The properties of both compounds and elements, especially the relative masses of the different elements which made up a compound, were studied extensively in the eighteenth century. Correspondence between chemists, formalized by publication of experimental results in the journals of scientific societies, led to the compilation of tables of composition and properties of pure compounds. It remained for the Quaker schoolmaster John Dalton to place these painstakingly collected measurements and empirical laws on the theoretical basis known as the atomic theory of matter.
The understanding of the nature of matter as molecules was first set down by John Dalton (1766-1844). Dalton lived in Manchester, England, and from 1793 earned his living as a private tutor there. In the course of his studies and teaching, as well as in discussions with other members of the Literary and Philosophical Society of Manchester, he developed the ideas which led to his formulation of the atomic theory of matter in 1805.
Quantitative chemical measurements, especially those of mass, had given rise to several empirical laws which preceded the theoretical basis given to them by John Dalton. Dalton was familiar with these laws and also with the laws which had been discovered to describe the behavior of gases as we shall see in later sections. It was to explain these empirical laws that he developed the atomic theory of matter.
Probably the best-known and most widely accepted of the emprical laws known to Dalton was the law of conservation of (total) mass, which had been used by A. L. Lavoisier in 1798. It was probably first assumed, then verified by a few experiments, although it can only be truly verified by very precise experiments using sealed systems such as were carried out some fifty to one hundred years later. It had generally been accepted by scientists, however, even prior to the work of Lavoisier. The law of conservation of mass can be stated as follows: The total mass of the reactants in any chemical reaction is exactly equal to the total mass of the products.
The second empirical law known to Dalton was the law of constant composition of compounds, most ably stated by J. L. Proust (1755 - 1826) in his publications in J. Physique (Madrid) in 1802 - 1808. This law, clearly established by the long controversy between J. L. Proust and C. Berthollet, was a statement of the observed fact that any pure substance has a fixed composition in terms of the chemical elements.
For example, there are three compounds known which contain only iron and oxygen. They have different chemical compositions and different physical properties, but for each compound the properties, and the composition, do not differ between samples obtained from Australia to Zanzibar or, indeed, from the earth or the moon. In other words, the mass ratios or percentage composition by mass of any compound is characteristic of that compound and cannot be altered without changing the other properties which are also characteristic of the compound. The law of constant composition of compounds can be stated as follows: For every pure compound the mass ratios of its constituent elements are constant.
The empirical law of multiple proportions was apparently developed by Dalton himself around 1804. Again, it was probably assumed, then verified by a few experiments; it became public in 1807. The law of multiple proportions can be stated as follows: When any two elements are observed to form more than one compound between them, the mass ratios in one compound will be related to the mass ratios in the other in the proportions of small whole numbers.
The determination of the empirical mass ratios in compounds should ideally be done by decomposition to the elements (elemental analysis). However, in the eighteenth and early nineteenth centuries it was often impossible to effect the decomposition quantitatively, so many of the ratios were determined by synthesis from pure elements. Prior to the development of electrolysis, decomposition of water to hydrogen plus oxygen was not possible, but preparation or synthesis of water from hydrogen and oxygen was possible. When hydrogen was burned in excess oxygen water of constant composition was always evolved. The water could be weighed directly by absorption, the hydrogen used could be obtained by difference, and the oxygen used was established assuming conservation of mass. It can then be calculated that water is 11.19% hydrogen and 88.81% oxygen by mass. Dalton assumed that hydrogen should be assigned an atomic mass of one because it was the lightest element known. The mass of oxygen is then 88.81/11.19 = x/1, where x is the atomic mass of oxygen, and the value of 7.937 obtained was rounded to eight.
The decomposition of Fe2O3 was likewise impractical, but the reaction (rusting or oxidation) of iron in the presence of excess oxygen is simple and yields a compound of constant composition. The original iron can be weighed, as can the resulting iron oxide. The oxygen can be obtained by difference since conservation of mass is assumed. The compound has the constant composition Fe 69.9%, O 30.1% by mass. The obviously different compound FeO cannot be prepared by synthesis from the elements but must be tackled by analysis: reduction with hydrogen, yielding water and iron metal. From the mass of the original iron oxide, the mass of the water produced, and the mass of the resulting iron, the composition by mass of FeO is calculated to be 72.7% Fe and 27.3% oxygen. The atomic mass of iron is then 72.7/27.3 = x/8, or 27.9 units.
The units of atomic mass are now called unified atomic mass units, abbreviated u or amu. Biochemists often call them daltons in honor of John Dalton.