Chalmers, Alan, "Atomism from the 17th to the 20th Century", The Stanford Encyclopedia of Philosophy  (Winter 2014 Edition), Edward N. Zalta (ed.): <http://plato.stanford.edu/archives/win2014/entries/atomism-modern/>.

4. Chemical Atomism in the Nineteenth Century

4.1 Dalton's Atomism…Dalton was able to take for granted assumptions that had become central to chemistry since the work of Lavoisier. Chemical compounds were understood as arising through the combination of chemical elements, substances that cannot be broken down into something simpler by chemical means. The weight of each element was understood to be preserved in chemical reactions. By the time Dalton (1808) made his first contributions to chemistry the law of constant composition of compounds could be added to this. [Joseph-Louis] Proust had done much to substantiate experimentally the claim that the relative weights of elements making up a chemical compound remain constant independent of its mode of preparation, its temperature and its state.

The key assumption of Dalton's chemical atomism is that chemical elements are composed of ‘ultimate particles’ or atoms. The least part of a chemical compound is assumed to be made up of characteristic combinations of atoms of the component elements. Dalton called these ‘compound atoms’. According to Dalton, all atoms of a given substance whether simple or compound, are alike in shape, weight and any other particular. This much already entails the law of constant proportions. Although Dalton himself resisted the move, Berzelius was able to show how Dalton's theory can be conveniently portrayed by representing the composition of compounds in terms of elements by chemical formulae in the way that has since become commonplace. Hereafter this device is employed using modern conventions rather than any of the various ones used by Berzelius and his contemporaries,

As Dalton stressed, once the chemical atomic theory is accepted, the promise is opened up of determining the relative weights of atoms by measuring the relative weights of elements in compounds. If an atom of element A combines with an atom of element B to form a compound atom of compound AB, then the relative weights of A and B in the compound as measured in the laboratory will be equal to the relative weights of atoms of A and B.

However, there is a serious under-determination of relative atomic weights by measurements of combining weights in the laboratory. If the compound atom in our example were A2B rather than AB then the relative atomic weight of B would be twice what it would be if the formula were AB. Dalton himself attempted to resolve this problem with a simplicity assumption. Formulae were always to take the simplest form compatible with the empirical data. If there was only one compound of A and B known then it was assumed to be AB, whilst if there were two then a more complicated compound, A2B or AB2 became necessary. As is illustrated by the latter example, as well as the problem of the truth of the simplicity assumption there was the problem of its ambiguity. Chemical atomists were to struggle for several decades with various solutions to the problem of arriving at definitive formulae and relative atomic weights, as we shall see.

This deficiency of Dalton's atomism aside, links were forged between it and experimentally determined combining weights that went beyond the law of constant proportions to include the laws of multiple and reciprocal proportions. If two elements combine together in more than one way to form compounds, as is the case with the various oxides of nitrogen and carbon, for example, then Daltonian atomism predicts that the weights of one of the elements in each compound, relative to a fixed weight of the second, will bear simple integral ratios to each other. This is the law of multiple proportions, predicted by Dalton and soon confirmed by a range of experiments. Daltonian atomism also predicts that if the weights of elements A and B that combine with a fixed weight of element C are x and y respectively, then if A and B combine to form a compound then the relative weights of A and B in the compound will be in the ration x:y or some simple multiple of it. This law was also confirmed by experiment.

There is a further component that needs to be added to the content of early atomic chemistry, although it did not originate with Dalton, who in fact did not fully embrace it. Gay Lussac discovered experimentally that when gases combine chemically they do so in volumes that bear an integral ratio to each other and to the volume of the resulting compound if gaseous, provided that all volumes are estimated at the same temperature and pressure.

For instance, one volume of oxygen combines with two volumes of hydrogen to form two volumes of steam. If one accepts atomism, this implies that there are some whole-number ratios between the numbers per unit volume of atoms of various gaseous elements at the same temperature. Following suggestions made by Avogadro and Ampere early in the second decade of the nineteenth century, many chemists assumed that equal volumes of gases contain equal numbers of ‘atoms’, with the important implication that relative weights of atoms could be established by comparing vapour densities.

As Dalton clearly saw, this can only be maintained at the expense of admitting that ‘atoms’ can be split. The measured volumes involved in the formation of water, for example, entail that, if equal volumes contain equal numbers of atoms then a water ‘atom’ must contain half of an oxygen ‘atom’. The resolution of these problems required a clear distinction between atoms of a chemical substance and molecules of a gas, the grounds for which became available only later in the century. This problem aside, the empirical fact that gases combine in volumes that are in simple ratios to each other became a central component of chemistry, although it should be noted that at the time Gay Lussac proposed his law, only a small number of gases were known to chemists. The situation was to change with the development of organic chemistry in the next few decades.

4.2 The Status of Daltonian Chemistry

…If Dalton's atoms were regarded as ontologically basic, then there needed to be as many kinds of matter as there are chemical elements. Further, atoms of each element needed to possess a range of characteristic properties to account for chemical combination as well as physical aggregation and other physical properties. As a philosophical theory of the ultimate nature of material reality, Daltonian atomism was not a serious contender and was not treated as such. A

more significant issue is the status of Daltonian chemistry as an experimental science. To what extent was Daltonian chemistry borne out by and able to fruitfully guide experiment?

A basic issue concerning the empirical statues of Daltonian atomism was already pinpointed in an early exchange between Dalton (1814) and Berzelius (1815). Dalton was keen to present himself as the Newton of chemistry. In his view, just as Newton had explained Kepler’s laws with his new mechanics, so he, Dalton, had explained the laws of proportion with his atomism. Without atomism the joint truth of the three laws of proportion is a mystery. Berzelius questioned the experimental grounds for assuming anything stronger than the laws of proportion, since, he argued, all of the chemistry could be accommodated by the latter. That is, nothing testable by the chemistry of the time follows from Dalton's atomic theory that does not follow from the laws of proportion plus the experimental law of combining volumes for gases.

Berzelius (1814) expressed his version of Daltonian chemistry using formulae. Dalton had pictured atoms as spheres and compound atoms as characteristic arrangements of spheres. Berzelius claimed that the two methods were equivalent but that his method was superior because it was less hypothetical…His point was that the testable empirical content of the two theories were equivalent as far as the chemistry of the time was concerned, but that his version was less speculative because it did not require a commitment to atoms. The symbols in Berzelian formulae can be interpreted as representing combing weights or volumes without a commitment to atoms.

A Daltonian atomist will typically take the hydrogen atom as a standard of weight and the atomic weight of any other element will represent the weight of an atom of that element relative to the weight of the hydrogen atom. On such an interpretation the formula H2O represents two atoms of hydrogen combined with one of oxygen. But, more in keeping with the weight determinations that are carried out in the laboratory, it is possible to interpret atomic weights and formulae in a more empirical way. Any sample of hydrogen whatever can be taken as the standard, and the atomic weight of a second element will be determined by the weight of that element which combines with it. The formula H2O then represents the fact that water contains two atomic weights of hydrogen for every one of oxygen. Of course, determining atomic weights and formulae requires some decision to solve the under-determination problem, but that is the case whether one commits to atoms or not.

Berzelius was right to point out that as far as being supported by and serving to guide the chemistry of the time was concerned, his formulation using formulae served as well as Dalton's formulation without committing to atomism. What follows from this will depend on one's stand on confirmation and explanation in science. A strong-minded empiricist might conclude from Berzelius’s observation that Dalton's atomism had no place in the chemistry of the time. Others might agree with Dalton that the mere fact that Dalton's theory could explain the laws of proportion in a way that no available rival theory could constituted a legitimate argument for it in spite of the lack of evidence independent of combining weights and volumes. Atomism could be defended on the grounds that attempts to articulate and improve it might well fruitfully guide experiment in the future and lead to evidence for it that went beyond combining weights and

volumes. But such articulations would clearly require properties to be ascribed to atoms in addition to their weight.

Berzelius himself took this latter option. He developed an atomic theory that attributed the combination of atoms in compounds to electrostatic attractions. He developed a ‘dualist’ theory to bring order to compounds involving several types of molecules. For instance, he represented copper sulphate as (CuO + SO3). Here electropositive copper combines with electronegative oxygen but in a way that leaves the combination slightly electropositive, whereas electropositive sulphur combines with oxygen in a way that leaves the combination slightly electronegative. The residual charges of the ‘radicals’ as they became known could then account for their combination to form copper sulphate.

Berzelius's conjectures about the electrical nature of chemical combination owed their plausibility to the phenomenon of electrolysis, and especially the laws governing it discovered by Faraday, which linked the weights of chemicals deposited in electrolysis to chemical equivalents. But evidence for the details of his atomistic theory independent of the evidence for the experimental laws that the theory was designed to support was still lacking. Contemporaries of Berzelius proposed other atomic theories to explain electrical properties of matter. Ampère proposed electrical currents in atoms to explain magnetism and Poisson showed how electrostatic induction could be explained by assuming atomic dipoles. In each of these cases some new hypothesis was added to atomism for which there was no evidence independent of the phenomenon explained…

4.3 Progress in Organic Chemistry Using Chemical Formulae

The period from the third to the sixth decades of the nineteenth century witnessed spectacular advances in the area of organic chemistry and it is uncontroversial to observe that these advances were facilitated by the use of chemical formulae. Inorganic chemistry differs from organic chemistry insofar as the former involves simple arrangements of a large number of elements whereas organic chemistry involves complicated arrangements of just a few elements, mainly carbon, hydrogen, oxygen and to a lesser extent, nitrogen.

It was soon to become apparent that the specification of the proportions of the elements in an organic compound was not sufficient to identify it or to give an adequate reflection of its properties. Progress became possible when the arrangements of the symbols representing the elements in formulae were deployed to reflect chemical properties. The historical details of the various ways in which chemical properties were represented by arrangements of symbols are complex. (For details see Rocke (1984) and Klein (2003)). Here we abstract from those details to illustrate the kinds of moves that were made.

The simplest formula representing the composition of acetic acid is CH2O using modern atomic weights. This formula cannot accommodate the fact that, in the laboratory, the hydrogen in acetic acid can be replaced by chlorine in four distinct ways yielding four distinct chemical compounds. Three of those compounds are acids that have properties very similar to acetic acid, and in which the relative weights of chlorine vary as 1:2:3. The fourth compound has the properties of a salt

rather than an acid. These experimental facts can be captured in a formula by doubling the numbers and rearranging the symbols, so that we have C2H4O2, rearranged to read C2H3O2H. The experimental facts can now readily be understood in terms of the substitution of one or more of the hydrogens by chlorine, with the three chloro-acetic acids represented as C2H2ClO2, C2HCl2O2H and C2Cl3O2H and the salt, acetyl chloride, as C2H3O2Cl. Such formulae came to be known as ‘rational formulae’ as distinct from the ‘empirical formula’ CH2O. Representing the replacement of one element in a compound by another in the laboratory by the replacement of one symbol by another in a chemical formula became a standard and productive device that was to eventually yield the concept of valency in the 1860s. (Oxygen has a valency of two because two hydrogens need to be substituted for each oxygen.)

Other devices employed to fashion rational formulae involved the notion of a radical, a grouping of elements that persisted through a range of chemical changes so that they play a role in organic chemistry akin to that of elements in inorganic chemistry. Series of compounds could be understood in terms of additions, for example to the methyl radical, CH3, or to the ethyl radical, C2H5, and so on. ‘Homologous series’ of compounds could be formed by repeatedly adding CH2 to the formulae for such radicals so that the properties, and indeed the existence, of complex compounds could be predicted by analogy with simpler ones. Another productive move involved the increasing recognition that the action of acids needed to be understood in terms of the replacement of hydrogen. Polybasic acids were recognised as producing two or more series of salts depending on whether one, two or more hydrogens are replaced. Yet another important move involved the demand that rational formulae capture certain asymmetric compounds, such as methyl ethyl ether, CH3C2H5O, as distinct from methyl ether, (CH3)2O, and ethyl ether, (C2H5)2O. By 1860, the idea of tetravalent carbon atoms that could combine together in chains was added. By that stage, the demand that rational formulae reflect a wide range of chemical properties had resulted in a set of formulae that was more or less unique. The under-determination problem that had blocked the way to the establishment of unique formulae and atomic weights had been solved by chemical means.

4.4 Implications of Organic Chemistry for Atomism

The previous section was deliberately written in a way that does not involve a commitment to atomism. It is possible to understand the project of adapting rational formulae so that they adequately reflect chemical properties by interpreting the symbols as representing combining weights or volumes as Berzelius had already observed in his early debates with Dalton…

A number of chemists involved in the early advances of organic chemistry who did adopt atomism expressed their ontological commitment to ‘chemical atoms’. In doing so they distinguished their theories from those brands of physical atomism that were in the tradition of mechanical or Newtonian atomism and which sought to explain phenomena in general, and chemistry in particular, by reference to a few physical properties of atoms. Chemical atoms had more in common with natural minima insofar as they were presupposed to have properties characteristic of the substances they were atoms of. Chemical atomism lent itself to the idea that

it was developments in chemistry that were to indicate which properties were to be attributed to chemical atoms, as exemplified in the path that led to the property ‘valency’…

Dalton's atomism had given a line on just one property of atoms, their relative weight. But it is quite clear that they needed far richer properties to play [their] presumed role in chemistry. It was to be developments in chemistry, and later physics, that were to give further clues about what properties to ascribe to atoms. (We have seen how chemists came to ascribe the property of valency to them.) There was no viable atomistic theory of chemistry in the nineteenth century that was such that chemical properties could be deduced from it…

The emergence of unique atomic weights and the structural formulae that organic chemistry had yielded by the 1860s were to prove vital ingredients for the case for atomism that could eventually be made. But there are reasons to be wary of the claim that atomism was responsible for the rise of organic chemistry and the extent to which the achievement improved the case for atomism needs to be elaborated with more caution that is typically the case…

There is a case for claiming that correct atomic weights were the outcome of, rather than a precondition for, progress in organic chemistry prior to 1860. After all, the majority of the formulae productively involved in that dramatic progress were the wrong formulae from a modern point of view! For instance, use of homologous series to project properties of lower hydrocarbons on to higher ones are not affected if the number of carbon atoms in the correct formulae are doubled, which results from taking 6 as the relative atomic weight of carbon, as many of the contemporary organic chemists did.