THE ORIGINS OF CHEMISTRY AND MODERN ATOMISM
Since modern atomic theory along with celestial mechanics represent the two most significant theoretical developments in the physical sciences that changed our conception of the modern world from the ancient Aristotelian model to the modern mechanistic one, the latter requires a separate chapter. Recall that it was the ancient Greeks who first endeavored to understand the universe in a more empirical-rationalistic manner to replace the earlier mythical or theogonic interpretations. This required describing the primal elements from which everything arose, along with explaining how the diversity of nature came to be from this primal state.
Though it was Empedocles’ conception of the four elements of fire, air, earth, and water as primary that was adopted by Aristotle, which prevailed throughout most of the past, modern classical science reinstated the atomic theory of Leucippus and Democritus, along with the theory of infinite particles composing the universe introduced by Anaxagoras and adopted by Epicurus and Lucretius. Thus it was natural philosophers like Mersenne, Galileo, Gassendi, Descartes, Boyle, Locke, and Newton who revived the atomic or particle theory in the seventeenth century by adopting the corpuscular-mechanistic framework, though the conception at that time was still entirely speculative and elementary.
Although as early as the third century BCE Anaxagoras had declared that basic particles were infinitely divisible, when Newton adopted the corpuscular theory as the basic physical reality these particles were still mainly defined in terms of the Democratean primary qualities of solidity, shape, indivisibility, and motion (although Epicurus claimed they were composed of an inseparable minima), along with the more recent additions of mass, momentum, inertia, and gravitational attraction. Although the pseudosciences of alchemy and astrology were still pursued, the former by such distinguished natural philosophers as Boyle and Newton, they would soon be eclipsed by advances in modern classical science whose superior methodology led to the discovery of more elementary particles such as the electron and proton and an explanation of chemical compounds and reactions according to their exact molecular components, structures, and properties, rather than by God’s will.
As was true of the transformation of the former notion of the celestial world to the modern conception of a gravity driven planetary and stellar universe according to mathematically defined astronomical laws, this new atomic and particle physics also would require a radical conceptual revision. Though not the first to use the balance to weigh the exact quantities of the reagents and products of chemical reactions, Antoine Laurent Lavoisier is considered the father of modern chemistry owing to his precise weighing of the components of combustion and oxidation that enabled him to determine that oxygen was a gas facilitating combustion, thereby refuting the prevailing phlogiston theory that postulated a fire-like element within combustible bodies. English chemist John Dalton similarly is regarded as the founder of modern atomism based on his discovery that natural elements like water, gases like carbon dioxide, and chemical compounds like sulfuric acid have a molecular structure that can be analyzed into specific atoms that combine according to simple numerical ratios according to their numbers: H2O, CO2 and H2 SO4 respectively.
As religiously and rationally significant as was the transformation of the conception of a heavenly or celestial cosmology to a natural physical universe, for most of us, except for weather predictions and hurricanes and tornados, it is somewhat remote from our daily lives. This, however, is not true of the empirical sciences such as physics, chemistry, biology, physiology, medicine, engineering, etc. It is these sciences in particular that have radically changed our lives from what they were before the advent of science.
It was the overthrow of the theory that combustion was due to the expelling of phlogiston and replaced with the burning of oxygen that is usually credited with having been the major factor in the development of chemistry. It began with German physician Johann Joachim Becher’s claim that combustion involved the burning off of the “fatty earth” described in his treatise Physicae subterraneae in 1669. Then German chemist George Ernest Stahl, in his book Fundamenta Chymiae (Fundamental Chemistry) in 1723, renamed Becher’s terra pinguis “phlogiston,” claiming that it was “the matter and principle of fire,” though not fire as such. According to the phlogiston theory certain substances, like wood, charcoal, and phosphorus contain large amounts of this “inflammable principle” that they give off when heated that is combustion.
Then the Swedish chemist Carl Wilhelm Scheele, in a book translated as Chemical Treatise on Air and Fire published in 1777, reported his discovery that air consisted of two components: one highly flammable that he called “Fire air” and the other inflammable designated “Foul air,” the first later renamed “oxygen” and the second “hydrogen.” Though he detected a flammable substance in the air he did not investigate it. Like Scheele, Joseph Priestley in different experiments noticed that when substances are burned in air and the residue and the air are carefully weighed the residue usually gained weight while the volume of air decreased, contrary to the phlogiston theory that claimed the burning material gives off phlogiston and thus should weigh less, while the air gaining the phlogiston should weigh more.
Priestley’s description of the transformation of calces (the residue of a burnt mineral) into metals is just one of many examples.
For seeing the metal to be actually revived, and that in considerable quantity, at the same time that the air was diminished, I could not doubt, but that the calx was actually imbibing something from the air; and from its affects in making the calx into metal, it could be no other than that to which chemists had unanimously given the name of phlogiston.49
Thus Priestley is credited with discovering that it was “something from the air,” a “new air,” that caused the combustion, but as his final word “phlogiston” indicates, he was so committed to the phlogiston theory that he “concluded that the new gas must contain little or no phlogiston, and hence he called it dephlogisticated air” (pp. 126–27), which meant air that is free from phlogiston or the element of inflammability.
And so the honor of explaining the significance of the discovery is attributed to Lavoisier. After many failures to explain the process of combustion, it was at a dinner meeting in Paris with Priestley in 1774 that the solution occurred to him. Priestley “told Lavoisier at dinner of his discovery of dephlogisticated air, saying he ‘had gotten it from precip [of mercurius calcinatus] per se and also red lead’; whereupon, he says, ‘all the company . . . expressed great surprise’” (pp. 126–27). What caused the surprise was that the so-called dephlogisticated air produced by heating mercury oxide had properties the opposite of carbon dioxide produced by heating charcoal: it supported burning and respiration and did not combine with lime and alkalis. Repeating the experiment Lavoisier obtained a gas purer than ordinary air which convinced him that while measuring the components of chemical reactions is crucial to chemistry, so is choosing the right experiments.
He read two papers describing his experiments on the oxide of mercury titled “On the Nature of the Principle which Combines with Metals during Calcination and Increases their Weight” before the French Academy of Sciences, the first on Easter 1775 and the second on August 8, 1778. Having initially decided that the gas produced in Priestley’s experiment though purer than common air was still a form of common air, when he learned of Priestley’s later experiment showing that when reacting with nitrous oxide it was more soluble in water than common air, he concluded that while it was a constituent of common air it was not identical to it! He thus considered it a gas that was absorbed in the conversion of metals to calces or oxides when burnt in air and emitted when the oxides themselves were heated. He named the new gas “oxygene.” According to chemist J. R. Partington:
In 1782 Lavoisier says Condorcet had proposed the name “vital air” for pure air, but in a memoir received in 1777 . . . and published in 1781, entitled “General considerations on the nature of acids and on the principles composing them”, Lavoisier called the base of pure air the “acidifying principle” or “oxigine principle” (principe oxigine), which he latter changed to “oxygene”[. . .] . (pp. 131–32)
The publication of Lavoisier’s Traité de Chimie (Treatise on Chemistry) in 1789 established the superiority of the explanation involving oxygen over that of phlogiston. This not only overthrew the phlogiston theory, it brought about a revolution in chemistry. It no longer was assumed that common substances such as air and water were irreducible, but indicated they were compounded of more basic elements that opened up a whole new world of research. Although Priestley never gave up the phlogiston theory himself, in his last book he graciously acknowledged Lavoisier’s contribution.
There have been few, if any, revolutions in science so great, so sudden, and so general, as the prevalence of what is now usually termed the new system of chemistry, or that of the Antiphlogistons. . . . Though there had been some who occasionally expressed doubts of the existence of such a principle as that of phlogiston, nothing had been advanced that could have laid the foundation of another system before the labors of Mr. Lavoisier and his friends. . . .50 (italics in original)
Yet as consequential as the chemical revolution has been, there was an impending revolution even more effective in transforming the conception of physical reality, namely, the reconstruction of the atomic theory. The success in chemistry of the experimental identification of oxygen and explanation of its function in combustion convinced scientists of the possibility of discovering the inner corpuscular elements of all ordinary substances, such as air, water, acids, metals, etc., along with the properties that could explain their combinations and reactions.
Thus while the intellectual and technological levels at the time of Leucippus and Democritus were insufficient to promote advances in the atomic theory, that was no longer true. Newton’s belief stated earlier “that God in the Beginning form’d Matter in solid, massy, hard, impenetrable, movable Particles, of such Sizes and Figures . . . as most conduced to the End for which he form’d them,” though a misrepresentation of their origin and properties, had finally been vindicated.
In 1787 Lavoisier, Claude Louis Berthollet, Guyton de Morveau, and Antoine François de Fourcroy published a book entitled Méthode de Nomenclature Chimique (The Method of Chemical Nomenclature) that presented the first modern list of elements based on recent experimental discoveries. Then in 1799 Joseph Louis Proust introduced “the law of constant proportions” stating that any sample of a compound or molecular substance, such as salt, always contains its constituents, sodium and chlorine, in fixed ratios by weight reinforcing the belief in the constancy of the reagents and the regularity of the reactions. Yet despite the considerable experimental evidence that substances were composed of more basic elements in ratios determined by their weights, there still was no scientific explanation as to why or how. And since Newton’s belief that they were caused by God was no longer adequate the search began for an explanation, another indication of the transformation in the conception of the external world and how to investigate it and how to understand it.
The person credited with initiating the explanation, John Dalton, like many of his predecessors, was a most unlikely candidate. Born to a Quaker family in the tiny, rustic Village of Eaglesfield, England (1766–1844), his father was a cottage weaver while his mother supplemented their income by selling writing materials. Unable to enter a private school, he attended the village schools in the neighborhood but had acquired a sufficient background that he was able to teach in the village school from the early age of twelve to fourteen. While teaching there he was fortunate to meet a wealthy Quaker named Elihu Robinson who, along with being educated in natural philosophy, especially meteorology, corresponded with Benjamin Franklin.
Having noticed Dalton’s mathematical aptitude when he won a dispute in mathematics, Robinson began tutoring him in mathematics with Dalton always appreciating the kindnesses, good advice, and intellectual awakening that Robinson and his cultured wife had contributed to his early development. Then, when he was fifteen, he moved to Kendal to become assistant in a boarding school rising to the position of principal. During his spare time he studied Latin, Greek, French, mathematics, and natural philosophy.
During the twelve years he lived in Kendal he made the acquaintance of a more unusual benefactor, another Quaker by the name of John Gough who, despite being blind and suffering from epilepsy, owing to his wealthy, well-educated, and intellectual family was able to acquire a sound knowledge of the classics, physics, mathematics, botany, and zoology. Though nine years older than Dalton and considerably more advanced in his studies, when he learned of their common interests and Dalton’s intellectual aptitudes, Gough became his close friend and academic mentor. His family having an excellent library and an extensive collection of scientific instruments, he shared these with Dalton who, in gratitude, served as his reader and amanuensis. As a result, Dalton became well schooled in “mechanics, [Newton’s] fluxions, algebra, geometry, chemistry (including some French chemical writings), astronomy and meteorology . . .”51 (brackets added).
Because of this close intellectual relationship, when Dr. Barnes of New College in Manchester wrote to Gough in 1793 (who had become a widely respected mathematician) seeking his suggestions in filling a position of professor of mathematics and natural philosophy at New College, Gough unselfishly recommended Dalton for the position, even though it would mean severing their very close association. When the position of tutor at New College was offered to him, Dalton readily accepted, partially because of his dissatisfaction with his teaching at Kendal and also because he foresaw a more promising future in Manchester, which was confirmed when he later described his life there as “very happy and fulfilling.”
Moving to Manchester he was immediately welcomed by the eminent “Mancunians,” as the patricians of Manchester were called, and elected to the prestigious Manchester Literary and Philosophical Society the following year. As author Elizabeth C. Patterson states:
The association which Dalton began with the Manchester Literary and Philosophical Society in 1794 was to continue until his death in 1844. During this half century the Society would play a central role in his life and he in its. Before it he read one hundred and seventeen papers, of which fifty-two were printed. For forty-four years he served as an officer—first as Secretary, then as Vice-President, and as President. To think of either—the Society or the man—is to think of the other. (pp. 59–60)
It was hearing the lectures and witnessing the experiments of English physician Dr. Thomas Garnett at New College that aroused his interest in molecular chemistry.
His early investigations and publications were centered on meteorology, including the nature of water vapor and the composition of the air and whether its components consisted of a mixture, a chemical compound, or some other structure. Then in a series of four essays he presented his research conclusions regarding gases, meteorology, and chemistry, in the last essay declaring that he had independently discovered Jacques Charles’s gas law that all gases at constant temperature will, with the same increase in temperature, expand equally. In tribute, Patterson declares that the “wealth of material in these four essays is extraordinary. Even today they are hailed as ‘epoch-making’ and as ‘laying the foundations for modern physical meteorology’” (p. 94).
It apparently was these initial experiments of the solubility of gases in water (similar to those of Robert Boyle) that led to his crucial insight that each element was composed of characteristic atoms that would be possible to distinguish by their atomic weights. The first explicit statement of this is in a paper he read to the Literary and Philosophical Society on October 21, 1803, entitled “On the Absorption of Gases by Water and Other Liquids.”
The greatest difficulty attending the mechanical hypothesis arises from different gases observing different laws. Why does water not admit its bulk of every kind of gas alike [i.e., why are they not equally soluble in water]? This question I have duly considered, and though I am not yet able to satisfy myself completely, I am nearly persuaded that the circumstance depends upon the weight and number of the ultimate particles of the several gases. . . . An enquiry into the relative weights of the ultimate particles of bodies is a subject, as far as I know, entirely new; I have lately been prosecuting this enquiry with remarkable success.52 (brackets in the original)
While other chemists were only investigating the relative weights in which the components of substances combine, such as the densities or weights of hydrogen and oxygen composing water, Dalton was the first to attempt to determine the relative weights of the components themselves. Appended to the paper was a “Table of the relative weights of the ultimate particles of the gaseous and other bodies,” while the paper itself presented the basic features of his atomic theory at the time.
After giving a series of lectures in Edinburgh and Glasgow that were highly praised—which must have been very gratifying considering his humble origins—he began writing his great work, a New System of Chemical Philosophy, published in 1808. The chapter “On Chemical Synthesis” was particularly significant because it explicitly states his original thesis that every sample of a basic substance such as water contains “ultimate particles” that always “are perfectly alike in weight, figure, etc.” When one considers the various possibilities as to how the ratios of the constituent particles could be construed, one begins to appreciate the complexity of the problem he faced. The key, he believed, lay in determining the individual atomic weights of the elements composing the substances. However, one can only claim that water is H2O rather than HO2 or H3O4 if one knows not only their atomic weights, but also in what numerical proportion they combine. As he summarized the challenge:
In all chemical instigations, it has justly been considered an important object to ascertain the relative weights of the simples which constitute a compound. But unfortunately the enquiry has terminated here; whereas from the relative weights in the mass, the relative weights of the ultimate particles or atoms of the bodies might have been inferred, from which their number and weight in various other compounds would appear, in order to assist and guide future investigations, and to correct their results. Now it is one great object of this work, to show the importance and advantage of ascertaining the relative weights of the ultimate particles, both of simple and compound bodies, the number of simple elementary particles which constitute one compound particle, and the number of less compound particles which enter into the formation of one more compound particle. (p. 229)
Frank Greenaway states Dalton’s “rules of greatest simplicity” as his guide to achieving this.
1 atom of A + 1 atom of B = 1 [compound] atom of C, binary.
1 atom of A + 2 atoms of B = 1 [compound] atom of D, ternary.
2 atoms of A + 1 atom of B = [compound] atom of E, ternary.
1 atom of A + 3 atoms of B = 1 [compound] atom F, quaternary.
3 atoms of A + 1 atom of B = 1 atom of G quaternary. etc. etc.53
Greenaway then presents the general rules that would guide Dalton in his “investigations respecting chemical synthesis”: that is, how the elements of compounds were arranged.
The composition of any substance must be constant (Law of Constant Composition). If two elements A and B combine to form more than two compounds then the various weights of A which combine with a fixed weight of B bear a simple ratio to one another (Law of Multiple Proportions). If two elements A and B combine separately with a third element C, then the weights of A and B which combine with a fixed weight of C bear a simple ratio to each other (Law of Reciprocal Proportion or Law of Equivalents). (p. 133)
He then assigned symbols to each of twenty elements with their atomic weights relative to hydrogen taken as 1 that he presented as a table of ELEMENTS in his New System of Chemical Philosophy. Although his symbols were later replaced by the Swedish chemist Jöns Jakob Berzelius (1779–1848) to those with which we are now familiar, he took the initial letter or letters from the Greek or Latin names or those rendered in English to designate the element: for example, “Fe (from the Latin ferrum)” to designate iron and H for hydrogen, O for oxygen, and Cl for chlorine, along with using superscript numerals to indicate the number of elements in a molecular compound such as water (H2O).54 It was Dalton who attempted the first systematic classification of atomic elements according to their atomic weights.
As Berzelius wrote when he first learned of Dalton’s atomic hypothesis: “supposing Dalton’s hypothesis be found correct, we should have to look upon it as the greatest advance that chemistry has ever yet made in its development into a science.”55 Later, after having read his New System of Chemical Philosophy, he wrote to Dalton that the “theory of multiple proportions is a mystery but for the Atomic Hypothesis, and as far as I have been able to judge, all the results so far obtained have contributed to justify this hypothesis” (p. 249). This was a vindication, finally, of the ancient theory of atoms introduced by Leucippus and Democritus that has proven so successful.
Yet despite Dalton’s achievements, his rules of atomic composition were still suppositional based on atomic weights that did not provide the exact ratios in which the elements combine to form compound substances. In 1809, the year after Dalton published his great work, French chemist Joseph Gay-Lussac published a classic paper, “Memoir on the Combination of Gaseous Substances with Each Other,” presenting his discovery known as “Gay-Lussac’s Law of Combining Volumes” that measured the ratios of combining gases by their volumes, rather than their weights as Dalton had done, and also provided more exact proportions thus allowing a more precise determination of the ratios of the combinations of the atoms in their molecular structures, such as CO2 or NH3. As quoted by Leonard K. Nash:
Thus it appears evident to me that gases always combine in the simplest proportions when they act on one another; and we have seen in . . . all the preceding examples that the ratio of combination is 1 to 1, 1 to 2, or 1 to 3. It is very important to observe that in considering weights there is no simple and finite [integral] relation between the elements of any one compound. . . . Gases, on the contrary, in whatever proportions they may combine, always give rise to compounds whose elements by volume are multiples of each other. (p. 260)
Gay-Lussac believed that his exact measurements of the integral ratios of combining gases were “very favorable” to Dalton’s rules of combining weights, adding empirical support. Yet Dalton had strong objections, especially to Gay-Lussac’s assumption that equal volumes of all gases under the same conditions contain equal numbers of atoms, instead maintaining that the different solubilities in water of the same volume of various gases implied that they were composed of atoms of different sizes and thus the same volume of different gases under the same conditions of temperature and pressure could not contain the same number of atoms.
If the density, mass, or weight of a substance is defined by the quantity of mass per unit volume, and the mass itself consists of the number and size of the particles composing the substance, then if two volumes of gas are equal but differ in density then this implies that either the number of particles are different or their sizes are or both. Thus Dalton concluded “that there are different numbers of particles in equal volumes of different gases was powerfully supported by experimental data on gaseous densities and combining volumes in gaseous reactions” (p. 266).
In addition he proposed as a maxim that the atoms in different gases vary in size: “That every species of pure elastic fluid has its particles globular and all of a size; but that no two species agree in the size of their particles, the pressure and temperature being the same” (p. 267). Thus Dalton attributed Gay-Lussac’s claim that his Law of Combining Gases provided a more exact method for calculating the number of atoms in a substance to the inexactness of his experiments, despite the fact that Gay-Lussac had the reputation of being one of the most exacting scientists. Yet there was no agreement and so the search for additional evidence went on.
In an effort to resolve the problem and also to combine both Dalton’s law of combining weights and Gay-Lussac’s law of combining volumes, Italian chemist Amedeo Avogadro in 1811 published his “Essay on a Manner of Determining the Relative Masses of the Elementary Molecules of Bodies, and the Proportions in which They Enter into These Compounds” in the Journal de Physique (Journal of Physics). As he stated, if the simple or elementary gases rather than being monatomic or composed of single atoms were composite: if the “‘particles’ present in the gaseous elements do not consist of the individual atoms of the elements but of groups of atoms of the same element joined in a single molecule of that element” (p. 284), then the anomalies can be eliminated. Thus the original polyatomic atoms of the combining gases could separate and recombine into such proportions as to maintain a constant number in every volume of gas.
There were, however, three weaknesses in the theory. First, there was as yet no empirical evidence to support the theory. Second, even if the smallest particles of the volumes of combining gases were polyatomic there was no way of determining in what proportion they combined nor the nature of the combining force. Third, there was a conflict in assuming that the force binding the polyatomic substances was attractive and the fact that the gas pressure was repulsive. So the challenge remained to find a method to determine the exact number and ratios of the elements of compound or molecular particles.
In the succeeding years much of the research was directed at trying to solve this problem. When Count Alessandro Volta, an Italian physicist, in 1800 developed a voltaic pile or electrochemical battery that produced a continuous electric current several physicists who believed that the force binding the polyatomic structure might be electrical realized that Volta’s electric current might decompose them into their components. Then in 1807 Sir Humphry Davy of the Royal Institution declared:
If chemical union be of the [electrical] nature which I have ventured to suppose, however strong the natural electrical energies of the elements of bodies may be, yet there is every probability of a limit to their strength; whereas the powers of our artificial instruments seem capable of indefinite increase . . . [Consequently, we may] hope that the new [electrical] method of analysis may lead us to the discovery of the true elements of bodies. (p. 296)
The theoretical assumption was that if the monatomic particles in polyatomic elements such as ammonia or water have contrasting electrical charges causing them to bind, then connecting two terminals or electrodes to them that in turn were connected to the oppositely charged poles of Volta’s battery with the charges strong enough, they would overcome the binding power of the charged particles thereby attracting them to the oppositely charged terminals. Today this would be called “decomposition by electrolysis.” Within a year Volta had confirmed his theory by decomposing alkali metals by electrolysis.
But it was J. J. Berzelius, previously mentioned in connection with the symbolic naming of the elements, who discovered in the electrolysis of water that when two oppositely charged electrodes connected to the opposite terminals of a battery were inserted into the water it decomposed into the elements of oxygen and hydrogen owing to their being attracted to the oppositely charged electrodes. Finding this to be true of other compounds and inferring that it was due to the opposite charges of the monatomic elements, he proposed the dualistic classification of “electropositive” and “electronegative.” Moreover, since electrical experiments had shown that like charges repel while opposite charges attract he inferred, as did Davy, that the stability of the elements could be explained by the attraction of their opposite charges which, when neutralized in polyatomic substances, left them uncharged.
Following his successful electrolysis of water, with the aid of Gay-Lussac’s law of combing volumes, Berzelius was able to designate the correct molecular structure of compound substances using superscript numerals (later changed to subscripts) such as water (H2O), ammonia (NH3) nitrous oxide (N2O) formed from gaseous elements. Though restricted to gases, his procedure was a significant advance because he was able to derive more precise atomic weights of the separate elements composing polyatomic substances. As a result, he spent more than a decade measuring the combing weights of the elements forming the compounds and publishing the results in tables in 1814, 1818, and 1826. And as indicated, having assigned symbols to each monatomic element based on the initial letter or letters of its assigned name, he then listed what he found to be the atomic structure of the polyatomic or compound substances with superscript numerals, indicating the ratios of their elements followed by what he had determined to be the compound substances’ atomic weights.
Though this was a great improvement over the table of ELEMENTS published by Dalton, it still was not a sure method for determining the precise atomic weights of most elementary atoms nor the exact molecular structure of compound elements. The next major effort was the specific heat method of Alexis Petit and Pierre Dulong presented in a paper to the French Academy of Science on April 12, 1819. Aware that despite the advances in determining the exact proportions of the monatomic elements composing the polyatomic substances the results were still based on somewhat arbitrary principles, Petit and Dulong believed that discovering additional exact properties of the combining elements would enable a more precise computation. Having devised an exact experimental method for determining the specific heats of various elements (defined at the time as the degree of heat required to raise the temperature of the weight of a given substance by one degree relative to that of water), they believed they had found such a property. Known as the Petit and Dulong Law that the monatomic structures of all the compound substances have exactly the same capacity for heat, this enabled them to calculate the approximate values of the atomic weights. And since the molecular formulas are determined by the interrelation of atomic weights and combining weights this provided another way of determining the molecular structures. In their judgment, “Whatever may be the final opinion adopted with regard to this relation, it can henceforth serve as a control of the results of chemical analysis. In certain cases it may even offer the most exact method of arriving at information about the proportions of certain combinations” (p. 307).
Using the Petit Dulong Law and his own method based on gas densities and their combining volumes, along with the analogous behavior of the elements in chemical reactions and in the structures of crystals, according to Nash, Berzelius by
judicious selections from among the various possibilities . . . he had, by 1840, arrived at atomic weights and molecular formulas that are in most cases in excellent agreement with those we now accept as correct. But alas, by this time a flood tide of skepticism was already lapping around the foundation of the atomic theory, and Berzelius’ fine work did not receive the attention it deserved. (p. 309)
There was a discrepancy between Avogadro’s belief that the relative weights of the heterogeneous particles could be inferred from the gas densities and the Petit Dulong Law that could not be applied to the gaseous elements.
Then in 1827 the French chemist J. B. A. Dumas devised a method for determining the gas densities at much higher temperatures that enabled him to study a much greater variety of substances at that higher temperature which in turn allowed him to reconcile the combining weights data derived from the Petit-Dulong Law with the relative gas densities obtained from Avogadro’s hypothesis that equal volumes of gases do not contain equal numbers of particles. Briefly, the reconciliation could be achieved if, in addition to agreeing there were
polyatomic molecules of the elements . . . it would now have to be further conceded that the polyatomic molecules of the different elements contain different numbers of the respective atoms . . . add[ing] the inability to explain why the molecules of different elements contain different numbers of their respective atoms. (p. 312; brackets added)
It took a little more than a quarter of a century before this phase of atomic physics reached a resolution by the Italian chemist Stanislao Cannizzaro. During that time there continued to be new discoveries despite the prevailing skepticism to accepting the truth of the atomic theory, such as the kinetic explanation of gas pressure as due to the mobility of dispersed particles in the gas and new evidence to support Avogadro’s theory of polyatomic particles. Accepting Avogadro’s law that equal volumes of similar gases contain the same number of particles of which some were polyatomic, Cannizzaro concluded that it should not be assumed that equal volumes of different gases contain the same number of basic particles. But if that were true, the weight of the atoms could not be inferred unless it was known how many atoms were contained in the volume, a near impossibility.
Thus he introduced a different procedure that involved weighing the densities of various compound gases containing the same element. Knowing the densities per unit volume of a number of compound gases containing that element, the weight of the element could be determined by what fraction of the weight of the compound was due to that element. Beginning with the smallest ratio, he found that in succeeding weightier compounds the ratios of that element were always in whole numbers. He then realized that he could calculate the relative weights of the elementary particles by comparing their ratios in the weights of the various compounds.
If the elementary compound contains one atom of that element this would give the atomic weight of that element. Then following Berzelius’ convention that established the atomic weight of hydrogen as the standard of 1, the weights of the other elements relative to hydrogen could be assigned: carbon as 12, oxygen as 16, sulfur as 32, and so forth. As these atomic weights agreed with those derived by the method of specific heats used by Petit and Dulong, this provided strong confirmation of the theory of atomic weights. Thus a half century later thanks to the efforts of preceding experimentalists, Cannizzaro proved Dalton’s belief in 1808 of “‘the importance and advantage of ascertaining the relative weights of the ultimate particles of both simple and compound bodies’” (pp. 318–19).
About two decades after Cannizzaro’s generally accepted determination of the atomic weights, in 1848 his research culminated in the independent publication respectively of the Periodic Law by Julius Lothar Meyer and the Periodic Table by Dmitri Ivanovich Mendeleev. Mendeleev’s table was published in Russian in April 1869 and though Meyer’s paper containing his Periodic Law was dated December 1869, it was not published in Germany until 1870. In a Faraday Lecture given to the Fellows of the Chemical Society of the Royal Institution in 1889 Mendeleev gave a succinct but comprehensive summary of what had been achieved up to that time and what could be anticipated in the future.
1. The elements, if arranged according to their atomic weights, exhibit an evident periodicity of properties.
2. Elements which are similar as regards their chemical properties have atomic weights which are either of nearly the same value (e.g., platinum, iridium, osmium) or which increase regularly (e.g., potassium, rubidium, caesium).
3. The arrangement of the elements, or of groups of elements, in the order of their atomic weights, corresponds to their so-called valences [the combining power of an element] as well as, to some extent, to their distinctive chemical properties—as is apparent, among other series, in that of lithium, beryllium, barium, carbon, nitrogen, oxygen, and iron (brackets added).
4. The elements which are the most widely diffused have small atomic weights.
5. The magnitude of the atomic weight determines the character of the element, just as the magnitude of the molecule determines the character of a compound.
6. We must expect the discovery of many yet unknown elements—for example, elements analogous to aluminum and silicon, whose atomic weight would be between 65 and 75.
7. The atomic weight of an element may sometimes be amended by a knowledge of those of the contiguous elements. Thus, the atomic weight of tellurium must lie between 123 and 126, and cannot be 128.
8. Certain characteristic properties of the elements can be foretold from their atomic weights.56
This enabled Mendeleev in 1895 to present his Periodic Table consisting of two columns, one vertical and the other horizontal. Both columns listed the elements under the headings given by Berzelius, with the vertical column presenting them according to their atomic weights and chemical properties while the horizontal column listed them in increasing numerals according to their atomic numbers as determined by the number of protons in their nucleus, beginning with hydrogen as 1 since it contains 1 proton. As now written, it appears as H but in a water molecule as H2O (with the subscript 2) because it contains two elements of hydrogen.
What a great progress had been made by chemists since Lavoisier’s explicit recognition of oxygen as the combustable gas and the resurrection of the ancient atomic theory of Democritus and Leucippus by Dalton culminating in Mendeleev’s Periodic Table. We now await the discovery of new elements and their molecular structure in compounds along with the interior structure of the atom.