Atomic Theory.

Atomic Theory. What is known as the atomic theory is a theory as to the nature of the ultimate particles of matter, which, supported as it is by both chemical and physical evidence, has been of great service in the explanation of chemical facts, as well as in the progress of scientific chemistry.

Theoretical speculations as to the nature of the constitution of matter date from the earliest times of philosophy, but the gradual development of the atomic theory into its present form is owing to the accumulation of chemical and physical facts during a period of about a century. Matter has long been regarded as not being continuous, but as possessing a grained structure—i.e. as being made up of extremely minute particles. These particles, or groups of such particles, are supposed to be arranged in any substance at a certain average distance from one another, such average distance depending not only upon the physical state of the substance—i.e. whether solid, liquid, or gas—but also upon the temperature and pressure.

The honour of having first formulated an atomic theory based upon experimental evidence, obtained from his own investigations and from those of his predecessors and contemporaries, falls to Dalton. Prior to Dalton's first publication, at the beginning of this century, of his contributions to the subject, a good deal of investigation had been made leading up to the establishment of a rational atomic theory, although it had not been fully recognised as such at the time. It was known that chemical compounds contained their elements in certain definite and fixed proportions. Moreover, by finding the quantities of various acids required to neutralise a given quantity of a particular base, and of various bases required to neutralise a given quantity of a particular acid, it had been ascertained that numbers could be assigned to each acid and base which represented quantities which were chemically equivalent. Thus a quantity of an acid, A, combined with a quantity of a base, B, to form a neutral salt, AB; the same quantity of A combined with a quantity of another base, C, to form a neutral salt, AC; and the same quantity of B combined with another acid, D, to form a neutral salt, DB. The important point which had been established was, that the quantity of acid D in DB, and that of base C in AC, combined to form another neutral salt, DC. It was also known that the quantities of various metals dissolved by the same weight of an acid were capable of uniting with the same weights of oxygen. Further, it was known that several metals and other elements formed more than one compound with oxygen, the proportion of the latter being different in each. Dalton showed that when elements united with each other in two different proportions, these proportions were related to each other in a very simple way. Thus, he showed that a given weight of carbon united with a certain proportion of oxygen to form carbonic oxide, and with just twice as much to form carbonic acid; also that in olefiant gas and in marsh-gas the relation of the proportion of hydrogen for a given weight of carbon was as one is to two. Other examples were also known to him, notably in the case of compounds of oxygen and nitrogen, which showed similar simple relations. This discovery of Dalton's was what has been subsequently known as the law of multiple proportions, and, to explain it, Dalton reverted to the atomic hypothesis, which assumed that matter consists of atoms of different weights, those of the same element being all of the same weight.

Dalton regarded carbonic oxide as a compound of 1 atom of carbon and 1 of oxygen, and carbonic acid as a compound of 1 atom of carbon and 2 of oxygen; olefiant gas as a compound of 1 atom of carbon and 1 of hydrogen, and marsh-gas as a compound of 1 atom of carbon and 2 of hydrogen. He further assigned atomic weights or relative weights of the atoms, based upon the results of his own experiments, to each of the elements he had examined, assuming the atomic weight of hydrogen to be unity, as hydrogen combined with other elements in smaller proportion by weight than any other element. Although Dalton's atomic weights were far from accurate, yet his theory was sound that the observed law of multiple proportions could be satisfactorily explained by the atomic hypothesis.

Dalton further introduced a system of chemical notation, in which the atoms were represented by symbols.

Quickly following on Dalton's discoveries, was the discovery by Gay-Lussac of the simple relations of the volumes of gases which combine to form new compounds, and of the volume of the new gas produced. Thus Gay-Lussac proved that 2 volumes of hydrogen combined with 1 volume of oxygen to form 2 volumes of water vapour (all measured at the same pressure and temperature—these being of course such that the whole of the water formed was in the gaseous state); that 1 volume of hydrogen combined with 1 volume of chlorine to form 2 volumes of hydrochloric acid gas; and that similar simple relations existed in other cases. Gay-Lussac also noted that the numbers representing the atomic weights of elementary gases were simply related to the specific gravities of these gases. Indeed this follows directly from a consideration of his law of the simple relation of volumes, for, if 1 volume of hydrogen combines with 1 of chlorine to form hydrochloric acid, and if the gases combine atom for atom, which seems to be the simplest assumption, it is obvious that the weights of hydrogen and of chlorine in a given volume, at the same pressure and temperature, must be proportional to the atomic weights of these gases. In short, the two sets of numbers are identical if hydrogen be assumed as unity for both atomic weights and specific gravities. But further, it is also manifest that the specific gravity of the resulting gas must bear a simple relation to the specific gravities of its component gases.

Up to this period, although attempts had been made, notably by Dalton and Berzelius, to affix atomic weights to the elements, it had not been possible to do so with much certainty. An important step, however, towards definitely fixing the atomic weights of many elements was made by the experiments and observations of Dulong and Petit. These investigators called attention to the relation between the specific heats of the elements and their atomic weights. Generalising from their experiments on a number of the elements, Dulong and Petit concluded that what are now known as the atomic heats of all elements (that is their specific heats multiplied by their atomic weights) were approximately identical. This conclusion might also be stated otherwise—viz. that the capacity for heat of elementary atoms is identical. Some notable exceptions to this generalisation were subsequently found, but the law, as it is called, of Dulong and Petit, is substantially correct. To take some examples, the specific heat of potassium multiplied by its atomic weight ($166\; \backslash times\; 39$), gives as product 6.5. In the case of iron, the number is $6.3\; (112\; \backslash times\; 56)$, and of mercury $6.4\; (.0319\; \backslash times\; 200)$. These numbers are very nearly constant, but all are not quite so close. Lothar Meyer states (Moderne Theorien der Chemie, 4th ed. p. 93) that although most atomic heats lie between 6.1 and 6.5, some are as low as 5.2 and others as high as 6.9. The fact that the atomic weights used by Dulong and Petit were different from those now commonly adopted, does not alter the general conclusions to be drawn from their observations. The constant number representing the atomic heat was, of course, different from that given above. It has not been possible to determine directly the atomic heats of all the elements, but the atomic heats of a good many have been estimated by indirect means from the molecular heats (that is molecular weights $\backslash times$ specific heats) of their compounds, for it has been found to be at least approximately true that the molecular heat of a compound substance is equal to the sum of the atomic heats of the atoms which it contains. In other words, the product of molecular weight $\backslash times$ specific heat, divided by the number of atoms in the molecule, gives a constant number.

The variations in atomic heats of various elements are, it will be noticed, considerable, but still it is sufficiently clear that an atomic heat determination is a most valuable assistance in fixing which of two or three numbers, each of which would sufficiently express the chemical relations of an element, should be adopted to represent its atomic weight. Other considerations have led in some cases to the fixing of a certain atomic weight for an element, but with the exception of what is known as Avogadro's Law, these need not be particularised here. Avogadro's Law is of the first importance, although its bearing was not recognised until many years after its promulgation, which precedes historically that of the law of Dulong and Petit.

Avogadro distinguished between elementary atoms, or the smallest indivisible particles of an element, and molecules, or the smallest portion of a substance, possessing all the properties of the substance. His molecules are hence groups of 2 or more atoms, each group being capable of a separate existence. Avogadro's Law is based upon some considerations of the physical properties of gases, which he held received their simplest explanation by the assumption that a given volume of any gas, whether elementary or compound, contains the same number of molecules as the same volume of any other gas when measured at the same pressure and temperature. This law is in complete accord with the dynamical theory of gases. See article GAS.

If now Avogadro's Law be true, the relative densities of gases must represent their relative molecular weights whether the gases be elementary or compound. Hence we are led to suppose that the molecules of some elementary gases, as of compound gases, consist of 2 or more atoms. Thus 2 volumes of hydrogen unite with 1 of oxygen to form 2 volumes of water vapour. Now, if 2 volumes of hydrogen contain twice as many molecules as 1 volume of oxygen, and the resultant 2 volumes of water vapour contain as many molecules as the original 2 volumes of hydrogen, it is obvious that each molecule of oxygen must be split into 2, and that the molecule of oxygen must consist of at least 2 atoms. Again, 1 volume of hydrogen unites with 1 volume of chlorine to form 2 volumes of hydrochloric acid gas. The number of molecules at the end hence remains unchanged, but whereas each molecule then consists of hydrogen and chlorine, it is obvious that before the union the molecules must have consisted of hydrogen and hydrogen and of chlorine and chlorine respectively—i.e. each molecule of hydrogen and of chlorine must have consisted of at least 2 atoms. The molecules of a few metallic vapours consist only of single atoms. This is the case with mercury and zinc. The molecules of some non-metallic elements consist of more than 2 atoms. For instance, molecules of phosphorus and of arsenic consist of 4 atoms.

The laws of Dulong and Petit, and of Avogadro, constitute the main grounds for fixing the atomic weights as at present used by chemists. It would not be possible within the limits of such an article as this to enter into the varied evidence, derived from chemical properties of substances, for assigning particular atomic weights to certain elements. Some examples of the kind of evidence must suffice. For many years, and indeed until comparatively recently, the atomic weight of oxygen was all but universally stated as 8, compared with that of hydrogen as unity, instead of 16 as at present. Adopting this weight, the formula of water becomes HO, this being the simplest expression for the proportion of 1 part by weight of hydrogen to 8 parts by weight of oxygen. Now, we know that hydrogen and chlorine combine in the proportions by weight of 1 to 35·4 to form hydrochloric acid, and further, that not less than 35·4 parts of chlorine ever take the place of 1 part of hydrogen in combination. We further know a compound of potassium, oxygen, and hydrogen (potassium hydroxide or caustic potash), which contains these elements in the proportions by weight of 39, 16, and 1, respectively.

Marsh-gas contains carbon and hydrogen, and we know that it must contain at least 4 atoms of hydrogen, because, by suitable means, the hydrogen which it contains can be removed in four successive stages and its place taken by chlorine. In performing the first of these removals, or replacements as they are called, we find that 35·4 parts by weight of chlorine take the place of 1 part by weight of hydrogen. We do not know any means by which 8 parts of oxygen can be got to take the place of these 35·4 parts of chlorine, nor is there any instance known in the whole range of chemistry where 8 parts of oxygen take the place of that quantity of chlorine which combines with 1 part by weight of hydrogen.

On treating the first chlorine derivative of marsh-gas with caustic potash, we get potassium chloride, and the place of the chlorine is taken by 16 parts by weight of oxygen and 1 part by weight of hydrogen. The second chlorine derivative of marsh-gas (that in which half of the total hydrogen, or 2 atoms, is replaced by chlorine) has 2 parts by weight of hydrogen removed, and its place taken by 35·4 × 2 parts by weight of chlorine. Corresponding to this, however, there is an oxygen compound in which 35·4 × 2 parts of chlorine are replaced by 16 parts of oxygen. Further, the next chlorine derivative of marsh-gas has an oxygen compound corresponding to it, which, however, has the place of the 35·4 × 3 parts of chlorine, taken by 16 × 2 parts of oxygen and 1 part of hydrogen. Finally, the last chlorine derivative of marsh-gas has a corresponding oxygen compound in which the place of the 35·4 × 4 parts of chlorine is taken by 16 × 2 parts of oxygen.

It will be noticed that only in those cases where an even multiple of 35·4 parts of chlorine exists, is the place of the latter taken by oxygen alone. In the other cases chlorine is replaced by oxygen and hydrogen. Thus we see that we do not get 1 atom of oxygen, weighing 8, simply taking the place of 1 of hydrogen, weighing 1, but 16 parts by weight of oxygen take the place of 2 of hydrogen. These facts can all be easily explained by assuming the atomic weight of oxygen to be 16, when the formula for water becomes H₂O, and the other facts mentioned above may be expressed symbolically as under :

Marsh-gas has the composition CH₄. Its first and succeeding chlorine derivatives are respectively CH₃Cl, CH₂Cl₂, CHCl₃, CCl₄; and the oxygen compounds corresponding are CH₃OH, CH₂O, CHOOH, CO₂.

Represented by graphic formulæ, these are :

Such is an example of the nature of the very varied chemical evidence for the fixing of a particular atomic weight. Below is a list of the elements at present known with certainty, and of their atomic weights as fixed by the various kinds of evidence obtained by very numerous, and in many cases, varied experiments. The old atomic weights are given in the last column for comparison. In most cases these are just half the new atomic weights (see the list below). The new numbers are now all but universally employed by chemists. See also ARGON.

It has been stated in the foregoing that 1 atom of oxygen is generally regarded as combining with 2 of hydrogen to form water, and that 1 atom of chlorine combines with 1 of hydrogen to form hydrochloric acid. These are not the only variations in the number of hydrogen atoms which unite with 1 atom of another element to form a compound. Thus in ammonia we have 1 atom of nitrogen united with 3 of hydrogen; in marsh-gas, 1 of carbon united with 4 of hydrogen. Similarly, in many metallic chlorides the number of atoms of chlorine for 1 of metal varies, as for instance, in KCl, CaCl₂, BiCl₃, SnCl₄, &c. Again, the oxides corresponding to these chlorides are respectively K₂O, CaO, Bi₂O₃, SnO₂, &c.; while we know other compounds in which 1 atom of potassium exists in combination with 1 of bromine or iodine; 1 atom of calcium with 1 of sulphur or 2 of fluorine; 1 atom of antimony with 3 of hydrogen; or 2 of antimony with 3 of sulphur, and so forth. From these examples it will be clear that the atoms of the elements possess different exchange values in replacing one another in chemical compounds. This exchange value of an element is called its valency. 1 atom of chlorine combines with 1 of hydrogen or 1 of potassium, and these three elements are said to be monovalent. Calcium, oxygen, and sulphur are divalent; nitrogen (in ammonia at least), antimony, and bismuth, trivalent; carbon and tin, tetravalent, &c. Some elements appear to have more than 1 valency, as for instance, phosphorus, which forms 2 chlorine compounds, PCl₃ and PCl₅. Much controversy has taken place as to fixing the valency of such elements, some chemists contending that valency is invariable, and that where we find in some compounds an apparent valency less than the highest known, it is to be explained by supposing a part of the combining power to remain unsatisfied. It seems simplest to assume that the valency of an element may differ in different sets of compounds. The amount of importance attached to the valency of an element, and to the question as to whether or not the valency of an element is variable, is not now so great as it was formerly.

What is known as the equivalent of an element is that proportion of an element which is capable of taking the place of 1 atom of a monovalent element. Thus, 39 parts of potassium combine with 35.4 of chlorine to form potassium chloride. The equivalent quantity of calcium is not 40 parts (for that quantity of calcium combines with 35.4 × 2 parts of chlorine) but $\backslash frac\{40\}\{2\}\; =\; 20$ parts. 20 is therefore the equivalent weight of calcium. Faraday has shown that those quantities of various elements are chemically equivalent which are separated from various electrolytes by the passage through them of the same quantity of electricity. See ELECTRICITY, GALVANISM.

When the elements are arranged in progressive order of their atomic weights, a certain regularity of succession of chemically analogous elements is observed. Attention was drawn to this by Newlands, and the matter was more fully developed by Lothar Meyer and especially by Mendeleëff, who formulated the Periodic Law which states that the properties of an element are a function of its atomic weight. The regularity of the recurrence of the cycles or periods is so striking that the appearance of gaps in the list led to the prediction not only of the existence of numerous undiscovered elements, but of the probable properties these elements would be found to possess when discovered. These predictions have been very fully verified in the case of two or three recently discovered elements, notably gallium and germanium.

A list of the elements (see ARGON) arranged according to the periodic system is given below :

In this table each horizontal line forms a period or cycle of the elements, and the various groups arranged in columns contain elements which ex- hibit chemical analogies and similarities which are often very striking. The elements placed at the left of the columns possess greater resemblances to each other than to those placed at the right, and vice versa. Those elements included under Group VIII. are exceptional. In each series where they occur we find a set of three elements which Mendeleëff calls transition elements, as they occur between the end of one and the beginning of the next long period or term of two series. Numerous gaps still exist in the table, indicating the places which, doubtless, elements which are still to be discovered will hereafter occupy.

A marked periodicity is noticed in the atomic volumes of the elements (that is, in the quotients of the specific gravities of the solids divided by their atomic weights) in each of Mendeleëff's series. This subject, however, cannot be pursued here, as its discussion would lead too far. For further information concerning it and the periodic law generally, the reader is referred to some of the larger modern text-books of chemistry. See also the article CHEMISTRY in this work.