Suggestion of a Viewpoint Change for the Classification Criteria of

Mar 22, 2017 - The four representative models that define oxidation–reduction reactions are often used differently in different situations or contex...
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Suggestion of a Viewpoint Change for the Classification Criteria of Redox Reactions Seoung-Hey Paik, Sungki Kim, and Kihyang Kim* Department of Chemistry Education, Korea National University of Education, Chungju City 28173, Republic of Korea ABSTRACT: The four representative models that define oxidation− reduction reactions are often used differently in different situations or contexts in chemistry textbooks. Although integrated models have been suggested to overcome the confusion caused by this, they have not been successful. We therefore aim to interpret the causes for difficulties in the attempted integration of the various epistemological perspectives. The four key models, namely, the oxygen, hydrogen, electron, and oxidation state (or number) models, focus on the movement of matter. We, however, suggest the use of the Goodstein method, where the relative changes in the electronegativity orders in redox reactions should be employed rather than the properties of the reactants or products themselves. This requires a change of viewpoint from matter to process. We also suggest that the confusion of classifying redox reactions may occur from the viewpoint of matter. Therefore, new interpretations of redox reactions based on the relationship between reactants and products are suggested, considering the process viewpoint as a valuable criterion. While many researchers have tried to solve this problem by suggesting new models to classify redox reactions, we instead suggest a novel interpretation of the various redox models. KEYWORDS: High School/Introductory Chemistry, History/Philosophy, Misconceptions/Discrepant Events, Textbooks/Reference Books, Oxidation/Reduction

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INTRODUCTION Many students, including high school students, undergraduate students, and graduate students, and even teachers, have difficulties in understanding redox reactions.1−7 Such difficulties typically arise because of the presentation of several different models for oxidation−reduction reactions in textbooks. In general, the common models used to describe redox reactions are the oxygen model, the hydrogen model, the electron model, and the oxidation state model.8 Indeed, these models are used in different ways depending on the situation, and the methods for applying the various models also differ in each textbook.9 In inorganic chemistry, authors often exclusively use the electron and oxidation number models, while the oxygen and hydrogen models are employed in organic chemistry and the hydrogen model and alternative representations are mainly used in biochemistry. However, the authors of such textbooks do not comment on the reasons for using different models in each case. Although they use two or more models simultaneously, they do not explain their reasoning behind the selection of each model. This can result in many students facing difficulties in understanding redox reactions.10,11 In South Korea, redox reactions are commonly taught in both high schools and colleges. As in other countries, various models are used to explain redox reactions on a case-by-case basis, while only partially showing the relationships among them.12 In Korea, most high school students understand redox reactions on the basis of the electron model, while most © 2017 American Chemical Society and Division of Chemical Education, Inc.

teachers and preservice teachers understand redox reactions according to the oxidation state model. This discord in understanding between the students and teachers may also contribute to the difficulties in teaching and learning. However, since no single model can explain all oxidation−reduction reactions, teaching methods tend to simultaneously include a range of different models. Although one study has insisted that presenting various models helps develop students’ learning,13 many other studies argue that the use of different models can hinder students’ understanding of redox reactions. In particular, one study found that the models represented in textbooks or by teachers are hybrid models comprising elements of different historical models, which are often treated as if they constituted a coherent whole. As such, these models are not based on the historical development of the concept.14 Therefore, we should either dispose of the methods that simultaneously use two or more models or at least honestly recognize that none of the models can successfully explain redox reactions.15 Received: August 5, 2016 Revised: February 15, 2017 Published: March 22, 2017 563

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CURRENT MODELS FOR TEACHING REDOX REACTIONS In this section, we briefly describe the representative models for explaining redox reactions, while the limitations of each model will be discussed in the following section. First, the oxygen model describes oxidation as the gain of an oxygen atom in an organic substrate, while the hydrogen model describes oxidation as the loss of a hydrogen atom from an organic substrate. In contrast, the electron model describes oxidation as the complete net removal of one or more electrons from a molecular entity, and the oxidation state model describes oxidation as an increase in the oxidation number of any atom in any substrate. The historical development of the redox reaction models can be represented as shown in Figure 1.

shown in reaction 2, where the carbonyl carbon (CO) in formaldehyde loses an oxygen atom but is not reduced: CH 2O + NH3 → CH 2NH + H 2O

(2)

This reaction is also an acid−base reaction, in which formaldehyde reacts as an acid and ammonia reacts as a base. The key point of these examples is that the gain or loss of hydrogen or oxygen atoms does not always indicate that a reaction is a redox reaction.17

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LIMITATIONS OF THE ELECTRON MODEL The most popular teaching model presented in textbooks and by teachers is the electron model; however, many studies have indicated problems associated with this model.1,15,18 This is partly the case because the electron model has limited validity in reactions involving only elements and simple ions or ionic compounds.18 In addition, many researchers have pointed out that electrons are not always exchanged literally in redox reactions, thus resulting in further confusion.1,15,18 For example, when hydrochloric acid dissociates by reacting with water (see reaction 3), a shared pair of electrons is moved completely to the chlorine atom; however, this reaction is an acid−base reaction and not a redox reaction. HCl(aq) + H 2O(l) → H3O+(aq) + Cl−(aq)

(3)

A number of academics have therefore proposed that instructors stop explaining redox reactions using the electron model, which is a common suggestion in textbooks.1,18−20

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Figure 1. Historical development of the different redox reaction models.

The majority of researchers currently agree upon the use of the oxidation state model.16,17 For example, Vitz (ref 16, p 400) insisted that “Oxidation and reduction should be defined in terms of changes in oxidation state. This allows us to discuss the overall processes without a knowledge of their mechanism or details of electronic structure.” Vitz16 also states that the oxidation state model would promote a more unified view of organic and inorganic chemistry and that it would allow a more realistic concept of the electronic structure. However, the various methods available for calculating oxidation states are also not unified. According to Jurowski et al.,21 the definitions and rules proposed by Pauling, the IUPAC, Jørgensen, Calzaferri and Gupta, Atkins, etc., differ from one another. In addition, Holder et al.22 stated that “No two textbooks have exactly the same set of rules.” The most common and accepted methods for calculating the oxidation numbers of species are the rules proposed by IUPAC and Pauling, and these are presented in the majority of general chemistry textbooks. However, each of these sets of rules has its own limitations, as will now be discussed.5,7,20,23−25

To date, various studies have questioned whether too many models exist to describe redox reactions,16,17 although a more serious problem is perhaps that students must be aware that the different redox reaction models can be appropriate in different contexts. However, students who are just beginning to learn oxidation−reduction reactions will find it difficult to understand what is considered “appropriate in different contexts”. In fact, even experts encounter difficulties in these situations. Thus, to address these difficulties, many chemists have tried to suggest various alternative methods; however, no agreement has yet been reached.

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LIMITATIONS OF THE HYDROGEN AND OXYGEN MODELS In organic chemistry, the hydrogen and oxygen models are commonly used to describe redox reactions. However, these models have limitations, as many redox reactions do not involve hydrogen and oxygen. In addition, systems such as acid−base reactions involve oxygen or hydrogen transfer but are not considered redox reactions. One such example is shown below in reaction 1, where nitrogen loses hydrogen but is not oxidized: NH3 + B(OH)3 → NB + 3H 2O

LIMITATIONS OF THE OXIDATION STATE MODELS

Limitations of the IUPAC Rule

(1)

According to IUPAC, oxidation states can be calculated as follows (ref 26, p 1097):

Instead, this is an acid−base reaction, in which ammonia reacts as an acid and hydroxide reacts as a base. A further example is 564

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Figure 2. Steps for determining the oxidation number based on Pauling’s rule.

Oxidation state is a general term defined to mean the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules: (1) the oxidation state of a free element (uncombined element) is zero; (2) for a simple (monatomic) ion, the oxidation state is equal to the net charge on the ion; (3) hydrogen has an oxidation state of 1 and oxygen has an oxidation state of −2 when they are present in most compounds (exceptions to this are that hydrogen has an oxidation state of −1 in hydrides of active metals, e.g. LiH, and oxygen has an oxidation state of −1 in peroxides, e.g. H2O2); (4) the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. For example, the oxidation states of sulfur in H2S, S8 (elemental sulfur), SO2, SO3, and H2SO4 are, respectively: −2, 0, +4, +6, and +6. The higher the oxidation state of a given atom, the greater is its degree of oxidation; the lower the oxidation state, the greater is its degree of reduction. However, the IUPAC rules have three important limitations. First, the calculation procedure is too complex. Several researchers have insisted that these sets of calculations can be rather complex (up to 16 rules to construct computer-driven expert systems), but in any case, they are incomplete and suffer from an increasing number of exceptions.22,23,27−29 To help students understand the complexity of redox reactions, a number of researchers therefore attempted to develop interactive artificial intelligence software.28,29 However, these attempts were not successful because the model itself, which was the basis for the software development, was approximate and also displayed limitations. Indeed, in science it is rather troubling to recognize that these definitions and rules are approximate or have limitations.30 The second limitation of the IUPAC rules was mentioned previously, namely, that there are many exceptions to the rules. For example, hydrogen normally has an oxidation state of 1, while oxygen normally has an oxidation state of −2; however, in LiH, hydrogen has an oxidation state of −1, and in H2O2, oxygen has an oxidation state of −1. Finally, the third limitation is the problems encountered when attempting to calculate the oxidation number of a compound without H or O (e.g., CN−), as the oxidation numbers of other atoms are normally determined

after assigning the oxidation numbers of H and O. Therefore, for the IUPAC method to be useful, the compound of interest should contain either H or O. Limitations of Pauling’s Rule

Pauling31 described the following definition of oxidation number in his book General Chemistry published in 1947 (ref 31, p 173): “The oxidation number of an atom is a number that represents the electric charge that the atom would have if the electrons in a compound were assigned to the atoms in a conventional way.” This way of assignment can be described for atoms in known structures of covalent compounds as follows (ref 31, p 173): The oxidation number of an atom in an elementary substance is zero. The oxidation state of a monatomic ion in an ionic substance is equal to its electric charge. In a covalent compound with a known structure, the oxidation state of each atom is the charge remaining on the atom when each shared electron pair is assigned completely to the more electronegative of the two atoms sharing it. An electron pair shared by two atoms of the same element is usually split between them. In this method, there are no exceptions, and the rules are simpler than in the IUPAC system. In addition, Pauling’s rules explain why the oxidation number of O is −1 in H2O2 and why the oxidation number of H is −1 in LiH. However, this method has four key limitations. First, students must know the Lewis structure of the compound of interest; this is not a trivial task for a lower-level learner. If the Lewis structure of the compound is unknown, calculation of the oxidation number is difficult. Second, large amounts of information are required to verify whether a reaction is in fact a redox reaction. For example, students should follow the steps outlined in Figure 2 to determine the oxidation number of C in acetic acid (CH3COOH). After construction of the Lewis structure of the compound (Figure 2, step 1), students require further information regarding the electronegativity of each atom and the number of valence electrons, and they must also be familiar with the oxidation number equation (oxidation number = number of valence electrons − number of assigned electrons) to solve steps 2 and 3. Pauling’s rules are therefore rather difficult to learn. 565

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H and O) are fixed in advance, so the oxidation number is a property of matter. In Pauling’s rules, the oxidation number is determined on the assumption that the electron pair is associated with the atom with the higher electronegativity. This assumption means that the oxidation number is based on a property (electronegativity) of matter (atom) and thus is also a property of matter. Therefore, the hydrogen, oxygen, electron, and oxidation number models can be classified on the basis of the viewpoint of matter.

In addition, misunderstandings can occur in the imaginary assignment of covalent electron pairs. This method calculates the oxidation number of an atom assuming that the more electronegative atom obtains the covalent electron pair from the bond in a manner similar to an ionic compound. In fact, there is actually a gap between the real movement of electron density and the oxidation number of an atom. For example, the carbon atom of CH4 is more electronegative than the hydrogen atoms according to Pauling’s rules (Figure 3), and therefore,

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A NEW INTERPRETATION OF THE CAUSES OF CONFUSION To address the inherent problems of determining the oxidation state of an atom, Goodstein proposed a new approach (ref 32, p 455): “It is proposed that an oxidation−reduction reaction be defined as a reaction in which a change in the relative order of atomic electronegativities takes place between a given atom and the atoms to which it is bonded prior and subsequent to reaction”. In this method, a redox reaction can be defined on the basis of a process viewpoint, which focuses on the relationship between the reactant and the product. Goodstein defined a redox reaction as one in which the relative order of atomic electronegativities in bonding changes during the reaction. This model can therefore be classified on the basis of the viewpoint of process. Goodstein’s definition can thus be interpreted as outlined below when considering the three types of relationship for the bonding of an atom: (I) Bonding with an atom of higher electronegativity (II) Bonding with an atom of equal electronegativity (III) Bonding with an atom of lower electronegativity (including lone-pair electrons) Therefore, atoms in any reaction could belong to one of the following combinations: (I) ⇄ (I), (I) ⇄ (II), (I) ⇄ (III), (II) ⇄ (II), (II) ⇄ (III), or (III) ⇄ (III). These combinations can then be classified into two types: In the first, there is no change in the relative order of atomic electronegativities in bonding (type A), and in the second, there is a change in the relative order of atomic electronegativities in bonding (type B). Reactions of the form (I) ⇄ (I), (II) ⇄ (II), and (III) ⇄ (III) would be classified as type A, while those of the form (I) ⇄ (II), (I) ⇄ (III), and (II) ⇄ (III) would be classified as type B. Only type B reactions are redox reactions according to Goodstein. As such, Goodstein’s model has the benefit of easily defining redox reactions with no exceptions. In reaction 1, we saw that the N−H bond of the reactant breaks and a new N−B bond is formed in the product. This

Figure 3. Assigning a covalent electron pair to the carbon atom in CH4 according to Pauling’s rules.

the oxidation number of the carbon atom is −4. However, according to Pauling’s rules, this is an irrationally low oxidation state for an organic compound.27 This can be explained by considering the actual electronegativities of the two atoms, which are 2.5 for carbon and 2.1 for hydrogen. These values indicate that the electronegativity difference between the two atoms is small, and thus, determination of the oxidation number using Pauling’s rules has limitations when explaining the reality of covalent bonds. The final limitation of Pauling’s rules regards the ambiguities that may arise in determining oxidation numbers when the compound has more than one Lewis structure. For example, resonance structures can provide more than one Lewis structure for a compound, and therefore, the oxidation number should be considered an average of the values for the different Lewis structures. As shown in Figure 4a, the oxidation number of carbon is −1.2 (−6/5). This assumes that the five carbon atoms of C5H5 have the same Lewis structure. However, Figure 4b illustrates an example in which three different Lewis structures can exist; in this case, the oxidation number determined for carbon differs according to the Lewis structure selected. However, it should be noted that the representation of different resonance structures does not indicate a redox reaction. These observations allow us to conclude that both the IUPAC rules and Pauling’s rules have a common limitation in that they consider the oxidation number to be a property of matter. In this case, matter refers to certain atoms or electrons. In the IUPAC rules, the oxidation numbers of the atoms (e.g.,

Figure 4. Resonance structures for two compounds. 566

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Figure 5. Acid−base reaction (ref 34, p 71).

reaction is therefore classified as (III) ⇄ (III) with respect to nitrogen because the electronegativities decrease in the order N > B > H, and therefore, the reaction is not a redox reaction. Furthermore, in reaction 2, the CO bond of the reactant becomes a CN bond in the product. In this case, the electronegativities decrease in the order O > N > C, and this reaction can be classified as (I) ⇄ (I), thus clearly indicating that this reaction is also not a redox reaction. Moreover, in reaction 3 is classified as (I) ⇄ (I) because the H−Cl bond of the reactant changes to a H−O bond (the electronegativities decrease in the order O > Cl > H), and thus, this reaction is also not a redox reaction. Let us now consider the following reaction: NH3 + BF3 → H3N − BF3

that it can clearly distinguish between acid−base reactions and redox reactions. Although acid−base reactions and redox reactions were not clearly distinguished when the development of the oxidation concept began, acid−base definitions have since been gradually refined, and the Lewis model is now universally accepted. This model was originally proposed by Lewis in 1916, and the extended acid−base definition of coordinate bonds was suggested in 1923, which included the concept of electronpair donors and acceptors. Because the Lewis model is related to electrons, this model obscured the boundary between acid− base reactions and redox reactions, especially in the field of organic chemistry. However, Goodstein’s definition clearly classifies acid−base reactions and redox reactions separately. For example, the reaction shown in Figure 5 could be misclassified as a redox reaction using the oxygen model because the oxygen atom is attached to the aldehydic carbon atom. However, this reaction is actually an acid−base reaction.33 To easily solve this problem, we could introduce the process viewpoint, where the carbon atom that is bonded to oxygen in the reactant changes and becomes bonded to another oxygen atom. Therefore, the order of electronegativities on the carbon atom does not change after the reaction. Of course, the proposed method also has a number of limitations. For example, it is difficult to decide which of the compounds in a redox reaction is the oxidant, which is the reductant, which is the oxidizing agent, and which is the reducing agent. In addition, it will still be necessary to use either the IUPAC rules or Pauling’s rules to address problems related to inorganic chemistry. Furthermore, we must know the connectivity of the atoms in a molecule to apply this method. The use of oxidation numbers in mechanistic organic chemistry will also remain unchanged by the introduction of the Goodstein method once the reaction is determined to be redox rather than acid−base in nature. However, as identified earlier, one of the major problems faced by students is determining whether a reaction is a redox reaction. We therefore propose that they use the Goodstein method to make this determination, and if they conclude that the reaction is a redox reaction, they are faced with a variety of methods to determine the oxidation state of each atom based on, usually, either Pauling’s rules or the IUPAC rules. Despite the issues related to the latter of these rules, it remains the most commonly used in practice. Finally, if we interpret the redox reaction from the process viewpoint, we should focus on the reaction process itself rather than the initial state of the reaction. If a reaction does not occur, it is meaningless to calculate the oxidation number in advance. We could thus examine the change in the electronegativity order of the atoms during the reaction process and determine whether it is a redox reaction.

(4)

In reaction 4, the N atom of the reactant bears a lone pair, which changes to a N−B bond in the product. We assume that the lone pair is classified as (III). Therefore, the electronegativities decrease in the order N > B > lone pair, and this reaction is classified as (III) ⇄ (III), which is, again, not a redox reaction. However, let us consider another reaction: 2H 2 + O2 → 2H 2O

(5)

In reaction 5, hydrogen belongs to type (II) → (I) because the H−H bond of the reactant breaks to give a new H−O bond in the product, and the electronegativities decrease in the order O > H. In contrast, oxygen belongs to type (II) → (III) because the O−O bond of the reactant changes to a H−O bond in the product, and the electronegativities increase in the order H < O. The relative order of the electronegativities of the bond in the reaction changed with respect to the hydrogen atom, so this reaction is a redox reaction. Previously, a number of researchers have proposed new models for defining redox reactions; here we propose a new interpretation for the various diverse models. We can interpret the meaning of arbitrary and ambiguous representations of oxidation state models as trials to assign definite oxidation numbers to individual atoms without considering the process of redox reactions. If we understand redox reactions from the process viewpoint, the ambiguous and arbitrary criteria of the diverse models to define oxidation−reduction reactions are meaningless.

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SUGGESTIONS FOR AN AGREEMENT In this study, we wish to provide a new interpretation for the oxidation−reduction process from an epistemological view of science concepts and to determine why an agreement has not been reached despite continuous suggestions. To address this issue, we suggest that the viewpoint of matter for assigning particular characteristics to an atom or electron should be replaced by a viewpoint of process that focuses on the change in electronegativity order during the reaction. This suggestion can overcome the limitations of the IUPAC rules and Pauling’s rules and make it easier to define oxidation−reduction reactions. Moreover, the process viewpoint is advantageous in

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. 567

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ORCID

(23) Steinborn, D. The Concept of Oxidation States in Metal Complexes. J. Chem. Educ. 2004, 81, 1148−1154. (24) Swinehart, D. F. More on Oxidation Numbers. J. Chem. Educ. 1952, 29, 284−285. (25) Jensen, W. B. Oxidation States versus Oxidation Numbers. J. Chem. Educ. 2011, 88, 1599−1600. (26) IUPAC Compendium of Chemical Terminology - the Gold Book. http://goldbook.iupac.org/ (accessed January 2017). (27) Calzaferri, G. Oxidation Numbers. J. Chem. Educ. 1999, 76, 362−363. (28) Birk, J. P. The Computer as Student: An Application of Artificial Intelligence. J. Chem. Educ. 1992, 69, 294. (29) Eggert, A. A.; Middlecamp, C.; Kean, E. An Oxidation Number Assignment Expert for CHEMPROF. J. Chem. Inf. Model. 1990, 30, 181−187. (30) Gupta, V.; Ganegoda, H.; Engelhard, M. H.; Terry, J.; Linford, M. R. Assigning Oxidation States to Organic Compounds via Predictions from X-ray Photoelectron Spectroscopy: A Discussion of Approaches and Recommended Improvements. J. Chem. Educ. 2014, 91, 232−238. (31) Pauling, L. General Chemistry; Freeman: San Francisco, CA, 1947; p 173. (32) Goodstein, M. P. Interpretation of Oxidation−Reduction. J. Chem. Educ. 1970, 47, 452−457. (33) Smith, D. W. Valence, Covalence, Hypervalence, Oxidation State, and Coordination Number. J. Chem. Educ. 2005, 82, 1202− 1204. (34) Eubanks, I. D. Preparing for Your ACS Examination in Organic Chemistry, The Official Guide, 9th ed.; American Chemical Society: Washington, DC, 2009; p 71.

Seoung-Hey Paik: 0000-0002-0393-4533 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This research was supported by the Basic Science Research Program through the National Research Foundation of Korea (NRF) (2016R1A2B4012991; 2016S1A5B6913974).

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REFERENCES

(1) Anselme, J.-P. Understanding Oxidation-Reduction in Organic Chemistry. J. Chem. Educ. 1997, 74, 69−72. (2) De Jong, O.; Acampo, J.; Verdonk, A. Problems in Teaching the Topic of Redox Reactions: Actions and Conceptions of Chemistry Teachers. J. Res. Sci. Teach. 1995, 32, 1097−1110. (3) Garnett, P. G.; Garnett, P. J.; Hackling, M. W. Students’ Alternative Conceptions in Chemistry: A Review of Research and Implications for Teaching and Learning. Stud. Sci. Educ. 1995, 25, 69− 96. (4) Garnett, P. J.; Treagust, D. F. Conceptual Difficulties Experienced by Senior High School Students of Electrochemistry: Electric Circuits and Oxidation-Reduction Equations. J. Res. Sci. Teach. 1992, 29, 121− 142. (5) Kauffman, J. M. Simple Method for Determination of Oxidation Numbers of Atoms in Compounds. J. Chem. Educ. 1986, 63, 474−475. (6) Ö sterlund, L.-L.; Ekborg, M. Students’ Understanding of Redox Reactions in Three Situations. NorDiNa 2009, 5, 115−127. (7) Woolf, A. A. Oxidation Numbers and Their Limitation. J. Chem. Educ. 1988, 65, 45−46. (8) Ringnes, V. Oxidation−ReductionLearning Difficulties and Choice of Redox Models. Sch. Sci. Rev. 1995, 77, 74−78. (9) Ö sterlund, L.-L.; Berg, A.; Ekborg, M. Redox Models in Chemistry Textbooks for the Upper Secondary School: Friend or Foe? Chem. Educ. Res. Pract. 2010, 11, 182−192. (10) Menzek, A. A New Approach to Understanding OxidationReduction of Compounds in Organic Chemistry. J. Chem. Educ. 2002, 79, 700−702. (11) Swinehart, D. F. More on Oxidation Numbers. J. Chem. Educ. 1952, 29, 284−285. (12) Yoon, H.; Park, H.-N. Analysis of Textbooks and Conception of High School Students, Pre-service Teachers and Chemistry Teachers on Expansion of Oxidation−Reduction Definition. Sch. Sci. J. 2012, 6, 103−108. (13) Hall, W. T. Oxidation-Reduction Reaction. J. Chem. Educ. 1929, 6, 479−485. (14) Justi, R.; Gilbert, J. A Cause of Ahistorical Science Teaching: Use of Hybrid Models. Sci. Educ. 1999, 83, 163−177. (15) Sisler, H. H.; VanderWerf, C. A. Oxidation-Reduction: An Example of Chemical Sophistry. J. Chem. Educ. 1980, 57, 42−44. (16) Vitz, E. Redox Redux: Recommendations for Improving Textbook and IUPAC Definitions. J. Chem. Educ. 2002, 79, 397−400. (17) Silverstein, T. P. Oxidation and Reduction: Too Many Definitions? J. Chem. Educ. 2011, 88, 279−281. (18) VanderWerf, C. A.; Davidson, A. W.; Sisler, H. H. OxidationReduction: A Re-evaluation. J. Chem. Educ. 1945, 22, 450−456. (19) VanderWerf, C. A. A Consistent Treatment of OxidationReduction. J. Chem. Educ. 1948, 25, 547−551. (20) Herron, J. D. What is Oxidation? J. Chem. Educ. 1975, 52, 51− 52. (21) Jurowski, K.; Krzeczkowska, M. K.; Jurowska, A. Approaches to Determining the Oxidation State of Nitrogen and Carbon Atoms in Organic Compounds for High School Students. J. Chem. Educ. 2015, 92, 1645−1652. (22) Holder, D. A.; Johnson, B. G.; Karol, P. J. A Consistent Set of Oxidation Number Rules for Intelligent Computer Tutoring. J. Chem. Educ. 2002, 79, 465−467. 568

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