TED ANKARA COLLEGE FOUNDATION HIGH SCHOOL EFFECT OF TEMPERATURE ON THE ACID DISSOCIATION CONSTANT OF ACETIC ACID Assessment Criteria: Full Investigation (D, DCP, CE) Session: May 2013 Candidate Name: Oğul Ersin ÜNER Candidate Number: D1129066 Centre Name: TED Ankara College High School, Ankara, TURKEY Date of Experiment: 10.04.2012 Instructor: Ms. Ayse Senay
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TED ANKARA COLLEGE FOUNDATION HIGH SCHOOL
EFFECT OF TEMPERATURE ON THE ACID
DISSOCIATION CONSTANT OF ACETIC ACID
Assessment Criteria: Full Investigation (D, DCP, CE)
Session: May 2013
Candidate Name: Oğul Ersin ÜNER
Candidate Number: D1129066
Centre Name: TED Ankara College High School, Ankara, TURKEY
Date of Experiment: 10.04.2012
Instructor: Ms. Ayse Senay
OĞUL ERSİN ÜNER D1129066
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BACKGROUND INFORMATION
What is Titration?
Titration determines the molarity of an unknown concentration of a solution. During the
process, the concentration of the solution is found by the addition of another substance which has a
known concentration. “The addition is stopped when the endpoint is reached.” 1
The colour of the
solution changes due to the presence of an indicator, a colourless solution which specifically changes
colour according to the pH of the medium.
In this titration experiment, the weak acid acetic acid, CH3COOH(aq), is titrated with the strong
base sodium hydroxide, NaOH(aq), in order to find the concentration of the acetic acid of unknown
concentration. The colour is changed by using the indicator phenolphthalein, which “is colourless in
acid solution and pink in alkaline solution”2.
What is Ka?
“An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization
constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium
constant for a chemical reaction known as dissociation in the context of acid-base reactions.”3 It can
only be written for weak acids, such as for acetic acid, which partially ionize in water. It can be
expressed as:
3
where:
Ka is the acid dissociation constant of the acid HA
[A-] denotes the concentration of the conjugate base of the acid HA
[H+] denotes the concentration of hydrogen ions.
[HA] denotes the concentration of the acid HA
What is Temperature and how does it affect the Ka of acids?
“Temperature is a physical property of matter that quantitatively expresses the common
notions of hot and cold. Objects of low temperature are cold, while various degrees of higher
Table 1: The initial and final readings on the burette containing the aqueous sodium hydroxide solution, the volume of the aqueous sodium hydroxide solution
obtained from the difference of the initial and final reading on the burette, the molarity of aqueous sodium hydroxide, volume of aqueous acetic acid solution
used, the temperature of the water bath, the room temperature and room pressure in each trial during titration.
TITRATION OF CH3COOH(aq) WITH NaOH(aq)
TRIAL
INITIAL READING
ON BURETTE
CONTAINING
NaOH(aq) SOLUTION
( 0.02 mL)
FINAL READING
ON BURETTE
CONTAINING
NaOH(aq) SOLUTION
( 0.02mL)
VOLUME
OF NaOH(aq)
SOLUTION
USED
( 0.04 mL)
VOLUME OF
CH3COOH(aq)
SOLUTION
USED
( 0.2 mL)
MOLARITY
OF NaOH(aq)
SOLUTION
(M)
TEMPERATURE
OF THE WATER
BATH/TITRATION
MIXTURE
( 0.20C)
NUMBER OF
PHENOLPHTHALEIN
SOLUTION DROPS
ROOM
TEMPERATURE
( 0.20C)
ROOM
PRESSURE
( 0.2 hPa)
1 30.00 1.10 28.90 30.0 1 20.2 3 19.5 1067.0
2 30.00 0.39 29.61 30.0 1 20.3 3 19.5 1067.0
3 30.00 0.11 29.89 30.0 1 20.1 3 19.5 1067.0
4 30.00 0.25 29.75 30.0 1 20.3 3 19.5 1067.0
5 30.00 0.34 29.66 30.0 1 20.0 3 19.5 1067.0
OĞUL ERSİN ÜNER D1129066
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DETERMINATION OF THE ACID DISSOCIATION CONSTANT OF CH3COOH(aq)
FOR 20.0°C ACETIC ACID SOLUTION
TRIAL
INITIAL TEMPERATURE
OF CH3COOH(aq)
SOLUTION
( 0.2 °C)
FINAL TEMPERATURE
OF CH3COOH(aq)
SOLUTION (JUST
BEFORE pH READING)
( 0.2 °C)
pH OF
CH3COOH(aq)
SOLUTION
(±0.01)
VOLUME OF
CH3COOH(aq)
SOLUTION USED
( 0.2 mL)
TEMPERATURE OF
THE WATER BATH
( 0.20C)
1 19.5 20.0 2.52 30.0 20.2
2 19.4 20.0 2.52 30.0 20.3
3 19.7 20.0 2.51 30.0 20.0
4 19.5 20.0 2.51 30.0 20.1
5 19.6 20.0 2.51 30.0 20.0
FOR 35.0°C ACETIC ACID SOLUTION
TRIAL
INITIAL TEMPERATURE
OF CH3COOH(aq)
SOLUTION
( 0.2 °C)
FINAL TEMPERATURE
OF CH3COOH(aq)
SOLUTION(JUST
BEFORE pH READING)
( 0.2 °C)
pH OF
CH3COOH(aq)
SOLUTION
(±0.01)
VOLUME OF
CH3COOH(aq)
SOLUTION USED
( 0.2 mL)
TEMPERATURE OF
THE WATER BATH
( 0.20C)
1 19.2 35.0 2.47 30.0 35.0
2 19.3 35.0 2.39 30.0 35.1
3 19.1 35.0 2.46 30.0 35.0
4 19.3 35.0 2.45 30.0 35.3
5 19.0 35.0 2.47 30.0 35.0
FOR 50.0°C ACETIC ACID SOLUTION
TRIAL
INITIAL TEMPERATURE
OF CH3COOH(aq)
SOLUTION
( 0.2 °C)
FINAL TEMPERATURE
OF CH3COOH(aq)
SOLUTION (JUST
BEFORE pH READING)
( 0.2 °C)
pH OF
CH3COOH(aq)
SOLUTION
(±0.01)
VOLUME OF
CH3COOH(aq)
SOLUTION USED
( 0.2 mL)
TEMPERATURE OF
THE WATER BATH
( 0.20C)
1 33.1 50.0 2.57 30.0 50.2
2 33.3 50.0 2.56 30.0 50.1
3 33.0 50.0 2.58 30.0 50.0
4 33.2 50.0 2.57 30.0 50.2
5 33.0 50.0 2.59 30.0 50.0
Table 2: The initial and final temperatures of the acetic acid solution, pH of the acetic acid, volume of
acetic acid solution used and the temperature of the water bath for 20.0, 35.0 and 50.0°C.
OĞUL ERSİN ÜNER D1129066
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Note: The solutions were prepared by the lab technician and therefore, the uncertainties for the
concentrations of the solutions could not be provided.
Qualitative Data:
CH3COOH(aq), phenolphthalein and NaOH(aq) were colourless before titration.
The indicator phenolphthalein did not change the colour of the mixture immediately when it was
added to the flask that contained acetic acid solution.
When NaOH(aq) started to drop into the flask with CH3COOH(aq) and phenolphthalein, the colour of
the solution started to turn into pink temporarily.
While swirling the flask, the temporary light pink colour observed faded away and the mixture of
CH3COOH(aq) and phenolphthalein became colourless again.
The light pink colour observed before did not fade away at the endpoint of titration for about 20.0
seconds.
The temperature of the water bath increased when NaOH(aq) was added to the flask of
CH3COOH(aq) and one could feel the heat liberated from the erlenmeyer flask by touching it, which
indicated that the reaction was exothermic.
Acetic acid had a strong vinegar odor that could be smelled from a distance.
Apparently, sodium hydroxide solution was odorless. A smell could not be detected from afar and
the solution could not be smelled closely due to safety reasons.
DATA PROCESSING
1. FINDING THE MOLARITY OF ACETIC ACID (ETHANOIC ACID) SOLUTION
USED
During titration, the end point, which is marked by the indicator by the change of color, is the
point where the stoichiometric coefficients of the titrant (sodium hydroxide) and the analyte (acetic
acid) are equal; “the amount of titrant is sufficient to fully neutralize or react with the analyte.”11
So
the number of moles of the acid and base are equal.
where n is the number of moles.
11 Whitney, W.D., Smith, B.E. (1911). "Titrimetry". The Century Dictionary and Cyclopedia. The Century co. pg. 6504. Print.
OĞUL ERSİN ÜNER D1129066
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In order to find the number of moles of acid and base at that point, the molarities should be multiplied
by the volume of the acid or base used.
where n is the number of moles, M is the molarity in molars and V is the volume in liters.
Balanced equations aid in observing stoichiometric coefficients and must be used to find the molarity
of the acetic acid solution. One should assume that the indicator changes color exactly at the
equivalence point, due to the fact that “some indicators change near the point.”12
The rightmost significant figures are shown in bold in every calculation.
Unless otherwise stated, all values given in tables are rounded according to their uncertainties.
The unrounded, rounded and average values are given in 2 Tables: At the end of the titration
process and at the end of the determination of acid dissociation constants in the results table.
Example for the First Trial of Titration:
THE TITRATION OF ACETIC ACID WITH SODIUM HYDROXIDE
X mol Y mol
Since there is 1 mol-1 mol stoichiometric ratio between acetic acid and sodium hydroxide in the
reaction equation above, the number of moles of sodium hydroxide and acetic acid are equal at the end
(The ionized part can be neglected due to the fact that it is so small that it would not have a major
effect on the calculations in determining the Ka of acetic acid.)
By conducting this method for the other trials, one can find the molarity of the acetate ions as:
Trial 2: M
Trial 3: M
Trial 4: M
Trial 5: M
ACETIC ACID SOLUTION AT 20.0°C
TRIAL pH OF ACETIC ACID SOLUTION
(± 0.01)
HYDROGEN ION CONCENTRATION IN
ACETIC ACID (M)
1 2.52
2 2.52
3 2.51
4 2.51
5 2.51
Table 4: The pH values of the acetic acid solutions found by a pH meter and the hydrogen ion
concentrations that are calculated from the pH values at 20.0°C in all five trials respectively.
Example for the First Trial of Determining the Acetate and Hydrogen Ion Concentrations at
35.0°C:
OĞUL ERSİN ÜNER D1129066
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THE PARTIAL IONIZATION OF ACETIC ACID
In: 1 M
Rxn: - Y M + Y M + Y M Ionized part : Y M
Eq: (1-Y) M Y M M
(The ionized part can be neglected due to the fact that it is so small that it would not have a major
effect on the calculations in determining the Ka of acetic acid.)
By conducting this method for the other trials, one can find the molarity of the acetate ions as:
Trial 2: M
Trial 3: M
Trial 4: M
Trial 5: M
ACETIC ACID SOLUTION AT 35.0°C
TRIAL pH OF ACETIC ACID SOLUTION
(± 0.01)
HYDROGEN ION CONCENTRATION IN
ACETIC ACID (M)
1 2.47
2 2.39
3 2.46
4 2.45
5 2.47
Table 5: The pH values of the acetic acid solutions found by a vernier and the hydrogen ion
concentrations that are calculated from the pH values at 35.0°C in all five trials respectively along
with their uncertainties.
Example for the First Trial of Determining the Acetate and Hydrogen Ion Concentrations at
50.0°C:
OĞUL ERSİN ÜNER D1129066
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THE PARTIAL IONIZATION OF ACETIC ACID
In: 1 M
Rxn: - Y M + Y M + Y M Ionized part : Y M
Eq: (1-Y) M Y M M
(The ionized part can be neglected due to the fact that it is so small that it would not have a major
effect on the calculations in determining the Ka of acetic acid.)
By conducting this method for the other trials, one can find the molarity of the acetate ions as:
Trial 2: M
Trial 3: M
Trial 4: M
Trial 5: M
ACETIC ACID SOLUTION AT 50.0°C
TRIAL pH OF ACETIC ACID SOLUTION
(± 0.01)
HYDROGEN ION CONCENTRATION IN
ACETIC ACID (M)
1 2.57
2 2.56
3 2.58
4 2.57
5 2.59
Table 6: The pH values of the acetic acid solutions found by a vernier and the hydrogen ion
concentrations that are calculated from the pH values at 50.0°C in all five trials respectively along
with their uncertainties.
OĞUL ERSİN ÜNER D1129066
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ACETIC ACID SOLUTION AT 20.0°C
TRIAL CONCENTRATIONS OF ACETATE AND HYDROGEN IONS IN ACETIC ACID SOLUTION
(M)
1
2
3
4
5
ACETIC ACID SOLUTION AT 35.0°C
TRIAL
CONCENTRATIONS OF ACETATE AND HYDROGEN IONS IN ACETIC ACID SOLUTION
(M)
1
2
3
4
5
ACETIC ACID SOLUTION AT 50.0°C
TRIAL
CONCENTRATIONS OF ACETATE AND HYDROGEN IONS IN ACETIC ACID SOLUTION
(M)
1
2
3
4
5
Table 7: The concentrations of acetate and hydrogen ions obtained from the pH of the solutions at
20.0°C, 35.0°C and 50.0°C respectively for all five trials.
3. FINDING THE ACID DISSOCIATION CONSTANT OF ACETIC ACID AT 20.0°C,
35.0°C and 50.0°C
Since the concentrations of acetate and hydrogen ions are known, the acid dissociation
constant at 20.0°C, 35.0°C and 50.0°C can be found. The constant for acetic acid can be expressed as:
where:
is the acid dissociation constant
OĞUL ERSİN ÜNER D1129066
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is the concentration of the acetate ions in the acetic acid solution
is the concentration of the hydrogen ions in the acetic acid solution
is the concentration of acetic acid.
Note that the average value and uncertainty of the molarity of will be used in the
calculations for determining the acid dissociation constants and their uncertainties.
Example for the First Trial of Determining the Acid Dissociation Constant at 20.0°C:
Average M
M
M
M
Uncertainty of the Acid Dissociation Constant of Acetic Acid for the First Trial:
(Percentage uncertainty of average ) + (Percentage uncertainty of pH of )
Consequently,
Comment [a5]: Unit?
Comment [a6]: The unrounded sum should first have been shown
OĞUL ERSİN ÜNER D1129066
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By conducting these steps for the other trials, one can find the acid dissociation constants and its
percentage uncertainties as:
Trial 2:
Trial 3:
Trial 4:
Trial 5:
Example for the First Trial of Determining the Acid Dissociation Constant at 35.0°C:
M
M
Uncertainty of the Acid Dissociation Constant of Acetic Acid for the First Trial:
(Percentage uncertainty of ) + (Percentage uncertainty of pH of )
OĞUL ERSİN ÜNER D1129066
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Consequently,
By conducting these steps for the other trials, one can find the acid dissociation constants and its
percentage uncertainties as:
Trial 2: M
Trial 3: M
Trial 4: M
Trial 5: M
Example for the First Trial of Determining the Acid Dissociation Constant at 50.0°C:
M
M
Uncertainty of the Acid Dissociation Constant of Acetic Acid for the First Trial:
(Percentage uncertainty of ) + (Percentage uncertainty of pH of )
OĞUL ERSİN ÜNER D1129066
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Consequently,
By conducting these steps for the other trials, one can find the acid dissociation constants and its
percentage uncertainties as:
Trial 2: M
Trial 3: M
Trial 4: M
Trial 5: M
OĞUL ERSİN ÜNER D1129066
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ACETIC ACID SOLUTION AT 20.0°C
TRIAL ACID DISSOCIATION CONSTANT OF ACETIC ACID
(M)
1
2
3
4
5
ACETIC ACID SOLUTION AT 35.0°C
TRIAL
ACID DISSOCIATION CONSTANT OF ACETIC ACID
(M)
1
2
3
4
5
ACETIC ACID SOLUTION AT 50.0°C
TRIAL
ACID DISSOCIATION CONSTANT OF ACETIC ACID
(M)
1
2
3
4
5
Table 8: The acid dissociation constants found in each trial at 20.0°C, 35.0°C and 50.0°C with their
percentage uncertainties.
Determining the Average Value of the Acid Dissociation Constants:
Note that unrounded values were used in this step so as to prevent imprecise calculations.
OĞUL ERSİN ÜNER D1129066
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Average Value of the Acid Dissociation Constant at 20.0°C:
M
Average Percentage Uncertainty:
Average Value of the Acid Dissociation Constant at 35.0°C:
M
Average Percentage Uncertainty:
OĞUL ERSİN ÜNER D1129066
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Average Value of the Acid Dissociation Constant at 50.0°C:
M
Average Percentage Uncertainty:
TEMPERATURE
OF ACETIC ACID AVERAGE ACID DISSOCIATION CONSTANTS OF ACETIC ACID (M)
20.0°C
35.0°C
50.0°C
Table 9: The average acid dissociation constants of the acetic acid solution found in different
temperatures in M with percentage uncertainties.
4. FINDING THE THEORETICAL ACID DISSOCIATION CONSTANTS OF
ACETIC ACID AT 20.0°C, 35.0°C and 50.0°C
The theoretical value of the acid dissociation constant of acetic acid at 20.0°C, 35°C and 50°C can
be found using Van’t Hoff’s equation.5
Note that the acid dissociation constants found by the Van’t Hoff’s equation will be the literature
values discussed in the Conclusion and Evaluation section.
OĞUL ERSİN ÜNER D1129066
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where:
KT2 is the equilibrium constants at T2, at 20.0oC (293K), 35.0°C (308 K) and 50.0°C (323 K).
KT1 is the equilibrium constant at T1, at 25.0oC (298K), which is 1.754x10
-5 M.
14
is the standard enthalpy change of the dissociation of acetic acid, which is - 385 J/mol. 15
R is the universal gas constant, which is 8.314 J mol-1
K-1
.16
Since the acid dissociation constant at 25.0°C is known, one can find the acid dissociation constant
at 20.0°C, 35.0°C and 50.0°C.
Determining the Theoretical Acid Dissociation Constant at 20.0°C:
M
Percentage Uncertainty of the Theoretical Acid Dissociation Constant at 20.0°C:
Determining the Theoretical Acid Dissociation Constant at 35.0°C:
14 Housecroft, C. E.; Sharpe, A. G. (2008). “Inorganic Chemistry” (3rd ed.). Chapter 6: Acids, Bases and Ions. Print. 15 Green, J., & Damji. (2008) “S. Chemistry”. Melton: IBID. Print. 16 Jensen, William B. (2003). "The Universal Gas Constant R". J. Chem. Educ. 80 (7): 731. Print.
OĞUL ERSİN ÜNER D1129066
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M
Percentage Uncertainty of the Theoretical Acid Dissociation Constant at 35.0°C:
Determining the Theoretical Acid Dissociation Constant at 50.0°C:
M
Percentage Uncertainty of the Theoretical Acid Dissociation Constant at 50.0°C:
°C
Comment [a7]: It is not clear how this one has 1 sig fig
OĞUL ERSİN ÜNER D1129066
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Table 10: Results of the calculations, concentration of acetic acid values obtained from titration with their percentage uncertainties and the acid dissociation
constants of acetic acid when the acetic acid solution is at 20.0°C, 35.0°C and 50.0°C, with their percentage uncertainties.