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A rational reconstruction ofthe origin of the covalentbond and its implications forgeneral chemistry textbooksMansoor NiazVersion of record first published: 20 Jul 2010.

To cite this article: Mansoor Niaz (2001): A rational reconstruction of the originof the covalent bond and its implications for general chemistry textbooks,International Journal of Science Education, 23:6, 623-641

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RESEARCH REPORT

A rational reconstruction of the origin of the covalentbond and its implications for general chemistrytextbooks

Mansoor Niaz, Chemistry Department, Universidad de Oriente, ApartadoPostal 90, Cumana, Estado Sucre, Venezuela 6101A, e-mail:mniaz@sucre.udo.edu.ve

The main objectives of this study are: (i) development of a perspective based on history and philosophyof science considerations (rational reconstruction) in order to understand the postulation of the covalentbond by Lewis; (ii) formulation of four criteria based on the perspective; and (iii) evaluation of 27textbooks based on the four criteria. Results obtained show that most textbooks lacked a history andphilosophy of science perspective and did not deal adequately with the following aspects: (i) Lewis’spostulation of the covalent bond in 1916 posed considerable conceptual difficulties; (ii) Lewis used thecubical atom (a hypothetical entity) in order to understand the sharing of electrons in the covalent bond(octet rule); (iii) sharing of electrons had to compete with the transfer of electrons (ionic bond) con-sidered to be the dominant paradigm until about 1920; (iv) postulation of the covalent bond (octet rule)was not an inductive generalization based on stability of the noble gases and the high dissociation energyof the covalent bonds; and (v) Pauli exclusion principle provides a theoretical explanation of the sharingof electrons, just as the cubical atom did previously. It is concluded that the development of the covalentbond does not follow the inductivist process, viz. experimental observations lead to scientific laws whichlater facilitate the elaboration of explanatory theories.

Introduction

Covalent bonding is considered to be a difficult topic for most freshman students.According to Gillespie et al. (1996), ‘The simplest approach to bonding is in termsof Lewis diagrams and the concept of the shared electron pair. This is as far as thistopic needs to be taken in the most elementary introductions’ (p. 622). Most text-books and chemistry teachers would agree with this—and this ‘consensus’ leadsprecisely to a dilemma. In an attempt to simplify the topic most textbooks presentrules (algorithms) for writing Lewis diagrams for covalent bonds, which are mem-orized by the students. A preliminary survey showed that textbooks devote almost5-10 pages in presenting these rules. Such presentations do not lead the studentstowards a conceptual understanding of the difference between covalent (sharing ofelectrons) and ionic (transfer of electrons) bonds. Recent research in chemicaleducation has shown an increasing interest in emphasizing conceptual understand-ing of a chemistry topic rather than memorizing algorithms (Nurrenbern andPickering 1987, Sawrey 1990, Niaz and Robinson 1992, 1993, Gabel 1993,Nakhleh and Mitchell 1993, Beal and Prescott 1994, Mason et al. 1994, Niaz1995, Lin et al. 1996, Noh and Scharmann 1997, Niaz 1998a).

International Journal of Science Education ISSN 0950-0693 print/ISSN 1464-5289 online # 2001 Taylor & Francis Ltdhttp://www.tandf.co.uk/journals

DOI: 10.1080/09500690010006491

INT. J. SCI. EDUC., 2001, VOL. 23, NO. 6, 623- 641

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Research in science education has recognized not only the importance of his-tory and philosophy of science but also its implications for textbooks (Hodson1988, Solomon 1991, Duschl 1994, Matthews 1994, 1998; McComas et al.1998). In spite of Kuhn’s (1970) advice to the contrary, there seems to be someconsensus that textbooks can facilitate conceptual understanding if students areasked to grapple with rival theories within a history and philosophy of scienceperspective (Siegel 1978, 1979, Brackenridge 1989, Kaufmann 1989, Burbulesand Linn 1991, Kovac 1991, Mahaffy 1992, Stinner 1992, Milne 1998, Niaz2000). Project 2061 (AAAS 1989, 1993) and the National Science EducationStandards (NRC 1996) in the USA, Science in the National Curriculum (NCC1988) in the UK and several other countries have also recognized the importanceof history and philosophy of science. However, in a recent study based on 23general chemistry textbooks, Niaz (1998b) found that only two textbooks men-tioned that Thomson’s cathode ray experiments were conducted against the back-drop of a conflicting (rival) framework, namely, cathode rays could have beencharged particles or waves in the ether. Again, only two textbooks described satis-factorily that Thomson determined mass-to-charge ratio to decide whether cath-ode rays were ions or a universal charged particle. This study attempts to continuethe evaluation of general chemistry textbooks started by Niaz (1998b), and most ofthe textbooks are the same, except that recent editions have been included.

A historical study of the development of the chemical bond shows that the ideaof sharing electrons (covalent bond) posed considerable conceptual constraints forthe scientists.

The main objectives of this study are: (i) development of a perspective basedon history and philosophy of science considerations (rational reconstruction) inorder to understand the postulation of the covalent (shared pair) bond by Lewis;(ii) formulation of criteria based on the perspective that could be useful in theevaluation of general chemistry textbooks; and (iii) evaluation of textbooks utiliz-ing the criteria based on the history and philosophy of science perspective.

A history and philosophy of science perspective

Lewis (1916) is generally considered to have presented the first satisfactory modelof the covalent (shared pair) bond based on the cubic atom in 1916. It is importantto note that the genesis of the cubic atom can be traced to an unpublished memor-andum written by Lewis in 1902 and recounted by him in the following terms,

In the year 1902 (while I was attempting to explain to an elementary class in chem-istry some of the ideas involved in the periodic law) becoming interested in the newtheory of the electron (Thomson’s discovery of the electron in 1897), and combiningthis idea with those which are implied in the periodic classification, I formed an ideaof the inner structure of the atom (model of the cubic atom) which, although it con-tained crudities, I have ever since regarded as representing essentially the arrange-ment of the electrons in the atom. (Lewis 1923: 29-30, emphasis added)

Lewis’s 1902 model of the cubic atom

Lewis (1916: 768) reproduced the following postulates of his 1902 theory of thecubical atom at length in his 1916 article:

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1. In every atom is an essential kernel which remains unaltered in all ordin-ary chemical changes and which possesses an excess of positive chargescorresponding in number to the ordinal number of the group in theperiodic table to which the element belongs;

2. The atom is composed of the kernel and an outer atom or shell, which inthe case of the neutral atom, contains negative electrons equal in numberto the excess of positive charges of the kernel, but the number of elec-trons in the shell may vary during change between 0 and 18;

3. The atom tends to hold an even number of electrons in the shell,and especially to hold eight electrons which are normally arranged sym-metrically at the eight corners of a cube;

4. Two atomic shells are mutually interpenetrable;5. Electrons may ordinarily pass with readiness from one position in the

outer shell to another. Nevertheless they are held in position by more orless rigid constraints, and these positions and the magnitude of the con-straints are determined by the nature of the atom and of such other atomsas are combined with it;

6. Electric forces between particles which are very close together donot obey the simple law of inverse squares which holds at greaterdistances.

After having presented the six postulates Lewis (1916) elaborated by providingfurther information. For example,

The first postulate deals with the two parts of the atom which correspond roughlywith the inner and outer rings of the Thomson atom. The kernel being that part of theatom which is unaltered by ordinary chemical change. . . . (p. 768)

Postulate 2 was illustrated by indicating how chlorine has eight electrons in theouter shell, while forming chlorides. Postulate 3 was the most striking and at thesame time controversial feature of Lewis’s theory, which led to the formulation ofthe ‘rule of eight’ or the ‘octet rule.’ The rule of eight as proposed by Lewisdiffered from the ‘law of octaves’ proposed by Newlands in 1865, according towhich the elements when listed in the order of increasing atomic weights in theperiodic table, the eighth element would be similar to the first. Lewis postulatedthat the eight electrons of an octet formed the eight corners of a cube, as thisprovided, ‘. . . the most stable condition for the atomic shell’ (p. 774). Thus thesingle bond was conceived of as two cubic atoms with a shared edge (pair ofelectrons, see figure 1) and the double bond as two cubes with a common face(see figure 2). Postulate 4 facilitated further the idea of sharing of electrons, ‘Thusan electron may form a part of the shell of two different atoms and cannot be saidto belong to either one exclusively’ (p. 772). Interestingly, in elaborating onPostulate 5, Lewis disagreed with Bohr’s (1913) model of the atom in the followingterms,

It seems to me far simpler to assume that an electron may be held in the atom in stableequilibrium in a series of different positions, each of which having definite con-straints, corresponds to a definite frequency of the electron, . . . (p. 773)

Postulate 6 reinforced a point of view that coincided with new developments in thestructure of the atom,

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. . . a study of the mathematical theory of the electron leads, I believe, irresistibly tothe conclusion that Coulomb’s law of inverse squares must fail at small distances.(p. 773)

Lewis’s model of the covalent bond in retrospect

In this section evidence is provided to show that Lewis’s theory of sharing elec-trons (covalent bond) had to compete with a rival theory, viz. transfer of electrons(ionic bond). From a philosophy of science perspective the rivalry between com-peting theories (paradigms/research programmes) is an integral part of scientificprogress. According to Lakatos (1970),

. . . research programmes have achieved complete monopoly only rarely and then onlyfor relatively short periods. . . . The history of science has been and should be a historyof competing research programmes . . . . (p. 155)

According to Kohler (1971), who has presented a detailed account of the originof Lewis’s ideas,

When it was first proposed, Lewis’s theory was completely out of tune with estab-lished belief. For nearly 20 years it had been almost universally believed that all bondswere formed by the complete transfer of one electron from one atom to another. Theparadigm was the ionic bond of Na‡ Cl¡ , and even the bonds in compounds such asmethane or hydrogen were believed to be polar, despite their lack of polar properties.From the standpoint of the polar theory the idea that two negative electrons couldattract each other or that two atoms could share electrons was absurd. (p. 344)

Rodebush (1928), a chemist reviewing the origin of the covalent bond in the late1920s, shared the same concern,

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Figure 1.

Figure 2.

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Since according to Coulomb’s law two electrons should exert a repulsion for eachother, the pairing of electrons seems at first glance to be a bizarre idea. In order toaccount for the peculiar behavior Lewis assumed the existence of a magnetic attrac-tion between the electrons. (pp. 513-514)

Lewis (1916) further clarified his attempt at building a theory of the atom,

In my original theory [1902] I considered the elements in the periodic table thus builtup, as if block by block, forming concentric cubes. (p. 769)

Later in the article Lewis recognizes that the cubic structure cannot represent thetriple bond and suggests its replacement by the tetrahedral atom (p. 780).

At this stage it is important to note that Thomson’s (1897) discovery of theelectron in 1897 and later publications (Thomson 1907) provided powerful argu-ments for the polar theory of the ionic bond. According to Thomson (1907),

For each valency bond established between two atoms the transference of one cor-puscle from the one atom to the other has taken place . . . . (p. 138)

Although Thomson accepted that overlapping of corpuscles could produce a non-polar bond in theory, he believed that in reality all bonds were polar bonds (p. 131).

One of the best known theories of electrovalence representing the polar ortho-doxy was proposed by Abegg (1904). According to this theory the only kind ofchemical affinity was electrostatic attraction, and all bonds, even in non-polarsymmetrical molecules, were electron transfer bonds.

Historical antecedents of the covalent bond

According to Kohler (1971), although sharing of electrons to form covalent bondsseemed shocking at first, few chemists have shown interest in the origin of theshared pair bond. In the previous sections it has been shown that the cubic atomwas important in the development of Lewis’s theory of the shared pair covalentbond. Lewis’s cubic atom was first conceived as a teaching device to illustrate theoctet rule and can be considered as ‘speculative’. Modern philosophy of science hasemphasized the importance of speculation as characteristic of scientific process(Kuhn 1970, Lakatos 1970). For example, Bohr’s use of the quantum postulatein his theory is considered to be speculative. Interestingly, Lewis (1919) in a letterto Robert Millikan complained,

. . . I could not find a soul sufficiently interested to hear the theory [1902 memoran-dum]. There was a great deal of research work being done at the university [Harvard],but, as I see it now, the spirit of research was dead.

First dissenting voices against the polar orthodoxy were raised in 1913 by Brayand Branch (1913), colleagues of Lewis at MIT. Bray and Branch (1913) objectedthat the polar theory had been extended far beyond its proper limits and suggested,

. . . there are two distinct types of union between atoms: polar, in which an electronhas passed from one atom to the other, and non-polar, in which there is no motion ofan electron. (p. 1443)

According to Kohler (1971),

By 1913 or so the polar theory completely dominated chemistry, and it did until it wasreplaced by Lewis’s theory in the early 1920s. (p. 355)

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After the lead of Bray and Branch, Lewis himself published two papers (1913,1916).

Arsem (1914), a student of Bray at MIT, was the next to publish with thefollowing introductory comments,

I have . . . ventured to present my views at this time with a feeling that they are inharmony with the present trend of scientific speculation’. (p. 1656)

Arsem (1914) presented his critique in the following terms,

There is a difficulty in accounting, on the basis of Thomson’s theory [1907], for theexistence of a hydrogen molecule made up of two positive atoms, or of a chlorinemolecule with two negative atoms, and it is hard to see why two neutral atoms or twoatoms having equal valency of the same sign should combine to form a stablemolecule. (p. 1658)

Arsem’s model of the hydrogen molecule was based on a single electron thatoscillated between the two atoms. The oscillating electron was an intrinsic partof both atoms at the same time. Although Arsem’s ‘half-electron bond’ was con-sidered to be flawed, it stopped just one step short of the two electron shared pairbond. Later Thomson (1914) himself changed and accepted that all bonds were notpolar bonds after all. Thomson conceived the non-polar bond as a tube of forcestretching from an electron in one atom to the nucleus of a second.

Parson (1915) was the next to contribute towards the development of thecovalent bond. Following the French physicist Langevin (1872-1946) he con-ceived the idea that the force responsible for chemical bonding was not electricalbut magnetic. Langevin had proposed in 1905 that the electron was a zero resist-ance electric circuit, i.e. an electro-magnet (magnetons). Parson’s atom consistedof a sphere of positive charge (Thomson 1907) in which the magnetons werearranged not in concentric rings but at the corners of cubic octets, reminiscentof Lewis’s 1902 cubical atom. Interestingly, Parson finished his manuscript whilespending a year at Berkeley with Lewis. Apparently, Parson had only two inter-views with Lewis, and in one of the interviews Lewis drew a cubic atom andremarked, ‘I once had the idea of a cube corresponding to the octave law’(Reproduced in Kohler 1971: 370).

Kohler’s reconstruction of Lewis’s theory of the covalent bond presentedabove contrasts with the interpretation of Rodebush (1928), a chemist who wasreviewing the literature on the development of the covalent bond. Rodebush con-siders Lewis’s shared pair covalent bond to have been induced from two empiricalfacts: (i) Moseley’s demonstration that helium, an inert gas, has a pair of electrons;and (ii) The fact that with a few exceptions the number of electrons in all com-pounds is even. According to Kohler (1971), ‘Rodebush’s empiricistic view revealswhat chemists at the time thought their science ought to be like’ (p. 371). Thisshows how Lewis’s work was incorrectly interpreted even in the late twenties.

Origin of the covalent bond: a Baconian inductive ascent?

At this stage it is instructive to study how philosophers of science have interpretedBohr’s model of the atom. According to Lakatos (1970),

. . . Bohr’s problem was not to explain Balmer and Paschen series, but to explain theparadoxical stability of the Rutherford atom. Moreover, Bohr had not even heard ofthese formulae before he wrote the first version of his paper. (p. 147)

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This version of the events has been corroborated by an extremely careful anddetailed study by Heilbron and Kuhn (1969). Interestingly, most textbooks con-sider Bohr’s major contribution to be the explanation of the Balmer and Paschenseries of the hydrogen line spectrum. On the contrary, philosophers of scienceconsider Bohr’s explanation of the paradoxical stability of the Rutherford modelof the atom as his major contribution (Heilbron and Kuhn, 1969; Lakatos, 1970).In a study based on 23 general chemistry textbooks, Niaz (1998b) found that veryfew textbooks considered the explanation of the paradoxical stability of theRutherford model of the atom as important and none of the textbooks interpretedthe quantization of the Rutherford model within a historical perspective.

Lakatos (1970) goes on to show the importance of this event in the history ofscience,

Since the Balmer and the Paschen series were known before 1913 [year of Bohr’s firstpublication], some historians present the story as an example of a Baconian ‘‘inductiveascent’’: (1) the chaos of spectrum lines, (2) an empirical law (Balmer), (3) the theor-etical explanation (Bohr). (p. 147)

A major premise of historians who follow the Baconian inductive ascent is thatscientific theories and laws are primarily driven by experimental observations.Furthermore, such empiricist interpretations consider scientific progress to bedichotomous, viz. experimental observations lead to scientific laws, which laterfacilitate the elaboration of explanatory theories.

A major thesis of this article is that the conceptualization of the covalent bondby chemists and textbooks approximates quite closely to a Baconian inductiveascent, according to the following stages:

1. The finding that diatomic molecules such as H2 and the hundreds ofcompounds found by organic chemists in the late nineteenth centurycould not be understood by the ionic bond (this corresponds to thechaos of spectrum lines before Balmer’s law).

2. Postulation of the non-polar shared pair covalent bond by Lewis as aninductive law or generalization. (This corresponds to Balmer’s empiricallaw for the hydrogen line spectrum.)

3. Theoretical explanation offered by quantum theory (Pauli exclusionprinciple) as to how two electrons (in spite of the repulsion) can occupythe same space. (This corresponds to Bohr’s explanation of the hydrogenline spectrum.) This aspect is corroborated by the following interpret-ation of the events by Rodebush (1928),

It will be recognized by the chemist however that Pauli’s rule [exclusion principle] isonly a short hand way of saying what Lewis has assumed for many years as the basis ofhis magnetochemical theory . . . of valence. If the electrons are paired in the atommagnetically, it is easy to see how two unpaired electrons in different atoms may becoupled magnetically and form the nonpolar bond. (p. 516)

The crux of the issue is that the inductivist interpretation construes Pauli’sexclusion principle as the theoretical explanation and ignores the fact that Lewis’scubic atom was crucial for his later explanation of the sharing of electrons. Thusscientific progress is characterized by a series of theories or models (plausibleexplanations), which vary in the degree to which they explain the experimentalfindings. In other words, science does not necessarily progress from experimentalfindings to scientific laws to theoretical explanations. According to Lakatos (1970:

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129) the conflict is not between theories and laws but rather between an interpret-ative and an explanatory theory. Blanco and Niaz (1997) found that manychemistry teachers and students consider progress in science to be characterizedby a ‘Baconian inductive ascent’, that is experimental findings !scientificlaws !theoretical explanations. Furthermore, Blanco and Niaz (1997) concludedthat this finding can provide, ‘A plausible blueprint for alternatives to the tradi-tional textbook treatment of progress in science’ (p. 228). The blueprint for analternative approach in the present case would be a textbook presentation empha-sizing the origin of the covalent bond as a product of conflicting or rival theories(models) for the explanation of bond formation. This shows that appropriate his-torical reconstructions can benefit students both by providing them with modelsfor alternative/rival approaches and by instilling in them a deeper conceptualunderstanding of the topic (cf. Chiappetta et al. 1991).

Criteria for the evaluation of general chemistry textbooks

Based on the history and philosophy of science perspective presented in theprevious section, here I present criteria for the evaluation of freshman/collegelevel introductory chemistry textbooks.

1. Lewis’s cubic atom as a theoretical device for understanding the sharing ofelectrons: Lewis’s cubic atom was based on his atomic theory based onpostulates formulated in 1902. The cubic atom was thus a theoreticaldevice that was later used for understanding the sharing of electrons(covalent bond) and provided the rationale for the octet rule. This cri-terion is based on the following references: Lewis (1916, 1923), Kohler(1971) and Jensen (1984). Following classifications were elaborated:

Satisfactory (S): treatment of the subject in the textbook is considered tobe satisfactory if it is briefly explained that Lewis (1916) used his modelof the cubic atom to explain the sharing of electrons and the octet rule.Mention (M): a simple mention of Lewis’s cubic atom.No-mention (N): no-mention of Lewis’s cubic atom.

2. Sharing of electrons (covalent bond) had to compete with the transfer ofelectrons (ionic bond): Lewis’s idea of sharing electrons (covalent bond)had to compete with the transfer of electrons (polar/ionic bond). Theorigin of the polar bond as the dominant paradigm in chemical combina-tion can be traced to Thomson’s discovery of the electron in 1897. By1913 the polar theory completely dominated chemistry, and it was in theearly 1920s that Lewis’s idea of sharing electrons became acceptable.This criterion is based on the following references: Thomson (1897,1907, 1914), Lewis (1916, 1923), Lakatos (1970) and Kohler (1971).Following classifications were elaborated:

Satisfactory (S): treatment of the subject is considered to be satisfactoryif the role of competing frameworks (polar/non-polar) is brieflydescribed.Mention (M): a simple mention of the competing frameworks.No-mention (N): no-mention of the competing frameworks.

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3. Covalent bond: inductive generalization/derived from the cubical atom: theobjective of this criterion (Rodebush 1928, Lakatos 1970; Kohler 1971) isto evaluate if the textbooks follow one of the following interpretationswith respect to the origin of the (shared pair) covalent bond:

Inductivist (I): Lewis’s covalent bond was an inductive generalizationbased on: stability of the noble gases or formation of the hydrogen mol-ecule leads to a lowering of the energy or Helium an inert gas has a pair ofelectrons or number of electrons in most compounds are even.Lakatosian (L): Lewis’s (shared pair) covalent bond was not inducedfrom experimental evidence but derived from the cubic atom.No-mention (N): textbook makes no-mention explicitly toeither of the two interpretations, presented above.

4. Pauli exclusion principle as an explanation of the sharing of electrons incovalent bonds: the objective of this criterion is to evaluate if textbooksconsider Pauli’s exclusion principle to provide an explanation of thesharing of electrons. This criterion is based on the following references:Pauli (1925), Rodebush (1928), Lakatos (1970) and Kohler (1971).Following classifications were elaborated:

Satisfactory (S): treatment of the subject in the textbook is considered tobe satisfactory if the role of Pauli exclusion principle is briefly described,in order to explain the covalent bond.Mention (M): a simple mention of Pauli exclusion principle, in the con-text of the covalent bond.No-mention (N): no-mention of Pauli exclusion principle.

In order to implement the criteria, a university chemistry professor with aPhD in inorganic chemistry and 25 years of teaching experience at both the fresh-man and higher levels, and the author applied the criteria separately to evaluatethree textbooks (selected randomly). It was found that both evaluators coincidedon all four criteria on two of the textbooks. On the third textbook the evaluatorscoincided on three of the four criteria. The point of disagreement referred tocriterion 3: In order to be classified as Inductivist (I), should the textbook referto all, some or one of the following reasons: stability of the noble gases, lowering ofenergy in the H2 molecule, He and most compounds have an even number ofelectrons. After some discussion it was decided that the textbook could be classi-fied as (I), if it referred to at least one of the above mentioned reasons. Similarly,an instructor with three years of teaching experience at the freshmen level was alsoasked to evaluate three textbooks. It was found that both evaluators coincided onall four criteria on two of the textbooks. On the third textbook the evaluatorscoincided on 3 of the 4 criteria. The point of disagreement referred to criterion4: in order to be classified as Satisfactory (S), should the textbook explicitly referto the Pauli Exclusion Principle. After some discussion it was decided that a text-book be classified as Mention (M) even if it does not explicitly refer to the PauliPrinciple but still mentions that the two electrons in the covalent bond must haveopposite spins. With this experience, the rest of the textbooks were then evaluatedby the author.

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Evaluation of chemistry textbooks: results and discussion

Criterion 1

Table 1 shows that only one textbook mentioned (M) Lewis’s cubic atom and threedescribed it satisfactorily (S). Following is an example of a satisfactory description:

Lewis assumed that the number of electrons in the outermost cube on an atom wasequal to the number of electrons lost when the atom formed positive ions . . . heassumed that each neutral atom had one more electron in the outermost cube thanthe atom immediately preceding it in the periodic table. Finally, he assumed it tookeight electrons—an octet—to complete a cube. Once an atom had an octet of electronsin its outermost cube, this cube became part of the core, or kernel, of electrons aboutwhich the next cube was built. (Bodner and Pardue 1989: 273)

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Table 1. Evaluation of general chemistry textbooks (covalent bond)based on a history and philosophy of science perspective.

Criteria

No. Textbook 1a 2b 3c 4d

1. Ander and Sonnessa 1968 N N I M2. Bodner and Pardue 1989 S N L N3. Brady and Holum 1981 N N I S4. Brady and Humiston 1996 N N I N5. Brown and LeMay 1988 N N I N6. Brown et al. 1998 N N I S7. Burns 1996 N N I N8. Chang 1981 N N I N9. Chang 1999 N N I S

10. Daub and Seese 1996 N N I N11. Dickerson et al. 1970 N N I S12. Ebbing 1997 N N I N13. Hein and Arena 1997 N N I N14. Holtzclaw and Robinson 1988 N N I N15. Joesten et al. 1991 N N I N16. Lippincott et al. 1977 N N I M17. Mahan and Myers 1990 S N L S18. Masterton et al. 1989 M N I M19. Mortimer 1983 N N I S20. Oxtoby et al. 1990 N N I N21. Pauling 1977 N N N M22. Petrucci 1977 S M N S23. Segal 1989 N N I S24. Stoker 1990 N N I N25. Whitten, Davis and Peck 1998 N N I M26. Wolfe 1988 N N I N27. Zumdahl 1990 N N I N

a Lewis’s cubic atom as a theoretical device for understanding the sharing of electrons.b Sharing of electrons (covalent bond) had to compete with the transfer of electrons (ionic bond).c Covalent bond: Inductive generalization/derived from the cubic atom.d Pauli exclusion principle as an explanation of the sharing of electrons in covalent bonds.(S ˆSatisfactory; M ˆMention; N ˆNo mention; I ˆ Inductivist; L ˆLakatosian).

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With this introduction Bodner and Pardue (1989) explain the formation of thecovalent bond in the following terms,

By 1916, Lewis had realized that there was another way atoms could combine toachieve an octet of valence electrons—they could share electrons. Two fluorineatoms, for example, could share a pair of electrons and thereby form a stable F2

molecule in which each atom had an octet of valence electrons. (p. 274)

The authors provide pictures of the individual cubes of fluorine comingtogether to share an edge and thus form the covalent bond in which the eightelectrons are oriented towards the corners of a cube. Furthermore, the authorsreproduce Lewis’s 1902 memo with hand drawings of the cubic atom that wasincluded by Lewis (1923) in his book on valence.

It is important to note that Bodner and Pardue (1989) go beyond by recogniz-ing the importance of Lewis’s contribution within a historical perspective, ‘Themagnitude of this achievement is underlined by the fact that this model was gen-erated only five years after Thomson’s discovery of the electron and nine yearsbefore Rutherford proposed the model of the atom . . .’ (p. 273).

Mahan and Myers (1990), the other textbook that gave a satisfactory descrip-tion, also included Lewis’s 1902 handwritten memo, and emphasized that it wasbased on conjecture, which later led him to formulate his valence theory. Anothertextbook (Chang 1981: 156) also reproduced the handwritten memo (about half apage) with no reference to the cubic atom nor as to what Lewis was trying to do.Interestingly, the recent edition (Chang 1999) of this textbook does not includeLewis’s handwritten memo. One interpretation of this (and others cannot be ruledout) can be that if a study similar to this had appeared earlier the author (Chang)could have found a reason for reproducing the memo. This shows that textbookslack a history and philosophy of science framework. According to Coppola andDaniels (1998), ‘When history does appear, it often does so in neatly isolated andeasily neglected textbook side bars . . .’ (p. 33). Similarly, Mahaffy (1992) in a studybased on five general chemistry textbooks concluded, ‘Unfortunately, the contem-porary texts examined in this study still often present history as simple chronol-ogy, interesting anecdote, or curiosity’ (p. 54).

Criterion 2

One of the textbooks mentioned (M) and none described satisfactorily (S) thatLewis’s idea of sharing of electrons (covalent bond) had to compete with thetransfer of electrons (ionic bond). It is important to note that a number of text-books start with the presentation of the covalent bond in terms that can be useful.Nevertheless, these authors do not interpret the origin of the covalent bond as arival research program, based on a history and philosophy of science perspective(Lakatos 1970). Following are examples of what constitute helpful preambles, butdo not facilitate the understanding of the covalent bond:

. . . when non-metals combine with other non-metals . . . In such instances, natureseeks a different way to lower the energy—electron sharing . . . This sharing of elec-trons is called the covalent bond. (Brady and Holum 1981: 219)

The vast majority of chemical substances do not have the characteristics of ionicmaterials; water, gasoline, banana peelings, hair, antifreeze, and plastic bags are ex-amples.. . . For the very large class of substances that do not behave like ionic sub-

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stances, we need a different model for the bonding between atoms. Lewis reasonedthat an atom might acquire a noble-gas electron configuration by sharing electronswith other atoms. A chemical bond formed by sharing a pair of electrons is called acovalent bond. (Brown and LeMay 1988: 233)

Although the existence of molecules was hypothesized as early as the seventeenthcentury, chemists had practically no idea of how and why molecules formed. Thefirst major breakthrough came in the 1910s, when Gilbert Lewis discovered the roleof electrons in chemical bond formation. In particular, he developed the importantidea that a chemical bond exists when a pair of electrons is shared between two atoms.(Chang 1981: 154)

Covalent bonds also result in a molecule lower in energy than the isolated atoms, butsince in a pure covalent bond there is equal sharing of electrons and a symmetriccharge distribution, a simple Coulomb’s law calculation cannot explain the energylowering as it does for ionic bonding. In order to understand many properties ofmolecules, such as their geometries and the strengths of the chemical bonds thathold the molecules together, we must investigate the nature of the covalent bond.(Segal 1989: 496, emphasis added) (Comment: Soon after this, the author, insteadof investigating the nature of the covalent bond, provides students with rules forwriting the Lewis electron dot structures.)

Examples from these four textbooks show that there is a common threadrunning through them, viz., there are two fundamentally different ways of formingchemical bonds. By emphasizing this the textbooks set the stage for the students tostart thinking. Nevertheless, instead of fostering a critical approach based on theproblematic nature of the covalent bond and its rivalry with the then-dominantionic bond, these textbooks credit Lewis with having discovered the ‘magical’solution—the sharing of electrons.

At this stage it is instructive to compare the helpful preambles (presentedabove) with examples of textbooks that present the difference between the twotypes of bonds as something natural and non-problematic:

Sodium chloride, NaCl, consists of Na‡ and Cl¡ ions united in a regular dispositionor crystal, formed by ionic bonds. The ionic bond results from the electrostatic attrac-tion between oppositely charged ions. A second class of chemical bond is the covalentbond. In a covalent bond, two atoms share valence electrons (outer shell electrons),that are attracted towards the positively charged centers of the two atoms. For ex-ample, chlorine gas consists of molecules of Cl2. (Ebbing 1997: 344, original empha-sis)

Another textbook first asked the following question, ‘what is it that determineswhether the interaction of two elements produces ions or molecules?’ (Stoker 1990:189). A few lines later the textbook provided the solution in the following terms:

It is useful to classify chemical attractive forces (chemical bonds) into two categories:ionic bonds and covalent bonds. An ionic bond is formed when one or more electronsare transferred from one atom or group of atoms to another . . . A covalent bond isformed when two atoms share one or more electron pairs between them. (Stoker1990: 189, original emphasis)

It is not far-fetched to suggest that if the sharing of electrons was considered tobe ‘bizarre’ and ‘absurd’ by the scientists (cf. Kohler 1971: 363), it could appearcounterintuitive to students as well. The controversial origin of the covalent bondand its rivalry with the ionic bond provides a good opportunity to illustrate howprogress in science is based on controversy and how established theories or ways ofthinking are difficult to change. Thus besides providing students detailed

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instructions for writing the Lewis structures (almost all textbooks do that), a briefreference to the historical details can facilitate conceptual understanding of thedifference between the two types of bonds.

Criterion 3

Table 1 shows that 23 textbooks consider the origin of the covalent bond to bean inductive (I) generalization. Following are some of the examples of suchpresentations:

This shell of eight appears highly stable, and . . . compound formation is limited to theless tightly bound valence electrons. Further confirmation, of the stability of theunderlying shell of eight electrons is obtained by the observation that neon andargon, each with eight electrons in the outer shell, are inert, like helium, and donot readily form compounds. (Lippincott et al. 1977: 67)

Because all noble gases (except He) have eight valence electrons, many atoms under-going reactions also end up with eight electrons. This observation has led to what isknown as the octet rule. (Brown and LeMay 1988: 225)

The simplest substance in which atoms are covalently bonded is the hydrogen mol-ecule, H2. A hydrogen atom has one electron in its 1s shell. The two electrons . . . fromthe two hydrogen atoms in a hydrogen molecule are shared by the two nuclei . . .These shared electrons spend most of their time in the region between the two nuclei,and the electrostatic attraction between the two positively charged nuclei and the twonegatively charged electrons holds the molecule together. The bond resulting fromthis attraction is very strong, as is evidenced by the large amount of energy required tobreak the covalent bonds in 1 mole of hydrogen molecules—436 kilojoules. (Holtzclawand Robinson 1988: 141-142)

These presentations are quite representative of most textbooks and showclearly that the octet rule is sustained by empirical evidence, viz., stability of thenoble gases and the dissociation energy of the covalent bond in the hydrogenmolecule.

The alternative interpretation (Lakatosian) would have emphasized the cubicatom—a hypothetical entity. No wonder, only two textbooks (Bodner and Pardue1989; Mahan and Myers 1990) presented the Lakatosian interpretation. Thesetextbooks trace the origin of the stability of the covalent bond to the cubic atomand go to considerable length to show that Lewis’s ideas developed slowly based onconjectures and were tentative. It is important to note that these two textbooks alsogo on to mention the stability of the noble gases and the high dissociation energy ofcovalent bonds. At first sight the difference between the two approaches may seemtrivial. Nevertheless, the inductivist approach considers all scientific findings to bedriven by experiment, which has been referred to by Gillespie (1997) as, ‘Puttingobservations first . . .’ (p. 485). Niaz (1999) has presented a critical appraisal of thisapproach. Similarly, Matthews (1994) traces the origins of such approaches in thehistory of science. The previous section illustrates how such presentations comequite close to what has been referred to as a ‘Baconian inductive ascent’ in thehistory and philosophy of science.

Interestingly, the octet rule itself shows the tentative nature of scientific pro-gress. Although most textbooks do point out its limitations a recent critique goesbeyond,

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My concern is that for many students the octet rule is more than a heuristic to dis-criminate stable from unstable species. . . . To present it as an explanatory frameworkfor understanding chemical reactions is both intellectually and pedagogically suspect.(Taber 1995: 84)

At this stage it is important to differentiate between the octet rule and thecubic atom. Lewis’s model of the cubic atom (and not the octet rule) was a theor-etical construct and led to the first tentative (conjectural) explanation of the cova-lent bond. A better explanation (‘progressive problemshift’ cf. Lakatos 1970) isprovided by Pauli exclusion principle (Pauli 1925), the subject of the next cri-terion.

Criterion 4

Table 1 shows that five textbooks mentioned (M) and eight described satisfactorily(S) Pauli Exclusion Principle as an explanation of the sharing of electrons (withopposite spin) in covalent bonds. Following are examples of a satisfactory description:

When two hydrogen atoms form a covalent bond, the atomic orbitals overlap in such away that the electron clouds reinforce each other in the region between the nuclei, andthere is an increased probability of finding an electron in this region. According toPauli exclusion principle, the two electrons of the bond must have opposite spins. Thestrength of the covalent bond comes from the attraction of the positively chargednuclei for the negative cloud of the bond. . . . (Mortimer 1983: 135)

By Pauli’s reasoning . . . The hydrogen molecule, H2, has two nuclei and two elec-trons. Both electrons can be accommodated in the [bonding] orbital if their spins arepaired, and a covalent electron pair bond is created. (Dickerson et al. 1970: 333)

This interpretation of a satisfactory description coincides with Gillespie et al.(1996), who consider the Pauli exclusion principle to be significant for freshmanchemistry students only if stated in terms that refer to the physical consequences ofthe principle, ‘Electrons with the same spin have a low probability of being closetogether and a high probability of being far apart, whereas there is no restriction onelectrons of opposite spin, which may come close together’ (p. 623).

Do textbooks follow a Baconian inductive ascent?

Presentations of some of the textbooks came quite close to what could be consid-ered as a Baconian inductive ascent, as the following two examples demonstrate:

Most of the reasons for matter bonding to matter . . . can be summarized by twoconcise notions: (i) Unlike charges attract; (ii) Electrons tend to exist in pairs.Couple these two ideas (one empirical, one theoretical) with the proximity require-ment that only the outer electrons of the atoms (the valence electrons) interact, and youhave the basic concepts that explain how atoms in over 10 million compounds bond toeach other . . . . (Joesten et al. 1991: 109, original emphasis)

In observing millions of stable compounds, chemists have learned that in almost allstable chemical compounds of the representative elements, all of the atoms have achieved anoble gas electron configuration. The importance of this observation cannot be over-stated. It forms the basis of our fundamental ideas about why and how atoms bond toeach other. (Zumdahl 1990: 373, original emphasis)

These are good examples of the first two steps of the Baconian inductiveascent, namely, chaos of experimental observations led to the postulation of the

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covalent bond. Nevertheless, these two textbooks did not go on to the third step—explanation of the covalent bond by Pauli’s exclusion principle.

Interestingly, of the eight textbooks that did consider Pauli’s exclusion prin-ciple as a satisfactory (S) explanation of the covalent bond, only two referred toLewis’s cubic atom. This shows that lack of a history and philosophy of scienceperspective perhaps leads the textbooks to recognize the role of Pauli’s exclusionprinciple and yet not even mention Lewis’s cubic atom.

Are recent textbooks different from those published earlier?

In order to respond to this question textbooks were divided into two groups: thosepublished before 1990 and those published since 1990. Table 1 shows that oncriterion 1, three textbooks described Lewis’s cubic atom satisfactorily (S), twowere published before 1990 (Petrucci 1977; Bodner and Pardue 1989) and one in1990 (Mahan and Myers 1990). On criterion 2, none of the textbooks describedsatisfactorily (S) that sharing of electrons (covalent bond) had to compete with thetransfer of electrons (ionic bond). One textbook made a mention (M) and waspublished before 1990 (Petrucci 1977). On criterion 3, only two textbooks had aLakatosian (L) interpretation, one before 1990 (Bodner and Pardue 1989) and theother in 1990 (Mahan and Myers 1990). On criterion 4, eight textbooks describedsatisfactorily (S) Pauli’s exclusion principle as an explanation of the sharing ofelectrons. Five of these were published before 1990 (Dickerson et al. 1970;Petrucci 1977; Brady and Holum 1981; Mortimer 1983; Segal 1989) and three in1990 or later (Mahan and Myers 1990, Brown et al. 1998; Chang 1999). Theseresults show that over the last 30 years textbooks have not changed much. Itappears that history and philosophy of science are still marginal to the interestsof most textbook authors.

This study also shows that an adequate (satisfactory) treatment of a criterion(topic) by an edition of a textbook that is out of print is still useful for the teacherand future textbook writers. It provides them material for developing classroomstrategies. Furthermore, although at the freshman level most teachers rely heavilyon the textbook, the importance of supplementary readings cannot be overlooked.An example of a course that used supplementary readings with respect to historyand philosophy of science is provided by Kovac (1991). Interestingly, one of thetopics covered by Kovac was the history of the covalent bond, and he used thearticle by Kohler (1971).

Conclusion: from Lewis’s cubic atom to Pauli’s exclusionprinciple

Lewis’s postulation of the cubic atom as a hypothetical entity shows that progressin science does not necessarily proceed from experimental observations to scien-tific laws and finally theoretical explanations.

It is concluded that most general chemistry textbooks lacked a history andphilosophy of science perspective and did not deal adequately with the followingaspects:

1. A reconstruction of the origin of the covalent bond enables us to under-stand how new ideas are resisted and also that the sharing of electrons

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had to compete with the transfer of electrons (ionic bond) considered tobe the dominant paradigm until about 1920.

2. Transition from Lewis’s cubic atom !Pauli’s exclusion principle !what next, provides an illustration of how scientific knowledge is tenta-tive. According to Project 2061, ‘The notion that scientific knowledge isalways subject to modification can be difficult for students to grasp. Itseems to oppose the certainty and truth popularly accorded to science,and runs counter to the yearning for certainty that is characteristic ofmost cultures, perhaps especially so among youth’ (AAAS 1993, p. 5).

3. Lewis’s postulation of the covalent bond (sharing electrons) in 1916posed considerable conceptual difficulties.

4. Postulation of the covalent bond (octet rule) was not an inductive gen-eralization based on stability of the noble gases and the high dissociationenergy of the covalent bonds.

5. Pauli’s exclusion principle provides a theoretical explanation of thesharing of electrons, just as the cubic atom did previously.

An important implication of this study that is relevant for future textbooks isthat Lewis’s theory of shared electron pair covalent bond when first proposed in1916 conflicted with the then-dominant paradigm of the ionic bond based on thetransfer of electrons. This conflict led to a rivalry between the two ways of con-ceptualizing bond formation and it was in the early 1920s that Lewis’s theory ofsharing electrons became acceptable.

Based on the analyses conducted in this study it is suggested that a discussionalong the following lines could facilitate conceptual understanding. Ionic bondsare formed by the actual transfer of electrons, which produces positively andnegatively charged ions. Formation of the ionic bond leads to a lowering of energy(stabilization) because of coulombic or electrostatic attraction between ions ofopposite charge. In this context, how can we explain the lowering of energywhen two electrons are shared to form a covalent bond? Apparently, the approachof two electrons having the same charge should produce repulsive forces and henceproduce destabilization. Thus it is not surprising that when first proposed the ideaof a covalent bond was considered to be ‘absurd’ and ‘bizarre’. Lewis and collea-gues were among the first chemists to support the rival theory (hypothesis/idea) ofcovalent bonding by the postulation of a model based on the cubic atom. Later thequantum theory provided further support to the theory of sharing electrons whenPauli introduced his exclusion principle.

Acknowledgments

Thanks are due to the four anonymous reviewers, Mar¹a Asuncion Rodr¹guez deAguirrezabala and Gustavo Liendo (Department of Chemistry, Universidad deOriente) for making valuable suggestions towards the improvement of the manu-script. Research reported here was supported by a grant from the Consejo deInvestigacion, Universidad de Oriente (Project No. CI-5-1004-0849/2001).

References

Abegg, R. (1904) Die valenz und das periodische system. Z. Anorg. Chem., 39, 330-380.

638 M. NIAZ

Dow

nloa

ded

by [

Rye

rson

Uni

vers

ity]

at 0

1:50

29

Apr

il 20

13

American Association for the Advancement of Science, AAAS (1989) Project 2061:Science For All Americans (Washington, DC: AAAS).

American Association for the Advancement of Science, AAAS (1993) Benchmarks forScience Literacy: Project 2061 (New York: Oxford University Press).

Ander, P. and Sonnessa, A. J. (1968) Principles of Chemistry (New York: Macmillan).Arsem, W. C. (1914) A theory of valence and molecular structure. Journal of American

Chemical Society, 36, 1655-1675.Beall, H. and Prescott, S. (1994). Concepts and calculations in chemistry teaching and

learning. Journal of Chemical Education, 71, 111-112.Blanco, R. and Niaz, M. (1997) Epistemological beliefs of students and teachers about

the nature of science: From ‘baconian inductive ascent’ to the ‘irrelevance’ of scien-tific laws. Instructional Science, 25, 203-231.

Bohr, N. (1913) On the constitution of atoms and molecules. Philosophical Magazine, 26, 1-25.

Bodner, G. M. and Pardue, H. L. (1989) Chemistry: An Experimental Science (New York:Wiley).

Brackenridge, J. B. (1989) Education in science, history of science, and the textbook—necessary vs. sufficient conditions. Interchange, 20, 71-80.

Brady, J. E. and Holum, J. R. (1981) Fundamentals of chemistry (New York: Wiley).Brady, J. E. and Humiston, G. E. (1996) General chemistry: Principles and structure

(Spanish) (New York: Wiley).Bray, W. C. and Branch, G. E. K. (1913) Valence and tautomerism. Journal of American

Chemical Society, 35, 1440-1447.Brown, T. L. and LeMay, H. E. (1988) Chemistry: The Central Science (4th edn)

(Englewood Cliffs, NJ: Prentice Hall).Brown, T. L., LeMay, H. E. and Bursten, B. E. (1998) Chemistry: The central science (7th

edn, Spanish) (Englewood Cliffs, NJ: Prentice Hall).Burbules, N. and Linn, M. C. (1991) Science education and philosophy of science:

congruence or contradiction? International Journal of Science Education, 13, 227-241.Burns, R. A. (1996) Fundamentals of Chemistry (2nd edn, Spanish) (Englewood Cliffs, NJ:

Prentice Hall).Chang, R. (1981) Chemistry (New York: Random House).Chang, R. (1999) Chemistry (6th edn, Spanish) (New York: McGraw-Hill).Chiappetta, E. L., Sethna, G. H. and Dillman, D. A. (1991) A quantitative analysis of

high school chemistry textbooks for scientific literacy themes and expository learningaids. Journal of Research in Science Teaching, 28, 939-951.

Coppola, B. P. and Daniels, D. S. (1998) Mea culpa: Formal education and the dis-integrated world. Science and Education, 7, 31-48.

Daub, G. W. and Seese, W. S. (1996) Basic Chemistry (7th edn, Spanish) (Englewood Cliffs,NJ: Prentice Hall).

Dickerson, R. E., Gray, H. B. and Haight, G. P. (1970) Chemical principles (New York:W. A. Benjamin).

Duschl, R. A. (1994) Research on the history and philosophy of science. In D. L. Gabel(ed.), Handbook of Research on Science Teaching (New York: Macmillan) 443-465.

Ebbing, D. D. (1997) General chemistry (5th edn, Spanish) (New York: McGraw-Hill).Gabel, D. L. (1993) Use of the particle nature of matter in developing conceptual under-

standing. Journal of Chemical Education, 70, 193-194.Gillespie, R. J. (1997) Reforming the general chemistry textbook. Journal of Chemical

Education, 74, 484-485.Gillespie, R. J., Spencer, J. N. and Moog, R. S. (1996) Part 2. Bonding and molecular

geometry without orbitals—The electron domain model. Journal of ChemicalEducation, 73, 622-627.

Heilbron, J. L. and Kuhn, T. (1969) The genesis of the Bohr atom. Historical Studies inthe Physical Sciences, 1, 211-290.

Hein, M. and Arena, S. (1997) Foundations of College Chemistry (Spanish) (Pacific Grove,CA: Brooks/Cole).

Hodson, D. (1988) Towards a philosophically more valid science curriculum. ScienceEducation, 72, 19-40.

THE ORIGIN AND IMPLICATIONS OF THE COVALENT BOND 639

Dow

nloa

ded

by [

Rye

rson

Uni

vers

ity]

at 0

1:50

29

Apr

il 20

13

Holtzclaw, H. F. and Robinson, W. R. (1988) General Chemistry (8th edn) (Lexington,MA: Heath).

Jensen, W. B. (1984) Abegg, Lewis, Langmuir, and the octet rule. Journal of ChemicalEducation, 61, 191-200.

Joesten, M. D., Johnston, D. O., Netterville, J. T. and Wood, J. L. (1991) World ofChemistry (Philadelphia: Saunders).

Kauffman, G.B. (1989) History in the chemistry curriculum. Interchange, 20, 81-94.Kohler, R. E. (1971) The origin of Lewis’s theory of the shared pair bond. Historical

Studies in the Physical Sciences, 3, 343-376.Kovac, J. (1991) A capstone experience in chemistry. Journal of Chemical Education, 68,

907-910.Kuhn, T. (1970) The Structure of Scientific Revolutions (2nd edn) (Chicago: University of

Chicago Press).Lakatos, I. (1970) Falsification and the methodology of scientific research programmes. In

I. Lakatos and A. Musgrave (eds) Criticism and the Growth of Knowledge (Cambridge,UK: Cambridge University Press) 91-195.

Lewis, G. N. (1913) Valence and tautomerism. Journal of American Chemical Society, 35,1448-1455.

Lewis, G. N. (1916) The atom and the molecule. Journal of American Chemical Society, 38,762-785.

Lewis, G. N. (1919) Letter to Robert Millikan (reproduced in Kohler, 1971).Lewis, G. N. (1923) Valence and the Structure of Atoms and Molecules (New York: Chemical

Catalog Co.).Lin, Q., Kirsch, P. and Turner, R. (1996) Numeric and conceptual understanding of

general chemistry at a minority institution. Journal of Chemical Education, 73,1003-1005.

Lippincott, W. T., Garrett, A. B. and Verhoek, F. H. (1977) Chemistry (3rd edn) (NewYork: Wiley).

Mahaffy, P. G. (1992) Chemistry in context: how is chemistry portrayed in the introduc-tory curriculum? Journal of Chemical Education, 69, 52-56.

Mahan, B. and Myers, R. J. (1990) University Chemistry (4th edn, Spanish) (Wilmington,DE: Addison-Wesley).

Mason, D., Shell, D. F. and Crawley, F. E. (1994) Differences in problem solving bynonscience majors in introductory chemistry on paired algorithmic-conceptual prob-lems. Journal of Research in Science Teaching, 34, 905-923.

Masterton, W. L., Slowinski, E. J. and Stanitski, C. L. (1989) Chemical Principles (6thedn, Spanish) (New York: McGraw-Hill).

Matthews, M. R. (1994) Science Teaching: The Role of History and Philosophy of Science(New York: Routledge).

Matthews, M. R. (1998) In defense of modest goals when teaching about the nature ofscience. Journal of Research in Science Teaching, 35, 161-174.

McComas, W. F., Almazroa, H. and Clough, M. P. (1998) The role and character of thenature of science in science education. Science and Education, 7, 511-532.

Milne, C. (1998) Philosophically correct science stories? Examining the implications ofheroic science stories for school science. Journal of Research in Science Teaching, 35,175-187.

Mortimer, C. E. (1983) Chemistry (5th edn) (Belmont, CA: Wadsworth).Nakhleh, M. and Mitchell, R. C. (1993) Concept learning versus problem solving. Journal

of Chemical Education, 70, 190-192.National Curriculum Council, NCC (1988) Science in the National Curriculum (York,

UK: NCC).National Research Council, NRC (1996) National Science Education Standards

(Washington, DC: National Academy Press).Niaz, M. (1995) Progressive transitions from algorithmic to conceptual understanding in

student ability to solve chemistry problems: a Lakatosian interpretation. ScienceEducation, 79, 19-36.

640 M. NIAZ

Dow

nloa

ded

by [

Rye

rson

Uni

vers

ity]

at 0

1:50

29

Apr

il 20

13

Niaz, M. (1998a) A Lakatosian conceptual change teaching strategy based on student abilityto build models with varying degrees of conceptual understanding of chemical equi-librium. Science and Education, 7, 107-127.

Niaz, M. (1998b) From cathode rays to alpha particles to quantum of action: a rationalreconstruction of structure of the atom and its implications for chemistry textbooks.Science Education, 82, 527-552.

Niaz, M. (1999) Should we put observations first? Journal of Chemical Education, 75,734.

Niaz, M. (2000) A rational reconstruction of the kinetic molecular theory of gases based onhistory and philosophy of science and its implications for chemistry textbooks.Instructional Science, 28, 23-50.

Niaz, M. and Robinson, W. R. (1992) From ‘algorithmic mode’ to ‘conceptual gestalt’ inunderstanding the behavior of gases: An epistemological perspective. Research inScience and Technological Education, 10, 53-64.

Niaz, M. and Robinson, W. R. (1993) Teaching algorithmic problem solving or conceptualunderstanding: role of developmental level, mental capacity, and cognitive style.Journal of Science Education and Technology, 2, 407-416.

Noh, T. and Scharmann, L. C. (1997) Instructional influence of a molecular-level pictorialpresentation of matter on students’ conceptions and problem-solving ability. Journalof Research in Science Teaching, 34, 199-217.

Nurrenbern, S. C. and Pickering, M. (1987) Concept learning versus problem solving: isthere a difference. Journal of Chemical Education, 64, 508-510.

Oxtoby, D. W., Nachtrieb, N. H. and Freeman, W. A. (1990) Chemistry: Science of Change(Philadelphia: Saunders).

Parson, A. L. (1915) A magneton theory of the structure of the atom. SmithsonianMiscellaneous Collections, 65, 1-80.

Pauli, W. (1925) Uber den Zusammenhang des abschlußes der elektronengruppen im atommit der komplexstruktur der spektren. Zeitschrift fur Physik, 31, 765-785.

Pauling, L. (1977) General Chemistry (10th edn, Spanish) (San Francisco: Freeman).Petrucci, R. H. (1977) General Chemistry: Principles and Modern Applications (Spanish)

(New York: Macmillan).Rodebush, W. H. (1928) The electron theory of valence. Chemical Review, 5, 509-531.Sawrey, B. E. (1990) Concept learning versus problem solving: Revisited. Journal of

Chemical Education, 67, 253-254.Segal, B. G. (1989) Chemistry: Experiment and Theory (2nd edn) (New York: Wiley).Siegel, H. (1978) Kuhn and Schwab on science texts and the goals of science education.

Educational Theory, 28, 302-309.Siegel, H. (1979) On the distortion of the history of science in science education. Science

Education, 63, 111-118.Solomon, J. (1991) Teaching about the nature of science in the British National

Curriculum. Science Education, 75, 95-103.Stinner, A. (1992) Science textbooks and science teaching: from logic to evidence. Science

Education, 76, 1-16.Stoker, H. S. (1990) Introduction to Chemical Principles (3rd edn) (New York: Macmillan).Taber, K. S. (1995) The octet rule—a pint in a quart pot? Education in Chemistry, 32, 84.Thomson, J. J. (1897) Cathode rays. Philosophical Magazine, 44, 293-316.Thomson, J. J. (1907) The Corpuscular Theory of Matter (London: Constable).Thomson, J. J. (1914) The forces between atoms and chemical affinity. Philosophical

Magazine, 27, 757-789.Whitten, K. W., Davis, R. E. and Peck, M. L. (1998). General Chemistry (5th. edn,

Spanish) (New York: McGraw-Hill).Wolfe, D. H. (1988) Introduction to College Chemistry (New York: McGraw-Hill).Zumdahl, S. S. (1990) Introductory Chemistry: A Foundation (Lexington, MA: Heath).

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nloa

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