Laboratory Experiment pubs.acs.org/jchemeduc
Undergraduate Analytical Chemistry Experiment: The Determination of Formation Constants for Acetate and Mono- and Dichloroacetate Salts of Primary, Secondary, and Tertiary Methyl- and Ethylamines Ronald P. D’Amelia,* Stephanie Chiang, Stephanie Pollut, and William F. Nirode* Department of Chemistry, Hofstra University, Hempstead, New York 11549, United States S Supporting Information *
ABSTRACT: The formation and the hydrolysis of organic salts produced by the titration of a 0.1 M solution of the following amines: methyl-, dimethyl-, trimethyl-, ethyl-, diethyl-, and triethylamine with a 0.1 M solution of acetic, chloroacetic, and dichloracetic acids are studied. The pKb of the amine and the pH at the end point were determined for each reaction from the corresponding titration curve. The experimentally determined values were in very good agreement with the literature and calculated values. The salt formation constants (Ksf) for acetate and mono- and dichloroacetate salts of primary, secondary, and tertiary methyl- and ethylamines were also determined. The conclusions were as follows: (1) the pH at the end point decreased as the strength of the titrating acid increased, (2) the stronger the amine base, the higher the pH at the end point, and (3) the Ksf increased with increasing acid and base strength. KEYWORDS: Second-Year Undergraduate, Analytical Chemistry, Laboratory Instruction, Hands-On Learning/Manipulatives, Acid−Base, Titration/Volumetric Analysis
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undergraduate chemistry laboratory classes),5,6 (2) determine the Ka of the selected acid, (3) determine the pKb from the titration curve, (4) determine the pH at the end point from the titration curve, (5) examine the titration curves and contrast these with strong acid−weak base and weak acid−strong base titrations, (6) calculate the salt formation constants (Ksf) for the reaction of an amine and a carboxylic acid, and (7) understand the effect of additional chloro substituents on the strength of the acid.
rganic salts can be formed by the reaction of a carboxylic acid and an amine base. The Bronsted-Lowry acid−base theory describes the mechanism as the carboxylic acid donating a proton to the amine base, thus producing a cationic amine− anionic carboxylate called an organic salt. The equilibrium constants of acid−base reactions have been reported by Thompson et al.1 Equilibrium constants from hydrolysis and acid−base reactions can be experimentally determined from basic titrimetric principles.2 The acidity and basicity of a solution produced by different organic salts depends on the strength of the acid and the base used. The pH of a solution produced during a reaction between an acid and a base varies. The solution is basic when Kb > Ka and acidic when Ka > Kb. If Ka and Kb are known, then reaction completeness or Ksaltformation, Ksf, can be determined. In the pharmaceutical industry, the Ksf is used to help optimize drug development. Salt forms of drugs have a large effect on the drugs’ quality, safety, and performance. The properties of salts affect pharmacological properties of the drug such as solubility, bioavailability, and stability. The importance of organic salts in pharmaceutical formulations provides a real-world application.3,4
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THEORY According to equilibrium rules when several balanced chemical reactions are algebraically added, the net equilibrium constant is the product of the individual reaction equilibrium constants (Knet = K1 × K2 × K3, etc.). The series of eqs 1−8 show the derivation of Ksf from the relevant equilibrium expressions RNH 2(aq) + H 2O(1) ⇌ RNH3+(aq) + OH‐(aq) K1 = Kb CH3COOH(aq) + H 2O(1) ⇌ CH3COO−(aq) + H3O+(aq)
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K 2 = Ka
LEARNING OBJECTIVES The purpose of this experiment is to develop an understanding of the principles of acid−base equilibria. By performing this experiment, students should learn how to (1) use the Vernier Logger Pro software and automatic Labquest titration system (Vernier and other similar devices have been used in © XXXX American Chemical Society and Division of Chemical Education, Inc.
(1)
(2)
H3O+(aq) + OH‐(aq) ⇌ 2H 2O(1) K3 = 1/K w
A
(3)
dx.doi.org/10.1021/ed400844y | J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
Laboratory Experiment
(the exact value was provided to 4 significant figures), and each standardized the given acid using a buret and indicator or pH electrode. The following week, each student used the Vernier Labquest system to titrate the assigned amine series using the standardized ∼0.1 M acid. After the titration curves were obtained, pH at the half neutralization point, the pH at the end point, and the volume of titrant were determined using the Vernier Labquest computer program. To determine the pH at the end point a graph of the first derivative curve of the original titration curve was obtained.
RNH 2(aq) + CH3COOH(aq) ⇌ RNH3(aq) + CH3COO−(aq) where K net = K1 × K 2 × K3 = (Kb × K a)/K w and K sf = K net (4)
From the ionization of the organic amine, using [OH−] = Kw/ [H3O+] and solving for [H3O+] ⎛ K ⎞ [RNH3+] [H3O+] = ⎜ w ⎟ ⎝ Kb ⎠ [RNH 2]
(5)
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From the ionization of CH3COOH [CH3COOH] [H3O ] = K a [CH3COO−]
HAZARDS All of the above reagents are corrosive, capable of causing severe burns, and flammable. They should all be handled with suitable gloves and in the hood for proper ventilation. All of the amine compounds have an unpleasant odor.
+
(6)
After combining eqs 5 and 6, it can be determined that at the end point, where [CH3COOH] = [RNH2] and [RNH3+] = [CH3COO−] the hydronium ion is then given by eq 7 1/2 ⎧ K ⎫ [H3O+] = ⎨K w × a ⎬ Kb ⎭ ⎩
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RESULTS AND DISCUSSION Titration of the amine base with the standardized organic acid allows for the determination of the end point from the titration curve and the corresponding first derivative plot (see the Supporting Information for an example). The theoretical pH at the end point can be calculated by using eq 7. Table 1 is a
(7)
The Ka of the organic acid can be determined from the initial pH and the initial concentration of the acid, [HA], and using eq 8 Ka =
[H3O+][A−] [HA]
Table 1. Summary of the Student Experimental pH at the End Point (EP) of 0.1 M Amine Titrated with 0.1 M Acida (8)
where [H3O+] = [A−].
Acetic Acid
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MATERIALS All reagents were purchased from Sigma-Aldrich and are all reagent grade chemicals. Methylamine and dimethylamine were 40 wt % solutions in water; trimethylamine was a 25 wt % solution in water; ethylamine was a 70 wt % solution in water; diethylamine and triethylamine were 99.5% pure. Glacial acetic acid, chloroacetic acid, and dichloroacetic acid were greater than 99% pure. Chloroacetic acid was a solid. A standardized 0.1 M solution of NaOH (exact value is stated to 4 significant figures) was provided.
Chloroacetic Acid
Dichloroacetic Acid
Amine
pH at EP
Theor pH
pH at EP
Theor pH
pH at EP
Theor pH
Methylamine Dimethylamine Trimethylamine Ethylamine Diethylamine Triethylamine
7.72 7.73 7.26 7.72 7.78 7.73
7.70 7.74 7.27 7.70 7.79 7.75
6.74 6.83 6.22 6.76 6.86 6.78
6.76 6.79 6.33 6.75 6.85 6.80
6.06 6.09 5.50 5.94 6.11 6.09
6.01 6.04 5.58 6.00 6.10 6.05
The % errors were calculated, and they averaged ±0.7% between the experimental pH and theoretical pH.
a
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comparison of the experimental pH and the theoretical pH at the end point. This data is compiled from a class of 8−12 students. As the number of chlorines increased, the acid strength increased. As the acid strength increased, the pH at the end point decreased. As the base strength increased, the pH at the end point increased.
APPARATUS The pH was measured using a Vernier Labquest system consisting of a Vernier computer interface connected to a Vernier pH sensor (PH-BTA) and Vernier drop counter (VDC-BTR) and a Dell PC. The data collection software used to analyze the pH−volume data obtained from the titrations was the Vernier Logger Pro 3 edition 3.6.1.
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Table 2. Student Values of pH at Half End Point, Experimental and Literature Values of pKb, and % Error
EXPERIMENTAL DETAILS This experiment was carried out by a class of 8−12 students in an undergraduate quantitative analysis chemistry laboratory over a two week period (totaling approximately 8−10 h). The students worked independently with specific acids and bases assigned by the instructor. They collated their data for the final analysis. During the first week, each student prepared a ∼0.1 M solution of an instructor-assigned acid (either acetic, chloroacetic, or dichloroacetic acid) and ∼0.1 M solutions of an instructor-assigned amine series (either the methyl-, dimethyl-, trimethylamine or the ethyl-, diethyl-, triethylamine series). Each student was given a standard 0.1 M solution of NaOH
Aminea
Exp pH at Half End point
Exp pKb
Litb pKb
% Errorc
Methylamine Dimethylamine Trimethylamine Ethylamine Diethylamine Triethylamine
10.63 10.76 9.79 10.61 10.87 10.72
3.37 3.24 4.21 3.39 3.13 3.28
3.34 3.27 4.20 3.35 3.16 3.25
0.90 −0.92 0.24 1.19 −0.95 0.92
a Amines were reacted with dichloroacetic acid. bLiterature values were obtained from the Handbook of Chemistry and Physics, 87th edition, 2006−2007. cThe % error between experimental and literature pKb values.
B
dx.doi.org/10.1021/ed400844y | J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
Laboratory Experiment
Table 3. Student Experimental and Theoretical Salt Formation Constants (Ksf) for Each of the Organic Acids with Each Aminea Acetic Acid Amine
Exp Ksf
Methylamine Dimethylamine Trimethylamine Ethylamine Diethylamine Triethylamine
× × × × × ×
7.47 1.01 1.08 7.12 1.30 9.19
+5
10 10+6 10+5 10+5 10+6 10+5
Chloroacetic acid Theor Ksfb
Exp Ksf
× × × × × ×
× × × × × ×
7.99 9.36 1.10 7.82 1.21 9.84
+5
10 10+5 10+5 10+5 10+6 10+5
5.76 7.76 8.32 5.49 9.98 7.08
+7
10 10+7 10+6 10+7 10+7 10+7
Dichloroacetic acid
Theor Ksfb
Exp Ksf
× × × × × ×
× × × × × ×
6.16 7.22 8.51 6.03 9.34 7.58
+7
10 10+7 10+6 10+7 10+7 10+7
1.91 2.57 2.76 1.82 3.31 2.35
Theor Ksfb +9
10 10+9 10+8 10+9 10+9 10+9
2.04 2.39 2.82 1.99 3.09 2.51
× × × × × ×
10+9 10+9 10+8 10+9 10+9 10+9
The % errors were calculated and they averaged ±5% between Ksf theoretical and Ksf experimental. bThe Ksf theoretical was calculated from Ka and Kb values obtained from Handbook of Chemistry and Physics, 87th edition, 2006−2007 a
instead of traditional trial and error methods. Students appreciated an alternative titration technique using the Vernier Labquest system and enjoyed the experiment based on the comments received.
Once the students determine the pH at the end point, they can determine the pH at half the volume of the end point. The pH at half the volume of the end point can be used to determine the pKb of the amine base by subtracting the pH from 14.00. The experimentally determined pKb values taken from the titrations are shown in Table 2. A comparison between the experimental pKb and the theoretical pKb is made with very low percent errors. Once the reactions for the organic acid and amine base are completed, students can algebraically determine the overall chemical reaction for the formation of the organic salt (see eqs 1−3). Using the experimentally determined Ka and Kb values, the Ksf for the organic salt can be calculated (see eq 4). The comparison of these calculated Ksf theoretical values and Ksf experimental values obtained from the titrations explained above are shown in Table 3.
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ASSOCIATED CONTENT
S Supporting Information *
Student handout and instructor notes. This material is available via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Authors
*E-mail:
[email protected]. *E-mail:
[email protected]. Notes
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The authors declare no competing financial interest.
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CONCLUSIONS This lab experiment promotes the students’ understanding of acid−base titrations and equilibrium reactions while introducing the calculation of Ksf. After doing the titrations and collecting all the necessary data, students draw conclusions from this experiment such as (1) the experimental pKb values for all the amines studied were in very good agreement with the literature, (2) the experimental pH values at the end point in the titration of the amines with acetic acid, chloroacetic acid, and dichloroacetic acid were in very good agreement with the theoretical values, (3) as the acid strength increased the pH at the end point decreased; dichloroacetic acid is the strongest acid and has the lowest pH at the end point, (4) as the base strength increased, the pH at the end point also increased; the base strength for the methylamine series is dimethyl-, methyl-, and trimethyl- amine from highest to lowest and the base strength for the ethylamine series diethyl-, triethyl- and ethyl- is from highest to lowest (see the Instructor Notes in the Supporting Information for the corresponding figures), (5) the experimental Ksf was in very good agreement with the theoretical Ksf, and (6) Ksf increased as the acid and base strengths increased. The students were evaluated based on their responses to the thought questions in the Student Handout and on the measured values and percent relative errors of the Kb and Ksf values from the selected amines titrated with the selected acid. The students used compiled class data to evaluate the other acids with their selected amines as well as the other set of amines. Most of the students said that they significantly improved their understanding of the principles of acid−base equilibrium and salt formation reactions, and they were able to make the connection from acid−base equilibria to calculate Ksf and how Ksf values can be used to screen drug candidates
ACKNOWLEDGMENTS This work was supported by a Hofstra HCLAS Faculty Research and Development Grant. REFERENCES
(1) Thompson, R. J. The Extent of Acid-Base Reactions. J. Chem. Educ. 1990, 67, 220−221. (2) Nyasulu, F.; McMills, L.; Barlag, R. Weak Acid Ionization Constants and the Determination of Weak Acid-Base Reaction Equilibrium Constants in the General Chemistry Laboratory. J. Chem. Educ. 2013, 90, 768−770. (3) Brittain, H. G. Strategy for the Prediction and Selection of Drug Substance Salt Forms. Pharm. Technol. 2007, 77−88. (4) Kumar, L.; Amin, A.; Bansal, A. K. Salt Selection in Drug Development. Pharm. Technol. 2008, 128−146. (5) Vannatta, M. W.; Richards-Babb, M.; Solomon, S. D. Personal Multifunctional Chemical Analysis System for Undergraduate Chemistry Laboratory Curricula. J. Chem. Educ. 2010, 87, 770−772. (6) Nyasulu, F.; Moehring, M.; Arthasery, P.; Barlag, R. Ka and Kb from pH and Conductivity Measurements: A General Chemistry Laboratory Exercise. J. Chem. Educ. 2011, 88, 640−642.
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dx.doi.org/10.1021/ed400844y | J. Chem. Educ. XXXX, XXX, XXX−XXX
