CHM 152
Lab 8: Titration curve of a Weak Acid
Last updated Jan. 2012
Introduction
A titration curve plots the pH of a solution as a second solution (the titrant) is slowly added, usually via a
burette. Titration curves of weak acid or base solutions using a strong base or acid, respectively, are
especially useful as it can allow one to experimentally determine the acid’s dissociation constant, K a.
A typically acid-base titration curve is shown below. The equivalence point is the pH where the acid has
been completely converted to its conjugate (HA  A-). On a titration curve, it’s located on the part of the
graph that turns sharply upward, at the halfway point of this slope (see figure below). It shouldn’t be
confused with the endpoint of a titration, which usually refers to the point where the indicator changes
color. When choosing an indicator for an acid-base titration, the indicator’s endpoint should match the
predicted equivalence point of the titration as closely as possible.
The pH of the equivalence point is primarily determined on the type of salt that is formed during the
reaction. For example, a weak acid and strong base usually react to produce a basic salt, so the equivalence
point is predicted to have pH>7.
A second point of interest is the midpoint, where the concentrations of the acid and conjugate base are
present in equal concentrations, [HA] = [A-], making the acid/conjugate ratio of 1. At this ratio, the
solution pH should be equal to the pKa of the weak acid. On a titration curve, the midpoint is the halfway
point between the initial pH and the point where it the slope increases sharply.
14
12
10
8
Equivalence
point
pH
Midpoint
6
4
2
0
0
10
20
30
40
50
60
Volume of base added, mL
In this lab you will determine the identity of an unknown weak acid by titrating it with a sodium hydroxide,
a strong base. You will then plot a titration curve to find the acid’s dissociation constant as well as its
equivalence point.
Procedure
Part I: Preparing the burette
1. Use a burette clamp attached to a ring stand to hold a 50.0mL burette
2. Add approximately 20mL of 1 M NaOH to the burette.
3. Open the stopcock for a couple of seconds to rinse out the tip, and then slowly pour the remaining
solution down the drain with running water, rotating to ensure that the entire burette is rinsed.
4. Repeat steps 2 and 3 if needed.
Part II: Calibrating the pH meter
1. Fill a 50mL beaker with 20-30 mL of pH = 7 buffer.
2. Carefully remove the probe of the pH meter from the storage bottle and place it in the buffer solution.
3. After turning on the pH meter, it should read “MEAS” (for measure mode) at the top of the screen.
Switch to calibration mode (CAL), using the “cal/meas” button.
4. The meter will automatically “guess” what the pH of the calibration solution is. After a second the
bottom number on the meter should read “7.00.” Press the “Enter” button to accept this value. The top
number should now also read 7.00
5. Switch the meter back to MEAS mode. If the top number changes slightly (0.01-0.05), this is normal.
The meter should now be calibrated and ready for use. Leave the probe in the buffer solution until you’re
ready to begin the titration.
Part III: Titration of an unknown weak acid
1. Fill the burette with 1M NaOH to the 0.0mL mark.
2. Use a graduated cylinder to measure 25.0 mL of your unknown acid and transfer it to a 250 mL beaker.
3. Place the pH meter in the acid solution and record the initial pH on your data sheet (“0.0 mL NaOH
added”).
4. Add 50.0 mL of NaOH to the acid solution, in the increments shown on the data sheet. After each
addition, swirl to mix and record the pH. If an addition is different than what’s listed (e.g., 11.0 mL instead
of 10.0 mL), note the actual volume added.
6. Repeat steps 1-5 for a second determination
Part IV: Preparing a Titration Curve
1. Using Excel, prepare a graph of pH over milliliters of NaOH added for each set of data. Use an scatter
plot with smooth lines and markers.
2. Format the y-axis to show major tick marks in increments of 1 pH unit and minor tick marks in
increments of 0.2 pH units (right-click axis and select “Format Axis”).
Waste Disposal
Solution waste can be poured slowly down the drain with running water. Rinse the burette twice with
approximately 20 mL of water, similar to what was done at the start of lab. Don’t forget to drain a small
amount through the stopcock each time to rinse the tip.
Name: _____________________________
Section: ________
Data
Unknown code __________
Volume NaOH added, ml
0.0 (initial pH)
pH
(Trial #1)
__________
pH
(Trial #2)
__________
5.0
__________
__________
10.0
__________
__________
15.0
__________
__________
20.0
__________
__________
22.0
__________
__________
24.0
__________
__________
26.0
__________
__________
28.0
__________
__________
30.0
__________
__________
32.0
__________
__________
36.0
__________
__________
40.0
__________
__________
45.0
__________
__________
50.0
__________
__________
Equivalence point
__________
__________
Average
__________
Midpoint
__________
__________
pKa
__________
__________
Ka
__________
__________
Average Ka
__________
Include your titration curves with your lab report. Label the location of the midpoint and equivalence
point.
Name: _____________________________
Section: ________
Post-lab Questions
1. Based on your results, determine the identity of your unknown from the following list. Circle your
choice.
Ka
Hydrogen sulfate ion
1.3 x 10-2
Formic acid
1.7 x 10-4
Acetic acid
1.8 x 10-5
Hydrogen sulfite ion
5.6. x 10-8
Boric acid
5.8 x 10-10
2. Based on your results, which of the following indicators would work best for this titration? Explain.
Indicator
Thymol blue
Bromophenol blue
Chlorophenol blue
Bromothylmol blue
Cresol red
Phenolphthalein
Color in acidic
solution
Red
Yellow
Yellow
Yellow
Yellow
Colorless
Color in basic
solution
Yellow
Bluish purple
Red
Blue
Red
Red
pH range of
color change
1.2 – 2.8
3.0 – 4.6
4.8 – 6.4
6.0 – 7.6
7.2 – 8.8
8.3 – 10.0
3. Given your answer to question 1, is the salt that’s formed during your titration consistent with your
observed equivalence point? Explain.
4. How would each of the following errors affect your calculated value of K a (incorrectly high, low, or no
effect)? In each case, explain your answer.
a. The pH meter was not calibrated before beginning the titration and all pH readings were erroneously
low.
b. A student failed to read the burette correctly and volume measurements were erroneously high.
Name: _____________________________
Section: ________
Pre-lab Questions
1. Define the following.
a. Equivalence point
b. Midpoint
c. Endpoint
2. Using the Henderson-Hasselbach equation (see textbook or lecture notes), explain why the pH of the
midpoint is predicted to be the same value as the weak acid’s pKa.
3. Calculate the dissociation constant of a weak acid that has a pK a of 5.4.
4. If you were analyzing a diprotic acid (H2A), how would the titration curve differ from the one shown in
the introduction section of this lab?
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