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Chemistry

Nineteenth-century Chemical Atoms

Nineteenth-century chemistry has often been described as a smooth sequence of discoveries that gradually established modern chemistry. Positivist historians were content with reporting landmark events such as John Dalton's atomic theory, Amedeo Avogadro's law, Dmitry Ivanovich Mendeleev's periodic system, and so on. In thus conveying a cumulative process, they not only distorted history but also tended to deprive past scientists of inner consistency because they failed to accept many of these "great discoveries."

One major feature of nineteenth-century chemistry is that the history of ideas did not follow the pathway that seems logical from a present-day perspective. For instance, the laws of chemical proportions did not follow from Lavoisier's program. Instead, they emerged from salt chemistry, more precisely from attempts to determine the weight or ponderous quantity of a base that could neutralize a definite quantity of acid. While the chemical revolution was drawing all attention toward Paris, two German chemists, Karl Friedrich Wenzel (1740–1793) and Jeremias Richter (1762–1807), were initiating a quantitative chemistry that they named stöichiometrie or stoichiometry (from the Greek stoicheion, or element). The main assumption that the properties of any substance depend upon the nature and proportion of its constituent elements opened up a research program for chemists whose key words were analysis and quantification (titration or dosage).

The field of stoichiometry was extended in 1802 when Joseph Louis Proust (1754–1826) applied Richter's notion of "equivalent," so far limited to reactions between acids and bases, to all combinations and formulated a general law: the relationships of the masses according to which two or several elements combine are fixed and not susceptible of continuous variations. While Berthollet questioned the generality of Proust's law of definite proportions, John Dalton (1766–1844), a professor at Manchester, made it the basis for his atomic hypothesis. He assumed that chemical combinations take place unit by unit, or atom by atom. He added a law of multiple proportions: when two elements form more than one compound, the weight proportions of the element that combines with a fixed proportion of another one are in a simple numerical ratio. Dalton's hypothesis rested on the assumption that atoms were solid and indivisible, that they were surrounded by an atmosphere of heat, that there were as many kinds of atoms as there were elements. Dalton's atoms were not the uniform and minute discrete units that structured all material bodies of ancient atomism. Rather, they were minimum and discrete units of chemical combination. And Dalton assumed that atoms would combine in the simplest way, that is, two atoms formed a binary compound. The main advantage of Dalton's hypothesis was that it allowed simple formulas. Instead of determining the composition of a body by percentages, chemists could express it in terms of constituent atoms thanks to the determination of atomic weights. Of course it was impossible to weigh individual atoms. Since Dalton could not determine this weight by using the neutrality of the compound like Richter, he elected a conventional standard. He chose hydrogen as the unit of reference. For instance, hydrogen and oxygen were known to form water, whose analysis gave the ratio 87.4 parts by weight of oxygen for 12.6 of hydrogen. If hydrogen has a weight equal to 1, the relative atomic weight of an atom of oxygen will be roughly 7. Shortly after Dalton's New System of Chemical Philosophy (1808), Joseph Gay-Lussac (1778–1850) announced that volumes of gas that combine with each other were in direct proportion and that the volume of the compound thus formed was also in direct proportion with the volume of the constituent gases. The volumetric proportions thus seemed to confirm Dalton's weight ratios.

To explain the convergence, both Amedeo Avogadro (1776–1856) in 1811 and André-Marie Ampère (1775–1836) in 1814 suggested that in the same conditions of temperature and pressure, equal volumes of gases contain the same number of molecules. In order to admit this simple hypothesis, however, both men had to admit a second hypothesis: that when two gases joined to form a compound, the integrating molecules should divide into two parts.

From a modern perspective this hypothesis was a big step forward because it suggested the distinction between atoms and molecules. Why, then, was it rejected in the 1830s and for several decades afterward by the most prominent chemists? The rejection of Avogadro's law is a classical topos for illustrating the opposition between presentist and historicist approaches. The alleged "blindness" of nineteenth-century chemists appears as a perfectly consistent attitude once one considers that Avogadro's hypothesis of molecules formed of two atoms of the same element stood in contradiction to the theoretical framework of chemical atomism. In Dalton's theory, such diatomic molecules were physically impossible because of the repulsion between the atmospheres of heat of two identical atoms, and they were theoretically impossible in the new electrochemical paradigm set up by Jöns Jacob Berzelius (1779–1848) on the basis of his experiments on electrolytic decomposition. Elements were defined by their electric polarity, and the intensity of the positive or negative charge determined the affinity between them. Molecules of two atoms of the same element were impossible because of the repulsion between two identical electrical charges.

The distinction between atom and molecule did not directly follow from Avogadro's law. Rather it was formulated by a young chemist, Auguste Laurent (1807–1853), trained in mineralogy and crystallography. He challenged Berzelius's electrochemical view. For Berzelius all compounds resulted from the electric attraction between two elements. This dualistic view was extended to organic compounds thanks to the notion of radical—for instance the benzoyl radical discovered by Leibig—a group of atoms that, like elements, persisted through reactions. In the 1830s Jean-Baptiste-André Dumas (1800–1884) prepared trichloracetic acid by the substitution of three atoms of chlorine to three atoms of hydrogen in acetic acid. His student Laurent noticed the similarity of properties between the two acids. Electronegative atoms of chlorine could replace electropositive atoms of hydrogen without changing the properties too much. He consequently developed a unitary theory that Dumas called "type theory" in 1838: in organic compounds there exist types that persist when elements are changed. This suggested that the properties of compounds depended more on the architecture of molecules than on the nature of constituent atoms. Type theory applied to the increasing crowd of organic compounds while inorganic chemistry was still ruled by dualism.

Charles Frédéric Gerhardt (1816–1856) attempted to re-unify chemistry by extending the type theory to all compounds, using analogy as a guide. Although he admitted that we do not know anything about the actual arrangement of atoms in a molecule, he defined three types—hydrogen, water, and ammonia—as an ideal taxonomic scheme. According to Gerhardt, all compounds derived from these three types.

The theory of type implicitly indicated that atoms of different elements had different combining powers or valences. Hydrogen, for example, has a valence of one, while oxygen has two and nitrogen three. It is worth emphasizing that although nineteenth-century organic chemists such as Gerhardt and August Kekulé (1829–1896) did not believe in the actual existence of atoms and molecules, they were able to invent structural formulas that allowed them to predict and create new compounds.

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