Page:The New International Encyclopædia 1st ed. v. 04.djvu/644

* CHEMISTRY. 562 CHEMISTRY. 32 parts of oxygen ; water of 2 parts of liyilrojjcn and 10 parts of oxygen; marsh gas of 12 parts of carbon and 4 parts of hydrogen; defiant gas (ethylene) of 24 parts of carbon and 4 parts of hydrogen, etc. The combining weights of the three elements niontionod are, approximately: carbon, 12; hydrogen, 1: oxygen, 16; and it is easy to see that in stating the composition of our compounds we have been able to use either these numliers or simple multiples of them. Dal- ton was led to the discovery of this law by the liypothesis according to ^vhi<•h all matter is made up of 'atoms' — i.e. of minute jiarticles incapable of further subdivision. The atomic theory, based on this hypothesis, comprises the following as- sumptions: The atoms of any given element are identical; the atoms of diflVrcnt elements are difl'erent and have difl'erent weights; by the force of diemical altinity several atoms may be held in combination, forming a particle, or 'molecule,' of a compound, and veiy large num- bers of molecules are necessai-y to make up even the smallest amounts of compounds which we are actually capable of handling. In accord- ance with these assumptions, let .1/ stand for the numljer of molecules making up a certain amount of some compound containing two ele- ments; and let the amount of the first element in it be o, the amount of the second element 6. Let, furtlicr. A stand for the weight of a single atom of one of the elements, and » stand for the number of such atoms contained in a single molecule of the compound. Then, evidently, the weight a of the first element in the compound equals .1/ X « X A. Similarly, if .1' stand for the weight of a single atom of the second ele- ment, and n' for the number of its atoms con- tained in a single molecule of the compound, then, evidently, the total weight b of the second element in M molecules equals if X n' X A'. We therefore have n: 6:: ^rnA: Mh'A' or, o: 6: : nA : n'A' In case a molecule of the compound should contain eipial numbers of atoms of the two ele- ments, then n = n', and hence n: 6:: A: A' But these proportions tell us that the weights (a and h) of the elements in a compound are j)roportioiial either to the weight-s of single atoms {A and A'), or to multiples (ii.l and n'A') of those weiglits. Thus the fundamental law of chemical composition follows as a direct consequence from, and is. hence, completely ex- plained by, the atomic hypothesis, without which it would be a mystery. The 'combining weights' mentioned above in connection with our state- ment of the law are seen, in the light of the hypothesis, to represent the relative weights of the atoms themselves, and are therefore termed atomic iicightfi. But while the fundamental assumptions of the atomic theory thus establish a general relation be- tween the quantitative composition of substances and the relative weights of atoms, they do not furnish a sufficient basis for determining these relative weights in an unequivocal manner. When we apply the above proportions to some given substance — say. water— for the purpose of determining the atomic weights of its ele- ments, we find ourselves compelled to make some additional assumption. Indeed, chemical analy- sis shows that w^ater contains 11.1 per cent, of hydrogen and 8S.il per cent, of oxygen. We therefore have (i:h = 11.1:88.9 = 1:8 (nearly) and hence itA : n'A' : : 1 : S, where .1 and .1' are, respectively, the weights of
 * ^ingle atoms of hydrogen and oxygen, while

ji and )i' are, respectively, the numbers of atoms of these elements in a molecule of water. What we are after is the ratio .1: .1' — i.e. the relative weights of single atoms; but this we evidently cannot find unless we assign some numerical value to the ratio n : n'. Dalton assumed that a molecule of water is made up of one atom of hydrogen and one of oxygen — i.e. n = ji' ^ 1, and therefore he found A: A':: 1: 8, i.e. an atom of oxygen is 8 times as heavy as an atom of hydrogen. (In reality, Dalton thus obtained, for the atomic cight of oxygen, the ligure (i : but this was due to liis imperfect knowl- edge of the proportion of hydrogen and oxygen in water). Dalton's assumption was quite arbitrary. But in subsequent years, as the substances kno«ii grew numerous and complex, diemists began to feel the want of some general theoretical prin- ciple which would render arbitrary, and heiue confusing, assumi)tions unnecessary. Then (iier- liardt and Cannizzaro enriclied Dalton's atomic theory by adding to it a jirinciple which had been eimnciated by Avogadro as early a.s 1811, but which had remained unemployed as long as it was not urgently needed. According to Avo- gadro, equal volumes of different gases contain e(|ual numbei-s of molecules if the temperatures and pressures of the ga.ses are the same. This theoretical principle, and its use in determining the atomic weights of the elements, have been expbiincd at some length under Atomic Wkiohts and .voG.m«)'.s Ri'le (qq.v.), and require no further discussion here. Suffice it to state that it forms part of the very foundation of the pres- ent atomic and molecular thcoiv. and that it is involved in the discussion of most, if not all, problems of modern chemistry, A still further addition was made to the fun- damental hypotheses of the atomic theory before it attained its maximum of possibilities. This last addition, grndually incorporated during the second half of the Xinetwnth Century, consists of certain assumptions concerning the combin- ing forces of atoms, the number of such forces ]icculiar to the atom of each element, and the directions in which those forces act. (See below, under Chemical Formulas : and see the articles V.mj;nct; Carhon Compounds; and Stereo- CUKMISTKY.) These assumptions, forming the so-called 'doctrine of valency.' were adopted mainly because of the necessity of explaining the isomerism of organic compounds — i.e. the fact that quite different compounds may have exactly the same composition. And it was mainly when fortified by these assumptions that the atomic theory enabled us to know compounds before they have actually been found in nature or in a chemical laboratory. "But," says Ostwald. "it seems as if the adapf> ability of the atomic hypothesis is near exhaus- tion, and it is well to realize that, according to the lesson repeatedly taught by the history of science, such an end is sooner or later inevit- able." In the future, he believes, chemists will