1 Abstract The experiment was conducted to identified the value of Ka of weak acid by two methods which were potentiometric titration and “half volume” methods. The Ka value is an indication of acid strength. The larger the value of the Ka, the stronger the acid. This value is characteristic of the acid and can be used to help identify an unknown acid. In part A by using potentiometric titration method, a burette was rinsed with tap water and followed by distilled water. Then,the burette was filled with sodium hydroxide solution. The tip of the burette was ensured filled completely with the solution. 10 mL of unknown acid was pipette in a beaker and water was added until it covered pH electrode. pH electrode was rinsed with distilled water before using it and the set up was adjusted so that the pH electrode was above stirring bar. Total volume of NaOH added was recorded in each titration. In part B which is “half volume” method, a sample of unknown acid was obtained and 100 mL of distilled water was measured with graduated cylinder and poured into 250 mL Erlenmeyer flask. The sample acid was dissolved and stirred with water. The solution prepared then was divided into two. The portion in the flask was titrated with phenolphthalein endpoint with NaOH solution. As endpoint approaches, NaOH was added drop by drop until the solution has permanent pink colour. The titrated solution was mixed with the other half of the acid and the pH of resulting solution was determined. The results of this experiment in part A for average value of pKa and Ka were 4.45 and 4.43x10 -5 . Theoretically, the pKa and Ka value of the unknown acid which is believed to be benzoic acid were 4.2 and 6.3x10 -5 . Therefore, an error of 5.95 % of pKa value and 29.68 % of Ka value was calculated in potentiometric titration while in “half volume” method, percentage error for Ka value was 100% .
Transcript
1
Abstract
The experiment was conducted to identified the value of Ka of weak acid by two methods which were potentiometric titration and “half volume” methods. The Ka value is an indication of acid strength. The larger the value of the Ka, the stronger the acid. This value is characteristic of the acid and can be used to help identify an unknown acid.
In part A by using potentiometric titration method, a burette was rinsed with tap water and followed by distilled water. Then,the burette was filled with sodium hydroxide solution. The tip of the burette was ensured filled completely with the solution. 10 mL of unknown acid was pipette in a beaker and water was added until it covered pH electrode. pH electrode was rinsed with distilled water before using it and the set up was adjusted so that the pH electrode was above stirring bar. Total volume of NaOH added was recorded in each titration.
In part B which is “half volume” method, a sample of unknown acid was obtained and 100 mL of distilled water was measured with graduated cylinder and poured into 250 mL Erlenmeyer flask. The sample acid was dissolved and stirred with water. The solution prepared then was divided into two. The portion in the flask was titrated with phenolphthalein endpoint with NaOH solution. As endpoint approaches, NaOH was added drop by drop until the solution has permanent pink colour. The titrated solution was mixed with the other half of the acid and the pH of resulting solution was determined.
The results of this experiment in part A for average value of pKa and Ka were 4.45 and 4.43x10-5 . Theoretically, the pKa and Ka value of the unknown acid which is believed to be benzoic acid were 4.2 and 6.3x10 -5 . Therefore, an error of 5.95 % of pKa value and 29.68 % of Ka value was calculated in potentiometric titration while in “half volume” method, percentage error for Ka value was 100% . This was due to some error that having during the experiment. Hence, by comparing the percentage error in part A and part B, it can be concluded that potentiometric titration method was more accurate since there was only slightly difference from the theoretical value compared to “half volume” method and the unknown acid was identified as benzoic acid.
2
INTRODUCTION
Acording to Bronsted-Lowry acid-base theory, the strength of an acid is related to
its ability to donate protons while the strength of base is related to its ability to
accept protons from another substances [1]. Acids and bases are often describe
as being “weak” or “strong”. Strong acids are electrolytes that ionise completely
to H+ or H3O+ ions in aqueous solution, while weak acids are electrolytes that
partially ionise to H+ or H3O+ ions in aqueous solution. For strong bases,
electrolytes are ionise completely to OH- ions in aqueous solution while weak
bases are electrolytes that partially ionise to OH- ions in aqueous solution [2].
The Ka is acid dissociation constant for weak acid. For acids, Ka values are
commanly used. The ionization of an acid can be shown by the following
equation:
HA (aq) ↔ H+ (aq) + A- (aq) (equation 1)
Since an equilibrium exists, an equilibrium constant, Ka, can be written as :
Ka = [H + ] [A - ] (equation 2)
[HA]
[A-] is the molar concentration of the conjugate base and [HA] is the molar
concentration of the weak acid. Ka can be calculated by using the initial
concentration of the acid and the initial pH of the solution. The initial pH gives the
[H+] which equals the [A-] in the initial weak acid solution. If the weak acid is only
slightly ionized, the [HA] is assumed to be approximately equal to its initial
concentration [1] . The larger the value of the Ka, the stronger the acid. This
value is characteristic of the acid and can be used to help identify an unknown
acid. A similar system exists for bases.
Two methods may be used to determined the Ka value. Both methods require the
use of pH meter. In the first method, a sample of acid is titrated with base. The
pH value are plotted vs. the volume of base added. The equivalence point is
3
determined from the graph [3] . At the equivalence point of a titration, exactly
enough OH- ions have been added to react with all of the HA molecules that were
originally present. Around this point, the pH of the mixture increases suddenly
and dramatically. The rapid increase in pH can be detected by using pH meter
[4] . Next, the volume of base halfway to the equivalence point is found, and the
pH at this volume is noted. The [H+] corresponding to this pH is equal to this Ka
for the acid. At a point halfway to the equivalence point, [H+] = [HA] = [A-] for a
monoprotic acid. Canceling out [A-] and [HA] in equation 2 will gives Ka = [H+] .
The second method for determining Ka values involves a “ half volume” method.
A solution of the acid is prepared and divided in half as accurately as possible.
One portion is titrated to its endpoint with phenolphthalein. The two portions are
then recombined, and the pH of the resulting solution is measured. Since half of
the has been titrated, [H+] = [HA] =[A]. Again, if [A-] and [HA] are canceled in
equation 2, ka = [H+]. The pH value of the combined solutions can be converted
to [H+] to give a Ka value.
4
Aims
a) The purpose of this experiment is to determine the Ka value for a weak acid by
two methods ; potentiometric titration and “half volume” methods.
b) To determine the end point of the titration from the titration curve.
c) To determine the acids ability to ionizes or produce ions when dissolved with
water.
d) To explore weak acid in equilibria system
5
Theory
Figure 1:A typical set up for using a pH meter to measure data for a titration
curve
Titration is a procedure for determining the concentration of a solution of
unknown concentration by titrating it with a solution of known concentration. In
the potentiometric titration, a burette is used to dispense a small, quantifiable
increment of solution of known concentration. A typical has the smallest
calibration unit of 0.1 mL , therefore volume dispense from the burette should be
estimated to the nearest 0,01 mL. The pH meter will be used to measure the pH
of the solution (Figure 1).
In this experiment, equivalence point occur when the equivalent amount
of acid and base has mixed stoichiometrically. At this moment, the equivalent
amount of acid and base present exactly neutralises one another. In other words,
6
the number of moles of H+ ion (from acid) equals to the number of moles of OH -
ion (from base) at the equivalence point. Neither the acid nor base is present in
excess and thus acid has completely reacted with base at the equivalence point.
The sudden change in the pH of the solution shows that the titration has reached
the equivalence point. pH in an aqueous solution is related to its hydrogen ion
concentration. Symbolically, the hydrogen ion concentration is written as [H3O+] .
pH is defined as the negative of the logarithm of the hydrogen ion concentration
as stated in equation 3.
pH = -log [H3o+] (equation 3)
pH scale is a method of expressing the acidity or basicity of a solutions.
When the pH value of the solution less than 7 (pH <7), the solution is said to be
acidic while if the pH of the solution is greater than 7 (pH >7), the solution is
considered as a basic. The solution is at neutral condition when its pH value is
equal to 7 (pH=7).
There are two general methods used to determined the Ka value of weak
acid which are potentiometric method and “half-volume” method. In this
experiment, we will be dealing with weak monoprotic acid. Based on Bronsted
Lowry, an acid is a proton donor whereas a base is a proton acceptor. This
portrays is a very important idea to understanding monoprotic and polyprotic
acids and bases since monoprotic, as a matter of fact, is basically referred to the
transfer of one proton. On the contrary, polyprotic corresponds to the transfer of
more than one proton. For the weak acid titrated with NaOH, a titration curve is
produced by plotting the pH of the acid solution versus the volume of NaOH
added. The equivalence point of titration is reached when all the weak acid (HA)
has completely reacted with NaOH. On the titration curve, the equivalence point
is read at the center of the region where pH increase sharply. The half-
equivalence point for the titration is reached when exactly one half of the based
required to completely neutralizes the acid has been added.
7
Figure 2: Acid-base titration curve of weak acid titrated with NaOH
At this half-equivalence point, concentration of the acid in the solution, [HA] is
equal to the concentration of its conjugate base, [A-].
[HA] = [A-] (equation 4)
Equation 2 can be simplified to yield equation 5,
Ka = [H3O+] (equation 5)
Taking the negative of the logarithm of each side of equation 5,equation 6 can be
obtained from derivation,
-log Ka = -log [H3O+]
pKa = pH (equation 6)
8
Equation 6 indicates that pKa for the acid is equal to the pH of solution at the half
equivalence point. The Ka of the acid is determined from the pKa values as
follows :
Ka = 10-pKa (equation 7)
By knowing the pH of the weak acid and initial weak acid concentration from he
pH of the acid solution, [HA], the H+ and A- ion concentration can be determined,
which is related to the pH of the solution by equation 8,
[H3O+] = 10-pH (equation 8)
By substituting [HA], [H3O+] and [A-] at equilibrium into equation 2, the value of Ka
can be calculated.
Table 1 : Ionization constant (Ka) for particular weak acids at 25 o C.
Acid Formula Ka
Acetic acid CH3COOH 1.8 X 10-5
Benzoic acid C6H5COOH 6.3 x 10-5
Carbonic acid H2CO3 4.2 x 10-4
Formic acid HCOOH 1.8 x 10-4
Hypochlorous acid HOCl 3.5 x 10-6
Dihydrogen phosphate ion H2PO4- 6.2 x 10-8
Hydrogen phosphate ion HPO42- 3.6 x 10-13
Hydrogen carbonate ion HCO3- 4.8 x 10-11
Nitrous acid HNO2 4.0 x 10-4
Phenol C6H6O 1.6 x 10-10
Potassium hydrogen phthalate KC8H5O4 5.1 x 10-6
9
Apparatus
In the experiment to determined the value of Ka of a weak acid, there are
some materials and apparatus that are use during experiment. The materials that
are use are standard base solution which is 0.1 M NaOH, buffer solutions which
has ability to maintain its pH when a small amount of strong base or strong acid
is added to it, phenolphthalein that act as indicator to indicate the equivalence
point of an acid-base titration by its colour change, acid sample and distilled
water. The apparatus that are use in carrying the experiment are pH meter. It is
an electronic device used for measuring the pH (acidity or alkalinity) of liquid.
Beaker also use as a liquid measuring container while burette is to delivers small
amounts of liquids slowly and precisely. Graduated cylinder is use to measure
liquid volume accurately while retort stand is to holds other equipment such as
burette. The other apparatus is magnetic stirrer to stir the solution by spin very
quickly in the immersed liquid.
10
Procedure
Part 1 : Determination the value of Ka of weak acid by potentiometric titration.
A burette was rinsed with tap water and followed by distilled water. The
burette is cleaned if it is drained without leaving drops of water behind. The
burette was washed with a dilute detergent solution after figure out it was not
clean enough and was rinsed by several times with tap water and with distilled
water. Some of the sodium hydroxide solution was poured into a beaker. A small
amount of sodium hydroxide was poured into the burette to let it drained through
into a “waste” beaker. The burette was filled with sodium hydroxide solution until
the tip of the burette had been completely filled. The sodium hydroxide solution
was drained to the 0.00 mL line or below.
10 mL of the unknown acid was pipette into a beaker and the water was
added until the end of the pH electrode was recovered. The initial reading of the
sodium hydroxide solution in the burette was recorded. The pH electrode was
rinsed and it is placed in the beaker with the acid. The beaker was then placed in
a stir plate and magnetic stirrer was inserted into the beaker to let it stirred the
solution. The set-up was adjusted, so that the burette containing sodium
hydroxide solution was over the beaker and the pH electrode was above the
stirring bar. The experiment began with the addition of small portion of the
sodium hydroxide to the acid sample. The total volume of sodium hydroxide
added and the value of the pH solution were recorded after each addition of
sodium hydroxide. The smaller portions of the sodium hydroxide was added
when the pH began to rise rapidly. The sodium hydroxide was continuedly added
until the pH was in the basic region and had leveled off. The titrated acid sample
in the beaker was discarded and the titration was repeated. The graph was
plotted with pH (y-axis) as a function of volume of base added (x-axis).
11
Part B : Determination of the value of Ka of a weak acid by “half-volume” method.
The pH meter and burette were assumed to be ready for used. A sample
of an unknown acid was obtained. 100 mL of distilled water was measured
out with a graduated cylinder and was poured into a 250 mL Erlenmeyer
flask. The sample of acid was dissolved in water and was stirred to mix
throughly. The solution that had been prepared was divided into two equal
portions by using graduated cylinder . The portion in the flask was titrated to
a phenolphthalein endpoint with sodium hydroxide solution. The sodium
hydroxide was slowly added while swirling the flask. The sodium hydroxide
solution was added drop by drop as the end point approaches until the
solution has a permanent pink colour. The titrated solution was mixed with
the other half of the acid solution and the pH of the resulting solution was
determined. The Ka of the unknown acid was calculated from the observed
pH.
12
Results
Part A : Potentiometric Titration
Initial burette reading : 50 mL
The data from the experiment for the pH value of unknown acid solution after the
addition of sodium hydroxide solution by titration was tabulated below :
Table 2 : pH of solution and the volume of NaOH (mL) added
Volume of NaOH (mL) pH (titration 1) pH (titration 2)
0 4.067 4.069
1 5.692 5.488
2 6.990 10.694
3 12.119 12.012
4 12.390 12.279
5 12.476 12.467
6 12.530 12.555
7 12.564 12.627
8 12.605 12.684
9 12.618 12.716
10 12.669 12.760
11 12.705 12.775
12 12.752 12.799
13 12.783 12.815
13
Figure 3: Graph of pH value of unknown solution after addition of some volume of
NaOH (mL)
Titration 1
Equivalence point on the graph = pH 9.5
Volume of NaOH correspond to equivalence point = 2.5 mL
Titration 2
Equivalence point on the graph = pH 8.3
14
Volume of NaOH correspond to equivalence point = 1.6 mL
Part B : “Half-volume” Method
Trial 1
Original pH of solution = 3.68
Volume of NaOH added to neutralize the solution = 10.85 mL
pH of solution at equivalence point = 8.40
pH after mixing the solution = 4.15 pH
Trial 2
Original pH of solution = 4.08
Volume of NaOH added to neutralize the solution = 11.60 mL
pH of solution at equivalence point = 6.85
pH after mixing the solution = 4.13
15
Calculations
Part A : Potentiometric Titration
Titration 1
pH at half equivalence point = 9.5
2
pH at half equivalence point = 4.75
From equation 6, pKa = pH
pKa = 4.75
Based on equation 3, pH = -log [H3O+]
4.75 = -log [H3O+]
[H3O+] = 1.78x10-5
From equation 5, Ka = [H3O+]
Ka = 1.78x10-5
Titration 2
pH at half equivalence point = 8.3
2
pH at half equivalence point = 4.15
From equation 6 pKa = pH
pKa = 4.15
Based on equation 3, pH = -log [H3O+]
4.15 = -log [H3O+]
[H3O+] = 7.08x10-5
16
From equation 5, Ka = [H3O+]
Ka = 7.08x10-5
Average of pKa = 4.75 + 4.15 = 4.45
2
Theoretical value of pKa of benzoic acid = 4.2
Percentge error of pKa = | 4.2-4.45 | x 100%
4.2
= 5.95%
Average dissociation constant of weak acid, Ka = (1.78x10 -5 )+(7.08x10 -5 )
2
=4.43x10-5
Theoretical value of Ka of benzoic acid = 6.3 x 10-5
Percentage error of Ka = | (6.3x10 -5 )-(4.43x10 -5 ) | x 100%
(6.3x10-5)
= 29.68%
Part B : “half volume” method
Based on reaction equation of the experiment
HA (aq) + H2O+ (l) H3O+ (aq) + A- (aq)
By considering the ‘ICE’ table
Components HA H3O+ A-
Initial concentration 0.1 0 0
Change in equilibrium -X +X +X
Concentration at 0.1- X X X
17
equilibrium
The ‘ICE’ table is used to calculate the value of Ka in the experiment Part B. The
initial pH of unknown acid was recorded without adding any bases solution in the
acid.
Trial 1:
pH value of unknown acid at equivalence point = 8.4
pH = - log [H3O+]
8.4 = - log [H3O+]
[H3O+] = 10 –Ph
= 10 -8.4
[H3O+] = 3.98x10-9
According to ICE table,
HA (aq)+ H2O (l)↔ A- (aq) + H3O+ (aq)
Initial 0.1 M 0 M 0 M
Change _ 3.98x10-9M + 3.98x10-9 + 3.98x10-9
Equilibrium 0.0999999996 M + 3.98x10-9 + 3.98x10-9
thus, Ka = (x)(x)
0.1-x
Ka = ( 3.98x10 -9 ) 2 0.0999999996
Ka = 1.58 x 10-16
Trial 2 :
18
pH value of unknown acid at equivalence point = 6.85
pH = - log [H3O+]
6.85 = - log [H3O+]
[H3O+] = 10 –Ph
= 10 -6.85
[H3O+] = 1.41x10-7
According to ICE table,
HA (aq)+ H2O (l)↔ A- (aq) + H3O+ (aq)
Initial 0.1 M 0 M 0 M
Change _ 1.41x10-7M + 1.41x10-7 + 1.41x10-7
Equilibrium 0.099999859 M + 1.41x10-7 + 1.41x10-7
thus, Ka = (x)(x)
1-x
Ka = ( 1.41x10 -7 ) 2 0.099999859
Ka = 1.98 x 10-13
Average of Ka of unknown acid = (1.98x10 -13 ) + (1.58x10 -16 )
2
= 9.91x10-14
Percentage error of Ka = (6.3x10 -5 )-(9.91x10 -14 ) | x 100%
(6.3x10-5)
= 100.00%
19
Discussion
A lot of information is required in order to assure that the identity of the
unknown acid is conclusive. After analyzing the results of the experiment as well
the Ka value, it is concluded to be remarkably similar to benzoic acid which is
weak acid. The strength of weak acid was determined by its value of Ka. The
stronger the acid, the higher the concentration of H+ or H3O+ ions, the lower the
value of pH, the larger the value of Ka, the smaller the value of pKa. Based on
figure 3, the graph shows that pH for titration 1 had increased significantly from
2mL until 3 mL of NaOH but a sudden change of pH at the next point, which from
3.1 mL to 13 mL of NaOH. From the graph, the equivalence point was then
calculated and for titration 1, its pH was 9.5 . In titration 2, pH of the solution was
sharply increased from 1 mL until 2 mL and its equivalence point was at pH 8.3.
This result was got by method of potentiometric titration. The equivalence point of
the titration was reached when all the unknown acid has completely reacted with
NaOH. On the titration curve, the equivalence point was read at the center of the
region where the pH increases sharply.
The unknown’s Ka value was calculated by dividing equivalence point with 2
and the results of the average unknown’s Ka value was 4.43x10-5 whereas
benzoic acid’s Ka value was 6.3x10-5 . The half-equivalence point for the titration
was reached when exactly one and a half of the base required to complete the
neutralization has been added. At this point, the concentration of the acid in the
solution is equal to the concentration of its conjugate base. As for pKa,the
average pKa value of the unknown acid was calculated to be 4.45 whereas
benzoic acid’s pKa was 4.2. For half volume method, the average unknown’s Ka
value was 9.91x10-14 .
20
The main objective of the experiment is to determined the Ka value of the
unknown monoprotic acid so as to identify the acid. However, after calculated the
percentage error of acid dissociation constant, Ka and pKa for both experiment in
part A and B, there were large value of percentage error after comparing the
theoretical value with the experimental value that obtained. In part A, percentage
error for pKa was 5.95 % while percentage error for Ka was 29.68 %. This was
different compared with part B. Its percentage error for Ka was 100.00 %. Based
on this result, it can be said that potentiometric titration method was more
accurate to determined Ka value of weak acid due to less percentage error
compared to “half-volume” method . A few mistakes or lack of awareness of the
precautions that must be considered when conducting the experiment may be
the reasons that lead to the erroneous calculations.
First error that might affect the calculated values was done during
weighing the pellets of sodium hydroxide. The weight of an empty beaker should
have been considered as well. Then, it must be substracted from the weight of
the beaker containing the pellets. Hence, the titration curves might not have the
accurate values as the concentration of sodium hydroxide was not perfectly 0.1
M. Second, that has been used has never given the definite readings, as the
values that are shown are always changing rapidly. Thus, one can never tell the
accurate readings of pH values. Therefore, these will also affect the titration
curve as well as pKa values which correspond to the pH values at half
equivalence point. Besides that, the equivalence point was not necessarily being
at pH 7 as it occurs just when the concentration of base reacted in solution.
Therefore, the final pH depends on the major species of ions left in the solution
after the reaction. In addition, the pH electrode might have come into contact with
magnetic stirrer. Therefore, the problem might be encountered during the
readings of pH values on the pH meter. Lastly, percentage error also can occur if
the burette did not rinsed with distilled water before using it the presence of air
bubbles in the burette which can cause error in the true value of NaOH used.
21
Conclusion
After careful consideration of all the possible results, the average
experimental values of pKa and Ka of the unknown acid in part A (potentiometric
titration method), were 4.45 and 4.43x10-5 whereas the average value of Ka of
unknown acid by using “half-volume” method was 9.91x10-14 . Theoretically, the
pKa and Ka values of the unknown monoprotic acid which was believed to be
benzoic acid were 4.2 and 6.3x10-5 . Therefore, there were an error during the
experiment was proceed. In part A, percentage error of pKa and Ka were 5.95 %
and 29.68 % while in part B, percentage error of Ka was 100.00 %. There was
slightly difference in percentage error of Ka in part A compared in part B.
Therefore, potentiometric method was more accurate method to determined and
identified the Ka of weak acid which was benzoic acid.
Recommendations
22
There are few recommendations and precautions that have to be considered
during the experiment in order to get an accurate value and readings data.
Firstly, the standard solution that was used should be a hundred percent pure
and stable at room temperature. Thus, it was more preferable to use a dried
standard material before weighing and diluted. Secondly, in order to be more
conclusive in identifying the unknown monoprotic acid, the molecular weight of
the acid should be considered as well. This then can be used to compare it with
the theoretical value of molecular weight of benzoic acid. Other than that, we also
have to ensure that there are no air bubbles trapped at the tip of the burette
during the filling of NaOH solution. We have to make sure our eye is
perpendicular to the scale to avoid from making parallax error. Lastly, rinsed the
burette with distilled water before filling NaOH solution to get an accurate results.
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