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Weak Acid Lab 1

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Determination of the First Dissociation Constant of Weak Acids
SCHA-266
Experiment 1
Apparatus:
pH meter with electrode
sodium hydroxide solution, about 0.1 M (standardized)
unknown solution of a weak acid (100mL of approx. solution)
25-mL pipet
buret, buret clamp, and ring stand
150-mL beaker
250-mL beakers (3)
standard buffer solutions, pH 4, 7 and 10
General Theory
According to the Brønsted-Lowry acid-base theory, the strength of an acid is related to its ability to
donate protons. All acid-base reactions are then competitions between bases of various strengths for
these protons. For example, the strong acid HCl reacts with water according to Equation [1]:

HCl(aq) + H2O(l)
H3O+(aq) + Cl-(aq)
[1]
This is a strong acid and is completely dissociated (in other words, 100 percent dissociated) in dilute
aqueous solution. Consequently, the [H3O+] concentration of a 0.1 M HCl solution is 0.1 M. Thus HCl is
a stronger acid than water and completely donates a proton to water to form H3O+.
By contrast, acetic acid, HC2H3O2 (abbreviated HOAc), is a weak acid and is only slightly dissociated, as
shown in Equation [2]:
H3O+(aq) + OAc-(aq)
H2O(l) + HOAc(aq)
[2]
Its acid dissociation constant, as shown by Equation [3], is therefore small:

Ka =
[ H 3 O ][ OAc

[ HOAc ]
]
= 1.8 x 10-5
[3]
Acetic acid only partially dissociates in aqueous solution, and an appreciable quantity of undissociated
acetic acid remains in solution.
For the general weak acid HA the dissociation reaction and dissociation expression are:
H3O+(aq) + A-(aq)
HA(aq) + H2O(l)

Ka =
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[4]

[ H 3 O ][ A ]
[5]
[ HA ]
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Weak Monoprotic Acid
Recall that pH is defined as:
–log [H3O+] = pH
[6]
Solving Equation [5] for [H3O+] and substituting this quantity into Equation [6] yields:
[H3O+] = Ka
[ HA ]
[7]

[A ]
–log [H3O+] = –log Ka – log
[ HA ]

[8]
[A ]
pH = pKa – log
[ HA ]

[9]
[A ]
where pKa = –log Ka
If we titrate the weak acid HA with a base, there will be a point in the titration at which the number of
equivalents of base is just one-half the number of equivalents of acid present (at ½ Ve). This is the point at
which 50 percent of the acid has been titrated to produce A- and 50 percent remains as HA. At this point
[HA] = [A-], the ratio [HA]/[A-] = 1, and log [HA]/[A-] = 0. Hence, at this point in a titration, that is, at
one-half the equivalence point, Equation [9] becomes
pH = pKa
@
½ Ve
[10]
By titrating a weak acid with a strong base and recording the pH versus the volume of base added, one
can determine the dissociation constant, Ka, of the weak acid. From the resultant titration curve we
obtain the dissociation constant, as explained in the following paragraph.
From the titration curve (Figure 1), we see that at the point denoted as half the equivalence point, where
[HA] = [A-], the pH is 4.3. Thus, from Equation [10], at this point pH = pKa, or
pKa = 4.30
–log Ka = 4.30
log Ka = – 4.30
Ka = 5.0 x 10-5
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Weak Monoprotic Acid
Equivalence point
12
8
6
½ equivalence point
pKa = 4.30
pH 
4
2
3.7 mL
7.4 mL
0
0
2
4
6
Volume of base added, mL 
8
10
12
Figure 1. Titration curve for the titration of the weak acid HA with a base.
(Your instructor will show you a graphical method for locating the equivalence point on your titration
curves.)
Operation of the pH Meter
In order to measure the pH during the course of the titration, we shall use an electronic instrument called
a pH meter. This device consists of a meter and two electrodes.
The main variations among different pH meters involve the positions of the control knobs and the types
of electrodes and electrode-mounting devices. The measurement of pH requires two electrodes: a sensing
electrode that is sensitive to [H3O+] concentrations, and a reference electrode. This is because the pH
meter is really just a voltmeter that measures the electrical potential of a solution.
The reference electrode is just an electrode that develops a known potential that is essentially
independent of the contents of the solution into which it is placed. The glass membrane electrode is
sensitive to the H3O+ concentration of the solution into which it is placed; its potential is a function of
[H3O+]. It operates by mobility of H3O+ ions in the glass membrane. This can be more precisely
described, but for our purposes here it is sufficient for you to understand that two electrodes are required.
These two electrodes are sometimes combined into an electrode called a combination electrode, which
appears to be a single electrode. The combination electrode, however, does contain both a reference and
a sensing electrode.
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Preliminary Operations with the pH meter
1. Obtain standard buffer solutions of known pH. (Try to conserve these. Use scintillation vials
to hold these solutions.)
2. Plug in the pH meter to line current. Our meter will not require a warm up time. It should be
left plugged in until you are completely finished with it however.
3. The meters used in the lab can read in units of mV or pH. Select the pH mode.
4. Prepare the electrodes. Make certain that the solution in the reference electrode extends well
above the internal electrode. If it does not, ask your instructor to fill it with saturated KCl
solution. Slide down the rubber collar on the reference electrode. Rinse the outsides well with
distilled water. (Our electrodes now are gel filled and require no maintenance but your solution
should be up to the white line which marks the reference junction)
5. Standardize the pH meter. Set the pH meter control to correct room temperature. Carefully
immerse the electrodes into the first buffer solution in a small beaker. You should select the two
buffers that will come close to the values you will measure (4 and 7 are a good starting pair).
Remember that the glass electrode is very fragile; it breaks easily!
Don’t touch the bottom of the beaker with the electrode surface!!
Set the function to “read” or “pH”. The meter will adjust itself to the value of the buffer you
have placed it in. When the meter indicates ready then hit yes, rinse the electrode with
distilled water and move it to the next buffer. When this value is obtained enter yes. Record
the slope that is briefly displayed. . The pH meter is now ready to measure pH.
RECORD ALL DATA DIRECTLY INTO YOUR NOTEBOOK!
Procedure
A. Determination of pKa of Unknown Acid.
With the aid of a pipet bulb, pipet a 25.00 mL aliquot of your unknown acid solution into a 250-mL
beaker and carefully immerse the previously rinsed electrode into this solution. Measure the pH of this
solution. Record the pH in your notebook. Begin your titration by adding 1 mL of your standardized
base from a buret and record the volume of titrant and pH. Repeat with successive additions of 1 mL of
base until you approach the end point; then add 0.1-mL increments of base and record the pH and
milliliters of NaOH added. For the first run plot your titration in your lab book as it progresses. When
the pH has gone through a major jump in value you may discontinue data collection. From these data
plot a titration curve of pH versus mL titrant added. Repeat the titration with two more 25-mL aliquots
of your unknown acid and plot the titration curves. From these curves calculate the ionization constant.
Time may be saved if the first titration is run with larger-volume increments of the titrant to locate an
approximate equivalence point; then the second and third titrations may be run with the small increments
indicated above. Once you know your approximate endpoint you might wish to collect higher density of
points in the Ve/2 region.
B. Concentration of Unknown Acid
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Using the volume at the equivalence point, calculate the concentration of the unknown acid, and record
this in your notebook.
Prelab Questions
Before beginning this experiment in the laboratory, you should be able to answer the following
questions:
1. Define Brønsted-Lowry acids and bases.
2. Why isn’t the pH at the equivalence point always equal to 7 in a neutralization titration? When would
it be 7?
3. What is the pKa of an acid whose Ka is 2.4 x 10-5?
4. What is a buffer solution?
5. The pH at one-half the equivalence point in an acid-base titration was found to be 5.8. What is the
value of Ka for this unknown acid?
6. If 30.15 mL of 0.0995 M NaOH is required to neutralize 0.180 g of an unknown acid, HA, what is the
molecular weight of the unknown acid?
Report
1. Prepare and abstract for this lab to include with your report.
2. Provide properly labeled Data Tables for the data you have collected for the experiment.
3. Provide Excel plots of the titrations, properly labeled as Figures.
4. Provide first and second derivative plots of the three titrations. See Table 11-3 (page 210) in
your text for how to construct these plots. These will be additional figures for your report.
5. Provide a Results Table with the endpoints and pKa values determined.
6. Give a brief Results and Discussion section. Propose an acid that might be your unknown and
explain the difference from the book pKa value.
7. The answers to the above questions should be in your report. They may be on your lab notebook
pages but they must be readable.
8. A copy of the notebook pages should be included. They should contain one plot. (as above)
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