Three Scientific Revolutions: How They Transformed Our Conceptions of Reality (17 page)

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 (H
²
O).
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 CO
₂
or NH
₃
. 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 (H
²
O), ammonia (NH
³
) nitrous oxide (N
²
O) 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.