Determination of Ionization Constant
Experiment #4
04/06/2023
Grant Rudnick
Alberto Perez and Kathleen Romero Martinez
CHEM 1225L-027
Experimental Procedure
1. Gather all required tools for the lab including a Lab Quest 2, a pH probe, a burette, a
ring stand, a 3-prong clamp, a stir plate, and a laptop. Make sure that a stir bar, stir bar
retriever, and Kim wipes are nearby. Also, obtain the unknown solution.
2. Using litmus paper, determine if the unknown solution is acidic or basic. Gently drop a
little bit onto the litmus paper and if it turns red, it is acidic, and if it turns blue it’s basic.
3. Start by performing a dirty titration.
a. Pick a titrant for the titration; if the sample is acidic, use NaOH as the titrant, and
if the sample is basic, use HCl as the titrant (0.07 M).
b. Pick an indicator for the titration, if the sample is acidic, use phenolphthalein,
and if the sample is basic, use bromothymol blue.
c. Add ~10 mL of the unknown to a beaker placed below the burette.
d. Prepare ~15 mL of the titrant in the burette. Use a funnel lifted slightly above the
top of the burette and be careful not to spill.
e. Place about 3 drops of the indicator into the burette with the unknown.
f. Begin the dirty titration and slowly add the titrant to the solution.
g. Record approximately the milliliters of titrant needed for the equilibrium point
to be reached.
h. If the mL used is less than 15 before the change occurred, use a less
concentrated titrant. If it is more than 15, use a more concentrated titrant.
4. Collect the following chemicals from the fume hood: pH 7 buffer, pH 4 buffer (only if the
sample is acidic), and pH 10 buffer (only if the sample is basic).
5. Prepare the burette on the ring stand and attach the pH probe to the stand to be used.
6. Calibrate the pH electrode using the pH buffers using the following steps:
a. Open LoggerPro and go to “Open Folder”. Click on “Advanced Chemistry with
Vernier” and then “Experiment 7A”. The border should turn yellow.
b. Clean the tip of the probe with DI water and blot with a Kim wipe (do NOT wipe).
c. On LoggerPro, go to experiment, calibrate, pH. Make sure that it’s set to 2-point
calibration, then click “calibrate now”.
d. Place the pH electrode into the pH 7 buffer and gently swirl and shake the
beaker. When the voltage stabilizes, enter pH 7. Keep the data.
e. Clean the tip of the probe.
f. Place the pH electrode in the pH 4 or 7 buffer and gently shake it. Then when the
voltage stabilizes, enter pH 4 or 10, respectively.
g. If the pH reads the correct pH of the buffer is in it, then the calibration was
successful.
7. Run the titration using the following steps:
a. Add 20 mL of the unknown to the beaker.
b. Add the sir bar, turn on the stir plate, and add enough DI water so the pH probe
is covered with liquid. Make sure that the stir bar doesn’t hit the probe.
c. Fill the burette with ~30 mL of the titrant and record the starting point.
d. Collect the first data at 0 mL of the titrant. Always press “keep” so the data is
cumulative.
e. Add titrant at ~1 mL increments. As the graph gets closer to the equivalence
point, make smaller increments for a more accurate graph.
f. Record all the values in the notebook too in case of any digital accidents.
8. Repeat the titration once more.
9. Using the first and second derivative graphs, the equivalence point can be determined.
10. Clean up the lab station.
Data Analysis:
When unknown A was tested with litmus paper, the paper turned blue, indicating a base.
A dirty titration was completed which demonstrated that it took 18 mL of 0.07 M HCl to titrate
15 mL of the unknown, so the concentration of HCl needed to be increased to 0.12 M.
Unknown + 0.12 HCl
Titration 1
First derivative
Second Derivative
The x-intercept of the second derivative graph occurs at 19.17 mL, so the first endpoint is at
19.17 mL. The half neutralization point occurs at half of the volume of the endpoint, so the half
neutralization point for titration one will be at 9.59 mL.
At the half neutralization point, pH is equal to pKa. The pH at 9.59 mL is 9.42.
(1) 𝐾𝑎 = 10−𝑝𝐾𝑎
(2) p𝐾𝑏 = 14 − 𝑝𝐾𝑎
Using equation 1,
𝐾𝑎 = 10−9.42 = 3.80 × 10−10
And using equation 2,
p𝐾𝑏 = 14 − 9.42 = 4.58
𝐾𝑏 = 10−4.57 = 2.63 × 10−5
Titration 2
First derivative
Unknown + 0.12 HCl
Second Derivative
The x-intercept of the second derivative graph occurs at 16.77 mL, so the first endpoint is at
16.77 mL. The half neutralization point occurs at half of the volume of the endpoint, so the half
neutralization point for titration one will be at 8.38 mL.
At the half neutralization point, pH is equal to pKa. The pH at 8.38 mL is 9.43.
Using equation 1,
𝐾𝑎 = 10−9.43 = 3.72 × 10−10
And using equation 2,
p𝐾𝑏 = 14 − 9.43 = 4.57
𝐾𝑏 = 10−4.57 = 2.69 × 10−5
The average Kb value is equal to
(2.69+2.63)×10−5
2
= 𝟐. 𝟔𝟔 × 𝟏𝟎−𝟓
Compared to the known bases, Ammonium hydroxide (Kb = 1.78x10^-5), aniline (Kb =
7.41x10^-10), and ethylamine (Kb = 4.47x10^-4), the unknown A is ammonium hydroxide.
To calculate the molarity of the unknown A, the moles of titrant used must be compared to the
moles of unknown. To find the moles of titrant used:
Trial 1: (0.01917 L)(0.12 M) = 0.0023 moles of titrant
Trial 2: ( 0.01677 L)(0.12 M) = 0.0020 moles of titrant
At the endpoint, the moles of the titrant are equal to the moles of the sample. Therefore, to
find the molarity of the sample, the moles of titrant (and unknown) are divided by the volume
of the sample:
Trail 1: (0.0023 moles)/(0.020 L) = 0.115 M sample
Trail 2: (0.0020 moles)/(0.020 L) = 0.101 M sample
Average molarity = (0.115 + 0.101) / 2 = 0.108 M
Summary of Experimental Results:
For unknown A, the sample was determined to be a base. The litmus turned blue when in
contact with the sample, indicating a base. In the dirty titration, 18 mL of 0.07 M HCl titrant was
used, and in the two trials, 16.77 and 19.17 mL of 0.12 mL HCl titrant were used to reach the
endpoint. In the dirty titration, 15 mL of unknown A was used and in the real titrations, 20 mL
of the unknown was used. The titration curves show a strong dip in pH around the endpoints,
and halfway back from the endpoint is the half-neutralization point, where pH = pKa.
Because the pH at the equivalence point was found to be 9.42 and 9.43, the Ka was calculated
with (1) to be 3.80 × 10−10 and 3.72 × 10−10 . Rearranging and using (2), the average Kb was
found to be 2.66 × 10−5 .
Experimental Results:
Discuss your selected molarity of titrant and the preparation of the sample. Was it a correct
decision? Was the choice of your volume interval appropriate?
The first titrant that was used was 0.07 M HCl. After about 20 mL of titrant was added, the trial
was given up because the ratio was too little, and the process was taking too long. In the dirty
titration, it was discovered that 17 mL of titrant was needed to reach the equivalence point, so
from that, it should have been determined that the concentration of titrant should have been
increased. For the last two trials, 0.12 M HCl was used as titrant and the reaction took place at
the correct rate.
The volume of titrant added provided good data for the graphs. Initially, 1 mL increments of
titrant were used, and as the graph approached the equivalence point, the increments were
decreased to 0.5 mL, 0.3 mL, and even 0.2 mL. This provided a pretty good image of what was
going on around the equivalence point while also maintaining a good amount of certainty in the
collected data.
Data and Observations: