Molecule. The smallest mass of any substance which can exist in the free state without modification, through subdivision, of its chemical constitution or of its physical properties. Molecules are held to be of definite size and mass; and any piece of apparently homogeneous matter is held to be made up of molecules or granules which conjointly make up the aggregate mass. That matter, apparently homogeneous, is really granular in this sense, appears firstly from the following physical considerations: (1) Compressibility and porosity show that matter does not fill space; (2) dispersion of light by the prism would be impossible if the glass were homogeneous; (3) chemical combination is attended with the absorption or evolution of limited amounts of heat; (4) soap films require energy to stretch them, and if they could be attenuated beyond a certain limit the energy applied in stretching them would be more than sufficient to volatilise them; (5) the kinetic theory of gases, when numerically worked out, shows the number of molecules in a given space to be limited (see MATTER). These considerations lead to the conclusion that it is impos- sible to have a sheet of matter thinner than from cm. to cm. thick; and that a figure of this order represents the diameter of the space occupied by one molecule. From the chemical point of view, the laws of definite proportion and of multiple proportion find their best explanation in the atomic theory; but the chemical atom, which is active, is not the same thing as the physical molecule, which is chemically somewhat inert; as we see, for example, in the contrast afforded by the mutual combination of oxygen and hydrogen when in the nascent state, or when exposed to high temperature, and their want of combination when simply mixed as free gases at ordinary temperatures.
The kinetic theory of gases—that gases consist of molecules in a continual condition of mutual collision and rebound—leads to the theorem (Avogadro's Law) that all gases contain within the same volume, under equal conditions of temperature and pressure, the same number of molecules; and if we assume each molecule to be made up of the masses of a group of chemical atoms sufficient to make up the smallest possible free mass of the chemical substance, then this law is in perfect accord with the observations of Dulong and Petit on the relative densities of gases, which are proportional to the weights of the molecules, as derived from the known atomic weights or combining masses of the elements (see CHEMISTRY).
For this reason it has been usual to ascertain the molecular weight of substances by observation of their gas or vapour density, the molecular weight of hydrogen being reckoned as 2, or that of oxygen as 32. Thus, hydrochloric acid gas, which is 18.25 times as heavy as hydrogen gas, has a molecular weight 36.5, which is equal to the sum of the atomic weights of chlorine (= 35.5) and of hydrogen (= 1). Hydrogen itself has a molecular weight = 2, not = 1, because its molecule is held to be made up of two atoms, each of which has an atomic weight equal to unity; and so with other elementary gaseous or volatile substances. There are abnormalities, however; sulphur vapour has so great a density as to show its molecule to consist, below a red heat, of six atoms, while at higher temperatures it consists of two; iodine vapour at lower temperatures has two, at higher one atom in the molecule; mercury, cadmium, zinc, potassium, and sodium molecules contain one atom each; chlorine and bromine molecules partially break up into single atoms when heated. The molecules of liquids are probably compound; those of solids are almost certainly so, as allotropic modifications and variations in crystalline form show, though the vapours, produced from these various forms, if the solids be volatile, are identical. Other and more convenient means of measuring molecular weights have recently become known. These are—(1) a solid dissolved in a liquid lowers the vapour pressure (or 'vapour tension') of that liquid by an amount which bears to the original vapour pressure at the particular temperature of experiment the same ratio as the number of molecules of the dissolved substance does to the total number of molecules in the solution (Raoult); (2) the freezing-point of a solution is lowered in a similar ratio (Raoult); (3) osmotic pressure (see OSMOSE) is the same as that which would be exercised by the substance dissolved if it were transformed into a 'gas' and made to occupy alone the space occupied by the solution (Van t'Hoff); whence the specific gravity of the 'gas' can be ascertained and the molecular weight of the dissolved substance computed. In all these cases there are abnormalities observed, which are due to dissociation or breaking up of the molecules; these are specially observed in the case of salts and other electrolytic substances. A solution of common salt, NaCl, for example, contains (Arrhenius) very little combined chloride of sodium, the particles of which have split up into monatomic sub-molecules of sodium and chlorine, heavily charged with opposite electricities; and the electrolytic conductivities afford additional means of measuring the proportion of the salt which has thus dissociated. Those acids which thus dissociate most into active sub-molecules are the most active. Heat has also a powerful action in breaking up molecules, and causing either dissociation, as in the case of acetic acid, whose molecules are complex until its vapour is strongly heated; or chemical decomposition, as in many well known reactions. Those compounds whose formation is attended with the evolution of heat have molecules which are generally most stable at low temperatures; those which absorb heat during their formation are, as a rule, most stable at high temperatures. When a substance is heated, the energy imparted to the molecules is not only spent in giving them motion relatively to one another, but also, and in many instances to a still greater extent, in giving the molecules themselves movements of elastic vibration and of rotation; and this disturbed condition affects the stability of the molecule. How the molecules are built up from their constituent molecules in each particular substance and its isomers is a matter now being diligently investigated; the graphic formulæ to be found in all the text-books of chemistry, and which have been of enormous utility in the systematisation and discovery of organic compounds, are beginning to give place to 'stereochemical' formulæ, in which the attempt is made to represent the tridimensional relations of the atoms within the molecule. The forces between molecules vary with their mutual proximity (see MATTER); and many of the observed anomalies in the obedience of Gases (q.v.) to Boyle's Law under varying pressures and temperatures can be reduced to order by taking instead of , the whole volume occupied by the gas, a term , the free space within that volume unoccupied by the molecules of the gas; and instead of , the external pressure, a term where represents a force of mutual attraction of the molecules. See Ostwald's General Chemistry (1890).