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Solubility Rules: Three Suggestions for Improved Understanding Bob Blake Department of Chemistry and Biochemistry, Texas Tech University, Lubbock, TX 79409;
[email protected] The subject of the solubility of ionic compounds in water is discussed in general chemistry textbooks, typically within the first several chapters (1). Precipitation reactions are suitable for an introduction to chemical reactions because they are easy to observe. Once a simple understanding of solubility is achieved, the course of precipitation reactions is also fairly easy to predict. Unfortunately, many textbooks are plagued with at least one of the following problems in their discussion of solubility and precipitation reactions: a. Dualistic (2) treatment of solubility—substances are only classified as soluble or insoluble, with no mention of the range of solubility values b. Imprecise or inappropriate definition of solubility c. An entirely memorization-based approach to solubility
Student understanding of solubility and precipitation reactions can be improved if: a. More solubility data were included early in the discussion b. A precise, appropriate definition of solubility were provided c. Solubility were explained from simple physical principles
Context Almost ten years ago, I was preparing the solutions required for a demonstration of the formation of a silver mirror. Unfortunately, I was unable to find a bottle of silver(I) nitrate. Instead, I found a bottle of silver(I) sulfate. This was exciting for me because I remembered from my general chemistry textbook that all sulfates are soluble, with the exception
Table 1. Solubility of Common Inorganic Sulfate Compounds Formula
Solubility/(g/100 cc) 0.00246
0.00010
PbSO4
0.00425
0.00014
SrSO4
0.0113
0.0006
Hg2SO4
0.06
0.0012
CaSO4
0.209
0.015
Ag2SO4
0.57
0.018
Na2SO4
4.76
0.335
26.0
2.16
NOTE: Data from reference 3. Solubility measured at 25 ⬚C.
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Which of the following ions would not precipitate sulfate? a. Pb2+
b. Hg22+
c. Ca2+
d. Sr2+
Prior to writing the question, I diligently checked the solubility rules in the textbook to make sure that there would be one and only one correct answer. To my surprise, almost immediately upon posting the question, several of my students started yelling out that all of the choices would precipitate sulfate. Their reason for thinking that mercury(I) would precipitate sulfate is that they remembered that I told them that mercury(I) would precipitate sulfate, despite what our textbook indicated. This stirred mixed emotions because I was simultaneously happy that my students were paying attention and learning from lecture, but disappointed that there may have been an important omission from the solubility rules in our textbook. To clear this issue up for me and my students, I looked up the solubility values (Table 1) in the CRC Handbook of Chemistry and Physics (3). Clearly, if calcium sulfate is listed as insoluble, then mercury(I) sulfate should also be listed as insoluble. Dualism
Solubility/(mol/L)
BaSO4
MgSO4
of Pb2+, Ca2+, Sr2+, and Ba2+. Thus, I would be able to perform the demonstration despite my lack of silver(I) nitrate. To my surprise and disappointment, silver(I) sulfate did not dissolve. I walked to my office and checked the solubility rules to find that my memory was accurate, but the solubility rules were not. This story has served me well in my teaching of general chemistry. It highlights the reality that the world is too complex to be summarized with a short list of generalizations, and introduces the idea that even a textbook can contain errors. I was emphasizing these points to my class. I also went on to say that I would guess that mercury(I) sulfate might also be insoluble. Despite my lack of confidence in the solubility rules of any given textbook of general chemistry, I told my students “for the purposes of predicting precipitation reactions in this class, we will use the solubility rules in our textbook (1h)”. Later in the term, I posed the following multiple-choice question in class:
One of the biggest challenges in teaching science is to help dualistic thinkers (2) progress to a higher level of thinking and problem solving. Unfortunately, this development is difficult and often hampered by simplistic discussions that reduce complex phenomena to dualistic systems without reference to the inherent complexity of the situation. All chemists know that there are degrees of solubility, yet they do not necessarily present that information to students while discussing solubility in the context of precipitation reactions. Presenting data to students such as Table 1 would help avoid generating the misconception that substances are either soluble or insoluble.
Journal of Chemical Education • Vol. 80 No. 11 November 2003 • JChemEd.chem.wisc.edu
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Definitions One popular general chemistry textbook states: “A substance is said to be soluble if a fair amount of it visibly dissolves when added to water”(1d). This definition does not specify the quantity of water used and gives a very subjective criterion for the quantity of solid that dissolves to classify a solid as soluble. Even so, more than one modern textbook presents a definition that is equally imprecise. A classic article on solubility generalizations (4) defines the term soluble as meaning “greater than 5 g of salt [dissolves] in 100 mL of water”. This is a more useful definition because it states a specific criterion for the quantity of substance to be used and details the specific quantity of water in which it must dissolve to be classified as soluble. If we wish to create a dualistic rating system for a complex phenomenon, it is important to have a clear, specific criterion for decision making. Since the textbooks use solubility rules for predicting the outcome of mixing two solutions, the criterion for solubility should be useful for this purpose. A problem with some modern definitions is that they are inappropriate for their stated purpose. Solubility rules are generally given in the context of predicting the outcome of mixing two solutions. If a combination of cation and anion are insoluble, then the salt formed from those anions will be expected to precipitate. A solubility limit of 0.001 M is not generally useful for this purpose because of the typical concentrations of solutions found in most chemical laboratories. The solubility of calcium hydroxide is 0.185 g兾100 mL (0.025 M), thus most textbooks consider it to be one of the soluble hydroxides. If a 0.1 M NaOH solution were mixed with a 0.05 M CaCl2 solution, what would a student be expected to predict? Based on the solubility rules in many textbooks, all combinations of cation and anion in the resulting solution would produce soluble salts, thus the “correct” prediction would be that there would be no reaction. Despite this prediction, calcium hydroxide will precipitate from this solution because of its limited solubility. For the purposes of predicting whether a precipitation reaction would occur or not, solubility limits of 0.01 M or 0.001 M are inappropriate unless students are only to be asked about very dilute solutions. In the context of predicting whether a solid precipitates from a solution of ions or not, a criterion for solubility between 1 g and 5 g per 100 mL of water seems to be more appropriate. Physical Principles Another critical aspect of undergraduate education is the conceptual understanding of topics, rather than an educational experience that stresses memorization and algorithmic learning. Unfortunately, in the treatment of solubility, most textbooks treat solubility by asking students to memorize a list of rules. This type of learning is not conducive to a deeper understanding and conditions students to think that learning is equivalent to memorization. The consideration of one physical principle along with a basic understanding of dissociation can help students understand one important aspect of solubility. The attraction of oppositely charged ions is proportional to the magnitude of the charges of those ions. During dissociation, oppositely charged ions in the solid phase
are separated from each other. This would suggest that if a salt were composed of highly charged ions, it would not be soluble, but salts composed of ions with lower charges would be soluble. A single guideline, based on these simple rules, can serve as a starting point for the understanding of solubility. Any salt involving a +1 cation or a ᎑1 anion is likely to be soluble. This one guideline predicts many of the traditional rules (halides, nitrates, acetates, and perchlorates are soluble, while sulfides, carbonates, and phosphates are not) and most of the exceptions to these rules (ammonium and alkali metal salts are generally soluble, while calcium, mercury(I), strontium, lead, and barium sulfates are not). Learning a more complex system of rules gives greater accuracy, but with far more work. Chemically significant examples like the insolubility of silver halides (except silver fluoride) certainly should be brought to the attention of students because they are likely to encounter the use of silver(I) solutions as a test for halides in the laboratory. The importance of other factors that are very important in the overall solubility of solids should also be discussed, including the interaction of the dissolved species with the solvent, and entropy. Detailed, quantitative discussions of these issues can wait until later in the textbook. Errors and Omissions Several texts list the rule that all sulfates are soluble except salts of Ba2+, Pb2+, Sr2+, and Ca2+ (1). If calcium sulfate is listed as insoluble (solubility of 0.209 g兾100 mL, 0.015 M), then calcium hydroxide (solubility of 0.185 g兾100 mL, 0.025 M) and mercury(I) sulfate (solubility of 0.06 g兾mL, 0.0012 M) should also be characterized as insoluble. Some textbooks state that all sulfides are insoluble except those of ammonium, group I metals, and group II metals (1). According to the CRC Handbook of Chemistry and Physics (3), magnesium sulfide and barium sulfide decompose in solution, strontium sulfide is insoluble, and the solubility of calcium sulfide is 0.021 g兾100mL of water. Thus, none of the group II sulfides are soluble without decomposition. Conclusion In general, modern textbooks are complete and accurate, but a detailed examination of the discussions of solubility related to precipitation reactions reveals obvious areas for improvement. These improvements include the correction of factual errors, the removal of ill-defined dualistic discussions, and the introduction of a conceptual understanding of solubility. It is hoped that these considerations will lead to greater student understanding and a less memorization-based approach to learning. Literature Cited 1 (a) Silberberg, M. S. Chemistry: The Molecular Nature of Matter and Change, 3rd ed.; McGraw Hill: New York, 2003; pp 138–140. (b) Brown, T. L.; LeMay, H. E.; Bursten, B. E.; Burdge, J. R. Chemistry: The Central Science, 9th ed.; Prentice Hall: Upper Saddle River, NJ, 2003; pp 117–121. (c) Zumdahl, S. S.; Zumdahl, S. A. Chemistry, 5th ed.; Houghton Mifflin Company: Boston, MA, 2000; pp 147–152. (d)
JChemEd.chem.wisc.edu • Vol. 80 No. 11 November 2003 • Journal of Chemical Education
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Information • Textbooks • Media • Resources Chang, R. Chemistry, 7th ed.; McGraw Hill: New York, 2002; pp 108–111. (e) McMurray, J.; Fay, R. C. Chemistry, 2nd ed.; Prentice Hall: Upper Saddle River, NJ, 1998; pp 121–123. (f ) Jones, L.; Atkins, P. Chemistry: Molecules, Matter and Change, 4th ed.; W. H. Freeman and Company: New York, 2000; pp 102–104. (g) Ebbing, D. D.; Gammon, S. D. General Chemistry, 7th ed.; Houghton Mifflin Company: Boston, MA, 2002; pp 140–142. (h) Silberberg, M. S. Chemistry: The
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Molecular Nature of Matter and Change, 2nd ed.; McGraw Hill; New York, 2000; p 141. 2. Perry, W. G. Forms of Intellectual and Ethical Development in the College Years: A Scheme; Jossey-Bass Publishers: San Francisco, CA, 1998. 3. CRC Handbook of Chemistry and Physics, 62nd ed.; Weast, R. C., Astle, M. J. Eds.; CRC Press, Inc.: Boca Raton, FL, 1981. 4. Heimerzheim, C. J. J. Chem Educ. 1941, 18, 377.
Journal of Chemical Education • Vol. 80 No. 11 November 2003 • JChemEd.chem.wisc.edu
