Molecule

MOLECULE (from the Latin molecula - little particle), the smallest particle of a substance that retains all its chemical properties. The molecule is the limit of the fragmentation of a substance as such: upon attempts at further subdivision, molecules must divide into their constituent parts—atoms (see Atom), which possess different chemical properties. Therefore, the chemical formula of a molecule is at the same time the chemical formula of any quantity of this substance, which consists of an aggregate of a huge number of identical molecules. Over the recent decades, as a result of a significant expansion of knowledge about the finest structure of matter, the concept of the molecule has changed and been refined somewhat: only in gases do molecules possess individual existence and are located at huge (compared to their size) distances from each other; in liquids they combine ("associate") with each other into more or less large and stable aggregates; in solid (crystalline) bodies they are brought so close together that the boundaries between individual molecules are erased and in many cases a regular alternation of atoms (or ions) is obtained, which are difficult to attribute to specific molecules. However, for gases and for (diluted) solutions, the concept of a molecule has fully retained its meaning. Hereafter, speaking of molecules, one must keep in mind, mainly, these two forms of substance. For them, Avogadro's law holds true, according to which equal volumes of various substances in a gaseous or vaporous state (at equal pressure and temperature) contain an equal number of molecules. At atmospheric pressure (760 mm) and 0°, a volume equal to 22.4 liters contains such a quantity in grams of any substance that is equal to its molecular weight: 2.0156 g of H₂, 32.000 g of O₂, etc. This volume is called gram-molecular. The number of molecules in this volume [or at other (liquid and solid) states of matter, in a weight quantity of it equal (in grams) to the molecular weight (this quantity is designated by the abbreviation "mole")] is according to Avogadro's law the same for all substances, i.e., it represents a universal constant. Usually, it is called Avogadro's number and is designated by the letter N. Sometimes it is called Loschmidt's number and is designated by the letter N_L. At present, there exists a whole series of independent ways of determining N, yielding very close results. The most accurate of these methods give the value N = 6.06·10²³. Dividing this huge number by the volume of a mole of a substance in a condensed state (liquid or solid), one can find the upper limit of the dimensions of the space occupied by a single molecule. Other methods also make it possible to determine, admittedly with less accuracy than N, the dimensions of molecules or, rather, the spheres of their action. They are different for different substances, increasing with the complexity of the chemical composition, but on average their order of magnitude is 10^-7 cm. It is understandable that at such dimensions molecules are invisible even in the best microscope or ultramicroscope, and all our information about molecules can be obtained only by indirect methods. The mass of an individual molecule is very small. It can be found with high accuracy for each substance by dividing the mole by Avogadro's number. Thus, there has been obtained for hydrogen 0.324·10^-23 g, for helium 0.639·10^-23 g, and so on. As indicated above, molecules consist of even smaller structural units—atoms. Therefore, the question arises about the structure of molecules, about how atoms are located within molecules, by what forces they are bound to each other, what energy must be expended for their separation, and so on. This question is of particular importance for chemistry, since it reduces to the question of the structure and formation of chemical compounds. The problem of the structure of complex molecules constitutes the most important part of modern organic chemistry and was developed during the 19th century with great completeness. As for simpler molecules, predominantly of inorganic compounds, their structure and formation have begun to be clarified only recently in connection with the theories of atomic structure that have formed a new chapter of physical chemistry, termed in recent years chemical physics. It turned out that the bond between atoms in a molecule can be different. In this sense, all molecules can be divided into two large groups: heteropolar and homeopolar molecules. Between the extreme representatives of both groups there exist a large number of gradual transitions, and it is not always possible to determine to which type the molecule of a given compound belongs. However, typical heteropolar molecules, such as for example NaCl, differ sharply from typical homeopolar molecules, such as e.g. H₂. The former consist of atoms or groups of opposite chemical character, furthermore charged with electricity of opposite sign, i.e., of ions, for example Na⁺ and Cl^-, Ca⁺⁺ and CO₃^--, etc. The latter, on the contrary, consist of identical or close in character atoms or groups devoid of electrical charge: H₂, N₂, O₂, Cl₂, C₂H₆, etc. In molecules of the first kind, the bond between components is carried out apparently by attractive forces of an electrostatic character, i.e., by the attraction of their opposite charges; the nature of the bonding forces in homeopolar molecules is less clear, although it likewise can probably be reduced to forces of an electrical character. The picture of the formation of heteropolar molecules from the standpoint of the electronic theory of atomic structure was given by Kossel: upon the formation of a molecule, for example NaCl, the valence electron of Na is pulled over to Cl, imparting a negative charge to the latter and leaving Na positively charged. The resulting ions Na⁺ and Cl^- mutually attract each other electrostatically, with each of them having eight electrons in its outer shell (see Atom), just like the atoms of the nearest noble gases. Proceeding along this path, Kossel explained the formation not only of simple heteropolar compounds of the NaCl type, but also of significantly more complex coordination compounds (see), and with the aid of energy considerations it was possible to calculate the coordination number (see Coordination theory). In typical homeopolar molecules, a complete transfer of electrons from some components of the molecule to others probably does not occur. It is believed that here the electrons effecting the chemical bond do not leave the atom from which they originated, but become common to the connected components, surrounding their nuclei with their orbits and spending part of the time in the sphere of action of each of the components of the molecule. Finally, in molecules of an intermediate type, called semipolar (Lux), the valence electrons are common to the components of the molecule, but are shifted to a greater or lesser degree toward one of them. Regarding the structure of molecules, tremendous successes have been achieved in recent years through the study of the optical and electrical properties of molecules. Molecular spectra are significantly more complex than the line spectra of atoms: they consist of bands which, with powerful instruments, are resolved into a series of the finest, very close-lying lines. The origin of these bands is explained in the following way. Whereas in atoms spectral lines appear only as a result of the jumps of electrons from outer orbits to inner ones, in molecules, besides these abrupt changes in state and energy store, vibrations of the components of the molecule relative to each other and rotations of the molecule as a whole are possible. All these three types of motion, according to the principles of quantum theory (see), can be accompanied not by continuous, but only by strictly defined discrete changes in the energy store of the molecules. Each of them under certain conditions can be accompanied by the emission of a separate spectral line. Changes in rotational motion are associated with very small changes in energy, small quanta, and the lines corresponding to them lie in the far infrared (see Infrared rays) region. Changes in vibrational motion lead to significantly (100–200 times) larger energy changes, and the corresponding lines are located in the short infrared part of the spectrum. Electron jumps in molecules cause the appearance of lines in the visible and ultraviolet parts of the spectrum. These three series of motions are interconnected, and e.g. the vibrational motion of the components of a diatomic molecule along the line connecting their centers always also causes a simultaneous change in rotational motion. This is why the appearance of a vibrational spectrum line always causes the appearance of a series of rotational spectrum lines located near the first one at very small distances and merging with it in weak optical instruments into a single spectral band of the vibration-rotation spectrum. The jump of an electron excites, besides the main line, a whole series of lines of the vibrational and rotational spectrum merging with the first into a broad band in the visible or ultraviolet part of the emission or absorption spectrum (see Absorption). A detailed study of molecular spectra has led to the precise determination of the length (distance between atoms) of certain diatomic molecules, their moment of inertia, and important energy quantities, such as e.g. the heat of dissociation of many molecules into atoms. Important results have also been yielded by the study of the electrical properties of molecules. Polar molecules, consisting of components that are heterogeneous in an electrical respect, are completely neutral only at distances large compared to their length.

At small distances from their ends (poles), the asymmetrical arrangement of electrical charges within them begins to manifest: positive charges predominate at one pole, and negative charges at the other. The distance between the centers of gravity (poles) of both types of charges, multiplied by the magnitude of these charges, yields the dipole moment of such a molecule, which is called a dipole. The asymmetrical arrangement of opposite charges, imparting a dipolar character to the molecule, is permanently inherent to some molecules (permanent dipole moment), while in others it can be caused by external influences, for example, by placing the molecule in an electric field (induced dipole moment). The study of the permanent dipole moments of molecules, conducted primarily by Debye (P. Debye) and his school, led to the determination of the length of dipolar molecules and the spatial arrangement of their constituent atoms. Furthermore, it proved possible to explain the phenomenon of association in liquids and vapors (see Liquids), deviations in the behavior of binary mixtures from simple mixing laws, the nature of the so-called Van der Waals forces of attraction between molecules, and so forth. At the same time, it was clarified which parts of the molecule (groups or radicals of complex compounds) determine its polar properties, to assign a specific dipole moment to each such polar group, and by additive summation of the latter (taking into account the spatial arrangement of individual groups) to calculate the dipole moments of complex molecules in close agreement with experimental data.

The induced dipole moment can appear in a molecule that does not possess a permanent dipole moment, under the influence of light falling upon it, which constitutes a rapidly alternating electromagnetic field. The appearance of this moment affects the refraction of light when it passes through such polarizable molecules. The study of molecular refraction R = ^{n2 - 1}/_{n² + 2} ^M/_q (where n is the refractive index, M is the molecular weight, and q is the density) of a series of compounds led to the discovery of additively compounding atomic and ionic refractions of individual molecular components and to the elucidation of the structure of certain molecules. When applied to polar molecules of salt-like compounds, the study of refraction led to the discovery of the deformation of the electronic shells of ions in chemical compounds, which allowed Fajans (K. Fajans), primarily, to explain a whole range of specific properties of various salts (vapor pressure, boiling and melting points, solubility, color, and many others).

In recent years, a new physical method for investigating molecules has been discovered, which has already yielded very interesting results. This method, discovered almost simultaneously by Raman (Raman) in India and Mandelstam and Landsberg in Moscow, boils down to obtaining scattering spectra of the studied substances, in which, alongside the usual lines of the incident light, a number of satellites appear, i.e., lines of greater and lesser wavelength.

The X-ray analysis method has also proved very useful for studying molecular structure, allowing us to ascertain the spatial arrangement and absolute distances between atoms comprising the molecules of complex compounds such as naphthalene, hexamethylenetetramine, and others.

Dipole moments of individual groups in complex molecules

Bond or radical

C-O...C-Cl

C-NO₃ \OH

Dipole moment μ · 10¹⁸

0,4

1,6

2,3

0,7

1,5