Activity-Based Approach For Teaching Aqueous Solubility, Energy, and Entropy
Laura Eisen †, Nadia Marano *‡, and Samantha Glazier ‡
† Department of Chemistry, The George Washington University, Washington, DC 20007, United States
‡ Department of Chemistry, St. Lawrence University, Canton, New York 13617, United States
J. Chem. Educ., 2014, 91 (4), pp 484–491
DOI: 10.1021/ed4005563
Publication Date (Web): February 21, 2014
Copyright © 2014 The American Chemical Society and Division of Chemical Education, Inc.
Supporting Information (17 pages)
In this article, the authors present an activity-based approach in which students make observations to learn about the solubility of molecular and ionic substances. The students also receive guidance in using thermodynamic data to explain the observed solubility. The authors propose to incorporate energy minimization and entropy maximization principles in the discussion of solubility. They believe that even though entropy is not fully covered till later in the course, it can be introduced early on even though students may not understand all aspects of it. They point to research evidence that shows that repetition and discussion in multiple contexts help in student understanding and retention of concepts.
The activity they created focuses on the following concepts (taken verbatim from the article):
• Polar substances generally dissolve in polar solvents such as water, whereas nonpolar substances do not.
• Some soluble substances dissolve exothermically and others endothermically.
• The solubility of ionic compounds varies with charge density.
The activities created by the authors were based on their experience using the ACS General Chemistry textbook and feedback from workshops they have held. Table 1 in the article gives a good but brief summary of the topics, the observations students make, and the focus of discussion for each of the five activities. One of the things they want to emphasize is the improving observation/measurement skills and using data to reach conclusions.
The Supporting Information documents include the student worksheets along with the instruction guide. In Activity 1, the students carry out solubility tests for urea and ethanol in water and observe any temperature change to deduce whether the dissolving process is endothermic or exothermic. The focus of discussion is to show that the dissolving process can be endothermic or exothermic thus showing that energy minimization is not the only factor promoting solubility.
In Activity 2, the students look at structural properties that affect the solubility of molecular compounds in water. To appreciate the structure, students first have to build models of methanol, butanol, and hexane, predict their solubility in water, and then carry out the tests for actual observations.
In Activity 3, the students look at the solubility of ionic substances in water. They are asked to predict any temperature change in the dissolving process based on a table of lattice and hydration energies that are then actually tested.
In Activity 4, the students carry out microscale solubility/precipitation tests by mixing aqueous solutions of different ionic compounds. They are then asked to identify the soluble and insoluble compounds. From their data, they are asked to come up with some generalizations. As homework, students further work on coming up with a hypothesis to explain some general observations, write net ionic equations, and make predictions about other ionic compounds based on what they learned.
In the last activity, activity 5, students are asked to calculate thermodynamic values for change in enthalpy, entropy, and Gibbs free energy. They are then asked to look for patterns to explain solubility properties of the compounds.
Some of the more quantitative and thermodynamic aspect of these activities require that studens undertake these activities later on in the 1st semester, after they have learned about energy diagrams and the stepwise process for dissolution (after Chapter 11 in 1A which is the second to the last chapter covered in the first semester at LPC). Activity 5 would have to be done at the 1B level after students learn about entropy and free energy.
Two sections of the article are devoted to the authors’ discussion of the role of energy and the role of entropy in the dissolving process. In the role of energy, while the enthalpy change (sum of lattice energies and the hydration energies) can illustrate what ionic compounds dissolve exothermically or endothermically based on the overall enthalpy change, there is no discernible correlation between the enthalpy change and solubility (see data for multiply charged ions). The authors consider entropic considerations in understanding solubility that they are strongly advocating that entropy be introduced at this early stage even at the qualitative level. Even with large energy requirements for dissolution, some compounds dissolve readily because the large energy input required is compensated for by an increase in entropy. The authors do note also that water reorganization needs to be accounted for when looking at entropy changes (not just the mixing process which increases entropy). This is a subtle factor that makes all the difference in cases when two things are not soluble in one another. For example, in the example of 1-butanol showing no or very little solubility in water despite its hydrogen-bonding capability, the authors (based on cited papers) point to the “decrease in entropy that arises from a shell of water molecules arranged in a fixed orientation around the nonpolar hydrocarbon end of 1-butanol” that makes it insoluble in water. The decrease in entropy of water is more than the increase in mixing entropy. The movement of water around the solute becomes too restricted compared to more degrees of freedom of motion in pure water. See Figure 3.
The reorganization and resulting lower entropy of water comes up again in understanding why ionic compounds with multiply charged ions are less likely soluble. Because of their high charge densities, these ions have both high (more negative) lattice and hydration energies so energy minimization alone can not account for why they are insoluble. The high hydrate energies imply that the water molecules are strongly held by the ions and are there again restricted in their movements which leads to lower entropy compared to when they are in pure form. Exceptions to this noted by the authors are the insolubility of AgCl containing monovalent ions (explained by the high net dissolution enthalpy) and of sulfates soluble (explained by highly favorable net enthalpy of dissolution change).
The correlation of solubility to favorable entropy changes and ultimately to free energy changes are treated in detail in Activity 5, usually appropriate only during the second semester after completing the thermodynamic sections of the 1-year General Chemistry course. In this activity, students will discover, as noted earlier, that most of the exceptions to generalization that monovalent ion compounds are more likely soluble than those containing multivalent ions are due to net enthalpy change values. In further integrating other structural correlations, students are asked to predict the properties (both solubility and thermodynamic) of the compound LiF thus invoking size effects and charge densities as important factors in predicting and understanding the behavior of particles in water. See Figure 5.
In the last section of the article, the authors share some new insights from current research on the behavior of water when pure and when in contact with solutes. As they have noted previously in the article, water ordering has been observed to occur when in contact with small ions and with urea although the exact nature of this ordering is still not clear. Solutes like butanol and Mg2+ are kosmotropes, their interaction with water causing water molecules to become more ordered, whereas urea and Na+ are chaotropes, their interactions with water causing more disorder. The interaction that fixes water in a single orientation is strong dipole-dipole as some research has shown. They also noted that solvation energies can be affected by the sign of similarly-charged and similarly-sized ions because of the asymmetric nature of the charge distribution in a water molecule: F- has a negative solvation energy while Rb+ has a positive solvation energy. The negative charge can get closer to the positive surface of the water molecule on the H atoms; there is a longer interaction distance between a cation and the negative oxygen.
With nonpolar molecules, modelling shows that water packs tightly around the nonpolar molecules resulting in smaller cavities for solutes, less likely solvation, and decreased entropy for water.
In addition to the exposure of students to integrating various concepts encountered at different points during the semester and development of their observation skills and abilities to see patterns and generalizations, the authors end with this note as another reason why this approach is an improvement over the traditional method: Research in all of the major subdisciplines of chemistry relies on entropy and energy to explain such diverse topics as protein folding, water-filled carbon nanotubes, activated complexes in mechanisms, and binding reactions. Textbooks necessarily evolve with current research, and we feel that the substantial body of literature on the importance of entropy in explaining chemical phenomena justifies a parallel shift in our teaching.
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