Demonstration pubs.acs.org/jchemeduc
Are Aqueous Solutions of Amphiprotic Anions Acidic, Basic, or Neutral? A Demonstration with Common pH Indicators Jervee M. Punzalan and Voltaire G. Organo* Department of Physical Sciences and Mathematics, College of Arts and Sciences, University of the Philippines, Manila 1000, Philippines S Supporting Information *
ABSTRACT: A simple, fast, and microscale qualitative demonstration illustrating the acidic and basic nature of amphiprotic anions suitable for senior high school or undergraduate general chemistry classes is presented. The demonstration involves frequently used pH indicators and salt solutions in general chemistry laboratory classes. This will provide an excellent avenue to engage students in hydrolysis and equilibrium concepts involving polyprotic anions in a more visually appealing approach. This demonstration can be completed in less than 10 min provided all reagents have been prepared prior to experiment. KEYWORDS: High School/Introductory Chemistry, First-Year Undergraduate/General, Demonstrations, Hands-On Learning/Manipulatives, Inquiry-Based/Discovery Learning, Aqueous Solution Chemistry, Acids/Bases, Dyes/Pigments, Equilibrium
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INTRODUCTION The acid−base concept is an essential topic in a chemistry course. From the chemical reactions in our cells, to enzymes, to pH-related behavior of drugs, to commercial products and cleaning agents, the importance of having a background in acid−base chemistry cannot be overemphasized. However, most students find it challenging to understand several topics related to the acid−base concept. Several reports published in this Journal include difficult topics such as relationships among strong and weak acids and bases, hydrolysis, equilibrium, and amphoterism.1−3 In a study done by Paik,4 a tabulated literature review of the types of and reasons for students’ difficulties in understanding acid−base concepts was presented. This includes identifying acids and bases and their relative strengths. It was found that students often predict acids or bases by simply looking for the presence of H+ or OH− in the formula. This has been attributed to the lack of skills and knowledge on conjugate acid−base pair concepts.4 To address such difficulties, researchers have devised a plot that depicts relationships between strong and weak acids and bases and their conjugates.1 Others have suggested methods like using conjugate acid−base charts to visualize strengths of acids and bases.5 One of the more problematic acid−base concept deals with the nature of amphiprotic anions. There is a general perception © XXXX American Chemical Society and Division of Chemical Education, Inc.
that anions are basic species since most examples of base hydrolysis taught in general chemistry involve anions.6 However, amphiprotic anions can also act as Bronsted− Lowry acids in aqueous solution. On the other hand, since amphiprotic anions such as HCO3− and HPO42− contain H+ in their formula, students may regard such anions as acids only. They may even find it difficult to determine whether salts containing amphiprotic anions such as HCO3− and HPO42− will result in acidic, basic, or neutral solutions. Predicting the acid−base properties of these amphiprotic anions is not as direct as any other anion. This requires knowledge of both ionization constants Ka and Kb. If we consider an aqueous solution of NaHCO3, we find that it can undergo two equilibrium reactions:7 base hydrolysis HCO3−(aq) + H 2O(l) ⇌ H 2CO3(aq) + OH−(aq) Kw Kb = 2.4 × 10−8 = K a,H2CO3 Received: September 15, 2016 Revised: May 8, 2017
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DOI: 10.1021/acs.jchemed.6b00711 J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
Demonstration
Figure 1. (A) Addition of Thymol blue indicator and (B) neutral, acidic, and basic forms of Thymol blue indicator.
solutions. Use extra caution in handling the following compounds due to their associated hazards: potassium carbonate, sodium hydrogen sulfate, methyl orange, bromocresol green, and alizarin are all irritants; sodium hydroxide, hydrochloric acid, and sodium phosphate are all corrosive; congo red and phenolphthalein are suspected carcinogens.
acid ionization HCO3−(aq) + H 2O(l) ⇌ CO3(aq)2 − + H3O+(aq) K a = 4.8 × 10−11
Here, it can be inferred that HCO3− will produce a basic solution since its Kb > Ka. It is important to note that in an aqueous solution of NaHCO3, while the acid−base conjugate pair (CO32− and HCO3−) species are both present, the solution predominantly contains HCO3−. Thus, hydrolysis of HCO3− producing more OH− ions in solution outweighs formation of H3O+ ions via ionization of HCO3−. Educators and students alike often use the above equilibrium expressions and ionization constants to understand the nature of amphiprotic anions.1,5,8 However, there is a need for concrete experimental evidence to support such abstract concepts. A simple and visual demonstration based on color changes may enhance learning and appreciation of theoretical viewpoints. In this demonstration, we use common pH indicators to determine whether sodium or potassium salts containing amphiprotic anions will produce acidic, basic, or neutral solutions. This approach can aid students integrate concepts such as base hydrolysis and acid ionization of amphiprotic anions, including acid/base strength.
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DEMONSTRATION
This demonstration can be presented in laboratory classes to complement lectures on acid−base equilibria particularly after introducing concepts such as Ka, Kb, hydrolysis, amphiprotic anions, and polyprotic acid−base system. This is also suitable for introducing students to common pH indicators and their corresponding colors in various pH ranges. Both indicators and aqueous salt solutions are prepared prior to demonstration. Indicator solutions are prepared as described in the Acid Base Indicators section of the CRC Handbook,9 which can also be found in the Supporting Information. Students are first tasked to mix 1−3 drops of indicator with 8−10 drops of distilled water, 0.1 M HCl, and 0.1 M NaOH in spot plates. These will serve as reference solutions to determine the colors of the indicator in neutral, strongly acidic, and strongly basic solutions, respectively. An example of this preparation using thymol blue as indicator is shown in Figure 1. Students are then instructed to predict whether the anions will produce an acidic, basic, or neutral solution by applying what they have learned from lectures. To test their predictions, they are given eight sets of 12-well spot plates where, in each well, 1−3 drops of selected indicator are mixed with 8−10 drops of anion solution. (See Supporting Information for sample setup.) Students then compare the observed colors to that of the reference solutions. They can also be asked to classify the given anions into strongly acidic, slightly acidic, slightly basic, and strongly basic. The pH of anion solutions can be measured to support their observations. Finally, the class is prompted to explain their observations based on acid−base equilibrium principles. It is worth noting that not all indicators used are found to differentiate anionic conjugate acid−base pairs. Thymol blue, phenolphthalein and alizarin are useful in differentiating CO32− from HCO3−. On the other hand, thymol blue, congo red,
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MATERIALS AND CHEMICALS All salt solutions and pH indicators are prepared using analytical grade reagents and distilled water. Potassium salts of CO32−, HCO3−, and HPO42− and sodium salts PO43−, H2PO4−, SO42−, and HSO4− are prepared at a concentration of 0.1 M (aq). All pH indicator solutions which can be used in this demonstration are listed in the Supporting Information. Other materials and equipment include a pH meter, spot plates, and droppers.
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HAZARDS The low concentrations and quantities of reagents used in this demonstration minimized the hazards and waste disposal concerns. However, it is still imperative to wear safety goggles, gloves, and appropriate clothing at all times. In addition, qualified personnel in the laboratory must prepare the indicator B
DOI: 10.1021/acs.jchemed.6b00711 J. Chem. Educ. XXXX, XXX, XXX−XXX
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Table 1. Comparison of Color Changes of Reference, Carbonate, and Hydrogen Carbonate Solutions under Selected pH Indicators
Table 2. Comparison of Color Changes of Reference, Sulfate, and Hydrogen Sulfate Solutions under Selected pH Indicators
true at low H+ concentration (high pH) where the In− color predominates.11 For example, thymol blue exhibits a color change from red at pH < 1.2 to yellow at pH 2.8−8 and to blue at pH > 9.6. At pH between 1.2 and 2.8, the resulting color is somewhere between gradients of red to yellow, while at pH between 8 and 9.6, gradients of yellow to blue can be observed. pH indicators are therefore useful in determining whether salts containing amphiprotic anions will produce acidic, basic, or neutral solutions. The pH indicators and anions involved in this demonstration are frequently used in high school level general chemistry classes. Depending on the available reagents in the laboratory, any sodium or potassium salt may be used as the anion source. This will avoid certain drawbacks of using alkaline earth salts such as insolubility or sparing solubility in water as well as susceptibility to thermal decomposition.12 In addition, this will also ensure that the cations will not interfere in color change since potassium or sodium ions lack any acidic/basic property in water. For comparison and baseline, color changes of the three reference solutions under various pH indicators were noted. Solutions of carbonate, sulfate, and phosphate anions as well as their respective amphiprotic anions upon addition of indicators
methyl orange, bromophenol blue, bromocresol green, methyl red, and bromothymol blue can be used for SO42− and HSO4−. Lastly, thymol blue, bromocresol green, methyl red, bromothymol blue, phenol red phenolphthalein and alizarin can differentiate phosphates. This poses a challenge to students in selecting the most useful indicators for analysis.
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DISCUSSION Acidic or basic solutions are usually determined in general chemistry classroom experiments using pH indicators. These are weak acids or bases that change color depending on the pH.10 In general, the observed color changes are consequences of protonation or deprotonation of an indicator. Consider the following hypothetical chemical equation at equilibrium where HIn and In− represent the protonated and deprotonated forms of the indicator, respectively. Hln ⇌ H+ +
(color 1)
ln−
(color 2)
Since these species differ in color, one can readily observe a change in color of the solution depending on the predominating form, protonated or deprotonated. At a high H + concentration (low pH), the equilibrium shifts to the left, favoring the appearance of HIn with color 1. The opposite is C
DOI: 10.1021/acs.jchemed.6b00711 J. Chem. Educ. XXXX, XXX, XXX−XXX
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Table 3. Comparison of Color Changes of Reference, Phosphate, Hydrogen Phosphate, and Dihydrogen Phosphate Solutions under Selected pH Indicators
For aqueous solutions of SO42− and HSO4− ions, the resulting color changes indicate that HSO4− is more acidic than SO42− (Table 2). The resulting color of HSO4− solution is similar to that of HCl while SO42− resembles the color of dH2O for all selected indicators. We can assume that HCl and HSO4− have approximately the same pH value of 1.0. In the same way, the pH of SO42− may be close to the pH of dH2O. The acidic nature of HSO4−(aq) can be explained by considering the two possible equilibria of HSO4− in aqueous solution:
were observed for changes in color. Results are summarized in Tables 1−3. All reference solutions were consistent with the expected results except for bromothymol blue subjected to 0.1 M HCl, which turned orange instead of the expected yellow color. The exposure of bromothymol blue to an extremely acidic pH possibly led to the appearance of orange color instead of the expected yellow color. This was verified by subjecting bromothymol blue to varying HCl concentrations. (See Supporting Information.) Table 1 shows the color comparisons of selected indicators in the presence of CO32− and HCO3− anions with respect to reference solutions. On the basis of the colors, both CO32− and HCO3− are basic in aqueous solution. However, it can be observed that CO32− is more basic than HCO3−. In thymol blue, for example, the blue color of CO32− is similar to that of NaOH solution. This implies that CO32− solution has a pH close to that of NaOH solution. On the other hand, HCO3− solution produced a green color, which is an intermediate of yellow and blue colors. This indicates that the pH of the HCO3− solution falls between the pH of dH2O (pH 7.20) and NaOH solution (pH 13.87). Phenolphthalein and alizarin also exhibited similar solution behaviors. These results can be accounted for by the difference in Kb values of the anions. CO32− is known to have a higher Kb of 2.1 × 10−4 and thus is more basic than HCO3− with Kb of 2.4 × 10−8. Furthermore, the basic property of HCO3− can be explained by the difference in the Ka value compared to its Kb. As previously shown, the Ka of HCO3− is less than the Kb. Hence, more OH− ions are produced by the hydrolysis of HCO3− compared to the H3O+ produced by the ionization of HCO3−. It is also noted that the difference in color between CO32− and HCO3− will only be visible if these solutions are freshly prepared. It is therefore necessary to observe proper storage of these solutions since the carbonate−bicarbonate system is unstable especially upon exposure to air.13
base hydrolysis HSO4 −(aq) + H 2O(l) ⇌ H 2SO4 (aq) + OH−(aq) Kw Kb = 4.2 × 10−21 = K a,H2SO4
acid ionization HSO4 −(aq) + H 2O(l) ⇌ SO4 2 −(aq) + H3O+(aq) K a = 1.3 × 10−2
Since the Ka value outweighs the Kb of HSO4−, more H3O+ is produced in solution, thus making the solution acidic. On the other hand, relatively small amounts of OH− ions are produced from the hydrolysis of SO42− due to its small Kb value of 7.7 × 10−13. This value is even smaller compared to the Kb value of the more basic CO32−. This may be the reason why the color of SO42− in thymol blue is yellow compared to that of CO32− which is blue. Table 3 shows the color comparisons of selected indicators in the presence of PO43−, HPO42− and H2PO4− anions. On the basis of the color changes, it can be inferred that PO43− is strongly basic, and HPO42− is slightly basic while H2PO4− is slightly acidic in aqueous solution. In thymol blue, for example, the blue color of PO43− is similar to that of NaOH solution while the green color of HPO42− is an intermediate of yellow and blue colors of dH2O and NaOH solutions, respectively. On the other hand, the yellow color of H2PO4− is similar to the D
DOI: 10.1021/acs.jchemed.6b00711 J. Chem. Educ. XXXX, XXX, XXX−XXX
Journal of Chemical Education
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color of dH2O. However, the color of H2PO4− in other indicators such as bromocresol green, methyl red, bromothymol blue, phenol red, and alizarin falls in between the colors of HCl and dH2O which indicates that H2PO4− is slightly acidic. On the basis of the ionization constant for PO43− (Kb= 2.1 × 10−2), more OH− ions are produced resulting in a basic solution. For HPO42−, two equilibrium systems are possible:
Demonstration
ASSOCIATED CONTENT
* Supporting Information S
The Supporting Information is available on the ACS Publications website at DOI: 10.1021/acs.jchemed.6b00711. Experimental details including sample setup, pH information for indicators and anion solutions, protocols, and color change information (PDF, DOCX)
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base hydrolysis HPO4 2 −(aq) + H 2O(l) ⇌ H 2PO4 −(aq) + OH−(aq) Kw Kb = 1.6 × 10−7 = K a,H2PO4−
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
Voltaire G. Organo: 0000-0001-9272-4433
acid ionization
Notes
HPO4 2 −(aq) + H 2O(l) ⇌ PO4 3 −(aq) + H3O+(aq)
The authors declare no competing financial interest.
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−13
K a = 4.8 × 10
However, Kb is much greater than Ka, which results in a basic solution but with a lower pH compared to PO43− ion solution. Finally, H2PO4− can also undergo two equilibrium reactions: base hydrolysis H 2PO4 −(aq) + H 2O(l) ⇌ H3PO4 (aq) + OH−(aq) Kb = 1.3 × 10−12 =
REFERENCES
(1) Adcock, J. Teaching Bronsted-Lowry Acid-Base Theory in a Direct Comprehensive Way. J. Chem. Educ. 2001, 78 (11), 1495. (2) Shamai, R.; Stavy, R. Teaching an Introductory Course in Qualitative Analysis in Order to Enhance Learning General Chemistry. J. Chem. Educ. 1986, 63 (8), 707. (3) Tesh, K. The Use of Potassium Alum in Demonstrating Amphoterism. J. Chem. Educ. 1992, 69 (7), 573. (4) Paik, S. Understanding the Relationship Among Arrhenius, Brønsted−Lowry, and Lewis Theories. J. Chem. Educ. 2015, 92 (9), 1484−1489. (5) Treptow, R. The Conjugate Acid-Base Chart. J. Chem. Educ. 1986, 63 (11), 938. (6) Sattsangi, P. Microscale Procedure for Inorganic Qualitative Analysis with Emphasis on Writing Equations: Chemical Fingerprinting Applied to then-Bottle Problem of Matching Samples with Their Formulas. J. Chem. Educ. 2014, 91 (9), 1393−1400. (7) Chang, R. General Chemistry, 1st ed.; McGraw-Hill: Boston, 2008; pp 548, 557−558, 561−562. (8) Davidson, D. Amphoteric Molecules, Ions and Salts. J. Chem. Educ. 1955, 32 (11), 550. (9) Lide, D.; Frederikse, H. CRC Handbook of Chemistry and Physics; CRC Press: Boca Raton, FL, 1995; Vol. 1995−1996. (10) Brenner, R.; Hess, K.; Morford, J. Understanding Electrophoresis Through the Investigation of Size, Shape, and Charge of pH Indicators. J. Chem. Educ. 2015, 92 (10), 1705−1708. (11) Silva, C.; Pereira, R.; Sabadini, E. Color Changes in Indicator Solutions. An Intriguing and Elucidative General Chemistry Experiment. J. Chem. Educ. 2001, 78 (7), 939. (12) Mittal, A. Chemistry; A.P.H. Publ. Corp: New Delhi, 2007; pp 29−30. (13) Sodium Bicarbonate Chemistry; Integrated Biomedical Technology, Inc.: Elkhart, IN, 2003.
Kw K a,H3PO4
acid ionization H 2PO4 −(aq) + H 2O(l) ⇌ HPO4 2 −(aq) + H3O+(aq) K a = 6.2 × 10−8
In this case, ionization of H2PO4− to generate H3O+ ions predominates over hydrolysis of H2PO4− to produce OH− ions, thus making this solution acidic.
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CONCLUSION Color changes of pH indicators affected by ionization or hydrolysis of amphiprotic anions can help students determine whether the anions will produce acidic, basic, or neutral aqueous solutions. With this demonstration, they also get to explain color changes using acid−base equilibrium concepts. Moreover, educators are presented with an innovative approach in teaching topics such as equilibrium, acid−base conjugates, and pH indicators through visual representation of chemical reactions.
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LIMITATIONS Indicators respond to the OH− and H+ ions produced by anions in aqueous solution. In this study, it is assumed that only hydrolysis or ionization of amphiprotic anions lead to significant OH− and H+ quantities thereby altering the pH and consequently leading to a color change. The cations used in this experiment, potassium and sodium, were assumed to have negligible effects on pH and hence have no role in causing color change. Finally, the tests are neither specific nor selective to a certain anion but good for a given concentration of OH− and H+. E
DOI: 10.1021/acs.jchemed.6b00711 J. Chem. Educ. XXXX, XXX, XXX−XXX
