, in Appendix I, presents
typical data obtained from a potentiometric titration.
The Reference Electrode
Most commonly, the reference electrode is the silver/silver chloride electrode. The
potential is based on the following equilibrium:
AgCl(s) + e →
←
Ag(s) + Cl-
(aq)
The half cell is:
Ag[(AgCl(Sat'd),KCl(xM)]
NOTE: Electrodes respond to the activity of the electroactive species in solution. However, as a practical
matter it is more convenient experimentally to use concentration. For this reason the discussion in this laboratory
experiment will be made in terms of concentration rather than the more correct activities.

2. The practical version of this electrode is a silver wire dipping into a saturated solution of
KCl; when fabricated this way its electrode potential is 0.199V (vs. Normal Hydrogen Electrode,
NHE) @ 25o
C. The potential is a function of temperature and the concentration of KCl in the
solution. Such an electrode is comparatively rugged, reliable, and inexpensive.
The Indicator Electrode
The heart of the glass electrode is a thin glass membrane, specially fabricated to preferentially
exchange H3O+
. The outside of the membrane is in contact with the analyte solution containing
the unknown [H3O+
]. The inside of the membrane contacts a hydrochloric acid solution of fixed
concentration. A silver wire, coated with AgCl dips into this solution; the other end of the wire
is connected to the measuring device. A combination glass electrode with silver/silver chloride
reference may be represented as shown in Figure 1 below:
FIG. 1. pH Electrode
The Mechanism of the Response
A change in hydronium ion concentration causes a change in composition of the glass membrane
due to an ion exchange process involving the solution and the membrane (see textbook for
details). A corresponding change in membrane potential, proportional to pH, is what is
measured. All other potentials are constant. In effect the membrane potential (variable) is
measured against two fixed potentials, the external reference and the internal reference, both
Ag/AgCl reference electrodes.
Potential difference is measured using a high impedance potentiometer. This high resistance
dictates a very small current flow.
Each glass electrode is different, due to the difficulty of reproducing the glass membrane; it is,
therefore, necessary to standardize the meter and electrode against at least two solutions of
accurately known pH. Such standard buffer solutions are available from many different
manufacturers.

3. PART I
DETERMINATION OF THE % ACETYLSALICYLIC ACID IN ASPIRIN
INTRODUCTION
The purpose of the experiment is to introduce the student to the use of electrochemical
instrumentation in analytical chemistry, specifically, potentiometric end-point detection using
derivative techniques.
In this part of the experiment you will use a pH meter to determine the active ingredient in an
aspirin tablet. The aspirin tablet will be dissolved directly in an aqueous-ethanol solution and
titrated with previously standardized NaOH. The change in pH will be monitored with the pH
meter throughout the titration. (See Appendix I). The equivalence point volume in the titration
will be determined graphically from the titration curve and by the first derivative methods (see
Appendix II). From these results the concentration of acetylsalicylic acid in the tablet will be
determined.
PROCEDURE
Accurately weigh the tablet in a 150 mL beaker. Drop onto the tablet two drops of deionized
water and wait a few minutes. The tablet should crumble. To the tablet in the beaker now add
15 mL of 95% ethanol and stir the tablet with a glass rod until it dissolves. The white solid that
dissolves slowly is starch and it may not dissolve completely. Titrate this solution using
the procedure in Appendix I.
CALCULATION OF RESULTS
Calculate the % acetylsalicylic acid in each sample, the average value, the absolute deviation,
and the relative average deviation in ppt and the confidence interval at a suitable confidence
level. Assuming the value supplied on the label is the true value, calculate the absolute and
relative error of your determination.
PART II
DETERMINATION OF THE % ACETIC ACID IN VINEGAR
INTRODUCTION
By law, any substance bottled and sold as vinegar must have a minimum acetic acid
concentration of 4.0%. In this experiment you will analyze one of a number of commercial
vinegars, using a potentiometric titration, to determine % acetic acid. You will note that wine
vinegar, cider vinegar, etc. must be titrated potentiometrically as their color precludes the use of
a visual indicator.

4. PROCEDURE
Pipet 5.00 ml aliquots of vinegar into an Erlenmeyer flask and dilute with 40–50 mL of
deionized water. Titrate using the procedure in Appendix I.
CALCULATION OF RESULTS
Calculate % acetic acid in each sample, the average value, the absolute deviation, and relative
average deviation in ppt and the confidence interval at a suitable confidence level. Assuming the
value for % acetic acid supplied by the manufacturer is 4.00%, calculate the absolute and relative
error of your determination.
PART III
DETERMINATION OF CARBONATE IN A SAMPLE
INTRODUCTION
Carbonate is a dibasic anion and which when titrated with a strong acid gives two distinct
equivalence points. The relevant chemical reactions are as follows.
CO3
2-
(aq) + H3O+
(aq) →
←
HCO3
-
(aq)+ H2O
HCO3
-
(aq) + H3O+
(aq) →
←
CO2 (g) + 2H2O
The carbonate system is ideally suited for use in a potentiometric titration to demonstrate a
titration curve with two inflection points corresponding to the two equivalence points. In this
experiment you will titrate a sample of sodium carbonate.
PROCEDURE
Accurately weigh one sample (about 0.2 g each) of sodium carbonate, which has been
previously dried in an oven for 1 2
1
hours, into three 200-mL beakers. Dissolve each sample with
in approximately 50 mL of deionized water. Titrate each sample using the procedure in
Appendix I. Note that in this case your titrant will be your standardized 0.1 M HCl rather than
the standardized 0.1 M NaOH.
CALCULATION OF RESULTS
Calculate the % sodium carbonate in each sample, the average value, the absolute
deviation, the relative average deviation in parts per thousand, the standard deviation and the
confidence interval at an appropriate confidence level. Assuming that your sample of sodium
carbonate was pure, determine if there is a determinate error in the method.
APPENDIX I
Standardize your pH meter using steps 1 to 14. Steps 15 to 16 describe the titration procedure
using the standardized pH meter.
1. Plug your pH meter in “Press STANDBY To Resume” will appear on screen.

5. 2. Press “Standby”.
3. A screen will appear which will say either, ‘Measure pH’, ‘Measure mV’, or ‘Measure
rel. mV’. Select appropriate screen for your measurements by toggling through ‘Mode’.
For pH Measurements Only.
4. Select screen ‘Measure pH’ by toggling ‘Mode’.
To Standardize pH meter for pH measurements.
5. If on ‘Measure pH’ screen under ‘BUFFER’ the numbers 4, 7, 10, etc. appear, they must
be removed as these are previous standardizations. If no numbers appear, proceed to Step
7.
6. To remove numbers of previous standardizations, press ‘Setup’ twice and then ‘enter’.
To Standardize pH meter using ‘two buffer method’.
7. Clean and dry two 50mL beakers. Into each beaker put ~ 25mLs of one of the two buffer
solutions. Into one of the beakers, place a small magnetic stir bar and place on a stirrer
hot plate. Adjust the stirrer until the magnetic stir bar is turning slowly (~ 1 turn a
second.)
8. Wash the pH electrode off with DI water and dry carefully using a kimiwipe. Place the
pH electrode into the electrode stand attached to the pH meter and using the stand
carefully lower the electrode into the beaker until the electrode tip is ~ 1 cm below the
surface. Make sure the magnetic stir bar does not hit the pH electrode. (If it does, add
more buffer solution and raise the pH electrode.)
9. Press ‘ Std.’ key, pause, and press ‘Std.’ key again.
10. The pH value of your buffer should appear on the screen under ‘BUFFER’.
11. Remove pH electrode from buffer, wash with DI water, and dry carefully with a
kimiwipe.
12. Place the second beaker with the second buffer solution on the stirrer–hot plate. Transfer
the magnetic stir bar using a ‘wand’ (wash and dry the stir bar before transfer) into the
second beaker.
13. Standardize with the second buffer as in steps 8–10.
14. If the standardization has worked ‘Good Electrode’ should appear on the screen. If it
does not, consult your instructor.
Making a pH measurement.

6. 15. After standardization wash and dry the pH electrode using DI water. Place the pH
electrode into the solution to be measured together with a magnetic stir bar. The pH
reading will appear on the screen. Allow the value to stabilize before reading. Record
the initial pH of the solution.
16. For Acidic Solutions
Titrate with standardized ~ 0.1M NaOH. Add the NaOH, approximately one
milliliter at a time, until the pH increases to about 4.5. Be sure to record the volume and
pH after each addition. After each addition of base allow the pH meter to stabilize before
recording the pH. (If the pH has not stabilized after ~ 1 minute consult your instructor).
Continue addition of NaOH using 0.2 mL increments until the pH increases to about 5.0.
(It is not necessary to record volume pH values after each addition. This approach allows
you to “feel” your way to the equivalence point region). Once the pH exceeds 5.0 make
an accurate volume reading and then use the following sequence.
(a) Add four drops of titrant to the beaker; record the pH but not the volume.
(b) Repeat step (a) until the pH increases to a value above 10.5.
(c) Record the volume which corresponds to the last four drop increment. Continue the
titration using 1.0 mL increments until the pH change corresponding to each addition
is less than 0.1 pH unit.
For Basic Solutions
Titrate with standardized ~ 0.1M HCl. Add the HCl approximately one milliliter
at a time, until the pH decreases to about 9.5. Be sure to record the volume and pH after
each addition. After each addition of acid allow the pH meter to stabilize before
recording the pH. (If the pH has not stabilized after ~ 1 minute consult your instructor).
Continue addition of HCl using 0.2 mL increments until the pH decreases to about 9.00.
(It is not necessary to record volume pH values after each addition. This approach allows
you to “feel” your way to the equivalence point region.) Once the pH has decreased
below 9.0 make an accurate volume reading and then use the following sequence.
(a) Add four drops of titrant to the beaker; record the pH but not the volume.
(b) Repeat step (a) until the pH has decreased to below a value of 3.5.
(c) Record the volume which corresponds to the last four drop increment. Continue the
titration using 1.0 mL increments until the pH change corresponding to each addition
is less than 0.1 pH unit.
17. For the next titration repeat steps 15 and 16.
APPENDIX II
The following data (Table I) was obtained from a potentiometric titration when a 0.4077 g
aspirin tablet dissolved in a water/ethanol mixture was titrated with a 0.1273 M NaOH solution.
You should use this as a guide for your calculation of the results of the potentiometric titration.
You should enter the data in a spreadsheet (for example Excel) and graph the results as you
collect them. Similarly, set up the spreadsheet to calculate the first rivative function.

7. TABLE I
Volume of NaOH (mL), pH of Solution, Corrected Volumes,
∆pH/∆V for the above Titration
Volume (mL) NaOH pH Corrected Volume (mL)
∆pH/∆
V
0.00 2.79
.
.
.
.
.
.
.
.
.
.
12.32 4.54
13.32 4.98
1344 0.48
4 drops (13.55mL) 5.01
13.67 0.74
4 drops (13.78 mL) 5.18
13.90 1.00
4 drops (14.01 mL) 5.41
14.13 2.49
4 drops (14.24 mL) 5.96
14.36 17.26
4 drops (14.47 mL) 9.93
15.59 2.35
4 drops (14.70 mL)
10.4
7 14.80 0.61
4 drops (14.90 mL)
10.6
6
15.84
11.1
1
16.87
11.3
8
17.83
11.4
3
18.86
11.5
0

8. 0.00 5.00 10.00 15.00 20.00 25.00
2.00
4.00
6.00
8.00
10.00
12.00
FIG. 2. Titration Curve (pH vs. mL titrant added) for Acetylsalicicylic Acid titrated with
0.1273M NaOH
Volume Titrant Added (mL)
Fig. I. Shows a plot of pH vs. mL of Titrant added. The equilavalence point can be estimated
from the inflection point in the titration curve.

9. 13.0 13.2 13.4 13.6 13.8 14.0 14.2 14.4 14.6 14.8 15.0
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00
16.00
18.00
FIG. 3. First Devirative ∆pH/∆V vs. V (corrected volume of titrant Added (mL))
Corrected Volume of Titrant added (mL)
Fig. 2. Shows a plot of ∆pH/∆V vs. v. corrected volume of titrant added (mL). The equivalence
point can be estimated from the peak in the curve which corresponds to the point of inflection in
Fig. 1.
∆pH/∆V

10. 13.0 13.2 13.4 13.6 13.8 14.0 14.2 14.4 14.6 14.8 15.0
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00
16.00
18.00
FIG. 3. First Devirative ∆pH/∆V vs. V (corrected volume of titrant Added (mL))
Corrected Volume of Titrant added (mL)
Fig. 2. Shows a plot of ∆pH/∆V vs. v. corrected volume of titrant added (mL). The equivalence
point can be estimated from the peak in the curve which corresponds to the point of inflection in
Fig. 1.
∆pH/∆V