Successful learning of many aspects in physiology depends on a meaningful understanding of fundamental chemistry concepts. Two conceptual diagnostic questions measured student understanding of the chemical equilibrium underlying calcium and phosphate homeostasis. One question assessed the ability to predict the change in phosphate concentration when calcium ions were added to a saturated calcium phosphate solution. Fifty-two percent of the students correctly predicted that the phosphate concentration would decrease in accord with the common ion effect. Forty-two percent of the students predicted that the phosphate concentration would not change. Written explanations showed that most students failed to evoke the idea of competing chemical equilibria. A second question assessed the predicted change in calcium concentration after solid calcium phosphate was added to a saturated solution. Only 11% of the students correctly predicted no change in calcium concentration; 86% of the students predicted an increase, and many based their prediction on a mistaken application of Le Chatelier's principle to heterogeneous equilibria. These results indicate that many students possess misconceptions about chemical equilibrium that may hamper understanding of the processes of calcium and phosphate homeostasis. Instructors can help students gain greater understanding of these physiochemical phenomena by adopting strategies that enable students achieve more accurate conceptions of chemical equilibria.
- chemical equilibrium
- conceptual diagnostic test
- skeletal physiology
successful learning of many aspects in physiology depends on a meaningful understanding of underlying chemistry concepts. However, it is well documented that undergraduates possess a variety of profound and persistent misconceptions about chemistry (4, 6, 9, 10). When students bring these misconceptions to the physiology classroom, it can adversely impact their ability to conceptualize the necessary chemical processes and thus can make the learning of physiology more difficult (7). Misconceptions about chemical equilibrium are particularly prominent (1, 2, 4–6, 10, 11, 14, 15) and may hinder the capability of students to correctly apply the general model of “mass action” (8) to a range of physiologically important chemical reactions. However, to the author's knowledge, there has been no empirical documentation of the prevalence of chemistry misconceptions among students of physiology.
Accurate understanding of the concept of equilibrium is particularly important for making sense of the chemistry underlying hormone-receptor interactions, oxygen handling, acid- base balance, electrolyte balance, and metabolism, among other physiological processes. For example, the relation of the skeleton to extracellular calcium and phosphate homeostasis can be represented by the relationship between solid calcium phosphate and calcium and phosphate ions in solution, since “Free ionized calcium in the extracellular fluid exists in chemical equilibrium with calcium present in the hydroxyapatite crystals of bone” (12). Therefore, misconceptions about heterogeneous equilibrium (the equilibrium between substances in more than one phase) may prevent students from correctly predicting the response of bone to disturbances in plasma calcium and phosphate concentrations.
This study was designed to examine the understanding of chemical equilibrium that students bring to their learning of calcium and phosphate homeostasis. The goal of this investigation was to uncover misconceptions that students have about the common ion effect or competing equilibria (where systems share two or more reactions with a common element) and about heterogeneous equilibria. Uncovering these misconception will help instructors to develop strategies to help students to better understand the physiochemical factors that influence extracellular calcium and phosphate balance.
Students were enrolled in the first semester of a two-semester, sophomore-level course in human anatomy and physiology and had all previously taken a full year of inorganic chemistry that used a common textbook (3). Most students were science majors (e.g., Biology, Biochemistry, and Biology Education majors). Responses were collected from students enrolled in two successive offerings of the course.
Assessment of student misconceptions.
Misconceptions were assessed by a conceptual diagnostic test that consisted of two questions (Fig. 1) The questions were created by the author to assess student understanding of the single ion effect and heterogeneous equilibrium as these concepts are applied to calcium and phosphate homeostasis. Each conceptual diagnostic question was structured as a two-tiered query following the recommendations of Treagust (13) and others (9, 14). The questions provided brief descriptions of alterations in some aspect of the chemical equilibrium, and students were asked to predict the direction of the resulting change in the concentration of one of the chemical species in the reaction by selecting a single response from a series of multiple-choice answers. Students were then asked to provide a short written explanation of their predictions.
Misconceptions were assessed in class at the beginning of the section of the course devoted to the physiology of skeletal system, where students learn about the processes of development, growth, and remodeling of bone tissue and the mechanisms of extracellular calcium and phosphate homeostasis. Students were given between 5 and 10 min to formulate their answers, and tests were collected after the last student had completed his or her responses. The assessment tool and protocol used in this project were approved by the Institutional Review Board for the Protection of Human Subjects at Niagara University. Informed consent was obtained from all participants in this study.
Analysis of student responses.
Answers to the multiple-choice questions were tabulated. Written answers were classified according to the explanatory categories shown in Tables 2 and 3 , which emerged from a critical reading of the student responses by the author. The construction of these categories was based on the identification of key phrases and causal mechanisms found in the student written explanations. In the absence of explicit statements of causality, characteristic phrases found in the student responses were used to interpret and categorize the written answers.
As shown in Table 1, in response to conceptual diagnostic test item 1, 52% of the students (34/66 students) correctly predicted that the concentration of phosphate would decrease upon the addition of calcium ions to the solution. A review of the written responses indicated that the majority of these students (21/34 students) explained their prediction by indicating that the reaction equilibrium would shift to a greater formation of calcium phosphate (see “decrease” in Table 2). However, Table 2 also shows that 29% of the students (10/34 students) predicted that the phosphate concentration would decrease because of a dilution effect, i.e., that the addition of calcium would cause a reduction in the concentration of phosphate. Thus, only 32% of the students (21/66 students) correctly explained their prediction of a decreased phosphate concentration on the basis of a change in the chemical equilibrium.
Forty-two percent of the students predicted that the concentration of phosphate would not change upon the addition of calcium ions to the solution (Table 1). The most common written response given for this prediction (12/28 students) emphasized the notion that the amount of phosphate had not changed (see “not change” in Table 2). Another frequent explanation for this prediction was the postulate that the calcium concentration has no influence on the concentration of phosphate (7/28 students). Approximately 6% of the students thought that the addition of calcium would cause an increase in the concentration of phosphate (Table 1). The major explanation implies that the phosphate concentration would increase to balance the chemical equation.
As shown in Table 1, in response to conceptual diagnostic test item 2, 85% of the students incorrectly predicted that the concentration of calcium will increase when solid calcium phosphate is added to the solution. The most common explanation for this prediction suggested that chemical equilibrium must be restored and, therefore, that the concentration of calcium ions must increase (24/57 students; see “increase, explanation 1” in Table 3). Another common category of explanation centered on the general notion that calcium was added to the solution and, therefore, that the concentration of calcium must increase (17/57 students; see “increase, explanation 2” in Table 3). Some of the written explanations in this category implied that the equilibrium was disturbed and, therefore, that more calcium was dissolved in the solution. However, other explanations appeared to indicate that students understood the increase in calcium concentration to reflect the increase in the total amount of calcium in the beaker. Another recognizable set of explanations expressed the notion of dissolution or dissociation without specifying a causal mechanism. Although this category might be considered akin to the suggestion of restoring equilibrium, the idea of equilibrium was not explicitly mentioned in the explanation. The notion of a balanced equation, the consideration of the stoichiometry of calcium phosphate, and the postulation that the solvent level remained constant were other, less frequently evoked explanations for why the concentration of calcium would increase.
Only 11% of the students correctly predicted that the concentration of calcium would not change when solid calcium phosphate was added (Table 1). There were a variety of explanations provided for this prediction (see Table 3). Two students emphasized the importance of the solid phase of the added compound in accounting for absence of a change in concentration. Three students pointed out that the addition of calcium and phosphate was “equal” and, therefore, that the concentration of calcium would not change. Other responses suggest the idea of chemical bonding or kinetics were concepts important to student explanations for their answers.
Student responses to the conceptual diagnostic test indicate that substantial percentages of students bring misconceptions about the chemical equilibrium between calcium, phosphate, and calcium phosphate in solution to their study of the homeostasis of these electrolytes. In particular, a considerable number of students failed to correctly predict and explain why the addition of calcium ions would cause a reduction in phosphate concentrations, and a large majority of students incorrectly indicated and justified why the addition of solid calcium phosphate to the solution would cause an increase in the calcium concentration. Both of these errors are diagnostic of misconceptions about the common ion effect in competing equilibrium and the effect of solids in heterogeneous equilibrium systems, respectively. These are topics typically covered in a first-year chemistry course (3). Moreover, these general categories of misconceptions have been reported to be common among students in advanced high school and undergraduate general chemistry courses (1, 5, 11, 15). Our study confirms these findings, extends the identification of chemistry misconceptions to heterogeneous equilibria in solution, and reports and analyzes individual student explanations for their predictions about changes in the equilibrium state.
Misconceptions about the common ion effect in competing equilibria.
In response to conceptual diagnostic test item 1, 52% of the students examined correctly predicted a decrease in phosphate concentration in response to increased calcium levels. Nevertheless, nearly a third of these students failed to provide a correct explanation for their prediction. Instead, these students indicated that the decrease in the concentration of phosphate occurred by dilution. By itself, this explanation is a plausible reason for a decrease in solute concentration, i.e., the addition of mass to a solution would be expected to cause an increase in its volume. Nevertheless, the explanation shows a naïve view of solution chemistry and indicates that these students were unable to recognize the common ion effect involved in the competing equilibria. Thus, a substantial number of students made a correct prediction about a state variable on the basis of an inadequate understanding of the underlying chemistry of the problem. This finding highlights the shortcoming of relying only on the predicted change in a system without having students explain the reasoning behind their predictions.
Students who incorrectly predicted that the concentration of phosphate would not change in response to the addition of calcium to the solution showed no consideration of chemical equilibrium. Instead, as indicated in their written explanations, some of these students explained their thinking by affirming a correct but irrelevant tenet of solution chemistry: that solute concentration can be increased by the addition of mass. In this view, since there was no addition of phosphate, there must be no change in phosphate concentration. In an alternate but related explanation, some students stated that the calcium concentration in the solution was unrelated to the phosphate concentration, again affirming an ignorance of the chemical equilibrium established between the components of the solution. A small fraction of students incorrectly predicted an increase in phosphate concentration upon the addition of calcium. Where explanations were interpretable, it is apparent that some of these students expressed ideas consistent a balancing notion of chemical equilibrium (5, 6, 10), i.e., an increased concentration of calcium should lead to a “compensatory” increase in phosphate concentration.
These results show that a substantial percentage of physiology students would not be able to correctly predict and interpret the changes in body fluid levels of physiologically relevant solutes where the common ion effect is necessary for determining the concentrations of these solutes under competing equilibrium. For example, with regard to body phosphate homeostasis, our results suggest that a considerable minority of students would predict that the addition of CaCl2 to body fluids would not alter the concentration of phosphate in the blood: a failure to activate prior knowledge of the common ion effect and apply it to the equilibrium condition. Alternately, these students would presumably make the complementary error when asked to predict the effect of hyperphosphatemia on plasma calcium concentrations. Even among students who would make a correct prediction in these circumstances, our results suggest that that a third of them would do so for the wrong reasons. The prevalence of this type of misconception has widespread physiological significance because the common ion effect plays a role in a number of physiochemical reactions that involve competing equilibria, such as the contribution of sodium bicarbonate to the pH of the extracellular fluid. Misconceptions about the common ion effect could thus limit the ability of a substantial number of students to make correct predictions about a variety of physiological processes that depend on a meaningful understanding of the concept of two or more simultaneous reactions with a mutual element.
Misconceptions about heterogeneous equilibrium.
Only 11% of the students correctly answered the conceptual diagnostic question about the addition of solid calcium phosphate to a saturated solution of calcium phosphate. This occurred despite the fact that all of the students had prior instruction in inorganic chemistry that informed them that the equilibrium is not affected by the amount of the pure solid (3). A few students explained their prediction by suggesting that the concentration of calcium will not change because calcium phosphate was added as a solid. However, none of the students clearly evoked the notion of a heterogeneous equilibrium where a solid salt has been added to a saturated solution of ions. At best, one student stated that calcium phosphate would remain a solid after addition to the solution, thereby implicating that no further dissociation of this compound would occur to cause an increase the concentration of calcium ions. Even here, there was no indication that the student was bringing to mind the notion of a dynamic equilibrium between a solid phase and a dissolved phase where the concentration of the solid remains constant as more mass is added.
A number of students evoked some idea of balance, i.e., that equal parts of calcium and phosphate were added as a solid and, therefore, that the concentration of calcium would not increase. This conclusion does not actually explain why the calcium concentration should not increase, but it is consistent with the widespread prevalence of student reasoning on equilibrium based on the notion of balance (5, 6, 10). Other explanations suggest that students were making sense of their prediction based on stoichiometry, kinetics, and bonding. Thus, a majority of the students who correctly predicted that the calcium concentration would not change after the addition of solid calcium phosphate made use of reasoning about chemistry that was either spurious or irrelevant.
The most prevalent misconception about the addition of solid calcium phosphate was the prediction that the concentration of calcium would increase. Many students explained this prediction by alluding to Le Chatelier's principle and the idea of a restoration of equilibrium. As stated in their chemistry textbook, Le Chatelier's principle describes the response of a system to a change where “the system shifts in equilibrium composition in a way that tends to counteract this change” (3). Thus, when students evoke the notion of “drive” or “want” or “tendency” or “shift” of the chemical reaction to restore equilibrium or to offset a change, they are expressing an understanding of Le Chatelier's principle whether they referred to this principle by name or not. The extent of these responses illustrates the widespread tendency for students to naïvely evoke Le Chatelier's principle in circumstances where conditions limit or contradict its application (1, 5, 11).
A less frequent category of explanation focused on the addition of calcium to the beaker without specifying a causal mechanism whereby the calcium concentration should increase. In some instances, students may have confused “concentration” with “amount” and thus must have thought that the question asked for a prediction about the change in total amount of calcium in the beaker. Fewer students alluded to a dissociation of calcium phosphate that was consistent with Le Chatelier's principle but without explicitly mentioning the concept of a shift in equilibrium. Thus, most of the instances where students predicted an increase in calcium concentration with the addition of calcium phosphate demonstrated that students held different misconceptions about the application of Le Chatelier's principle to heterogeneous equilibrium. This finding indicates that misconceptions associated with the concept of heterogeneous equilibrium are widespread among physiology students, and, as noted by Nakhleh (10), are resistant to prior instruction in inorganic chemistry.
These results suggest that many students would mistakenly apply Le Chatelier's principle to physiological circumstances where it is not appropriate, e.g., the contribution of the skeletal system to calcium and phosphate homeostasis. For example, using the reasoning given for their explanations of conceptual diagnostic test item 2, a large majority would predict that increased bone mass should cause an increase in the concentration of calcium ions in body fluids. Corresponding misconceptions could be inferred when students are asked to predict the effect of decreased bone mass. Furthermore, by the same reasoning, students would be expected to have difficulty making sense of the chemistry underlying the pathophysiological formation of other solid phase components, such as kidney stones or gall stones, that are at equilibrium with dissolved solutes in the body fluids.
Although there are only a few examples of solid phase equilibria in the body, misconceptions associated with heterogeneous equilibrium could also appear when students confront other sorts of related physiological phenomena. One category of analogous situations would involve the constant density of a fixed chemical species such as a receptor or an enzyme. For example, if the number of ligand-bound receptors in the body increases with an increased body mass and volume so that there is no corresponding increase in receptor density, then one should predict there would be no corresponding change in the concentration of the free ligand. However, when viewed from the perspective of the most common misconception about heterogeneous equilibrium, students might be expected to misuse the Le Chatelier's principle and conclude that the ligand concentration should increase with increased amounts of bound receptor. As each of these examples illustrate, student misconceptions about heterogeneous equilibria could have multiple ramifications concerning the meaningful understanding of physiochemical phenomena.
Suggestions for improvement.
Recognizing that the physiology teacher does not have the time to help students remediate all of the misconceptions about chemistry and chemical equilibrium that may appear in the classroom (4, 6, 9, 10), how can we best assist students to overcome the ones identified here and thus avoid their misapplication to physiological phenomena? As suggested by chemistry educators, the following approaches appear to have value in helping students come to a more accurate understanding of chemical equilibrium and, in particular, an understanding the role of the single ion effect in competing equilibria and the contribution of solids to heterogeneous equilibria in solution:
Provide students with a greater opportunity to review and relearn relevant principles of chemical equilibrium through the analysis of selected examples, particularly if they involve unusual or potentially discrepant events such as heterogeneous equilibrium (6, 9, 15). An evaluation of these examples in light of the appropriateness of Le Chatelier's principle could be done with the goal of properly applying the concepts learned to the interpretation of physiologically important reactions.
During a review of the chemical principles involved in the single ion effect and heterogeneous equilibrium, the instructor should help students clarify their microscopic view of the physiochemical system involved (5). This view should include 1) the particulate model for chemical change and 2) the kinetic behavior of interacting molecules to reinforce the dynamic nature of chemical equilibria (6, 10).
Instructors should model effective strategies for resolving crucial problems in physiologically relevant chemical equilibria that are associated with processes such as calcium and phosphate homeostasis. By their approach to analyzing problems that involve the single ion effect or heterogeneous equilibrium, instructors can illustrate the potentially discordant outcomes of taking inappropriate approaches to problem solving and help students realize that overemphasis on rote rules and algorithmic procedures can result in the misapplication of qualitative thinking tools such as Le Chatelier's principle (11).
Point out the limitations of Le Chatelier's principle with regard to heterogeneous equilibrium and/or replace it with more accurate representations of chemical equilibrium such as the laws of van't Hoff and the Equilibrium law (3, 6, 11, 15). While this later suggestion is widely advocated among chemistry educators, it is likely to involve more extensive instruction in chemical principles than most physiology instructors are willing to allocate time. Nevertheless, instructors need to help students avoid the tendency to view Le Chatelier's principle as the sole or universal approach to predicting changes in equilibria.
It is unavoidable that students will bring faulty understandings of chemistry to the physiology classroom that can hinder their ability to correctly learn various physiological mechanisms. It is the responsibility of the physiology teacher to be aware of these misconceptions and to use approaches to teaching and learning about chemistry that will help students avoid the potentially adverse outcomes. By doing so, the teacher will enable students to gain more accurate and useful mental models of the physical-chemical concepts that can lead to meaningful understanding of the corresponding physiological processes.
- Copyright © 2009 the American Physiological Society