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8/10/2019 identifying of unknown monoprotic acid
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4.0 THEORY
Our second experiment is identifying an unknown weak acid based on the Ka value obtained
from titration method and half volume method. We used these two methods in order to
identify the unknown acid and at the end of the experiment we had determined which method
is more accurate.
Theoretically, strength of and acid depends on its ability to lose proton (H +). Weak acids
usually will partially dissociated and release some of its hydrogen ion in the solution. Hence
its ability to donate proton is much lower than strong acid as it partially dissociated in water.
Therefore, these kind of acids have higher pKa than strong acids as strong acids completely
dissociated in water and release all the hydrogen atoms.
Ka, acid dissociation constant is defined asthe quantitative measure of the strength of and acid
in a solution. pKa, is equal to log10 Ka.Ka is the equilibrium constant for a chemical
reaction known as dissociation in the context of acid-base reactions. Symbolically,
equilibrium can be written as below :
HA (aq) + H2O (l) H3O+ (aq) + A- (aq)
Under such equilibrium conditions, the total concentrations of each species remain constant
even though the species in solution are constantly dissociating and recombining.
Monoprotic acid is an acid that donates only one proton or hydrogen atom per molecule to an
aqueous solution and will only have one equivalence point. As for polyprotic acid, literally it
corresponds to the transfer of more than one proton.
The ionization constant of weak monoprotic acid can be calculated as below :
Ka = [H+] [A-]
[HA]
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As we all know that the greater the value of Ka the higher the formation of H +is favoured so
that the pH of the solution become lower. Theoretically the Ka of weak acid varies between
1.810-16and 55.5. Acids that are lesser than that value are said to be weaker acids than water
itself.
There is one other way that can be used to identify an unknown acid by determining its
percent of dissociation which is symbolized as (alpha) and which it ranges from 0%
100%.Percent dissociation can be defined as :
= [A-]
[A-] + [HA]
As is not constant so it does not depends on the value of [HA]. Generally will increase as
[HA] decreases hence acids become stronger as they are diluted. Each proton will have Ka as
it said to be polyprotic. In other words percentage of ionization refers to the proportion of
neutral particles, such as those in a gas or aqueous solution which are later ionized into
charged particles.
The acid ionization constant, Ka can be determined from the titration curve. The pH values at
particular volume of sodium hydroxide being added can be obtained from the titration curve
so that the pKa values can be obtained from the plotted titration curve. As the data obtained,
two titration curves of pH versus volume of base added were plotted then the equivalence
point was obtained.
The pH values obtained from the titration curve can be used to calculate the acid ionization
constant hence the average value can be determined. Plus, the volume of base titrated at the at
the equivalence point is used to calculate the concentration of the unknown monoprotic acid.
The average is determined as well and the average acid ionization constant later will be used
to identify the unknown monoprotic acid.
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Apparatus and Materials
1. pH meter and electrode
2. 250 ml beaker
3.
40 ml of unknown acid
4. Distilled water
5.
Burette
6. Retort stand
7. Magnetic stir plate
8.
Magnetic stirrer
9. Graduated measuring cylinder
10.
NaOH solution
11.Erlenmeyer flask
Burette
Beaker
Unknown
acid
Magnetic stirrer
Magnetic plate
pH electrode
Retort stand
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Half volume method
1.
The burette was cleaned, rinsed and filled with NaOH solution.
2.
An unknown acid of 40 ml was measured. 100 ml of distilled water measured and
poured into a 250 ml of Erlenmeyer flask. The unknown acid was dissolved in
distilled water.
3. The solution we prepared in step 2 was divided into two parts using graduated
measuring cylinder.
4.
The solution prepared then titrated with NaOH until pH 12 was reached.
5. The titrated solution then was mixed with the solution which has not been titrated. The
value of Ka was once again determined from the [H+] value which was linked with the
pH of the solution.
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7.0 RESULTS
7.1 Titration of the unknown monoprotic acid with 0.1 M sodium hydroxide, NaOH.
Burette Reading Titration 1 Titration 2 Titration 3
0 3.76 3.66 3.711 3.90 3.81 3.86
2 4.06 3.96 4.01
3 4.21 4.08 4.11
4 4.32 4.19 4.20
5 4.42 4.29 4.40
6 4.52 4.40 4.50
7 4.62 4.49 4.55
8 4.70 4.56 4.60
9 4.79 4.65 4.71
10 4.87 4.72 4.8011 4.96 4.80 4.96
12 5.05 4.88 5.00
13 5.13 4.96 5.12
14 5.25 5.04 5.20
15 5.37 5.15 5.27
16 5.52 5.25 5.50
17 5.70 5.39 5.60
18 6.01 5.54 5.96
19 6.83 5.73 6.45
20 10.23 6.02 7.13
21 10.88 7.10 10.3622 11.02 10.01 10.80
23 11.11 10.52 11.10
24 11.19 10.76 11.19
25 11.24 10.89 11.23
26 11.29 11.01 11.25
27 11.36 11.08 11.32
28 11.40 11.14 11.41
29 11.42 11.19 11.45
30 11.45 11.23 11.47
31 11.47 11.27 11.4932 11.49 11.30 11.52
33 11.51 11.33 11.55
34 11.53 11.35 11.56
35 11.55 11.38 11.62
36 11.56 11.43 11.63
37 11.60 11.55 11.66
38 11.64 11.61 11.68
39 11.69 11.65 11.70
40 11.74 11.72 11.73
41 11.77 11.78 11.80
42 11.81 11.83 11.82
43 11.87 11.90 11.89
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44 11.92 11.92 11.90
45 11.96 11.95 11.93
46 12.01 11.99 12.02
Table 7.1
7.2 Half method titration of the unknown monoprotic acid with 0.1 M sodium
hydroxide, NaOH.
Burette Reading Titration 1 Titration 2 Titration 3
0 3.50 3.51 3.51
1 3.88 3.90 3.89
2 4.21 4.17 4.20
3 4.44 4.42 4.41
4 4.65 4.61 4.63
5 4.87 4.79 4.85
6 5.23 4.99 5.35
7 5.47 5.23 5.39
8 6.52 5.70 6.11
9 10.31 7.50 8.91
10 10.66 10.41 10.56
11 10.86 10.76 10.81
12 10.97 10.91 10.94
13 11.07 11.03 11.05
14 11.14 11.11 11.12
15 11.20 11.17 11.1916 11.26 11.23 11.24
17 11.31 11.28 11.30
18 11.35 11.32 11.33
19 11.39 11.36 11.37
20 11.42 11.39 11.40
21 11.45 11.43 11.44
22 11.49 11.46 11.47
23 11.52 11.51 11.50
24 11.66 11.63 11.64
25 11.79 11.75 11.7626 11.87 11.86 11.85
27 11.93 11.91 11.92
28 11.96 11.97 11.98
29 12.00 12.01 12.00
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7.3 Graph for the titration of the unknown monoprotic acid with 0.1 M sodium
hydroxide, NaOH.
7.3.1 Titration 1
0
2
4
6
8
10
12
14
0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46
pHmonoproticacid
Volume NaOH, ml
1) pH monoprotic acid Vs. Volume NaOH, ml
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7.3.2 Titration 2
0
2
4
6
8
10
12
14
0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46
pHm
onoproticacid
Volume NaOH, ml
2) pH monoprotic acid Vs. Volume NaOH, ml
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7.3.3 Titration 3
0
2
4
6
8
10
12
14
0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46
pHmonoproticacid
Volume NaOH, ml
3) pH monoprotic acid Vs. Volume NaOH, ml
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7.4 Graph for the half method titration of the unknown monoprotic acid with 0.1 M
sodium hydroxide, NaOH.
7.4.1 Titration 1
7.4.2 Titration 2
7.4.3 Titration 3
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8.0 CALCULATIONS
8.1 The titration of the unknown monoprotic acid with 0.1 M sodium hydroxide,
NaOH
8.1.1 Titration 1
By applying Henderson Hesselbalch Equation,
pKa = -log Ka + log
, since within the buffer region, the concentration of base reacted
with the concentration of acid was equal, so, log
= 0
Thus,
pKa = -log Ka
4.85 = -log Ka
Ka = -antilog 4.85
0
2
4
6
8
10
12
14
0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46
pHmonoproticacid
Volume NaOH, ml
1) pH monoprotic acid Vs. Volume NaOH, ml
19.50 ml
pKa = 4.85
9.75 ml
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= 1.41 x 10-5
8.1.2 Titration 2
Applying the Henderson Hesselbalch Equation
pKa = -log Ka
4.80 = -log Ka
Ka = -antilog 4.8
= 1.58 x 10-5
0
2
4
6
8
10
12
14
0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46
pHmonoproticacid
Volume NaOH, ml
2) pH monoprotic acid Vs. Volume NaOH, ml
21.40 ml
pKa = 4.80
10.70 ml
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8.1.3 Titration 3
Applying the Henderson Hesselbalch Equation
pKa = -log Ka
4.80 = -log Ka
Ka = -antilog 4.8
= 1.58 x 10-5
8.2 The half method titration of the unknown monoprotic acid with 0.1 M sodium
hydroxide, NaOH
8.2.1 Titration 1
Applying the Henderson Hesselbalch Equation
0
2
4
6
8
10
12
14
0 2 4 6 8 10 12 14 16 18 20 22 24 26 28 30 32 34 36 38 40 42 44 46
pHmono
proticacid
Volume NaOH, ml
3) pH monoprotic acid Vs. Volume NaOH, ml
20.50 ml
10.25 ml
pKa = 4.80
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pKa = -log Ka
4.65 = -log Ka
Ka = -antilog 4.65
= 2.24 x 10-5
8.2.2 Titration 2
Applying the Henderson Hesselbalch Equation
pKa = -log Ka
4.70 = -log Ka
Ka = -antilog 4.70
= 2.00 x 10-5
8.2.3 Titration 3
Applying the Henderson Hesselbalch Equation
pKa = -log Ka
4.70 = -log Ka
Ka = -antilog 4.70
= 2.00 x 10-5
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8.3 Titration table for the titration of the unknown monoprotic acid with 0.1 M
sodium hydroxide, NaOH
Titration
No.
Volume Of
NaOH at
equivalence
point, ml
pH at
equivalence
point
Volume of
NaOH
halfway to
equivalence
point, ml
pKa value Ka value
1 19.50 8.40 9.75 4.85 1.41 x 10-5
2 21.40 8.30 10.70 4.80 1.58 x 10-5
3 20.50 8.82 10.25 4.80 1.58 x 10-5
Average Ka 1.52 x 10-5
8.4 Titration table for half method titration of the unknown monoprotic acid with 0.1
M sodium hydroxide, NaOH
Titration
No.
Volume Of
NaOH at
equivalence
point, ml
pH at
equivalence
point
Volume of
NaOH
halfway to
equivalence
point, ml
pKa value Ka value
1 8.25 8.20 4.125 4.65 2.24 x 10-5
2 9.25 8.20 4.625 4.70 2.00 x 10-5
3 8.75 8.25 4.375 4.70 2.00 x 10-5
Average Ka 2.08 x 10-5
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9.0 Discussion
In this experiment, we have been assigned to carry out two types of method that focused on
the identifying of unknown monoprotic acid; titration curve method and half method. In the
expression, Ka is the acid dissociation constant. Strong acids typically dissociate completely,
and therefore would have a Ka value of greater than 1. Weak acids have Ka values much
smaller than 1 (typically less than 10-4).
For the first method, the titration of as much as 50 ml of NaOH was titrated into 20 ml of
unknown weak acid had been conducted until up to three titration. All of those titrations were
left to be until the pH value reached around 12. In the interest of identifying the unknown
weak acid that we worked on, we need to know the Ka value of the acid from each of the
three titration curves (D.C. Harris, 2004).
What we knew about the titration curve was that there was a buffer region that located right
along the way till to the equivalence point. Buffer region is a region that we added just
enough bases in the acidic solution that we have some conjugate base in the solution, and we
also have some acidic solution (D.C. Harris, 2004). The Ka value could be obtained from this
region.
In the simplest way, the Ka value from the titration curve can be obtained from the value of
pKa by applying the Henderson Hesselbalch Equation (A. Marie, 2014). pKa value in each
titration curve was exactly equal to the pH halfway to the equivalence point.
The pKa value had noticed to be 4.85, for titration 1, while titration 2, and 3 had recognized
the same pKa value of 4.80. As it went into the calculations, the Ka value of the acid had
accordingly became 1.41 x 10-5for titration 1 and the following titrations had seen an equal
Ka value of 1.58 x 10-5. The total average of these Kas was 1.52 x 10 -5. This value had not
been exactly equal but quite close to that of propanoic acid, C3H6O2(1.3 x 10-5) (Table of
Ka, 2010).
Meanwhile, from the half volume method, the pKa value of 4.65 for the titration 1was 4.65,
and for both of titration 2 and 3 the value was recorded to be 4.70. After using the Henderson
Hesselbalch Equation, the Ka value had been showing the corresponding value match the
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recognizable unknown acid. The Ka for first titration was 2.24 x 10 -5 while the second and the
third titration had seen an equal Ka value of 2.00 x 10-5. The average Ka for three titrations of
half volume method, was 2.08 x 10-5. So, from our sources table of Ka, the close value Ka of
unknown acid from our experiment of half volume method was falling closer toward the Ka
value of Crotonic acid (HC4H5O2).
From the both method, we identified the unknown acid that used in this experiment but we
got different weak acid. For first method the unknown acid identified as propanoic acid while
for the second method as a crotonic acid. The difference value pKa of first method and
second method just around 0.14. May errors preparation occur during the experiment caused a
difference in determining the unknown acid.
The deviation of the experimentally obtained average Ka value from any other close
Ka values from the Ka table might due to several errors that had incidentally associated
during the completion of this experiment. Mainly, it might due to the fact that the parallax
error had occurred. In a way, the volume on the burette was misread. This can be a parallax
mistake when the observer looking at the volumes on the incorrect angle. In addition, we did
not know the exact concentration of the NaOH we used in this experiment, as we had not
standardized the basic solution in prior to the execution of experiment. So, it could be that
error might have resulted from the uses of incorrect concentration of NaOH. Besides, it could
be that during the titration, there was several times where the solution had been transferred in
excess volumes by incorrectly leveling the scale on the burette.
10.0 Conclusion
The titration curves in this experiment have shown a consistent shape pattern for all
three titration of the both method with each of these curves have recognized slight differencesin pH value in the equivalence point. In the interest of identifying the monoprotic acid, the Ka
value of the acid was determined from the pHs that correspond to the volume of bases
halfway to the equivalence point from the titration curves. The pKa value had noticed to be
4.85, for titration 1, while titration 2, and 3 had recognized the same pKa value of 4.80.
Meanwhile, the Ka value of the acid had accordingly recognized to be 1.41 x 10 -5for titration
1 and the following titrations had seen an equal Ka value of 1.58 x 10 -5. The average value of
Ka was seen to be close to the value of Ka of propanoic acid, which was 1.30 x 10 -5from the
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theoretical Ka table. While the pKa of half method slightly differ from first method that
caused difference result in identification of monoprotic weak acid.
REFERENCES
A. Marrie. (2014).About.com Chemistry:Henderson Hesselbalch Equation and Example.
Retreived fromhttp://chemistry.about.com/od/acidsbase1/a/hendersonhasselbalch.htm
D.C. Harris. (2004).Determination of the Identify of an Unknown Weak Acid [Online PDF
Handout]. Retreived fromhttp://apbrwww5.apsu.edu/robertsonr/chem1110-
20/044%20Unknown%20Acid%20Ka%20MM.pdf
Table of Ka Values for Common Monoprotic Acids. (2015). Retrived from
http://www.bpc.edu/mathscience/chemistry/table_of_monoprotic_acids.html
P.Austin.(2012).Determining the Identity of an Unknown Weak Acid . State University
Department of Chemistry. Retrieved from
http://www.apsu.edu/sites/apsu.edu/files/chemistry/F12
Anonymous.Ka of an Unknown Acid. Retrieved April,22 2014 from
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http://web.utk.edu/~kcook/319S02/exp5m.pdfhttp://web.utk.edu/~kcook/319S02/exp5m.pdfhttp://web.utk.edu/~kcook/319S02/exp5m.pdfRecommended