ATOMIC WEIGHTS. Atomic weights have been defined as "the relative weights of the atoms of chemical elements referred to a common standard." This statement still serves as the simplest indication of the fundamental idea involved, although it now needs amplification. The concrete development of the idea was first effected in 1803 by John Dalton, an English chemist, when he converted the vague atomistic theory of the ancient Greeks into a highly valuable scientific asset by means of the concept of atomic weights. The chemical atomic theory thus initiated has been strengthened by modem investigation, and is to-day en trenched in a well-nigh impregnable position.
Atomic weights are quan tities of great practical and theoretical importance. They record the operation of the chemical law of definite combining propor tions ; hence they are the basis of quantitative chemical analysis, and are in everyday use throughout the world. Because of the parallelism between gravitational effect and inertia, they record also the relative masses of the atoms of the elements. They pos sess an extraordinary degree of definiteness, since the law of com bining proportions is one of the few known precise laws of the universe. Far deeper in meaning than the accidental astronomical "constants," such as the length of the day or the length of the year, the atomic weights of the simple elements and of the indi vidual isotopes (see IsoTOPES) stand out as among the peculiar and basic attributes of those 92 elementary substances of which everything is composed. Their interpretation is closely concerned with our inferences concerning the nature of matter.
Simple as the original concept of atomic weights seems to be, it nevertheless presents problems which are rather complex. For example, 22.997 grammes of sodium combine with 126.932 grammes of iodine to form sodium iodide. This ratio of the combining weights of these elements appears to be invariable. As Dalton pointed out, these weights must depend on the relative weights of the respective atoms; no other simple explanation is conceivable. There is in the experimental result, however, nothing which shows whether the sodium and iodine combine atom for atom, or whether one atom of sodium combines (for example) with two of iodine. Dalton himself perceived that this latter happening might in many cases occur ; indeed it is the essence of his Law of Multiple Proportions. There is now every reason to believe that in this particular case of sodium and iodine the atoms actually combine one to one and that the numbers given above represent really the relative weights of the atoms of sodium and iodine; but there are many less simple cases. For instance, 126-932 grammes of iodine combine with 20.035 grammes of calcium; here the latter number represents only half the atomic weight of calcium ; because every molecule of calcium iodide is believed on excellent evidence to contain two atoms of iodine for every atom of calcium (see VALENCY). Such a decision was beyond the reach of Dalton. It is based chiefly upon three subse quent discoveries to be briefly described.
In I 8 r 1 Count Amedeo Avogadro di Quaregna advanced the hypothesis, based upon Gay Lussac's Law of Volumes, that equal volumes of gases under like conditions of temperature and pressure contain the same number of mole cules, a molecule being defined as consisting usually of two or more atoms. This hypothesis (which has since been so amply confirmed as to become, in many minds, a statement of fact) fur nishes the most important means of deciding between the mul tiples or sub-multiples of the combining proportions which are to be taken as the atomic weights, because it fixes the molecular weights and formulas of volatile elements and compounds (see
