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In his article, Ronald J. Gillespie offered the idea of limiting Introductory Chemistry to six main topics: atom structure, the periodic law, chemical bonding, chemical reactions, chemical thermodynamics, and stereochemistry.

We presume that the introductory chemistry curriculum should be limited by an up-to-date explanation of the main chemical phenomena concerning chemical bonding, chemical reactions, the connection between these phenomena and between the physical and chemical properties and the structure of substances. The rest of the material, including physical and mathematical aspects, should be given only if it is meant to serve as a basis for the structure of a theory of chemical phenomena.

In what cases do we give the students additional material?

1) When it is vital for structuring a theory, say, in atom structure, or to reveal its layer-like structure.
2) When the material is important, during the illustration of the general precepts of a theory, for example, various types of bonds and/or chemical reactions.
3) When systemizing the material; for example, the Mendeleyev periodic table of elements.


With this in mind we will approach our aim.

"A thin book — if such things exist — containing guiding principles, basic knowledge, a thematic approach to general chemistry, would be prepared."  

We can even do as the above-mentioned authors of textbooks do — consider that the given book "is one of the products of the work of the Task Force." At least the time of its issue, the number of subjects touched upon, and most importantly — the data-driven models precede the theoretical discussions.

The types of models that we generally use are actually copies of model types given in the preface of the book Structure and Dynamics mentioned on the previous pages under the numerals of 1, 2, and 3.

Explanations concerning the layer-like structure of the atom and the method of defining the number of electrons in each layer on the basis of defining the ionization energy, the defining of valence on the basis of the data on the structure of stable compounds, The structure of models of chemical bonds on the basis of the comparison of the valence data and the data concerning atoms of the ionization energies, etc. are good examples of data-driven models.

Examples of models that reflect current understanding of chemical theories are well understood and therefore quite teachable as far as the theory of chemical bonding and chemical reactions are concerned.

The appearance of these models is conditioned by the fact that at the end of the twentieth century, according to Lippard (C&EN, Aug. 7, 2000), "as a result of the recent revolutionary changes in chemistry, older paradigms are giving way to new principles."

The paradigm, to the effect that the problem of chemical bonding can be solved only in the framework of quantum mechanics, was exchanged for a paradigm to the effect that this problem could be solved in the framework of the classical phenomenological approach. What is most important in education is to make these explanations really teachable

In 1868 Lotar Mayer announced: "Chemical phenomena must be treated as if they were problems of mechanics." This announcement serves as an epigraph to the chapter Rates and Mechanisms of Chemical Reactions issued by Chemical Principles in 1979. In textbooks there are visual aids for students illustrating drawings

where the interacting molecules approach each other at a great speed (A), then they collide (B) and break up(C).


Experimental facts that have been collected over the last fifty years lead to change in paradigms of chemical kinetics and catalysis. According to the new paradigm, chemical phenomena must be treated as problems relative to the change of the potential energy of the electrons in the outermost atomic shells. A comprehensible phenomenological explanation of chemical kinetics and catalysis phenomena in the framework of the new paradigms are suitable for introductory chemistry.

Gillespie not only indicated the minimal number of topics that introductory chemistry should be limited to, but he also indicated exactly what kind of material should be included in each of the six sections.

Our suggestion about the contents of the introductory chemistry course completely coincides with the main conclusions made by the Task Forces and mentioned in Spencer's and Gillespie's articles: 1) The exclusion of unteachable material; 2) The use of the core approach.

We support Gillespie's suggestion on how the Introductory Chemistry Curriculum should be limited. But generally, the curriculum that we suggest, differs from that suggested by Spencer and Gillespie in as much as our suggestion considers what Lippard called, "the results of the quiet chemical revolution," that occurred at the end of the twentieth century.

In accordance with the results of this revolution, Lippard suggested more than 10 points relative to grants among which were:    "We wish to evolve new theoretical approaches to understand chemical bonding and reactions and to test these theories via real chemical systems."

That is, in the course of the chemical revolution, the problem of chemical bonding and chemical reactions was solved, and the comprehensible methods of conveying these phenomena to the students have been found.

To compare the Gillespie suggestion with the newly offered explanation of the results of the researches carried out in 1990-2002, let's observe in greater detail, the contents of Gillespie's suggestions with those of the herewith-offered ones:

Atoms, Molecules, Ions
Researches conducted in 1990-2000 have shown that, indeed, we must limit the material suggested by Gillespie. The slight changes by which our explanations differ from those of Gillespie are relative to the fact that we use only ionization energy for defining the number of electrons in the outermost layer of the atom, and not the more modern data on photo-electronic spectroscopy. This is because the difference between electrons s, p, and d in the outermost layer, relative to the energy, is small and is obvious during photo-electronic spectroscopy, and is not obvious in the energy or in the chemical properties of bonds formed out of these electrons.
Chemical Bonds;
What Holds Atoms Together in Molecules and Crystals?

Great progress was gained in the interpretation and explanation of the physical essence of chemical bonding in the course of the quiet chemical revolution.

We offer the interpretation of chemical bonding in our manual Introductory Chemistry (Teaching Guide) on two levels that differ in depth of knowledge. Both levels of explanations, as already mentioned, are based upon the knowledge that the students received in their former period of education. The main difference between the old and the new explanations when teaching chemical bonding is that the latter answers the following questions:

1) How does the chemical bond form between the atoms?
2) What is the nature of the forces that bond atoms into molecules and molecules into substances?
3) Which interactions between the nuclei and the electrons are the most important ones?
4) Why does the thermal breaking of a bond require two times more energy than that, which is discharged during bond formation?
5) Why don't noble gases form covalent chemical bonds between themselves?
6) Why is it that the bonds between elements of group II are by one order of magnitude weaker than those between the elements of groups I and III, etc?

The main difference in the new explanation is the fact that it is phenomenological, that is, completely structured on the basis of the analysis and comparison of the experimental data.

Thus, the conclusion about the layer-like structure of the atoms and the number of electrons in the outermost shell of the atoms is made from the experimental data via the ionization energy of the atoms of elements of the 2nd and 3rd periods.     

All the concepts of the theory of chemical bonding are built on the basis of the comparison of experimental data relative to the number of electrons in the outermost shell and the data on the valence of the elements.

This comparison led to the introduction of a model for chemical bonding in a hydrogen molecule. According to this model, a covalent chemical bond is formed of two electrons rotating on a plane perpendicular to the axis and connecting the nuclei. Moreover, both bonding electrons enter the outermost shells of both bonding atoms. All these precepts are not mere declarations; they are based upon the above given comparisons.

In the case of a hydrogen molecule (two protons and two electrons), when presupposing that the usual electrostatic interactions exist between these four particles, the molecule's energy is calculated via an analytic (algebraic) solution. The energy value of a molecule, received in the course of the calculation, differs from the experimental result by less than 3 %, which proves that the defining forces in chemical bond formation are the electrostatic forces.

The attraction of the initial atoms' nuclei to the bonding electrons, situated between them, offers about 10% of the energy gain (the hydrogen atom's affinity to the electron is equal to 0.72 eV). The main electrostatic contribution is the increase of the nuclei's attractions to the electrons at the expense of the increase of the effective charges of the atoms being bonded, which occurs during chemical bond formation.

Equations compiled on the basis of a model allow answering the question: What is the optimal number of electrons necessary to form a bond? Or: Why is it that two electrons take part in bond formation, and not one or three? According to calculations, if one or three electrons took part in bond formation between two neutral atoms in the gas phase, there would be no energy gain, i.e., the molecule would not form.   

Then, as a result of calculations of the model for chemical bonding, it was shown that the bonding energy increases with the increase of the difference between the first ionization energies of the atoms being bonded. The calculated dependence of the bonding energy of one atom with others (dependence of the bonding energy on the first ionization energy of the second atom) coincided with the experimental results. This proved that in the course of the formation of ionic compounds of the NaCl type, there is no electron transition from Na to the halogen. A greater energy gain, in this case, is conditioned by a greater approach to the bonding electrons of the nucleus of an atom with greater first ionization energy. The distance between the nuclei of the Na and Cl atoms and the bonding electrons in the molecules is smaller than the distance between the nuclei and the electrons of the outermost shells in anions Na- and Cl-.

Gillespie's suggestion about limiting this section with the explanation of ionic and covalent bonding on the level of the Lewis rules, offered by the author in 1913, proved to be baseless in 1997.

These rules included such 'facts' as the wish of the atoms to build their shells up to the shells of the inert gases by a) transferring the electrons or b) sharing them among the atoms. Neither supposition had any physical sense. Even worse, the general supposition about the atoms wanting to build their shells up to the shells of the inert gases had an anthropo-physical sense. Lewis introduced this supposition with the aim of uniting the ionic and covalent bonds.

The first calculations of ionic bonding energy which were initiated out of the Lewis rules and the Kossel theory, that is, they presupposed that ionic bonding (for example, in NaCl) is formed at the expense of the attraction of all kinds of ions, as a result of the transition of electrons from sodium (Na) to chlorine (Cl), and differed from the experimentally defined bonding energy by less than 10 %. 

However, in 1970 Pimentel proved that these calculations are incorrect, and that the discrepancy between the calculation and the experimental results equaled to more than 80 %. That is, the Lewis-Kossel supposition concerning the formation of ionic bonds via their scheme, contradicted the calculation that proved the groundlessness of this supposition.

Analogously, considering a calculation as the basis of the electric nature of covalent bonding, as a result of the attraction of the nuclei to the bonding electrons, also led to results that differed from those of the experiments.

In the case of accounting for the hydrogen atoms' affinity energy to the electrons at the expense of the attraction of the nuclei to the electrons, the additional energy comprised 1.4 eV, which amounts to 30% of the bonding energy without considering the inter-nuclear repulsion. Together with the inter-nuclear repulsion, the energy gain at the expense of the nuclei's attraction to the bonding electrons amounts to 10 % of the bonding energy. That is, the readily accepted explanation, offered by Gillespie, argumentatively contradicted the scientific viewpoints on the nature of ionic and covalent bonding of the second half of the twentieth century.

To be more precise, the problem of the physical nature of chemical bonding was not yet solved by science. Therefore, this problem could only be included in the Introductory Chemistry course only as one that was not solved as yet. 

The Lewis Rules are an exception. The electron-deficit (BF3) and the electron-surplus compounds SF4, PF5, XeF4, etc., cause problems during the representation of structures such as SO2, NO2, C6H6, etc., and there are also problems when explaining the structure of stable compounds of the I3, I5, Br3 type, etc.

The existing textbooks, that limit the description of chemical bonding with the Lewis Rules, include exceptions to the Lewis Rules and solutions to the above-mentioned problems. Simultaneously, the structure of electron-deficit compounds is not explained, while the formation of stable electron-surplus compounds of the SF4 type is explained by the participation of d-orbitals in the bond formation. The structure of compounds (like I3) is explained, as a rule, only in monographs. The problem of writing such structures as SO2 is explained in the framework of the resonance theory. In other words, all the explanations are different and they are all quantum-chemical, that is, the students should be taught quantum chemistry, which is regarded as being unteachable.

In the framework of the new explanation of bonding energy such as homo-atomic (H2, Cl2) and hetero-polar (NaCl) both bonding electrons enter the outermost shells of the atoms to be bonded. That is, the number of electrons in the outermost shell of both atoms to be bonded is increased by one electron. Relatively, 8 electrons in the outermost shells of the atoms of the second and third periods is not at all the aim of the atom during stable molecule formation, but rather a limitation of the possibilities of the atom to connect (bond) other atoms. Correspondingly, the anthropo-physical component of the Lewis Rule is excluded.

The electron-deficit compounds are not exceptions to the theory. The theory, besides analyses of data on the valence of stable compounds, and the number of electrons in the outermost electronic layers of atoms of the second and third periods, also includes analyses of experimental data regarding electronic isomerization. These experimental data were received in the second half of the twentieth century.

It was proven experimentally that when one atom is bonded with analogous ones with various kinds of bonds (single, double, donor-acceptor, covalent, Van der Waals), electronic isomers turn into each other with at a great speed thanks to the rapid transition of the electrons. That is, the 50-year old quarrelsome question was solved: the structure of compounds like SO2, C6H6, I3, PCl5, SF4, and XeF4.

This explanation is the one and only one regarding all compounds, and it does not require the knowledge of quantum-chemical hypotheses and theories.

Explanations of exotic structural compounds in the framework of the Lewis Rules allowed declining the quantum-chemical explanations about the structure of these compounds.

The quantum-chemical explanations regarding compounds like SO2, C6H6 included such terms as resonance while quantum-chemical explanations of the structure of surplus compounds (SF4, PCl5) included such terms as d-orbital.

The introduction of the terms resonance and orbital was also required by the theory of valence bonds, which, in textbooks explained the valence rules. Analogously, the introduction of the term orbital was necessary in order to formulate the theory of molecular orbitals.

In order to explain the essence of the word orbital, today's Introductory Level Course of the General Chemistry textbook, introduces a description of the discovery made by Plank, Einstein, De Broglie, Bohr, and Schroedinger.

The authors of the discovery could not understand the physical contents of their discoveries— that is, the given material was regarded as unteachable.

Indeed, A.Einstein wrote: "All these fifty years of cautious brooding have brought me no nearer to the answer to the question:  'What are light quanta?'   Nowadays every Tom, Dick and Harry thinks he knows it, but he is mistaken."

N.Bohr  was known to have made remarks about quantum mechanics to the effect that "If  you think you understand it, that only shows that you don't  know  the  first  thing  about  it."

E.Schroedinger also left his message about quantum mechanics: "... it is not only practically inaccessible, but not even thinkable. Or, to be more precise, we can, of course, think it, but however we think it, it is wrong: not perhaps quite as meaningless as a  triangular circle,  but much  more so than a  winged lion."

"Nobody understands quantum theory," said R.Feynman, the greatest physicist of his generation, in 1980.

The Bohr Theory of hydrogen atom structure, given in textbooks, explains that the specter of the hydrogen atom has only a historic connection with the origin of the term to quantize energy, and further on, no mention is made in textbooks regarding the spectroscopy of atoms.

Thus, the exclusion of quantum-chemical explanations of chemical phenomena allows to: 1) exclude the explanations that are unteachable; and 2) exclude the descriptions of physical discoveries that have no direct cause-and-effect connections with chemical phenomena.

The general character of all the explanations of chemical phenomena in the existing textbooks— first of all, the quantum-chemical explanations ¾ does not leave any doubt in the students that at least such phenomena as wave properties of particles and delocalized electrons actually exist, that there are atoms and molecules in the orbitals where only electrons can be situated, and that these orbitals can hybridize.

The fact that quantum-chemical explanations, given in chemistry textbooks, are actually fabrications of the authors, has been admitted in one of the main textbooks "Quantum Chemistry" that were reissued five times from 1970 to 2004 by Ira Levine. Thus, in the first edition, 1970, on page 559 we read: "There is a tendency to consider 'configuration interaction', 'hybridization', etc. as real physical phenomena. Such concepts are only artifacts of the approximations used in the calculations. Even the concept of orbitals is but an approximation, and strictly speaking, orbitals do not exist."

In the fifth edition (2004) on page 609 we read: "The difference between the energy for the individual structure I and that found when all VB structures are included is the resonance energy of benzene. One says that benzene is "stabilized by resonance," but, of course, resonance is not a real phenomenon."

Such concepts as "configuration interaction," "resonance," "hybridization" and "exchange" are not real physical phenomena, but only artifacts of the approximations used in the calculations. Likewise, the concept of orbitals is but an approximation, and, strictly speaking, orbitals do not exist.

That is, in accordance with quantum chemistry, all the explanations based on such logic of the students as realistic physical phenomena accepted as orbitals, resonance, exchange, hybridization, etc., are related to false knowledge, which, as we all know, is worse than the lack of knowledge.  

        Molecular-shaped Geometry: Tri-Dimensional Chemistry

This is one of the most important sections of knowledge in chemistry. However, its explanation, even in Gillespie's simplified form, is still rather complicated as compared with the other sections of Introductory Chemistry. This is why in our Teaching Guide we have shifted the topic Three-Dimensional Chemical Structure to the third level of explanations and introduced our more simplified form of interpretation, instead of that of Gillespie.

Our interpretations contain cause-and-effect connections with the bonding theory that is given in each corresponding section.

Chemical Reactions

The changes in understanding and explaining chemical reactions in the course of the last chemical revolution took place on the level of changes in the paradigms. It was found that:

1) The interaction of saturated molecules proceeds, as a rule, not along the molecular mechanism, but through an intermediate formation of chemically activated particles (ions, radicals, conences, etc.). Active particles are formed of initial molecules in the course of dissociation that occurs as a result of thermal, photochemical, and other methods of increasing the energy, or, for example, dissolving and dissociating initial molecules in a solvent.

Previously, in the times of the old paradigm, it was supposed that a direct molecular interaction is a general mechanism for chemical reactions, while reactions via intermediate active particles were regarded as special cases.

2) The active particle interaction with saturated molecules proceeds not via the transition state (TS) situated at the top of the potential barrier, but via the intermediate compound (IC) situated in the minimum.

3) In the general case that lay at the bottom of the old paradigm, it was supposed that the interaction of molecules between themselves, and active particles with molecules, proceed in one stage. According to the new theory (i.e., new paradigm) the interaction of active particles with molecules proceeds in three stages: association, electronic isomerization, and dissociation.

4) The main driving force of chemical reactions, according to the old paradigm, was the kinetic energy of transiting initial molecules [theory of active collisions (TAC) and transition state theory (TST)], which is spent on breaking the old chemical bonds during the reaction. Here it was not clear exactly how, in the course of the chemical reactions that proceed at temperatures of 30° - 200° C, most of the chemical bonds break, though they require over 2,000° C thermal breaking.

According to the new paradigm, during the interaction of an active particle with a saturated molecule proceeding via the above indicated three stages, as a result of the electronic isomerization reaction, the strong old covalent bond turns into a weak Van der Waals bond (VWB) without any energy expenditure or energy discharge during isomerization. Thermal energy is spent mainly on the formation of extremely small (>0.01 %) active particles. Therefore, the visual pictures of the colliding saturated molecules, as a result of mechanical transition, commonly given in textbooks, are incorrect as far as the modern viewpoint is concerned.

5) In the framework of the mechanical approach, when describing the factors that define the reaction speed, only the influence of the temperature and the initial substance concentration were described. 

The question concerning the influence of nature of the reacting substances on the speed reaction was actually never raised.

In spite of this, according to experimental data, the chemical reaction speed is defined, first of all, by the chemical nature of the reacting substances. Thus, for example, radicals interact with saturated molecules by more than 10,000 times faster as compared with the interaction speed of the saturated molecules between themselves.

In the framework of the new approach, this point was not only mentioned as a main point, we managed to understand and explain the influence of the nature of the reacting substances on the speed of the chemical reactions.

6) In the framework of the old interpretation, the general explanation of catalysis turned out to be a 'black hole'. From 5 to 10 pages of the 800-page textbook are devoted to catalysis, though more than 90% of the chemical reactions are catalytic.

While in the framework of the new interpretation of the chemical reaction, the catalysis phenomenon becomes comprehensible and explainable at lessons of Introductory Chemistry relative to both general and special cases of catalytic reactions. Just as the theory of chemical bonding, the theory of chemical reactions is phenomenological.

The theory of chemical reactions is built on the basis of the analysis and comparison of experimental data via the interaction speeds of the active particles (ions, radicals, Lewis acids, Lewis bases, etc.) with saturated molecules.

The mechanism for the interaction of active particles with saturated molecules is based on experimental data relative to the electronic isomerization received from the biological, inorganic and organic chemistry.   

Energy and Entropy

Just like Spencer, Gillespie considers the interpretation of thermodynamics unteachable. He suggests that 'entropy' should be explained as a degree of disorder.

Gillespie offered to include this section in the Introductory Chemistry course because the energy and entropy allow answering the questions: 1) Why do some reactions take place and others do not?  2) Why do some reactions proceed to the end and others do not?          

In the course of the chemical revolution, it was found that reactions proceed when there are chemically active particles in the system, but not at all saturated molecules with a greater energy than that of the products received from the latter. That is, after answering the question how chemical reactions proceed, the answer to the fist question in the framework of thermodynamics is incorrect.

The second question: Why do some reactions proceed to the end and others do not? The answer to this question is obvious and teachable if we study not only the direct reactions, but also the reverse ones.

In Introductory Chemistry, this question, as well as the question about chemical kinetics, is of the same rank in chemistry, but this rank greatly prevails over the main questions included in Introductory Chemistry.

That is, this question, just as the question concerning chemical kinetics (reactions of various orders of magnitude, mechanisms and kinetics of various reactions) can be included in Physical Chemistry for Majors.

The inclusion of explanations of entropy as a degree of disorder in Introductory Chemistry, we think is a contentious point because of the following:

1) Out of the two explanations about entropy being discussed in literature — thermo-dynamic and statistic (Boltzmann) — the first is unteachable, while the second is mathematical, since it has no physical contents.     

2) The introduction of the notion 'entropy' (without the introduction of the notion 'free energy' and the equation ΔG = ΔH - TΔS) has no practical sense. According to Spencer, 'free energy' is unteachable in the framework of Introductory Chemistry; and according to Gillespie, 'thermodynamics' is a forbidding term for students, and besides, the equations on which formal thermodynamics is based, are even more forbidding.

The question concerning 'free energy' relative to chemistry (but not Introductory Chemistry) is a generalized (i.e., philosophical) question. After realizing the details of the formation and breaking of the chemical bond, the unproductive expenditure of energy during the thermal breaking of the bond that causes the decrease of the potential energy of the electrons (the increase of the distance between the bonding electrons and the nuclei) in the unbroken molecules, all the chemical questions can be explained without the introduction of the notions of 'free energy' and 'entropy'. 

What should be inserted in Introductory Chemistry besides the above mentioned by Gillespie and Spencer?

This is a question about the correlation of the macroscopic world of observations and microscopic world of atoms and molecules.           

A more complete answer to this question was received in the course of the chemical revolution. We described the connections between the parameters of atoms and molecules, and the physical and chemical properties of substances formed out of these atoms and molecules in the section 'Physical and Chemical Properties of Substances'.

In 1998 we have issued the book How Chemical Bonds Form and Chemical Reactions Proceed.  This book contains the results of all our work conducted in the years of 1982 to 1998. In the course of our work, we managed to answer the main questions dealing with chemical bonding and chemical reactions. That is, we have managed to get to the bottom of the physical essence of these phenomena and the cause-and- effect relations between the electronic structure of the atoms, chemical bonding, and chemical reactions.

Proper comprehension of the physical essence of phenomena, and the cause-and-effect relations between them was gained in the framework of the accepted classical (pre-quantum) phenomenological approach, which presupposed the possibility of a distinct (clear), but not quantum-chemical interpretations of such phenomena to the students

In 1998 we began work on structuring  General Chemistry, which will assist the chemistry teacher. Some of the chapters, including those concerning chemical bonding and chemical reactions were reported at conferences at the ACS twice a year (1999 to 2006 - more than 25 reports). The complete reports were published at our site.

The questions and discussions at the conference have shown that for 6 years our explanations concerning chemical bonding, kinetics, and catalysis had passed two of the three stages.

It is well-known that before a new concept takes the place of an old one, it must undergo three stages in the social and scientific conscience before it gets its common recognition.

The stages are as follows:
1)  That is impossible!
2) M-m-m ---- I guess you've got something there!
3)  Recognition of a theory by society. Such recognition  usually happens under such slogans as:This is a very well-known concept. This is the well-forgotten past!

The discussions of the reports at ACS conferences, the reviews on the book and on our site have shown that our interpretations of these chemical phenomena are comprehensible to a wide range of readers. As an example, here is an e-mail to us from a reader:

"I'm a research civil engineer and left chemistry behind when I was 18. At that time [1950] catalysis was presented as something of a black art. However, recently I needed to understand the phenomena of catalysis. So having tried various phrases in Google with the hope of finding a 101 explanation, I eventually hit upon "mechanism of catalysis" which turned up on your site.  Congratulations on a most interesting and clear exposition, which, even with my intermediate B.Sc. level of chemistry (and a modicum of subsequent treading), I had no difficulty in understanding and accepting as entirely logical.  

            Good luck in your future endeavors.  Frank Grimer"

In conclusion of this introduction, we would like to thank Professors J.N.Spencer and R.J.Gillespie for their articles, which stimulated our work in the realm of chemical education, and, what is most important — their efforts have defined the direction of further action including the contents of a new chemistry textbook Introductory Chemistry.

We are thankful to all the authors who took part in the thirty-year discussion, and especially to Professor S.J. Hawkes who compiled a compendium of proposals for changes in the Introductory Chemistry Curriculum and articles devoted to relative questions.

We are also thankful to the members of the Task Force for their contribution.

Chapter 1.Preface>>
James Spencer's viewpoint  >>
Ronald J. Gillespie's viewpoint >>   
Our viewpoint

Chapter 2. Aim of chemistry as a science >>