Teaching Chemical Equilibrium and Thermodynamics in

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Teaching Chemical Equilibrium and Thermodynamics in Undergraduate General Chemistry Classes -

Anil C ~anerjee' Regional College of Education, Mysore 570006, India

Teaching of chemical equilibrium and thermodynamics in high school and undereraduate chemistry classes alwayslhas been challengingVdueto the enormo& number of conceptual difficulties faced by students. Banejee and Power (1)reported the development and tryout of modules for teaching chemical equilibrium. Misconceptions of high school ( 2 5 ) and undergraduate students and school teachers (6) have been investigated. In a series of eight articles, Gordus (7) revisited chemical equilihrium principles. While some researchers (41,and most textbooks advocate a 'rate approach' to teaching equilibrium, others (6, 8-11) suggest a thermodynamic approach. Libby (12) suggested an equilibrium-kinetic approach for teaching introductory organic chemistry. In this paper, the conceptual difficulties of undergraduate students in some aspects of chemical equilibrium and thermodynamics are discussed along with teaching strategies for dealing with these difficulties. Methodology Sixty students enrolled in the third-semester general chemistry course of an eight-semester BScEd program a t the Regional College of Education a t Mysore were the subjects of this study. All these students had studied chemistry compulsorily for two years in senior secondary (Year 11 and 12) school level, and also in the first year of the BScEd course. The overall achievement of these students in chemistry ranged from good to very high. A20-item paper-pencil achievement test on various aspects of chemical equilibrium and thermodynamics was developed by the author and content validated by a group of school and college teachers. The third-semester general chemistry course of 16-weeksduration is on chemical equilibrium and thermodynamics. Two textbooks (8, 9)were used. The author taught the qualitative and quantitative aspects of gaseous, ionic, solubility, and aeid-base equilibria and, also the concepts, machinery, and applications of the first and second laws of thermodynamics, for 12 weeks through 36 lectures. Students also were exposed to the empirical-, kinetic-, and thermodynamic approaches for developing equilihrium concepts. However, the initial 12 weeks of teaching were not segmented, nor was a comparison of the effectiveness of the three approaches attempted. However, the analysis of conceptual difficulties of students arising out of the treatment of equilibrium through these three approaches was a maior aim of the study A test was eiven after 12 weeks to assess the conceptual understanding and problem-solvinr al~iliticsof studcn~s.The achievement test had items thacwould reflect the three approaches. However, the test items were not specifically and exclusively designed to test the effects of each type of approach. In this paper the word "approach" refers to the teaching approach that would produce the best results on a question. Analysis of the written responses and student interviews revealed widespread misconceptions and conceptual diffi-

'Present address: Regional Institute of Education, Ajmer 305004, India.

culties. Another four weeks ofthe course were spent on discussing and clarifying the conceptual difficulties, followed by a semester-end-test. The nature and level of this test was similar to the first test. Results and Discussion The concepts of equilibrium, equilibrium concentrations, and equilibrium constants were introduced. Data on equilibrium concentrations were provided for some reactions, and students solved problems on equilibrium expressions and equilibrium constants. After this, the students performed laboratory experiments to determine the equilibrium constants of some reactions, including that of hydrolysis of ethyl acetate. The conceptual problems of the students in these areas were revealed through their responses to the following two test items.

Item 1.If 1.0 x lo-' mol each of hydrogen and iodine gases are placed in a l-L flask at 448 "C with 2.0 x lo3 ma1 of hydrogen iodide, will more hydrogen iodide be formed (Ke4= 50.53)? Many students felt that the concentrations mentioned in the test item were equilibrium concentrations. Their conflict started when they found that their calculated equilibrium constant values were not matching with the given value. Students have the misconce~tionthat whenever equilibrium constant is given, all Eoncentrations mean equilibrium concentrations. Durinc discussion, the concept of reaction quotient or concentration ratio was brought in. Students found that the reaction quotient was less than the equilibrium constant. They now argued that the reaction had not reached equilibrium yet, and hence, more hydrogen iodide would be produced. Item 2. The equilibrium constant of the reaction:

Will there be any Agf(aq)ion left after the reaction? The common students' response was: 'No Ag+ (aq) ions, since equilibrium constant is very high'. It is interesting to note that students use the term equilibrium constant, but still feel that there will be no equilibrium when equilibrium constant is high. Misconceptions of students were widespread on rate and equilibrium. Some typical misconceptions were revealed through the following test item: Item 3. The reaction: CO(g)+ C12(g)= COC12(g)+heat is in equilibrium at 200% and 1atm. The temperature of the reaction is lowered to 150 "C, keeping the volume constant. When the system comes to another equilibrium: (a) the rate at which COC12(g)is formed will be (b) the equilibrium constant will be

Give responses as: (A)greater than-; (B)less than-; (C)same as first equilibrium; (Dl data insufficient for conclusion. Volume 72 Number 10 October 1995

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About 50% of students had the misconception that the equilibrium constant would be less and that the rate would be greater. Many students feel that, (i) when the temperature is decreased in a n exothermic reaction, the rate of the forward reaction increases, (ii) large values of equilibrium constant implies a very fast reaction (2,3, 6). Wheeler and Kass (13) found t h a t students failed to distinguish between rate and extent of a reaction. These miscouceptions largely originate d u e t o wrong o r overuse of t h e LeChAtelier ~ r i n c i ~to l e~ r e d i c rate t and extent of a reaction. I t shouid be iemekbered that the 'rate approach' is theoretically sound only for elementary reactions and thus is truly applicable for a limited number of reactions. I t also must be remembered t h a t equilibrium law depends on thermodynamics and not oi kinetics (8-11). Practical courses in senior secondary and undergraduate classes hardly have experiments measuring r a t e a t or near equilibrium. I n fact, most rate measurements presume the reverse reaction to be nvgligiblc. Despite evidence t o the contrary, most senior secondary and unrlrrpi~duntechen~istry text6ooks overemphasize =ate approach to equilibrium, and the LeChAtelier principle. LeChAtelier principle is vague and ambiguous, and the principle should be replaced by the v'ant Hoff laws on equilibrium (14,15). Thermodynamic approach seems to be the most logical way of discussing equilibrium because equilibrium laws depend on thermodynamics (1,8-11). The concepts of internal energy, enthalpy, entropy, Gibbs free energy, and thermodynamic equations interrelating these state functions were developed and a few problems solved in the class. The concepts of spontaneous and reversible reactions were broueht in. and the role of Gibbs free enerw and entropy in deciding direction of spontaneous reaction and state of equilibrium were discussed. However, the analysis of responses of students in the test indicated concentual difficulties in manv areas: e.%. .. . (a) on ecluilibrium in reactions which had vinually 'gone to complction'or 'do not occur': rb role of cutalvst on cqu~librium;~ C interprt!. I tation of the master equation AG A H - T&; (dl ~ i b b s free energy, spontaneity, reversibility, and equilibrium. According to Crosby (If?), many students believe that even though equilibrium reactions are reversible they still go to completion, other students think that the forward reaction goes to completion before the reverse reaction commences. Some of these misconceptions are revealed through discussions on test items 4-7.

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Item 4. For a reaction that has "gone to completion" or "does not occur",will there be any equilibridequilibrium constant? Explain.

Item 5. The equilibrium constant of the reaction N2(g)+ O&) = 2NOW at 25 "C is very low. (a) Will the reverse reaction be spontaneous? (b)What might be the reasons for formation of automotive smog? Suggest a way out to reduce this air pollution.

Students' common response to part (a) of this question was Yes, the reverse reaction will be spontaneous since the reverse reaction has high equilibrium constant.

But thev were not able to come out with suitable armments forpart (b) of the item, though they knew that automotive smog was due to nitric oxide and nitrogen dioxide. In the discussion class, we argued that though the reverse reaction was spontaneous, rate was very slow; hence, most of the nitric oxide formed went to the atmosphere. The role of a catalyst to speed UD the reverse reaction, and the development of a suitabie catalyst to reduce automotive smog formation were then discussed. Item 6. Calculate the Gibbs free energy change for the reaction at 27 "C given that enthalpy and entropy changes were -2808KJmol-' and +182.4JK1: C6H1206(g)+ 602(g)= 6C02(g)+ 6H20 How will you interpret the AG in terms of AH and AS?

Students could solve this simple problem using the master equation of thermodynamics: AG=AH-TAS Their interpretation was: High negative value of AH and positive value of TAS, make the right-hand side ofthe equation negative; hence, AG is negative and the reaction is spontaneous. Though the students were using the logic correctly, their interpretation was false. Most school and undergraduate textbooks also give this false interpretation (though it is a good rule of thumb). Atkins (7)specifically mentions this and presents a correct interpretation. The tendency to lower Gibbs free energy is solely a tendency toward greater overall entropy. Systems change spontaneously solely because that increases the entropy of the universe, not be-

Typical responses for this item was 'No'. During interviews, students argued that for reactions that had 'gone to completion' or 'do not occur', t h e r e w a s no equilibrium/equilibrium constant. During discussion sessions, the following concepts were elaborated (10). 1. The fundamental tendency in chemical systems is the tendency to attain equilibn'um. Hence, for all chemical reac-

tions there must be equilibrium/equilibrium constants. 2. This fundamental tendencv " aoolies .. also to reactions that 'go to campletian'or 'do not occur'. In these eases,the equilibrium constant has very large ar very small values. 3. In practice, kinetic barriers (activation energy) may frequently prevent attainment of equilibrium. l b clarify these statements, some examples of chemical reactions were cited; e.g., the reaction of carbon (graphite) with oxygen a t 1000 "C virtually 'goes to completion' while the same reaction virtually 'does not occur'at mom temperature due to kinetic barriers. In both cases, the reaction CcgraPhite) + Oz(g)= COz(g),has a n equilibrium constant of 10". 880

Journal of Chemical Education

(1)

A

extent of reaction

Gibbs free energy versus extent of reaction for A = B. s t u d e n t s ' graph - - -actual graph

cause they tend to lower energy. AG is a measure of the change in the entropy of the universe caused by the reaction. The equation AG = AH - TAS gives the impression that systems favor lower energy, but this is misleading. AS is the entropy of the system and, -AHIT is the entropy change of the surroundings. Total entropy tends toward maximum for spontaneous reactions.

Item 9. Predict the effectof increasing temperatureon the rate of the forward exothermic reaction A+B, in the reaction A = B at equilibrium? The most common response was 'the rate of the forward reaction will decrease'. As discussed earlier, students still have misconceptions relating to rate and extent of a reaction.

Item 7. Draw a graph of Gibbs free energy versus extent of reaction of the reaction A = B. Discuss and interpret the graph.

Item 10. Assign the thermodynamic reasons for the elevation of boiling point of a dilute solution.

Atypical student's graph and the actual graph are shown in the figure. During interviews, students revealed that Gibbs free energy would increase or decrease linearly to make the reaction spontaneous either in the direction A+B or B+A, depending on whether A (reactant) or B (product) had more Gibbs free energy to start with. The concept of equilibrium as the lowest free energy position of the system and a tendency toward maximizing the entropy of the universe had not been registered in the students' conceptual framework of mind. In the discussion class, I brought in the following points for development (8-10): 1. A reaction that is at equilibrium is a reversible reaction. 2. Chemical reactions spontaneously approach the equilihrim state from both the directions. Hence, the direction of spontaneous change is both A+B and B+A, because the equilibrium state has always a lower free energy than that of either the reactant(s)or praduet(s). 3. All reactions tend to attain equilibrium by maximizing the entropy of the universe. 4. The correct definition of equilibrium is: It is a state of a chemical system having lowest Gihbs free energy and highest total entropy, under given conditions of temperature and pressure.

Analvsis of resnonses in the semester-end-test was done p r i m a l y to assess the effectiveness of the four-week discussion session on clarification of student misconceptions. Responses of students to some of the test items of the semester-end-test are now discussed. Item 8. The eouilibrium constant (K.1 for the reaction N,OAe) " =2~0 ~ r e l i cthe t direction in which ~. ..d"e i i s0.36 at 373 rhr rrnrtmn will m o w ro rrarh rquilhnum ifwr srnn wlth 0 20 mnl of Y,0, g nnd 0.20 mol of i\'O?lg, In a 4.0-Lcontainer 7

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Most of the students could predict correctly that the reaction would proceed in the forward direction because the reaction quotient was less than the equilibrium constant.

Many students could explain this in terms of increase in entropy of the solution as the contributing factor for decrease in the chemical potential of the solution. This indicates that the students are able to interpret the relation AG = AH - TAS correctly. Item 11. Draw and explain the Gibbs free energy cunre for a reaction that has almost gone to completion. Many students drew the graph showing a free energy minimum almost near the nroduct side and also save the correct explanation a s discussed earlier. Analysis of responses in the semester-end-test indicated that students had less conceptual diff~cnltieson many concents of eauilibrium and thermodvnamics. This indicates t h k the four-week discussion sessLon was useful. However, manv students still had misconceotions relatins to rate and kquilibrium. Conceptual dific;lties diagnosG in this study should not be treated as vew specific to the nresent sample of the study. Such miscon~ep~ions are widespread among students and even teachers. According to constructivism, removal of misconceptions is a stupendous task that often needs alternative conceptual framework.

Literature Cited

4. ~ a c * l i n i , M .W:Garnett, P. J.Aust Sci. lbach. J . 1986.31.8-13, 5. Burrmiat. W.:HeiI&nen.H. J . Chem. Edue. 1990.67 . . 100&1003 6. ~ a n & e , A , d.Int