Ions

IONS (from the Greek ion-going, wandering), atoms or chemical radicals that carry electric charges.

-History. As first established by Faraday, the conduction of electric current in solutions is associated with the movement of material particles carrying electric charges. The substance conducting electric current-the electrolyte-decomposes into positively and negatively charged radicals, which, through the action of electrostatic forces, are attracted-the former to the cathode, the latter to the anode. Such atoms or atomic groups (radicals) moving in solution and carrying electric charges, Faraday called ions: positively charged ions (moving to the cathode)-cations, and negative ones-anions. Unlike metallic conductors, in which the propagation of electricity is not associated with the transfer and decomposition of matter, solutions of electrolytes received the name "conductors of the second kind." Faraday believed that only when an electric current is passed through a solution do the external electric forces cause some of the electrolyte molecules to split into ions. The founder of the theory of electrolytic dissociation, Arrhenius, on the basis of extensive experimental material, showed that a known portion of electrolyte molecules is constantly dissociated into ions regardless of whether the solution is conducting an electric current at that moment. This laid the foundation for the concept of the existence of free ions in solution as a stable state of matter. The degree of dissociation of an electrolyte, indicating what portion of its molecules breaks down into I., constitutes in Arrhenius' teaching the main quantity characterizing the participation of the electrolyte in a whole series of processes occurring in solutions. Further development of the modern theory of electrolytic dissociation and electrolyte activity was received in the research of Bjerrum, Debye, and Hückel and others. They showed that the activity of an electrolyte, besides being determined by the number of free I. resulting from its degree of dissociation, depends on the electrostatic interactions that arise among the ions themselves. The influence of these electrostatic interionic forces made it possible to explain many features of electrolyte solutions that did not fit within the framework of Arrhenius' classical theory. The creators of the ionic theory had no concrete concept of the structure of I. and of the way matter and charge are combined in it. Similarly, the basic property of I.-its astonishing chemical inertness compared with the corresponding neutral atom-did not receive sufficient explanation. Thus, sodium atoms react vigorously with water, decomposing it with the liberation of hydrogen; iodine gives a specific reaction with starch, etc. But a solution of NaJ, consisting of free I. of sodium and iodine, does not exhibit any of these reactions until the charge of its ions is destroyed (as occurs during electrolysis). These most important properties of ions could only be understood in the light of the modern theory of atomic structure (see). Structure of the ion. According to the Rutherford-Bohr theory, matter is constructed of positive and negative electric charges. The elementary positive charge is the proton, which has the mass of a hydrogen atom, whereas the free negative charge-the electron-has a mass 1,800 times smaller. The atom is built from an extremely small central positive nucleus, around which-in analogy to the planets moving around the sun-electrons revolve in a complex system of orbits. The atomic nucleus consists of protons or of a combination of protons with a smaller number of electrons. The number of positive charges of the nucleus (or the excess of positive charges over the number of intranuclear electrons) equals the number of electrons in the shell surrounding the nucleus. This number increases uniformly by one in the transition from H (atomic nucleus charge 1) to each subsequent element, according to the order they occupy in the periodic system (see). The electron shell surrounding the atomic nucleus consists of a series of successive layers, each of which contains a certain number of electrons. The outer layer can contain up to 8 electrons (with the exception of the first electron layer, directly adjacent to the nucleus; the maximum number of electrons in it is two). When the outer layer contains a full number of electrons, the atom acquires a completed structure and an extraordinarily stable electronic configuration, and accordingly-complete chemical inertness. These are the atoms of noble gases, whose chemical valence is zero. The transition to the next element of the periodic system (an alkali metal) means the addition of a new electron, located on a new outer electron layer. The continued building of the atom in subsequent elements ends only with a new stable combination of electrons of the next noble gas. According to Kossel, the electronic configuration of a noble gas (with an eight-electron outer layer) represents a stable state to which the atom of each element tends to transition. This transition occurs through the loss or capture of missing electrons. It occurs most readily in alkali metals and halogens, of which the first need only lose, and the second acquire, one electron to resemble the nearest noble gas. Similarly, in other elements, the number of electrons they must lose or gain to expose or complete the outer eight-electron layer equals the maximum number of positive or negative valences they exhibit. However, in doing so, the electrical neutrality of the atom, the initial equality of its positive and negative charges, is violated. The atom transforms into a positive or negative I., and the charge of the latter in sign and magnitude corresponds to the valence of the corresponding atom or radical. The electrostatic attraction of oppositely charged I. combines them into a heteropolar molecule. In media having, like water, a high dielectric constant, the action of electrostatic forces is weakened, and the heteropolar molecule again splits into its ions. Thus each I. has an electronic structure not of the atom from which it originated, but of the nearest noble gas. It differs from the latter only by its charge (and the ease with which, by losing it, it is transformed back into the original element). This ion structure fully explains its most important property, noted even by Arrhenius: the astonishing chemical inertness, which is a characteristic of the free I. in contrast to the atom into which it transforms when it loses its charge. Approaching the structure of the stable, chemically inert noble gas, ions differ from each other only in the magnitude and distribution of their electric charge, i.e., purely physical properties. For this reason they represent an object primarily of physical methods of research, an object of physical chemistry. Hydration and size of I. The most important physical properties of I. are its size and magnitude of electric charge. The ratio of these magnitudes also determines the density of charge, which is the greater, the smaller the size of the particle carrying a given charge. However, if we wanted to form an idea of their relative size based on the structure of I., on their electronic model, we would make a serious error. Ions Li^-, Na', K' etc. in water consist not only of the substances indicated, but also of a considerable amount of water molecules closely associated with them and moving together. The water molecule, like the molecule of many other substances, represents a dipole, at the opposite ends of which are concentrated unlike charges (at one pole the negative charge of oxygen, at the other the positive charge of hydrogen). Such dipoles orient themselves around a charged particle, attracting to it with their unlike pole. As a result, each ion in aqueous solution is hydrated, surrounded by a shell built of water molecules. The farther from the center, the less precise this orientation becomes, gradually passing into the chaotic distribution of free water molecules. Thus the hydration of I. is conditioned by their electric charge (Born). As a result of hydration, the size of I., as a particle moving independently, can increase considerably, and often ions with smaller atomic dimensions, as for example Li, even reach a greater size than I. formed from larger atoms, as K. From this follows another, no less paradoxical conclusion, of great importance for understanding certain problems of cell permeability: when a molecule breaks down into ions, the latter (together with their surrounding water shell!) can have larger dimensions than the molecule itself that dissociates them. Mobility of I. Some actions are characteristic of I. along with neutral molecules. Such is osmotic pressure, which depends only on the kinetic energy of the dissolved particles. Others are conditioned by the electric charge, which constitutes the difference between I. and a neutral molecule. Such a property is electrical conductivity. It is determined by the product of the number of ionic charges and the mobility of I.

Each I. moves in an electric field with a speed proportional to the force acting on it and inversely proportional to the resistance it encounters. If the potential difference is equal to one volt per 1 cm, then the speed of movement (in cm/sec at 18°) is expressed for several ions by the following figures: Cation U (cm/sec) Anion V (cm/sec) Na+ 33.0·10-1 3.5·10-1 Ag+ 4.6·10-1 5.7·10-1 NH4+ 6.75·10-1 6.7·10-1 OH- 18.2·10-1 Cl- 6.85·10-1 Br- 7.0·10-1 I- 6.95·10-1 NO3- 6.5·10-1 MnO4- 5.6·10-1 These differences in mobility, observed in I. when the forces acting on them are equal, indicate different resistance arising from friction against water. Since the latter is determined by the size of the moving body, the study of electrical conductivity and mobility of I. allows one to judge their size and hydration. The figures in the table presented show that due to hydration, the differences in the size of I. are greatly smoothed out and often do not correspond, either in magnitude or even in sign, to the difference in size of the corresponding atoms or radicals. Activity of I. According to the Arrhenius theory, each ion has a characteristic mobility for it, independent of its concentration or the presence of foreign substances in the solution, and the electrical conductivity given by the solution should be directly proportional to its ionic concentration. However, in this case, only the effect of the external electric field on I. was taken into account, and the electrostatic interactions between the I. themselves were completely disregarded. These interionic forces can be neglected at low ionic concentrations (e.g., in pure solutions of weak, little-dissociated electrolytes), when the average distance between ions is sufficiently large. But, as recent research has shown, as soon as the concentration increases and I. approach each other sufficiently, the forces of electrostatic attraction and repulsion that arise between them begin to play an increasingly large role. Under their influence, the mobility and electrical conductivity of ions decrease. Similarly, the kinetic energy of I. decreases, and consequently the osmotic pressure they produce. It turns out to be less than that which would be produced by an equal number of neutral molecules at the same temperature. For the same reason, the activity of I. decreases, their degree of participation in other processes occurring in solutions, particularly in chemical processes, which in turn does not remain without influence on the solubility of electrolytes. The decrease in activity is usually expressed by coefficients showing how many times the activity of I. has changed in a given solution compared to the value it would have at sufficiently great dilution, i.e., with the complete elimination of interionic forces. The coefficients obtained for osmotic pressure (fπ), for electrical conductivity (fκ), and for chemical activity (fa) have different values, but all of them decrease as the concentration of the electrolyte increases. Since the change in activity depends on electrostatic interactions between I. in the solution, it is obvious that it will occur to the same extent both when the concentration of a given electrolyte increases and when foreign ions are introduced into the solution. Physiological actions of I. Poisonous and protective action. The action of ions in the body may depend on the osmotic pressure which I. produce along with neutral molecules; much more often it is determined by the special properties of I. Besides the necessity of I. for building parts of the body or their direct influence on various life functions, besides this nutritive (in the broad sense) value of I., they also exert no less important 'protective action'. This protective action some I. exhibit in relation to others, balancing the harmful influence of which (at a certain quantitative ratio between both antagonists; see Ion antagonism). The mutual equilibration of I. when they influence the most diverse life functions leads to the fact that in the body in the vast majority of cases the decisive significance is not the absolute concentration of one or another I., but their quantitative ratio. Experiments by Loeb and a number of other researchers have shown the extraordinarily wide, universal distribution of the phenomena of antagonism described in the most diverse organisms. They are observed in both marine and freshwater animals, as well as in isolated tissues and organs of terrestrial animals and humans. Experiments by Osterhout and subsequent botanical research have shown that they are no less important for plant organisms. Particularly characteristic in this respect is that even the quantitative ratios between antagonist ions coincide in many cases. Thus, for animals and their isolated tissues, the optimal living environment is a salt mixture in which the ions of sodium, potassium, and calcium are approximately in the same proportion as in sea water: per 100 Na ions about 2 K and 2 Ca. In the same proportion these cations are contained in the blood and tissue lymph of vertebrates. - Influence of ions on permeability. The protective action of I. in many cases depends on their influence on cell permeability. Often each salt separately loosens the cell membrane, increases its permeability, and, penetrating into the cell, damages the protoplasm. When acting together, however, they leave the cell surface unchanged or compact it, thereby preventing each other from gaining access inside. Experiments by Osterhout showed, however, that in this case different salts influence permeability in completely different ways. Some of them, such as salts of sodium and other alkali metals, increase permeability from the very beginning. On the contrary, an excess of some other salts, especially calcium salts, first produces a strong compaction of the membrane and a decrease in its permeability. Only later can this reversible decrease in permeability be replaced by its increase as a result of irreversible damage to the cell membrane by an excess of Ca ions. The compacting influence of calcium salts probably explains the observations of Chiari and Januschke, according to which abundant introduction of calcium prevents the formation of exudates caused by poisoning by iodine compounds, diphtheria toxin, etc.; the loosened connection between endothelial cells is strengthened by calcium. Action on muscles. One of the first objects on which the physiological action of I. was studied was the heart of vertebrates. Ringer established that the heart of a frog can pulsate for a long time if small amounts of potassium and calcium salts (and also a little soda to maintain a slightly alkaline reaction) are added to the isotonic NaCl solution. Such a solution was named 'Ringer's solution'. As Locke showed, the vital activity of the heart (and other organs) of mammals can also be maintained for a long time in a similar solution, if only the salt content is somewhat increased corresponding to the higher osmotic pressure of their blood and tissues. According to Ringer, calcium causes systoles of the cardiac ventricle, potassium-diastoles, and their combination in a suitable proportion ensures the proper alternation of systoles and diastoles and gives a normal heart rhythm. Loeb, in his interpretation of these phenomena, assigned a much greater role to sodium ions, whose presence he considered necessary for contractions, while their harmful effect is equilibrated by potassium and calcium. This view received significant support in the experiments of Lingle on the heart of a turtle. When sodium chloride was replaced by an isotonic solution of a harmless nonelectrolyte, contractions—despite the presence of KCl and CaCl2—ceased. They resumed only in the case where the solution contained at least a certain minimum amount of sodium salt. Thus, sodium ions also turn out to be necessary for the contraction of cardiac muscle. The extremely complex question of the relationship between Na, K, and Ca ions in heart work, despite the large number of studies devoted to it, cannot yet be considered finally clarified.—Of other muscular formations, the greatest number of studies has been devoted to striated muscles. Potassium has a very strong depressing effect on them. But they are significantly less sensitive to the absence of Ca ions than cardiac muscle; for a fairly long time they retain excitability in a pure isotonic NaCl solution (although their indirect excitability from nerve quickly fades in this case). Loeb noticed that a skeletal muscle, placed in such a solution, after some time begins to contract rhythmically in it. According to Loeb, for the occurrence of these rhythmic contractions, a certain quantitative ratio between the concentrations of sodium and calcium ions (CNa/Ca) is necessary—a ratio higher than that in which these I. are found in the blood. In a pure NaCl solution, calcium ions diffuse out of the muscle tissue, and as soon as their content in the muscle falls below a certain limit, sodium ions lose the ability to cause contractions.

Similarly, the contraction of skeletal muscle does not occur if the surrounding solution, and consequently the muscle itself, contains too much calcium (or magnesium) salt. The rhythmic contractions that have already begun in a pure NaCl solution cease upon the addition of CaCl₂ in the concentration that it normally has in the blood or in Ringer's solution. As Loeb notes, 'only thanks to the calcium and magnesium salts of our blood, our skeletal muscles do not contract as continuously as the heart contracts.' This effect of ion concentration on the rhythmic contraction of skeletal muscles is of great importance for understanding the mechanism of muscle cramps observed under pathological conditions. Action on nerves. The physiological significance of the effect of I. on nerves is even greater than their direct action on muscles. Nerve excitability, like muscle excitability, depends on the concentration and ratio of cations. In the normal composition of blood, the calcium content is too high in relation to sodium for nerves, as for muscles. Loeb showed that solutions of salts that bind and precipitate calcium (oxalates, citrates, etc.) sharply increase the excitability of motor nerves. To increase the sensitivity of the nerve-muscle preparation to subsequent irritations, it is sufficient to immerse the nerve for a short time in an isotonic solution of one such salt, for example, sodium citrate; with a longer action of the latter on the nerve (for several minutes), muscle cramps begin. Chiari and Frohlich observed a completely analogous phenomenon in the whole organism. In the experiments they conducted, the injection of oxalates caused a sharp increase in the excitability of the sympathetic and in general the entire autonomic nervous system, creating a state resembling the clinical picture of infantile tetany. The introduction of calcium salts has, as expected, the opposite effect. It lowers normal excitability and returns to the previous level the excitability artificially increased by oxalates. Similarly, calcium I. have a braking effect on the nervous system excited by toxins or alkaloids. Thus, the injection of a CaCl₂ solution stops the convulsions caused in a frog by the introduction of strychnine sulfate. In infantile tetany and in tetany, calcium finds application as a means of reducing the excitability of the nervous system. Thus, a change in ion concentration is a general means for exciting nerves and muscles. The decisive factor here is not the absolute concentration of individual I., but their quantitative ratio, in particular the ratio of the sodium and potassium ions on the one hand, and calcium, as well as magnesium, on the other. The numerical value of this ratio of monovalent to divalent cations is often called the 'ion ratio' or 'Loeb's ion coefficient.' The value of the 'ion coefficient' of the blood is of great importance for characterizing the normal or pathological state of the excitability of the nervous system. Changes in ion concentration cause excitation not only in cases where the concentration of salts in the solution surrounding living tissue is directly changed. The research of Nernst showed that in various types of electrical stimulation, the action of the electric current amounts only to the movement of I. along the nerve, to a change in the concentration and ratio of I. in the nerve. Proceeding from the fact that electrical stimulation ultimately amounts to the action of I., Nernst was able to derive the quantitative laws of electrical stimulation. According to the research of Lazarev, in a similar way, light stimulation is directly caused by ionization occurring in photochemical processes. Summarizing the results established for chemical, electrical, and light stimulation, Lazarev came to the conclusion that changes in ion concentration generally represent a universal irritant of living tissue, lying at the basis of all types of excitation. Based on this, he developed a general ionic theory of excitation in a series of studies. However, if the action of ions is an adequate irritant for any living tissue in general, different nerves, like different muscles, show by no means the same relation to individual I. Howell drew attention to the similarity between the action of the vagus nerve on the heart and the action of the potassium ion. According to Zondek, who further developed this analogy, the action of the vagus nerve on the most diverse organs coincides with the action of the potassium ion, whereas for the sympathetic nerve there is a corresponding relationship with the calcium ion. Thus, a direct dependence is established between the balance of potassium and calcium ions in the organism and the constitutional features of vagotonia and sympathicotonia. Other actions of I. Many other processes of excitation, contraction, and secretion depend on the concentration of I., and in some cases the action of I. is strictly specific, while in other cases, I. that are chemically similar show similar biological effects. A good example of the specific action of ions is given by the experiments of Koltsov on the marine ciliate Zoothamnium: magnesium I. are necessary for the ciliary movement of its cilia, while calcium I. sharply accelerate the contraction of its stalk. This dependence appeared with such unchanging regularity even in the presence of such small amounts of the corresponding salts that it made it possible, by means of 'biological analysis,' to determine, for example, impurities of calcium and magnesium in various types of commercial table salt. According to the experiments of Spaeth, cations have a characteristic effect on the chromatophores of fish. In Fundulus, the brown pigment cells (melanophores) expand under the influence of sodium salts, while potassium salts cause the pigment grains to gather into clumps in the center of the cell; on the yellow pigment cells (xanthophores), sodium and potassium ions acted in the opposite way. Hamburger found that the addition of a small amount of calcium salt to an isotonic NaCl solution significantly enhances the phagocytosis of leukocytes. Similarly, O. Nikolaev recently established by experiments on an isolated salivary gland that an increase in the concentration of calcium ions strongly increases the secretion of saliva. Magnesium I. have a very peculiar effect; it was studied mainly by Meltzer and Auer. The injection of a sufficient amount of magnesium salts causes deep anesthesia and paralysis of the motor muscles. Smaller amounts of magnesium, which have a very weak direct effect, increase the animal's sensitivity to ether many times over. The introduction of calcium salts quickly eliminates the phenomena of magnesium narcosis. The few examples cited give a sufficient idea of the enormous significance of the physiological action of salt I. An even greater dependence is shown by every living organism on the concentration of hydrogen and hydroxyl ions. (For their influence, see Active reaction, Hydrogen ions.) Physiological and colloidal actions of I. A living organism is built from colloids, and the diverse physiological actions of I. must be entirely attributed to their influence on biocolloids. Indeed, colloid chemistry gives on simpler systems the most important of those regularities that biology establishes for the action of electrolytes on life processes. First of all, biocolloids, in particular proteins, belong to amphoteric colloids, the physico-chemical properties of which (character of electrolytic dissociation, sign and magnitude of charge, etc.) are decisively influenced by the reaction of the surrounding solution; such a dominant influence also belongs to it in life phenomena. The isoelectric point of most amphoteric biocolloids lies in weakly acidic reactions. Therefore, in the weakly alkaline or neutral environment usual for life, they are electronegative. According to Hardy's rule, the decisive influence on the properties of colloids is exerted by I. of the opposite sign; in the case of electronegative colloids, by cations. The dominant role played by cations in the action of salts on the most diverse life functions proves the full applicability of Hardy's rule in biology. Furthermore, in full accordance with the rule established for colloids by Schulze, the physiological activity of an electrolyte rapidly increases with the valence of its cation. However, even between I. of the same valence, more or less significant differences in their action on colloids are found. As Hofmeister first pointed out, on the basis of these differences, I. can be arranged in sequential series in the order of increasing or decreasing colloidal-chemical activity. Completely the same ion series, called the Hofmeister series (see Hofmeister series), have been established for many physiological actions of I. (for their toxicity, protective action, etc.). Even the protective action of I., the ability of some I. to balance the action of others, giving harmless physiological balanced solutions, finds a complete analogy in such an antagonistic effect of I. on certain colloidal systems.

In a number of cases, ions, instead of summing up their coagulating action on colloids, mutually weaken it, similarly to how they mutually suppress their poisonous action on living organisms. Thus, the physiological actions of ions are entirely determined by their influence on cellular and tissue colloids.