A traditional teaching pattern for chemical equilibria suggests that pre-16 students are introduced to a “two-way” reaction treated qualitatively
While more complex ideas such as calculation of equilibrium constants and the meaning for these feature in post-16 courses. Le Chatelier’s Principle (LCP) is introduced commonly at this stage to help students predict the direction of change in equilibrium position. The ideas associated with chemical equilibria are commonly regarded as among the most difficult to teach and learn in pre-university chemistry courses, so perhaps unsurprisingly the topic has received extensive attention from researchers keen to explore the development of students’ thinking about the key concepts involved. The key points are reviewed here.
Issues in learning about chemical equilibria
1. A “dynamic” equilibrium
The most basic principle students need to understand is that an equilibrium position implies molecules exchanging between two “sides” at the same rate. The “sides” may be two phases, for example, the distribution of iodine molecules between water and hexane, or two reactions, such as occurs in the formation of ammonia. The dynamic nature cannot be seen, but is implicit in the chemical events.
Maskill and Cachapuz (1989) used a word association test (WAT) to investigate students’ intuitive responses to the statement “the reactions were at equilibrium”. About 76% of 14- 15 year olds who had not received teaching about equilibrium strongly associated this with “static” and “balance”. Little change was observed post-teaching, as this student’s response illustrates:
“…the reaction is finished, it is stable, it will not react anymore unless you add something…” (p 67)
Gorodetsky and Gussarsky (1986, 1990) found similar reasoning among students aged 17- 18. Their earlier study used WATs in conjunction with a teacher-administered test. They found that only the highest achievers on this test broke the link between “dynamic” and “static” to make the associations between “dynamic”, “chemical equilibrium” and “reversibility” instead. The authors’ later work explored the impact of a teaching sequence on students’ thinking, comparing a control group who received no tuition with two groups receiving teaching to different depths. Their results indicated that the teaching resulted in links between “equilibrium” and “chemical equilibrium”, but also a slight increase in the association of “static” and “state of balance” to both these terms. These data suggest that the notion of a reaction in which continued unobservable change is occurring is counter to intuition, so many students find this difficult.
2. An equilibrium reaction involves two separate reactions
Experienced chemists consider the forward and reverse reactions part of the same chemical system. Students view the two reactions as separate and independent events. Early evidence for this came from Johnstone et al (1977), who report that 80% of 255 16-17 year old students have this view. These researchers suggest that the double-headed arrow used in equilibrium reactions contributes to the “two-sidedness” students perceive. One arrow, used in a reaction which goes to or near completion, emphasises one reaction, so two arrows implies two separate reactions. Additional evidence for this reasoning comes from several other workers. Gorodetsky and Gussarsky (1986) found this in one-third of 17-18 year old chemists. Cachapuz and Maskill (1989) used word association tests (WATs) with 14-15 year olds to reveal the same thinking. Banks (1997) tracked the developing understanding of a small group of post-16 chemists through a post-16 chemistry course and found further evidence.
3. Problems with Le Chatelier’s Principle
In 1888 Henri Le Chatelier devised a summary statement which could help chemists make qualitative predictions about changes in equilibrium position:
“If a system is at equilibrium, and a change is made in any of the conditions, then the system responds to counteract the change as much as possible” (Burton, et al, 1994, p 137)
Several workers have probed students’ ability to apply LCP to situations in which additional reagents are added to a closed system. Hackling and Garnett (1985) found that although about 40% of 17 year olds could apply the reasoning expected, a common misconception was to treat all substances in the reaction independently, rather than viewing the interactions between them. Bergquist and Heikkinen (1990) report some 19 year old chemists using an “oscillating” model, suggesting that when one change has occurred, another must follow immediately because the first position has altered. They report, with no precise percentages, that a common idea was the notion of the equilibrium being re-established only when all additional reagent was used up. These ideas reflect students applying a “two reactions” model for chemical equilibrium - in this latter case, if a reagent was added, then the forward reaction would continue to “use up” the extra material, while the reverse reaction remained unchanged.
The limitations of LCP also present problems. Wheeler and Kass (1978) noted that 95% of their ninety-nine 17-18 year old chemists misused LCP, not realising that it cannot be applied in all situations. Quilez-Pardo and Solaz-Portoles (1995) studied the responses of sixty-five teachers and 170 students to five situations in which LCP did not apply. Between 70-90% of students and around 70% of teachers used LCP in answering these questions, resulting frequently in incorrect predictions.
4. Calculating and using equilibrium constants
The value of K indicates the extent of a reaction and is calculated by applying the Equilibrium Law. The higher the value of K, the more complete the reaction. K is constant for a specific reaction at a defined temperature. A number of studies reveal students’ difficulties with these ideas. One difficulty reported by Hackling and Garnett (1985) is that about 50% of 17 year olds think that there is a simple arithmetic relationship between the concentrations of reactions and products at equilibrium, most commonly, that these are equal. The authors suggest that:-
“This misconception can probably be attributed to the considerable emphasis placed on reaction stoichiometry in introductory chemistry topics.” (p 211)
Students will be aware that chemical equations must be “balanced” and transfer this idea when they consider an equilibrium position. A second, given by about 20% in the same study, is that K increases when equilibrium is re-established after changing concentration of a reactant. Students argued that this would result in more product and hence a higher value.
Thirdly, Hackling and Garnett and Gorodetsky and Gussarsky (1986) found that many students did not appreciate the effect of temperature on K, demonstrating an inability to judge when K is constant, or when and how K changes. The proportions expressing these ideas decreased post-teaching. In a small-scale study using a context-led chemistry course, Banks (1997) revealed little change in students’ thinking about K, with many remaining uncertain about when K changed.
5 Confusing rate and chemical equilibrium
At equilibrium, the rates of the forward and reverse reactions are equal, resulting in the dynamic “no overall change” position. Although this appears quite straight-forward, the literature reveals several ways in which students confuse rate of reaction with chemical equilibrium ideas.
Hackling and Garnett’s (1985) post-teaching study with thirty 17-year old chemists revealed that about 25% thought the rate of the forward reaction would increase from the time reactants were mixed until equilibrium was established. This may reflect the perception of the forward and reverse reactions being separate events.
Cachapuz and Maskill (1989) and Hackling and Garnett (1985) find some students who consider concentrations of reactants and products are equal at equilibrium. These students may be directly confusing equality of rate and concentration.
Thirdly, Hackling and Garnett report that about 50% of students think that changing conditions results in an increase in the rate of the favoured reaction and a decrease in the rate of the other reaction. Banerjee (1991) found similar reasoning among 35% of undergraduate chemists and 49% of chemistry teachers. Some students (27%), extended this to the role of catalysts, suggesting that the rates of forward and reverse reactions would be affected differently, a finding corroborated by Gorodetsky and Gussarsky (1986). Finally, Banerjee reports (without figures) that both undergraduate chemists and high school teachers tend to associate a high K value with a very fast reaction.
Summary of key difficulties
1. Equilibria are static, not dynamic
Students mainly experience chemical reactions that appear to go to completion. When they meet a reaction which does not go to completion, but which has a reverse reaction occurring to a significant extent, it is unsurprising they think of the equilibrium position as being fixed. That is, once achieved, there is no movement of particles between the two “sides”. This is a version of the reaction being complete.
2. An equilibrium reaction comprises two separate reactions
Perhaps the next step for students in learning about equilibria is to recognise that two reactions are occurring, but to think of these as separate from each other. Research suggests that using the double-headed arrow may contribute to this.
3. Le Chatelier’s Principle is used as if it applies in every case
Le Chatelier’s Principle is taught widely in post-16 chemistry courses as it has gained a reputation as a useful tool for predicting changes to an equilibrium position under some circumstances. However, students who are taught this as the only strategy for considering how an equilibrium position is adjusted will not learn that this Principle does not always apply.
4. Rate and equilibria can be confused
There is some evidence suggesting that students perceive the rates of one reaction in an equilibrium system may alter, while another slows or remains constant. They have not grasped the notion that rate applies to the system as a whole, rather than the component reactions. This difficulty is related to students’ perception of two separate reactions.
1. Present a wider range of reactions to 11-16 year olds
Students need to experience a wider range of reactions than is demonstrated currently. We should not be afraid of showing situations which do not meet an expected “norm”, but rather use these as ways to challenge thinking and promote a wider perspective on chemical events than the one-way reactions in popular usage permit. Demonstrations of “unusual” reactions could form part of a teaching sequence designed to challenge students’ perceptions of chemical change, encouraging them to accept equilibria in a qualitative way.
2. Teach using equilibrium laws and the laws of van’t Hoff
The widespread application of LCP in post-16 chemistry deserves to be challenged and replaced with a much clearer, more accurate and essentially more honest approach to considering equilibrium problems. Banerjee (1991) is quite right to advocate teaching the laws of van’t Hoff, which are based on thermodynamics, use of the Equilibrium Law may be added. LCP is unnecessary and unhelpful. However, consider Treagust and Graeber (1999)’s study comparing two different approaches to teaching equilibrium. The effects on students’ learning of the Australian approach, featuring LCP and rates of reaction taught using analogies and the German, using the equilibrium law and analogies in only the final lesson of a teaching sequence were compared. The results showed no significant differences.
Students’ difficulties with the equilibrium constant deserve attention. There is a need to establish the mathematical relationship between the value of K and the concentrations of the reactants and products. Students need to experiment with figures to see for themselves that changing concentrations does not result in change to K. Once this is firmly established, students then need to work out why temperature affects K, but changing concentrations does not. Teachers need to introduce and explain the effects of changing temperature in relation to the enthalpy change for the reaction, but to prevent students from adding in rates ideas. Awareness that this might occur should help teachers be cautious and careful in the language used.
3. Use diagnostic tests to determine students’ understandings
Voska and Heikkenen (2000) have devised a “Test to Identify Students’ Conceptualisations” (TISC) about aspects of chemical equilibria, specifically, the application of LCP, constancy of the equilibrium constant and the effect of a catalyst. The test adopts a two-tier multiple choice approach. The authors suggest that although open-ended questions may asses students’ reasoning more accurately, the multiple-choice test does allow teachers to identify a range of misconceptions requiring remedy (p 171). Despite their limitations, diagnostic tests are likely to be useful in determining students’ starting points, their progress and change in thinking post-teaching.
For a full list of references used by Vanessa Kind in Beyond Appearances please click here
These resources have been taken from the book, Beyond appearances: students’ misconceptions about basic chemical ideas by Vanessa Kind.
Beyond Appearances: Students misconceptions about basic chemical ideas
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Students’ ideas about chemical equilibria