Book Reading Update - Fontana History of Chemistry Chap 15

Chapter 15: The Renaissance of Inorganic Chemistry

The author gives this chapter the same title as Ronald Nyholm's 1956 book, The Renaissance of Inorganic Chemistry, in which he defines inorganic chemistry as “the integrated study of the formation, composition, structure, and reaction of the chemical elements and their compounds, excepting most of those carbon”.

The author begins this detailed account of renewed vigor in inorganic chemistry research with a detailed summary of Alfred Werner's extensive development of coordination chemistry in the section titled "Werner’s New Ideas".  Alfred Werner, whose name is considered synonymous with coordination chemistry, did his first studies on stereoisomerism.  Frankland and Kekule’s theory of valency provided an explanation for the structures of carbon compounds and simpler inorganic compounds but there was still debate on whether elements have fixed valencies.  This was to loom largely in Werner's analysis of structure and bonding in coordination complexes. The author summarizes Werner's work on coordination chemistry with the following three comments:
1) Werner’s coordination of water directly to the metal “not only solved the problem of water of crystallization, but formed a bridge between the hydrate theories of Mendeleev and others and the new dissociation theory of Arrhenius.  Conductivity and dissociation were dependent on the nature of solvents because only those solvents which can combine with metal salts to form coordination radicals are able to allow electrolytic conduction”.  [I think what this means is that conductivity, which, itself, is dependent on the extent of dissociation, depends on the solvent because dissociation can only occur if solvent molecules are able to coordinate with the metal cation.]
2) It stimulated thinking on the spatial representation of complexes that have 6 or 4 particles surrounding the metal cation and consideration of stereochemistry (e.g. geometric isomerism) by inorganic chemists.  An insightful finding cited was the observation that because the platinum compound Pt(NH3)2Cl2 has no optical activity, it must have a square planar arrangement with two geometric isomers as is now taught.
3) Werner’s ideas on coordination complexes ‘demanded’ a “revolution in valence theory” for “how could a divalent copper atom bond six ammonia radicals to itself”?    In answer to this type of question, Werner mused, “It would also appear that the amount of residual charge – of surplus affinity – possessed by a radical after combination with others depends both on its own nature and that of the radical or radicals with which it becomes associated”.  “It was, he suggested, as o the metal were a positive sphere surrounded by a neutral shell of water or ammonia whose inner surface was rendered negative, and consequently, its outer surface positive.”  Werner further distinguished the valency of an atom from its “coordination number”:  the coordination number refers to the number of groups combining with a metal to form a complex and the valency the number of monovalent atoms that can bound directly to the atom.

Along with Werner's foundational discoveries and theories at the end of the 19th century, Nyholm and other coordination chemists attribute the renewed interest and research in inorganic chemistry to Sidgqwick's work and 'interpretation of coordination chemistry' in the 1920's.  In 1927, Nevil Vincent Sidgwick published “an influential reworking of Werner’s theory in terms of Bohr’s theory of spectra and atomic constitution and Lewis’ view on valency and chemical combination”.   In this text, he indentified three different ways that atoms of the elements link together in compounds and complexes:  polar or ionizable bonds between oppositely charged ions; nonpolar, covalent, shared electron pair of Lewis and Langmuir, and coordinate linkages of Werner which he also recognized as covalent, using the term ‘covalent link’ for the case in which a ligand donates a shared electron pair to the metal acceptor atom and defining the coordination number as the number of shared electron pairs donated.  In reconciling Werner’s coordination theory with traditional structural theory, Sidgwick pointed to Pauling’s statement that “most coordination compounds were resonance hybrids” (like sulfuric and nitric acid are) pointed to a solution involving a more complete theory on valency that recognizes a third type of linkage “a covalency in which both electrons are supplied by the same atom".  Mathematical analysis ultimately led to the determination of a ‘crystal-field stabilization energy’ arising from filling in d-orbitals with relatively lowered energy due to decreased electrostatic repulsion from the incoming ligand [the term crystal pertains to the first use of the analysis in crystallography].  Further analysis using both crystal theory and molecular orbital theory eventually led to the formulation of the ligand field theory now taught in General College Chemistry.

The author included sections on Australian and Japanese Chemistry pointing out that in both these countries, although culturally dissimilar, chemistry started out with a more utilitarian focus.  The disciplines of chemistry and physics were institutionalized in both countries in the 1870’s.  Many of the early pioneers were educated in Europe and, in Australia in addition, some were British.

In 1918, the Australian Chemical Institute was founded but due to long distances, there was not much interaction between the regional chapters.  Even as more chemistry research departments started to be developed, most of the senior posts went to British Chemists. The development of chemistry was further slowed by the difficulty of receiving communications from Europe.  By 1975, however, the first Australian chemist won the Nobel Prize for his work on the stereochemistry of enzyme-catalyzed reactions.

In Japan, between 1888 and 1930, doctorates in engineering and medicine outnumbered those in chemistry for which only 25% had a basic research focus.  The lack of a Japanese chemical vocabulary made lecturing in Japanese difficult.  Bureaucracy made it challenging for Japanese chemists showing academic promise to pursue further training and a pedagogical tradition of teaching as information transfer did not encourage innovation.  Furthermore, “a cultural hierarchical  fear of challenging an older peer or teacher also inhibited criticism and novelty and discouraged cooperate research”.  While the Tokyo Chemical Society was founded in 1878 by foreign teachers followed by a journal publication starting in 1880, it was not until 1972 that an English-language journal was first published.  The author featured some early leaders in Japanese chemistry who were educated in Europe. One of these was Ikeda Kibunae known for his discovery and patenting in 1908 of a method of producing monosodium glutamate from kelp which became the basis for Japan’s first major chemical industry.  Demand for trained engineers, managers, and skilled factory workers led to the expansion of the university system in the 1920’s (like in America and Australia).  Interest in fuels chemistry during the Second World War led to quantum mechanical studies of unsaturated hydrocarbons.  In 1950, Kenichi Fukui developed the frontier orbital theory of reactions that states that “the progress of reactions depends upon the geometry and relative energies of the highest recipient molecular orbital of one reactant and the lowest molecular orbital of the other”.  In 1981, Fukui and Roald Hoffman won the Nobel Prize for this work.

Coordination chemistry was particularly highlighted by the author in the further development of chemistry in Australia.  Eustace Turner’s work (he was British) and collaboration with an Australian colleague on the characterization of an arsenic coordination compound in 1921 put “Australian chemistry firmly on the world map of chemistry” and had been the ‘catalyst for Australia’s future reputation in coordination chemistry”.  [They also prepared an arsenic analogue of indole which they were going to call “arsole” which did not escape the censorious eye of the editor and was then changed to arsindole]  Burrows, with his students, further worked on synthesizing arsenic complexes of zinc, cadmium, mercury, and platinum.  David Mellor of Tasmania used x-ray crystallography to study structure of coordination compounds, including the coordination chemistry of palladium, and showed “how the magnetic properties of transition metals could be used to clarify their complicated stereochemistry”.

The author devotes the last section of the chapter on Nyholm’s Renaissance, owing to Nyholm's contribution to coordination chemistry.  Ronald Nyholm, was born in Broken Hill, Australia, a mining town whose streets were given chemical names, completed his Ph.D. with Ingold as supervisor.  He worked with Dwyer (Sydney Technical College)and  published papers on the complexes of Rhodium. Nyholm’s work showed that certain ligands can induce what historically has been thought of unusual valency states (e.g. nckel(IV)) determined from measurements of magnetic moments.  He was also able to induce higher coordination numbers (e.g. Mo(VII) and Ti(VIII)) using unusual ligands.  Working with Gillespie, they were able to rationalize the small deviations from predicted bonding angles in regular geometries (linear, trigonal, tetrahedron, etc.) by pointing out that electrons with similar spins are repelled more than those with opposite spins and that bonding pair-bonding pair repulsion is weaker than bonding pair-lone pair repulsion.  Becuase of his organization skills, Nyholm was able to make predictive use of patterns, anomalies, and gaps he discovered within the transition metal groups by “comparing, contrasting, and correlating the properties of transition metal complexes horizontally and vertically within their periodic table positions”.  An example of this is the deduction and later experimental preparation of transition metals that can form metal-metal bonding.  As the author outs it, “Like Ingold for organic chemistry, Nyholm perceived that inorganic chemistry would benefit from the use of large-scale instrumentation for mass spectrometry and spectrophotometry.  Indeed, it was Nyholm’s contention that ‘the impact of quantum mechanics and of modern physical methods of attack are the main reasons for the renaissance of inorganic chemistry, leading to the present period of rapid growth”.  In the realm of chemical education, Nyholm advocated for an undergraduate teaching curriculum that involves the integrated use of tools of chemical investigations:  “I believe that we should find more time to enable students to acquire these techniques of inorganic chemistry, as for instance, the handling of substances in the absence of air or moisture, the manipulation of gaseous substances, and reactions involving low or high temperatures. Finally, I am convinced that, in keeping with the new sense of purpose in inorganic chemistry, the maximum opportunity should be provided for the undergraduate to prepare a compound, to establish its purity by analysis, and to investigate as many of its properties by chemical and physical techniques as he [sic] is able to do. This means that we effectively illustrate the wholeness of modern chemistry, and I believe that we thereby develop a genuine enthusiasm for this subject at undergraduate level.”  In 1963, Nyholm encouraged further progress in chemistry teaching and research by recognizing that “chemistry’s boundaries were dissolving and overlapping with interesting areas of physics, biology, geology, and mathematics”.  He recognized the necessity of sharing expensive instrumentation and that undergraduates should be able to use them too.

Author's chapter conclusion in his own words on the Renaissance of Inorganic Chemistry: “The central place of inorganic chemistry is now unquestionable. The synthesis of extraordinary structures and insights into unexpected mechanisms continues. organometallic chemistry was, for example, transformed with the synthesis of the first sandwich compound, ferrocene, in 1957. The power of transition complexes, especially those of platinum, to catalyse important polymerization reactions has similarly transformed industrial chemistry since 1950. Far from being outplayed, inorganic chemistry in the second half of the twentieth century has proved an essential component of the understanding of biochemistry, analytical chemistry, catalysis, electrochemistry, mineralogy, crystallography, radiochemistry and in all industrial processes involving high temperatures, catalysis and semiconductors.”

Book Reading Update - Fontana History of Chemistry Chap 14

The title of this chapter, Structure and Mechanism in Organic Chemistry, is taken from the text written by Christopher Kelk Ingold, described by the author as the great ‘systematizer’ of 20th century organic chemistry and “the chemist who firmly established the importance of physical chemistry in understanding the subject”. The author defines what the phrase structure and mechanism encompasses in organic chemistry: “The phrase ‘structure and mechanism’ referred not only back to the classical chemical theory that the properties of organic compounds could be explained solely in terms of the behaviour of tetravalent carbon atoms, but to the integration of the whole of organic chemistry in terms of the physical understanding of the causes, mechanisms and effects of basic reaction types such as additions, eliminations and rearrangements. Emphasis was to be placed upon the class of reactions rather than on the reactions of individual compounds.”

The chapter is subdivided into sections detailing the evolution of thinking on structure - mechanism dominated by two factions:  Lapworth-Thiele-Robinson and Michael – Flurscheim – Vorlander.  The next section then summarizes salient components of these two schools of thought into a generalized electronic theory of organic reactions and structure, ascribing to Ingold much of the credit for providing a "coherent explanation of many chemical reactions outside of free radicals and certain rearrangements".  In the next section, the author describes the important contributions of kinetic studies to discover, elucidate, or confirm mechanisms of reactions.  This collaboration between what were traditionally considered tools of and concepts of physical chemistry and organic reaction analysis led to the recognition of the new field of organic chemistry.  The spread and progress of physical organic chemistry was not extensive i the beginning nor was it smooth with the author alluding to the not-so-subtle rift between practitioners and theorists of the two fields as a possible impediment.  The author then gives an account of  further attempts at refining chemists' understanding of aromaticity in view of new principles of structure and mechanism.  In the last section, attempts to describe and explain in precise language and conceptualization the formation of "ionic" intermediates ("with internal electrostatic arrangement") launched the chemists into an intense "non-classical ion debate" which contributed to further rehashing and, perhaps, refinement and extension of valence theory.

In conclusion, the author states,
"The Ingold revolution was to look for patterns in reactivity and to explain why reactions occurred in the way they did by relating reactivity to molecular structure. This mechanistic approach, like the periodic table in inorganic chemistry, simplified the task of the learner – but at the cost of demanding much more competence in mathematics than was had by Victorian chemists. Even so, by dint of specialization in quantum mechanics, chemists like Ingold and his followers were able to rely a good deal still on geometrical reasoning."


Some notes:

The Lapworth-Thiele-Robinson Tradition
Lapworth made one of the earlier attempts at classifying reactions on the basis of mechanisms in 1898.  According to this classification, tautomeric reactions involve the movement of a univalent atom or group while desmotropic reactions involve two labile atoms of groups exchanging positions through a Williamson-Kekule intermediate.  Lapworth’s insight on the key property of labile group was that its reactivity stems from its electropositive or electronegative character (ionic).  This was supported by experiments in which salts of predicted ionic intermediate were isolated, “building up evidence that ions played a crucial role in reaction mechanisms”.  He formulated a general theory of ‘alternative polarities’ in which atoms in a molecule are attributed a latent polarity that can be activated by a key atom.  The latent polarity can be activated at a distance (“action at a distance in chemistry”) wherein the reactivity on one site can be induced by the activation of another site (“reactivity at one site influenced or induced action at another”).  This theory of inductive effect was “fully enunciated” by Lapworth in 1920.

Johannes Thiele in 1899 developed the theory of residual valency after observing that addition of hydrogen atoms to unsaturated systems caused a shift in the positions of the double bonds, “as if not all the combining power of the carbon atoms at double bonds had been used”.  The author compares Thiele’s and Robinson’s take on the explanation for this observation:  “Whereas Thiele saw unsaturated atoms as possessing residual combining powers in addition to their normal bonding powers or valency value, Robinson preferred to think of the normal bond as splitting into two (or even fractional) partial valencies.”

The Michael – Flurscheim – Vorlander Tradition
Michael explained Markovnikov’s rule through polarization effects within a molecule that cause some carbon atoms to be more positive than others.  H would normally attach to the more negative carbon atom which is typically the C in the C=C that has more H. (Lapworth explained this by invoking the polar inductive effect of the incoming halide atom.) Bernard Flurscheim believed that any carbon chain would have strong and weak links and “if two primary groups in a molecule had a strong affinity for one another, any addenda would be weakly bonded”.  He explained benzene substitutions as the transmission within the cyclic chain of affinity demand resulting in changes in affinity.  (His ideas were considered a hybrid of Thiele’s and Michael’s and Werner’s “notion of affinity over and above the normal valency value.)  Daniel Vorlander’s explanation for substitution patterns in the benzene molecule pertains to different electrical charges arising between positive hydrogen atoms and negative nuclear atoms of the carbon chain.  He then argued that the rate of substitution depended on the “relative electrochemical contrast between the incoming substituent and the C atom at the ortho, meta, and para positions.

In the author’s words, similarities are clear between all these theories of “residual affinity, alternative polarities, alterations of affinity intensity, and electrical contrasts”.  “By 1920, therefore, organic chemists had a number of competing explanations for the mechanisms of reactions.  Each was ad hoc and ultimately dependent upon there being a satisfactory explanation for the induction, or cause, of a polar, or stronger or weaker affinity.”

The Electronic Theory of Organic Reactions
Upon learning of Lewis’ shared electron pair, Robinson recognized that what he and Lapworth have been referring to as “saturated valency was a shared electron pair, a latent valency a free pair, and a virtual valency an incomplete octet”.  In 1922, Robinson and Kermack published “An explanation of the property of induced polarity of atoms and an interpretation of the theory of partial valencies on an electronic basis” explaining that the basis of the ‘alternating effect’ is the “facile displacement of electrons in unsaturated systems”.  In this paper, Robinson had recognized the “fact that unsaturated atoms share more electrons in common than saturated atoms” and “there will be a greater mobility of electrons”.  Kermak and Robinson started the practice of using curved arrows (“the most important symbol in 20th century organic chemistry) to indicate the direction of electron mobility within bonded atoms.

Ingold presented experimental evidence (based on results of selective o m, p-substitution reactions) supporting Flurscheim’s prediction and contradicting Robinson’s and Lapworth’s prediction.  These were later found to be anomalous.  In 1925 Robinson presented a “coherent” and clarified version of his theory of electronic mechanisms: “ He now suggested that there were two different electronic mechanisms at work in aromatic and conjugative systems: changes in the ‘covalency functioning of electrons’ (the transfer of electron pairs), which he and Kermack had earlier represented by curved arrows, and another effect (which went some way towards a compromise with Flürscheim and Ingold) ‘due to electrostatic induction, the general effect requiring no changes in covalency’.”  Robinson would later accuse Ingold of having “appropriated the electrostatic and covalency shift mechanisms”.

Organizing the Structure of Organic Chemistry
Ingold’s writings in the Annual Reports published by the Chemical Society between 1924 and 1928 helped “strengthen the vocabulary of the new electronic theory of organic chemistry”.  Ingold’s final major generalization on tautomerism articulated in 1933: “the tautomeric effect represented the chemists’ best attempt to represent, or to capture on paper, the elusive structure of  a compound”, renaming the tautomeric effect the ‘mesomeric’ effect (in between) recognizing the limitations of classical structural chemistry.   In 1928, Ingold coined the terms ‘nucleophilic’ (electron pair donor) and ‘electrophilic’ (electron pair acceptor), leading to a generalized theory of aromatic substitution: “A nucleophilic (electron-releasing) group like alkyl reduced positive polarization in an aromatic molecule and encouraged electrophilic substituents into the ortho and para positions because of alternating polarity in the ring.  Conversely, an electrophilic group, with electron-withdrawing tendencies, like CCl3, strongly retarded ortho – para substitution, but permitted meta.”  Ingold’s principles provided a coherent explanation of many chemical reactions outside of free radicals and certain rearrangements.

Frank Whitmore presented a paper in 1932 proposing that all reactions proceeded through an electron deficient intermediate stabilized by 1) bonding to a nucleophile in the solvent, 2) losing a proton or other electrophilic group to form an unsaturated compound, and 3) a group or atom that is more electrophilic than another group, the less electrophilic group could cause an attached group to migrate with its electrons to form a new positive ion which is then stabilized by mechanisms 1 and 2.  This new electronic theory of reaction mechanisms assert that “electron-withdrawing groups would slow down aromatic substitution”.

The Kinetics of Mechanism
Some key events in the development of kinetics:
The basic theory of reaction kinetics was developed by van’t Hoff in 1884 (Etudes de dynamique).  In this publication, he classified reactions in terms of their molecularity (uni-, bi-, ter-, etc.) depending on how many concentrations are involved in the direct proportionality (linear, square or product of 2 concentrations, etc.). Ostwald distinguished between stoichiometric molecularity (number of reactants involved) and order of reaction (mathematical dependence on the power of the concentration or the product of concentrations).  Ferdinand Wilhelmy in 1850 concluded based on results of experiments on sugar inversion that rate of inversion is directly proportional to concentration.
Berthelot and Leon Saint-Gilles in 1862 showed that the rate of reaction between acetic acid and ethyl alcohol was proportion to the concentrations of both reactants and the reaction reaches equilibrium after a certain amount of time.

Hughes studied the kinetics of the halogenations of organic acids and ketones and, using van’t Hoff’s differential method to determine rate order, found that the rate was independent of the halogen concentration.  By 1935, based on kinetic studies that shed light on the possible mechanism, Hughes and Ingold have identified 4 different types of substitution and elimination reactions (SN2, SN1, E2, and E1, with each number indicating the number of molecules involved in the rate-limiting step), “codifying a great deal of organic chemistry”.  In addition, they studied the effects of solvent polarity, steric hindrance, catalysts, salts, and stereochemistry on rate-mechanism.  Their analysis was complemented by data from measurements of dipole moments, and dissociation constants; Hughes’ use of H isotope labeling; and thermodynamics to understand the activated state, all becoming standard tools for the study of organic reaction mechanisms. Not everyone was receptive and Ingold and Hughes received challenges to and criticisms of their kinetic method for determining mechanism.

The Spread of Physical Organic Chemistry
The term physical organic chemistry was first mentioned in a textbook authored by American Chemist Louis P. Hammett in 1940 (in his text Physical Organic Chemistry, Reaction Rates, Equilibria, and Mechanisms) In the preface to his text, Hammett “noted how it had been almost a point of honor with both physical and organic chemists to profess ignorance of the other’s fields [An undercurrent of this slight rivalry still existed when I went to graduate school where my research topic and method was considered as physical organic chemistry].  Monographs on the Physical Aspects of Organic Chemistry and Hammett’s text provided a physical chemistry approach to organic chemistry in British universities by the early 1940’s.  “Although some American organic chemists were happily using physical techniques such as polarimetry and electrolysis in the 1920’s, there was a general scorning of thermodynamics, kinetics, and the new quantum mechanics within the organic chemistry community.  No doubt there was a kernel of truth in the reputed physical chemists’ view of organic chemists as ‘grubby artisans engaged in an unsystematic search for new compounds’.  This was to change in the late 1930’s as Ingold’s mechanistic viewpoint infiltrated textbooks and research programmes.”  The author then went on to enumerate the relevant personalities in various universities and research centers that played significant roles in ensuring the sustainability and progress of physical organic chemistry in America that led to “eventual domination by Americans of the subject after 1945”.

Aromaticity
After Kekule deduced the structure of benzene, much research attention was given to discovering rules for predicting and theories for explaining substituent orientation in substitution reactions of benzene and its derivatives.  “The advent of electronic theories, of course, allowed Robinson, Ingold, and others to explain the rule and its defects in terms of the inductive and electromeric effects and, more fundamentally, aromaticity itself as a reflection of mesomerism.”  In 1925, Robinson “ascribed the uniquely ‘aromatic’ properties to benzene and its analogues to the six extra electrons that produced the ‘stable association which is responsible for the aromatic character’”.  In 1931, Huckel applied quantum mechanical calculations to show that electron density in pi orbitals above and below the plane of the hexagon provided the “mesomeric, or resonance, energy that gave the molecule its extra stability by not containing its electrons to simple alternating single and double bonds”.  Huckel also established what is now known as Huckel rule to predict aromatic properties in molecules that have 4n+2 (where n = 0, 1, 2, 3,…) pi electrons in a closed ring.  Huckel’s rule using concepts of pi electrons, orbitals, and bonding and understanding and explanation of excited states in spectroscopy reinforced the advantages of molecular orbital theory over Pauling’s valence bond theory.  In 1935, Kathleen Lonsdale confirmed the planar hexagonal structure of benzene using x-ray crystallography.
• By 1950, “aromaticity was understood as a property of any ‘cyclic compound with a large resonance energy where all its annular atoms take part in a single conjugated system’”.

The Non-Classical Ion Debate
• Michael J. S. Dewar published The Electronic Theory of Organic Chemistry in 1949.
• In 1969, Dewar published The Molecular Orbital Theory of Organic Chemistry in which he noted that any organic chemist who did not understand MO theory would be left “high and dry”.
• By the 1960’s, the inductive effect can be explained by MO theory as being due to the formation of polar sigma bond between carbon atoms in a chain and various substituent groups.  “The formation of a positive carbonium ion, or a negative carbanion, brought about the further polarization of pi orbitals, which interacted in turn with the pi electron orbitals of substituents to produce what Ingold had called the electromeric effect.”  Dewar preferred to call the intermediate a pi complex consisting of a transient covalent pi bonds wherein a single pair of electrons is shared between three nuclei.
• “Non-classical ion” referred to an intermediate formed containing an “internal electrostatic attachment”.   Dewar felt that this and other terms (bridges, synartetic ions…) were confusing and that the phenomena can more precisely be described using pi terminology and symbolism.  (In many cases, however, sigma bonding was involved and not pi).
• In the mid-1960’s NMR analysis was able to confirm the presence of non-classical carbocations in a concentrated solution prepared by George Olah using magic acid (intensely strong).
• The continuing debate on the structure and characteristic of these non-classical intermediates raged on but continued to contribute to the expansion and refinement of valence theory.

Conclusion
• The Ingold revolution was to look for patterns in reactivity and to explain why reactions occurred in the way they did by relating reactivity to molecular structure. This mechanistic approach, like the periodic table in inorganic chemistry, simplified the task of the learner – but at the cost of demanding much more competence in mathematics than was had by Victorian chemists. Even so, by dint of specialization in quantum mechanics, chemists like Ingold and his followers were able to rely a good deal still on geometrical reasoning.

Book Reading Update - Fontana History of Chemistry Chapter 13

The title of this chapter is the title of Linus Pauling's seminal text "The Nature of the Chemical Bond".  In 1931, Linus Pauling published the book based on a series of seven papers written for the Journal of the American Chemical Society and Journal of Chemical Physics “which was to transform the teaching of chemistry”.  The author notes the significance of this book:  “One of the undoubted reasons for the book’s success was that Pauling went out of his way to appease chemists’ fears that quantum mechanics was too difficult for them by emphasizing the book’s solid grounding in empirical data”.

The author begins this account of the development of bond theory with a description of the development of the "Lewis atom"  which was really an account of the early electronic theory of bonding developed by G. N Lewis from a series of observations and studies starting from Helmholtz claim of the constant charge of an atom and the electrical nature of chemical affinity based on Faraday's electrolytic work.  A partial source of this charge was determined from the 1897 discovery of the electron by J. J. Thomson in Cavendish Laboratory where he detected a negatively charged particle lighter than a hydrogen atom when carrying out experiments on ‘conduction of electricity through gases at low pressures’.  From these results, Thomson developed the “plum pudding model of negatively charged electrons being embedded in a sphere of positive charge”.  Chemical bonds were then interpreted as stemming from the transfer of electrons from an outer circle of electrons to the adjacent atom.  In 1904, Thomson developed the first electronic theory of valency: “the chemical bond was nothing more or less than simple electrostatic attraction.  A bond was formed when two atoms exchanged, or transferred, one or more electrons, the donor thereby becoming positively charged and the receiver, negatively charged”.  As early as 1902, Lewis had developed ideas of bonding (e.g. ‘cubic atoms’) based on observations of the stability of the rare gases (noble) being attributed to the presence of 8 electrons in the outermost sphere of Thomson’s plum pudding model.  In 1904, R. Abegg, formally recognized a ‘rule of eight’, pointing to a pattern in the formation  of compounds whose valencies (I think he meant here the ‘charge’ formed) of the individual atoms add up to 8, e.g. see table in book.  In 1910, Falk and Nelson published a paper entitled “The electron conception of valence” supporting Thomson’s electronic theory of valency in which bonds were formed by the transfer of electrons producing positively and negatively charged particles, introducing bonds as arrows representing the direction of electron transfer.  “Although their polarity theory did not explain the bonding of nonionizing molecules, much solid-phase chemistry seemed to support it.”  Further essays by other physical chemists “firmly established a polar theory of chemical bonding in America”.  Polarity theory was useful to inorganic chemistry but did not serve organic chemistry well, arguing against it on the basis of “the phenomenon of tautomerism, the low dielectric constants, and the poor conductivities of organic compounds”.  After 1913(?), Thomson switched from his plum pudding model to one of a “planetary-nuclear model of the atom” based on evidence from Rutherford’s studies,  concluding that there were two types of bonds (independently of the Americans), polar and nonpolar where the nonpolar bond can be described as “tubes of force anchoring electrons both to the nucleus and to an adjacent atom” with each bond equivalent to two ‘tubes of force’ or two electrons.

In 1916, hitting upon the shared electron pair theory of bonding, Lewis wrote:  “It seemed rather that the union of sodium and chlorine and the union of hydrogen and carbon must represent extreme types of a method of combination which ultimately would be found to be common to all kinds of compounds”.  The author interprets this as “where electrons were equally shared, the electron exhibited no polar properties, but if one atom took an unfair share, then the charge being unequally distributed polarity would be produced”. Irving Langmuir helped publicize Lewis’ theory of the shared electron pair bond contributing the term “covalent bond’ to replace nonpolar bond and “reconciling’ Lewis’ static electrons with Bohr’s dynamic electrons.  In 1923, Lewis published Valence and the Structure of Atoms and Molecules.  In it, he wrote to explain the reactivity of elements and compounds using concepts of dynamic electrons and electrical Coulomb forces of attractions and repulsion.  He used Bury’s and Bohr’s theory of electrons in shells shielding the nucleus and where the “orbit as whole” and not the position of an electron within the orbit that drives the chemical behavior.  In this theory, “two atoms may conform to the rule of eight or the octet rule not only by the transfer of electrons from one atom to another, but also by sharing one or more pairs of electrons.  The electrons which are held in common by two atoms may be considered to belong to the outer shell of both atoms”.  Complete ionization was then considered a special case of the new theory.  In this context, valency was then defined as the ‘number of electron pairs which an atom shares with another atom’.
Two other important ideas were covered in Lewis’ Valence book: 1) a broader definition of acids as electron pair acceptors and bases as electron pair donors; and 2) identification of hydrogen coordination between a carbon and oxygen atom by his student Maurice Huggins as a “hydrogen bond”

Lewis' electronic theory of bonding found a receptive audience in America and Britain but was slow to spread in France and Germany due to a number of political, cultural, and social factors.  In the meantime, however, the wheels of mathematical physics were actively turning in Germany.  One of the significant events related to the electronic theory is the determination of the number of electrons necessary to neutralize what can then be measured by x-ray spectroscopy (by correlation) as the number of positive charges in an atom.  In 1912, Max von Laue, an optical physicist, postulated that “if x-rays were shortwave radiation then, assuming that crystals arose from a periodic array of atoms comparable in dimensions to the wavelength of the x-radiation, the waves ought to be diffracted”.  Experiments to confirm the diffraction effects were then used by Bragg father and son to develop a mathematical method of crystal structure determination.  Harold Moseley used these results to develop the science of x-ray spectroscopy.  By measuring the wavelengths of characteristic reflections for each element, Moseley discovered a ‘unique atomic number’ that when arranged in sequence (with a few exceptions) followed the same order of the elements in periodic table according to atomic weight’.  Because the atomic number equaled the positive nuclear charge (not sure how this connection was made), the number of electrons in each required for a neutral overall charge can now be determined.  This significant finding prompted the author to note that “Rontgen’s discovery therefore, not only revolutionized orthopedic surgery, but helped to transform chemists’ ideas of the structure of materials and stimulated the physicists’ interpretation of the nature of the atom”.

Development of the quantum mechanical theory of electronic structure:
By the end of the 19th century, there was a clear bifurcation in fields of interests of chemists and physicists: chemists dealt with 90 or so elements of matter while physicists dealt with “a more nebulous mathematical world of energy and electromagnetic waves”.  The chemists’ material particles were discrete while energy was continuous.  In the early 1900’s, Max Planck’s developed his quantum theory based on his blackbody radiation studies that showed that, "at the atomic level, energy could not be emitted or absorbed continuously but only in small discrete steps he called quanta.  In this quantum theory, energy radiated from a lamp or heat source only appeared to radiate continuously because it was a smoothed or averaged effect of large numbers of quanta” (which Einstein called photons in 1905).  In 1913 Bohr developed his planetary model of the atom skirting the "paradox of classical physics that revolving electrons would lose energy continuously and therefore slowly collapse into the nucleus", by assuming that electrons revolved around the nucleus without radiating continuously.  Adopting Planck's quantum theory, he correlated observed lines in a substance's spectrum with the energy of the emitted or absorbed photon as an electron jumped from one orbit to another.  Good correlation was found with experimental data for hydrogen but "atoms with more than one electron proved more recalcitrant.”  In 1923 Bohr and Bury’s were able to correlate atomic electronic structure theory with spectroscopic data by refining Bohr’s planetary model with the addition of quantum numbers to define the ‘orbits’ and the magnetic and spin properties of electrons within the atom.  In this refined theory, electrons were assumed to be arranged into shells corresponding to the principal quantum number n.  Subshells were assumed to exist to explain the properties of the transition and rare-earth elements.  In 1924 Louis de Broglie’s thesis postulated a dual nature (corpuscular and wave-like) to electrons in analogy to Einstein’s postulate of the dual nature of light. (Davisson and Germer in 1927 found evidence for this upon observing electrons exhibiting diffraction patterns like x-rays by being diffracted by matter.)  Building upon deBroglie's thesis, in 1926, Erwin Schrodinger used wave mechanics and formulated a partial differential equation (now known as Schrodinger’s equation) whose solution/s were interpreted as a measure of the density of electron charge at each point and equivalent to a probability density for an electron being found at a particular point in the ‘wave cloud’.  (Pauling was to later interpret this electron cloud as the region where bonds were most likely to form.).  Heisenberg applied abstract matrix algebra to reformulate Planck’s treatment of discontinuous phenomena and this "‘new quantum mechanics’ fitted the known spectral phenomena even better than Bohr’s and Sommerfeld’s treatment and was soon to be found equivalent formally to the ‘wave mechanics’ treatment of Schrodinger.”

The Pauling Bond
Pauling’s first substantial involvement in developing a bond theory could be found in his doctoral dissertation (on x-ray crystallography study of crystal structures) where he noted  that the atomic radii of atoms were “squashed” in the direction in which they formed covalent bonds, a finding that would prove useful in studying the nature of the chemical bond between atoms.  In 1927, Heitler and London produced the first ever quantum mechanical treatment of a chemical system, calculating the energy of the hydrogen molecule where the two electrons are held together by two protons. As a result of this, they were able to come up with a ‘fairly’ simple expression for the hydrogen wave function that could be fitted into the Schrödinger equation, “the solution to which gave a binding energy impressively close to that obtained from spectroscopic studies”.   Pauling developed his valence bond model for chemical bonding based on his analysis of the Heitler-London treatment of the hydrogen molecule which they extended to the H2+ molecular ion.  According to J. H. Sturdivant, Pauling did this “to make the results of quantum mechanics accessible and familiar to chemists untrained in the new theoretical physics, and to fuse the results into the foundations of chemical theory”.  Chemical bonding was then explained qualitatively using quantum mechanical results as an overlapping of orbitals (probability density regions defined by solutions to SE) occupied by shared electrons of opposite spins (now taught as Pauling’s spin pairing theory).  Friedrich Hund formulated a ‘rule of maximum multiplicity’ (now taught as Hund’s rule) which stated that successive electrons occupy separate orbitals before pairing up with opposite spins (1926).  In a 1931 publication in JACS, to explain the observed tetrahedral geometry of tetravalent carbon, Pauling developed the concept of hybridization of s and p orbitals showing lowered energy overall for stable bond formation and resulting in a tetrahedral geometry of the resulting orbitals.  The author noted other contributions by Pauling (detailed in his Nature of Chemical Bond book): determination of the ‘electronegativities of compounds’ as the “difference between their energies if all the bonds were covalent and all were ionic”; correlation of diamagnetic and paramagnetic properties to the presence of opposed spins in the paired bonding electrons and the presence of unpaired electrons, respectively; determination of bond angles, ion radii using crystallography; thermodynamic calculation of the energy of crystals; interpretation of the structure of benzene and its reactions.  During his prolific career, Pauling found himself working on biochemical research: he determined that the bond between oxygen and hemoglobin is covalent giving arterial blood diamagnetic properties whereas venous blood has paramagnetic properties.  Between 1946-1949, he discovered that the disease anemia was due to a deformation of the hemoglobin.  In 1936, he developed a theory (along with Alfred Mirsky) that proteins are coiled chains of polypeptide units held together by hydrogen bonds and came up with the stable alpha helix structure for some proteins.

Molecular Orbital Theory
Developed by Spectroscopists Hund and Robert S. Mulliken in 1927-28,  the Hund-Mulliken interpretation attributed the more complex spectra of molecules to various rotational, vibrational, and oscillatory motions of nuclei and of electrons around the nuclei.  In this treatment of diatomic molecules, the electronic structure of the hydrogen molecule was thought of as being that of a helium atom that has separated rather than two atoms of hydrogen to take into account the more complex interaction between a 2-proton center and each electron.  This approach then describes the H-H bond as a ‘molecular orbital’ emerging from the 2 1s electrons.  The author summarized the three consequences of this model: it ‘codified and clarified’ spectroscopy; led to the concept of electron promotion (and concept of excited states) and correlated orbitals of diatomic molecules with those of a ‘united atom’ and those of a ‘separated atom’; and became an alternative to the valence bond theory of Pauling.  In the 1940’s-1950’s, MO theory began to gain more ground as a better model for dealing with polyatomic molecules. Part of the reason for the increasing relevance of the MO theory was its ‘simplicity’ relative to Pauling’s resonance concept (“by picturing the VB treatment as equivalent to saying that mathematically the bonds between two atoms have to be built up from the polar and covalent forms”) which was not well-received in some circles.  He defended it by saying that “his approach was an honest extension of structural theory in which, in any case all formulae were idealizations”.   In 1967, Pauling conceded and devoted a substantial section of his introduction to modern structural chemistry text to MO theory but still noted the need to use the simpler VB method for introductory teaching.  The author ended this chapter with the following final comment on the competition between MO and VB models:  “Paradowski goes to the heart of the matter in his comment on the rivalry between VB and MO 15 : In physics it is possible to develop a simple and detailed model to explain certain classes of phenomena, but chemistry is too complex to be fully explained by such simple theories. To explain chemical phenomena at the present time [1972], one needs several good models. But these ‘good’ models are more flagrantly models, i.e. they explain only a selection of data, and hence the need for several models. Depending upon the symbolic apparatus used, different truths emerge. Twenty years on, this assessment remains true. Theoretical chemistry is still a quirky empirical science based upon a Schrödinger equation that can hardly ever be solved.”

Book Reading Update: Fontana History of Chemistry Chapter 12

This chapter is titled "The Chemical News" as it describes the development and proliferation of two vehicles by which chemists communicated with each other and disseminated information: publications and societies.  The rapid expansion of chemistry after the 1840's is further highlighted by the proliferation of periodicals and societies. Liebig's statement in 1834 underscores the importance of these periodicals to chemists:  “Chemical literature is not to be found in books; it is contained in journals.  In books, the individual author’s opinion dominates and his judgment is without appeal.  However, in journals there is defence, there is justification, and because of the necessity of a mutual neutralization of opinions, we approach the common goal of science.”

Until the 1840’s, both British and Continental chemical news was published in general commercial journals, with commercial trade journals being distinct from journals published by societies. Many of these commercial journals, however, were “extremely quarrelsome organs and served more as factional weapons of propaganda than as journals containing serious chemical, medical, and pharmaceutical information.  They make entertaining reading and are important sources for the social historian of science…”  For serious historians of the development of chemistry as an academic and practical subject, the author recommends Taylor’s Philosophical Magazine and the Chemical Gazette.

The first national Chemical Society in Britain independent from the Royal Society of Science was formed in 1841.  The author offers several reasons for why it took longer than other national science societies:  many of the societies formed were specialized and there was no collaborative chemical projects to bring together the small interest groups and different classes of practitioners and philosophers existed and practitioners were further divided into different fields of interest.  [A brief outline of the development and branching of the chemical profession: Chemistry back then was divided into the pure and the applied.  The industrial revolution further reinforced the dichotomy between chemical philosophers interested in the nature of chemical phenomena and chemical practitioners.  Chemical practitioners were further divided into chemists and druggists and industrial chemists.Apothecaries became medical practitioners of pharmacy as distinguished from those who were involved with the commerce of drugs and chemicals.Chemical practitioners also have established professions in the field of analysis for other types of chemical industries such as dye making, alkali production, etc.  Chemistry teaching also became an established chemical profession for both practical and philosophical chemists as the medical education made chemical education a requirement and scientific and nonscientific societies availed of lectures in chemistry.  Chemists also became expert witnesses in court litigations and patent cases and advisors to governments and industries in matters requiring chemical knowledge such as water sanitation.]  In May 1841, the Chemical Society of London was founded and became a model for other chemical societies, its founding driven by collaborative philosophers and practitioners of chemistry of all classes and professions open to ideas coming from the continent.  The need for a qualifying organization to certify competent analysts drove the formation of the Institute of Chemistry, a splinter group to remove the qualifying process from member selection that can be manipulated and corrupted.  At this time there was an increasing need for competent analysts due to demands of quality control in industries like alkali manufacturing and gasoil production and as advisors to local municipal governments to enforce and industries to comply with new public health regulations on water sanitation, air pollution, etc.  Nevertheless, specialized societies begun to crop up again to protect their own interests: technologists, dyers and colorists, electrochemists, etc.  Even with the amalgamation n the 1970’s of the many specialized organizations for economy of scale, the “peculiar” British class system of no mixing between managers and workers kept the Royal Society of Chemistry separate from the Society of Chemical Industry.

The Chemical Periodical:  In April 1778, Lorenz Crell founded the first ever chemical periodical, Chemisches Journal in the “home of specialized journals”, Germany.  The financial and wide-distribution success of this journal was attributed by the author to "its formula of providing fresh chemical intelligence to middle-class doctors, apothecaries, and chemical manufacturers who had little access to the proceedings of learned societies and academies".  It ceased publication in 1804 due to its endorsement of phlogiston chemistry.  In 1789, Annales de chimie (‘the voice of Lavoisier chemistry’) was founded by Lavoisier, de Morveau, Berthollet, and Fourcroy.  This periodical followed the German periodical tradition of providing translation of German and English papers.  In 1790, in Italy, Annali di Chimica was edited by Ludovico Brugnatelli.  In many cases, national societies published journals in their own language and "gave chemical communications a far more insular, even nationalistic, appearance than they had had in earlier more urbane and transnational commercial periodicals.”  By 1900, Germany had become the world leader in publishing important chemical periodicals catering to a prolific group of chemists including Americans.  In addition, the German Chemical Society was also responsible for updating Beilstein, a multi-volume compendium of organic compounds.  “Neither Britain, France, nor America matched the variety of German literature until after the First World War…”

The founding of the American Chemical Society:  In 1874, H. Carrington Bolton of the Columbia School of Mines suggested the gathering of chemists at Priestley’s home to celebrate the 100th anniversary of his isolation of dephlogisticated air (oxygen).  In this gathering, Charles Chandler proposed the formation of a national chemical society, independent of the chemical section of the American Association for the Advancement of Science.  It was voted down but Chandler proceeded to form the New York Chemical Society in 1876.  He shortly after that declared it should be an American Chemical Society with John Draper as its first president.  Not everyone agreed and some local societies were formed in response to the perceived hegemony of east coast chemists.  The biggest step to solidarity came in 1893 when the Journal of the American Chemical Society moved its publication from New York to Philadelphia.  By 1908, ACS had some 3400 members, more than the German Chemical Society.

A section tribute to William Crookes, Chemical Editor of Chemical News:  Crookes was described by the author as a “scientist of considerable originality and a journalist with sound commercial sense”.  “Crookes brilliantly overcame the difficulty of catering at once for the purely scientific reader, the practical chemist, and a still larger class of reader who, with but imperfect knowledge of science , [sought] only for information which may be turned to account commercially”.  Chemical News was an independent commercial chemical journal that ran weekly almost 70 years “of continuous service to the chemical community”.

Book Reading Update - Fontana History of Chemistry Chapter 11

The title of this chapter (How to Teach Chemistry) is taken from the title of a textbook written by one of the people the author considers as being greatly influential in the evolution of the teaching of chemistry, Edward Frankland.  The chapter is about the development of the teaching of chemistry in Great Britain (and some in the US) in the 19th and early 20th centuries, from a purely didactic form to one that involved hands-on and discovery-based laboratory experimentation leading to the growth of laboratories.  It also gives an account of the “internationalization of approaches to chemistry teaching in the 1960’s”.

The two “path-breakers” in this effort were Edward Frankland and Henry Armstrong in England “who both built upon the German experience and precedent of Liebig’s great teaching laboratory at Giessen”. The author also mentions Cannizzaro as one who had significant contributions to chemistry teaching, “based upon a standard two-volume molecule derived from Avogadro’s hypothesis, that equal volumes of gases at the same temperature and pressure contained the same number of molecules”.  A.W. Hoffman was also mentioned for his reform of elementary teaching of chemistry and innovations on facilitating chemical calculations derived from chemical equations.

Frankland’s State-Sponsored Chemistry: Frankland believed in the importance of science education in the “future professionalization” of science.  Most of his vision was put into practice when he became chief examiner of the state Department of Science and Art.  There were two important facets to Frankland’s contributions as an educator:  his innovative teaching and organization of chemistry facts illustrated in his textbook Lecture Notes for Chemical Students which he wrote to assist students in paying full attention to the lecture rather than focusing on taking down notes.  The other was his influence on the teaching of chemistry through the development and improvement of examinations and his publication of the teaching textbook How to Teach Chemistry.   Frankland’s disappointment at the poor performance of students on practical chemistry concepts was the catalyst for his efforts as the Department of Science and Art chief examiner to devise ways to improve the teaching of practical chemistry:  securing state funding for lab facilities in science classes, requiring teachers to demonstrate experiments, and training chemistry teachers on how to do these demonstrations.

Frankland’s emphasis was on the more ‘economical’ method of lecture demonstrations that allowed students to ‘see’ the experiments.  His predecessor Henry Armstrong was to take this one step farther, successfully bringing into education the ability for students to actually ‘do’ experiments. Armstrong's objection to the didactic teaching-centered method of learning drove him to advocate for and implement the heuristic method, “the child-centered method of learning through experience and experiment”.  “‘Heuristic methods’, said Armstrong, ‘are methods which involve our placing students as far as possible in the attitude of the discoverer.’ What students found out for themselves was best remembered because then they truly understood. Moreover, because a student- investigator is genuinely interested and involved in finding a solution to a problem, they would learn about it more efficiently than if they were given the solution didactically.”  He subsequently designed a new syllabus for beginning chemistry that “proceeds ‘systematically from the known to the unknown’.   The author considers Armstrong's innovative heuristic method as a catalyst for the school laboratory movement: by 1902, the number of school laboratories in Great Britain went from 150 to over 1,000.

Twentieth-Century Developments in Teaching:   In the USA, the American Chemical Society played a big role in directing how chemistry was taught in schools.  In the 1920’s, high school chemistry was organized into inorganic, organic, and physical chemistry including laboratory analysis.  In the 1930’s, the ACS provided a set of minimum standards , implemented an accreditation process, and published a standardized set of examination questions that ensured uniformity in the teaching of chemistry across the nation.  “The correlation of mathematics and science, however, remained a particularly American problem.”  By the 20th century, more and more advanced mathematical knowledge became necessary to understand the new chemical concepts being discovered and subsequently included in the chemical education syllabus.  Responses from educators and chemists to bridge this gap include the addition of trigonometry and calculus in the British education system and the publication of mathematics textbooks for the study of science.  Furthermore, in America, the launching of the Russian space shuttle Sputnik triggered a movement to improve science education.  The Chemistry faction responded with two new movements, the Chemical Bond Approach and the CHEM Study program.  The revised curricula offered by both these programs were disseminated internationally.  In 1967, these two programs and the Nuffield curriculum program in Great Britain “all came to fruition” after international meetings and the resulting “attitude towards chemistry teaching can be summed up as follows: when teaching the chemistry of the atmosphere, it is no longer necessary to show students how to prepare oxygen and nitrogen; instead take them straight from gas cylinders just like any other chemical”.  It was later on determined that this evolved curriculum only worked best with the “brighter students”.  In the 1970’s, researchers also found that students brought their own pre- and misconceptions to the classroom and the role of chemical education became that of facilitating the “unlearning a prior view and perceiving it anew”.  Also, new findings on the psychology of learning and a  ‘more critical philosophy and sociology of science’ further undermined the discovery method and there was a new vision to create a “science for all” rather than a “science for the few”.

The last section in chapter focused on the evolution of the laboratory facility.  The primary focus given by Frankland and Armstrong on a strong laboratory component in chemical education in the 19th century and subsequent curriculum reforms in the 20th century placed a huge demand on laboratory facilities, previously only required by research programs since Liebig’s time.  The design specifications of chemical laboratories have always been driven by the following problems and needs:  artificial ventilation due to production of noxious and sometimes dangerous fumes and liquids, gas, water, and drains for discarding waste.  In 1862, passage of the Morrill Land Act, donating federal lands to create private or state engineering and agricultural laboratories in America, led to a major expansion in science laboratories, allowing Universities of Wisconsin, Illinois, Michigan, and Cornell among others to become centers of science teaching and research.  The cost of building and equipping a laboratory rose in the 20th century due to the introduction of new equipment.  Between the two world wars, there was a second wave of laboratory building and extension.  Post second world war expansion of British universities led to new design for labs and experimentation with new material.  Based on changing number of students and scientific curriculum and the need for flexibility, some of the considerations were:  need for air conditioning for mass spectrometry and spectroscopy and constant temperatures for chromatography; special constructions to protect from radioactive materials and hazardous chemicals; switching to longer peninsular surfaces from island benches for convenience and servicing; and redesign of the ‘fume cupboard’ which had the most specialized construction needs and was the “most awkward thing to deal with”.  Lastly, it is worth noting the author’s note on the current (at time of writing in 1996) and future state of laboratory design, some of which had been realized: “The indications during the last thirty years are that designers, influenced by the special precautions needed for safe research, the availability of highly efficient ventilation systems, and the fact that laboratories in most countries fall within the orbit of much health and safety legislation, foresee a situation in which all actual experiments will be performed in centrally placed fume cupboards or shielded glove- boxes, leaving ‘the open benches as places for assembling equipment and placing reagent bottles, spare glassware, notebooks, etc.’”

Book Reading Update - Fontana History of Chemistry Chapter 10

This chapter gave an account of the development of the ionic theory, primarily attributed to Arrhenius but, much like everything other theory or law in science, evolved into what it is today based on the work of many others.  The development and refinement and drama surrounding the ionic theory was the core of the emergence of the field of physical chemistry as the author contends.

When chemists first started referring to the field of physical chemistry , it was initially the study of the physical properties of chemical substances, in large part inspired by studies of optical activities observed for stereoisomers.  The author gave some examples of findings that would have fallen under the "physical chemistry" rubric such studies on correlations between boiling points and molecular size and complexity, refraction, polarimetry, specific heats and atomic weight, what became known as Hess' law, etc.

On the development of the ionic theory and Arrhenius' subsequent formulation of it had a big help from Raoult and van't Hoff.  The author began the chapter with a short story on how Arrhenius' dissertation thesis on the dissociation of particles in solution was not well-received.  His ideas, however, were supported by subsequent experiments carried out on properties of solutions.  Raoult carried out studies on the effects of particles on the freezing points of solvents and observed a relationship between the size of the FP depression and the number of solute particles relative to solvent particles. Van’t Hoff observed an analogy between osmotic pressure and gas laws as expressed by the mathematical relationship PV=kT.  With Arrhenius’ insight on dissociation, he was able to determine the physical significance of the proportional constant needed to make the equation work for acids, bases, and salts, as a measure of the dissociation of the solute particles.  Later experiments showed a good agreement between i values measured (van't Hoff in particular)  in the context of different physical processes supporting Arrhenius’ interpretation of this quantity as a measure of dissociation of the ions.  Essentially this series of events tied together point to the idea that the behavior of solutes in solutions was illuminated by how the properties of the solvent and solution have changed and the number of particles is the unifying concept (we now teach this as part of the section on colligative properties).  Arrhenius formally formulated his theory of ionic dissociation based on strong experimental support for particle dissociation in solutions (and contributions by others that made use of mathematical and qualitative analogies between particles in solutions and gas particles following gas laws).

In 1887, Ostwald formulated a dilution law after applying the electrolytic theory to acids, confirming Arrhenius’ finding that at infinite dilution the molecular conductivities of all acids were the same.  Ostwald used the ionic dissociation theory and applied it to acids and bases and was able to show that the hydrogen ion concentration in water is 10-7 M and established the hydrogen ion concentration as a measure of the relative strength of an acid using the equilibrium dissociation formula found for water.   Sorensen was to formulate the pH scale later on to remove the negative index.

The ionic theory met social ("discomfort" from organic chemists about this emerging field not tightly beholden to laboratory studies and political friction between Fance and Germany) and scientific objections.  A prominent one was brought up by Kahlenberg based on experiments he and his students conducted which showed non-correlation between dielectric constant measurements and dissociation of ions in solvents other than water, suggesting that solute-solvent interactions should be taken into account and that Arrhenius' ionic theory is defective.  Lewis and other prominent physical chemists of the time defended Arrhenius' ionic theory citing that "‘Perfection is rare in the science of chemistry. Our scientific theories do not spring full- armed from the brow of the creator. They are subject to slow and gradual growth, and we must candidly admit that the ionic theory in its growth has reached the ‘awkward age’. Instead, however, of judging it according to the standard of perfection, let us simply ask what it has accomplished, and what it may accomplish in scientific service."

By 1923, ionic theory was 'essentially complete" after refinements and corrections based on a mathematical model developed by Milner and simplified by Debye and Huckel were made based on solute ions – solvent interactions and “statistical interference of neighboring ions on ionic mobilities”  In 1927, Onsager added refinements to the theory for nonaqueous systems.

Book Reading Update - Fontana History of Chemistry Chap 9

As the author stated at the end of the chapter, "The periodic classification of elements is one of the greatest and most valuable generalizations in science.  Important in its late 19th century construction for the clarification it brought, and the guidance it gave to chemical research, it was able to adapt to the deeper analysis of the structure of matter revealed by nuclear physics in this century.”

Although there were similar, some concurrent attempts at organizing the elements, the majority of the effort and the most significant initiative and push are attributed to Mendeleev.  Mendeleev used as his primary organizing principle the fundamental property of each unique elements its mass.  In "sorting the elements", Mendeleev was guided by atomic weights (relative scale with C=8), valency, and isomorphism.  In February 1869, he came up with the law “elements placed according to the value of their atomic weights present a clear periodicity of properties”. It is notable that Mendeleev had such confidence in his organization that it warranted redoing of experiments when data did not conform to his prediction as with the case of the measured atomic weight of tellurium and its “misplacement” in Mendeleev’s periodic table.  His predictions ranged from spot on to luck guesses to incorrect outcomes.

The identification of the rare earths proved to be the most difficult to chemists who have devoted their work to the study of these elements.  The rare earths have very similar physical properties which made their resolution into different elements from a mixture and characterization very difficult.  These rare earth discoveries were befuddling to Mendeleev because of the seeming lack of periodicity.  There was debate as to where to place them within Mendeleev’s periodic table and it was suggested they should be placed in their own table until more research has been done.  In 1913, Moseley confirmed that there were exactly 14 rare earth elements.  It was not until Bohr’s electron shell arrangement and atomic number concept did it become clear that “the rare earth group formed a miniature periodic system in which the relationships of the main system are reflected”.

Because chemists were mainly using chemical methods to identify elements, measure weights, and observe properties, the inert gases were not discovered till later, principally by Ramsay who collaborated with Lord Rayleigh in the beginning.  By 1900, 5 of the 6 inert gases have been discovered but it was not until the 1960's that a chemical compound was synthesized using xenon.  Spectroscopy played a big role in the verification of these new elements.

Other methods of discovering the missing elements in Mendeleev's periodic table included manufacturing them in the lab and observing their productions in radioactive decays and fission processes.

As the author concluded in the end, after Lavoisier’s definition of the chemical element, chemists became convinced that matter was built up from these different substances.  Mendeleev sought to find a way to organize these elements using the fundamental nature of mass as his organizing principle, encouraged by observations in "organic chemistry of the dependence of physical properties on molecular weights". "Mendeleev’s system, in which atomic weight became the fundamental feature, or chemical standard, for identifying and defining elements, was sufficiently elastic to allow the prediction of missing elements, or the inclusion of whole new groups such as the noble gases and actinides.”  The discoveries of other elements relied on physicists’ kinetic theory rather than chemical methods.  Spectroscopy played a big role in the identification and verification of many of these, showing “how the explanation of the (periodic) law lay in a theory of atomic structure.  “Following Moseley’s concept of atomic number, and its correlation with Rutherford’s idea of a nuclear charge, it became possible to place a finite figure on the number of possible elements, and to redefine atomic weights as an averaging out of isotopic masses.”