Cuyamaca College Weak Acid Titration Experiment Lab Report

EXPERIMENT: Weak Acid Titration ExperimentIntroduction:
In this experiment you will investigate what happens when a weak acid is titrated using a
strong base (NaOH). Titration curves are often used to visualize what happens when an acid is
titrated with a base, they are a graph where pH is on the y-axis and volume of titrant is on the
x-axis. We can break down the titration curve of a weak acid with a strong base into four distinct
regions: (1), the initial point (a weak acid solution), (2) the buffer region, (3) at the equivalence
point and (4) after the equivalence point (excess base region). The following titration curve is an
example of when 20 mL of 0.1 M HC2H3O2 is titrated with 0.1 M NaOH.
Acetic Acid Titrated with NaOH
Excess Base Region
(weak acid has been completely neutralized)
14
Initial Point
(Only Weak Acid Present)
12
10
Equivalence Point
(only conj. base left in solution)
½ Equivalence Point
(pH = pKa)
pH
8
6
4
2
0
0
5
10
Buffer Region
(buffer formed in solution)
15
20
25
30
35
40
Volume of NaOH added (mL)
At the initial point, the pH of the solution is due only to the weak acid present in the
solution. Which means the pH is dependent on the Ka value of the weak acid and its
concentration. The pH can be solved for using an ICE table.
Once NaOH has been added there are two reactions involved: the neutralization reaction
of the weak acid and the NaOH, and the weak acid dissociation equilibrium. If the NaOH is the
limiting reactant in the neutralization reaction, the resulting solution is typically a buffer. Below
is the typical neutralization reaction:
𝐻𝐴(𝑎𝑞) + 𝑁𝑎𝑂𝐻(𝑎𝑞) → 𝐻2 𝑂(𝑙) + 𝑁𝑎𝐴(𝑎𝑞)
If calculating the pH in the buffer region, the first step is to determine the concentrations of the
weak acid (HA) and the conjugate base (A- or NaA) after the neutralization reaction. After that
the use of the Henderson-Hasselbach Equation to calculate pH, seen below.
[𝐴− ]
𝑝𝐻 = 𝑝𝐾𝑎 + log
[𝐻𝐴]
At the equivalence point, there are equal moles of the weak acid and NaOH initially. This
means that after the neutralization reaction there will be no NaOH or HA remaining. All that will
be left in the solution will be water (H2O) and the conjugate base (A- or NaA). This means that
the pH in the solution will be dependent on the Kb of the conjugate base and its concentration.
Finally, in the excess base region of the pH curve the NaOH becomes the excess reactant,
and the weak acid (HA) is the limiting reactant in the neutralization reaction. After the
neutralization reaction the solution will contain NaOH and weak conjugate base. In most cases
the weak base will not affect the pH in the solution to any great extent, as the NaOH will
contribute a significant amount of [OH-] to the solution. This will decrease the amount of OHformed from the weak base in solution. For this reason, the pH in the solution at this point can be
calculated from the concentration of NaOH in the solution after the reaction with the weak acid.
Experimental Procedure:
Experimental Procedure: Titration of an Unknown Weak Acid with NaOH
1. Using a clean 400 mL beaker obtain approximately 100 mL of standardized NaOH
solution, be sure to record the concentration of the NaOH.
2. Condition a buret, using water first and then again with NaOH solution. Your
instructor should briefly cover how to do this.
3. Once calibrated, fill the buret with NaOH solution and record the initial volume (it
should not be 0.00 mL for the initial).
4. Your lab instructor will assign an unknown acid to you, record this number. Using a
clean 50 mL beaker transfer about 25 mL of this unknown acid into the beaker.
Record the concentration from the container.
5. Condition a 10 mL pipette using water, and then using the unknown acid. Your
instructor should briefly cover how to do this.
6. Using the 10 mL pipette transfer 20 mL of unknown acid solution into a 125 mL
Erlenmeyer Flask. Add 2-3 drops of phenolphthalein indicator.
7. Using the buret, titrate the unknown acid solution in the Erlenmeyer flask with the
NaOH solution. Record the initial volume on the buret and the final volume after
reaching the end point.
8. Dispose of the solution in the Erlenmeyer flask in the waste container.
9. Next, we will perform a titration but without any indicator, instead we will measure
the pH at different points in the titration. Calibrate a pH meter and record the pH
meter number on your data sheet. Using the same beaker retrieve approximately
25 mL of unknown acid again. Then using the conditioned pipette, transfer 20 mL to
a clean 150 mL beaker. Steps 9 thru 13 will all be for the same titration.
10. Measure the pH of the solution in the beaker before adding any NaOH.
11. Next measure the pH after adding approximately 2 mL of NaOH, continue adding
NaOH at an increment of 2 mL and measuring the pH until you have added about
16 mL total. Make sure to record the pH and the actual volume off the buret
after EACH addition of NaOH.
12. After adding a total of about 16 mL of NaOH, decrease the increment you are
measuring at to 1 mL NaOH. Measure the pH at a 1 mL increment until you reach a
total volume of 24 mL. Make sure to record the pH and the actual volume off the
buret after EACH addition of NaOH.
13. After adding a total of about 24 mL of NaOH, increase the increment you are
measuring at back to 2 mL NaOH. Measure the pH at a 2 mL increment until you
reach a total volume of 34 mL. Make sure to record the pH and the actual volume
off the buret after EACH addition of NaOH.
Experimental Procedure: Excel Simulations
1. Using Microsoft excel, we will simulate titration curves for two acids.
2. The unknown acids tested in the first part of the experiment were either:
a. acetic acid, Ka = 1.75 x 10-5
b. potassium biphthalate, Ka = 3.91 x 10-6
3. Using the concentrations of NaOH and Unknown acid from the titrations done in the
first part of the experiment, calculate the pH at 0.00 mL of NaOH and at each
1.00 mL of NaOH added until reaching 35.00 mL, for each of the two acids above.
4. Start by making a column in excel titled Volume of NaOH added, and in the first row
add 0.00, followed by 1.00 mL, then 2.00 mL and so on until you reach 35.00 mL.
5. Include cells that contain the following information: The Ka value for the acid being
simulated, the concentration of NaOH and the concentration of unknown acid.
6. Use your knowledge of how to calculate the pH of a weak acid in each of the regions
of a weak acid strong base titration, to write formulas in excel that will calculate the
pH for each volume of NaOH added.
a. Region 1: Initial Point
b. Region 2: Buffer Region
c. Region 3: Equivalence Point
d. Region 4: Excess Base Region
7. Calculations must be done using an Excel sheet, a main goal of this
experiment/assignment is to help you become more familiar using Microsoft Excel.
8. Once both simulations have been created in Excel, you need to make a graph with
three data series in it. On one graph include the following series of data:
a. pH of acetic acid Vs. mL NaOH added (Simulation).
b. pH of potassium biphthalate Vs. mL of NaOH added (Simulation).
c. pH of unknown acid Vs. mL of NaOH added (from experiment).
9. The graph if done right, should allow you to clearly identify the unknown acid.
Name: _______________________
Experiment: Weak Acid Titration
Data and Calculations:
A
Equivalence Point Measurement:
1.98
Initial Buret Reading:
_____________
Unknown Number: _________
30.99
_____________
0.0984
Concentration of Acid: ___________
Measured Volume of NaOH
to Reach Eqv. Point:
_____________
Calculated Eqv. Point: ___________
Final Buret Reading:
Calculations for Equivalence Point:
0.0991
Concentration of NaOH: __________
51
pH Titration Curve Measurement:
1.50
Initial Buret Reading: _____________
pH Meter#: _________
Unknown Acid Identity: ___________________
d
Buret Reading
d
.50
_____________
_____________
6.55
Actual Volume
pH
Buret Reading
Actual Volume
____________
0
2.81
___________
_____________
33.29
____________
___________
____________
5.05
3.83
___________
34.31
_____________
____________
32.81
___________
_____________
8.35
6.851
____________
___________
4.06
37.19
_____________
____________
35.69
11.47
___________
_____________
10.31
____________
8.81
___________
4.06
_____________
____________
___________
_____________
12.32
____________
10.82
4.29
___________
_____________
____________
___________
_____________
14.57
____________
13.07
___________
4.31
_____________
____________
___________
16.23
_____________
14.73
____________
4.47
___________
_____________
____________
___________
18.32
_____________
____________
16.82
___________
4.60
_____________
____________
___________
20.22
_____________
____________
18.72
___________
4.67
_____________
____________
___________
22.45
_____________
20.95
____________
4.87
___________
_____________
____________
___________
26.30
_____________
24.8
____________
5.25
___________
_____________
____________
___________
28.51
_____________
27.01
____________
5.36
___________
_____________
____________
___________
29.18
_____________
27.68
____________
___________
5.38
_____________
____________
___________
30.33
_____________
____________
28.83
5.63
___________
_____________
____________
___________
31.41
_____________
29.9
____________
___________
5.96
_____________
____________
___________
31.79
pH
10.70
11.20
Post-Lab Assignment
1. Give an explanation why the Henderson-Hasselbalch equation can’t be used at the
beginning and at the equivalence point of the titration.
2. In the titration of a 35 mL of 0.345 M weak base being titrated by 0.465 M HCl
determine, the Kb for the weak base is 3.30 x 10-5:
a. The pH at the initial point.
pH___________
b. The pH after 12.3 mL of HCl has been added.
pH___________
c. The pH at the equivalence point.
pH___________
d. The pH after 38.4 mL of HCl has been added.
pH___________