We know from our study of the structure of hide, that it consists in its natural state of gelatinous fibres which are soft and swollen with water, and easy putrescible. When these are dried, they contract and adhere to each other, forming a hard and almost homogeneous mass, resembling in degree, a sheet of glue or gelatine. After the tanning process, the fibres are changed in character, though not in form; they no longer absorb water so freely, and in drying they do not adhere together, but remain detached and capable of independent movement. The leather is therefore porous, flexible, and opaque on account of the scattering of light from the surfaces of the fibres, although the individual fibres are translucent. At the same time, chemical changes have taken place which render the fibres incapable of ordinary putrefaction. Our first necessity, therefore, in the conversion of skin into leather is to dry the fibres without allowing them to adhere. This is accomplished in the most primitive mode of leather dressing, by mechanically working fatty substances into the skin as it slowly dries, so as to coat and isolate the fibres, which are loosened by kneading and stretching; while at the same time the fat forms a waterproof coating which prevents them from again absorbing the water which is necessary to putrefaction. Similar results may be produced by causing chemical changes in the fibres themselves, which render them insoluble in water, and consequently non-adhesive; and a sort of leather may even be made by merely replacing the water between the fibres with strong alcohol, in which they are insoluble, and which absorbs and withdraws the water from them, allowing them to shrink and harden, while preventing their adhesion. The merit of having first clearly seen and expressed these cardinal principles in leather production belongs to the now venerable Professor Knapp, who published in 1858 a short paper (Natur und Wesen der Gerberei und des Leders) which is a model of clear explanation and practical experiment. Knapp, however, deals mainly with the changes in the condition of the fibre which are necessary to convert it into leather, and not with their physical causes; and before we can explain the means by which these changes are brought about, we must be acquainted with certain facts and theories about solutions which have become much clearer since he wrote.
The particles (molecules) of all substances are drawn together by attractive forces somewhat of the same character as the attraction of gravitation which holds together the solar system, and which is the cause of weight. It is indeed even possible that these forces are identical. Like gravitation, these molecular attractions increase rapidly in intensity as the distance of the attracting bodies diminishes, so that in solids and liquids, where the molecules are near together, they are immensely powerful, while in gases and vapours they are barely perceptible. These attractions are opposed by the motion of heat, which takes the same part in molecular physics which the energy of planetary motion does in the solar system. In solids, the attractive forces hold the molecules rigidly in position, the motion of heat being limited to short vibrations round a fixed point, the effects of which are visible in the expansion caused by rising temperature. If the temperature is increased, most substances become liquid, a condition in which the particles can roll round each other, but are still held together by their mutual attractions, as the sun holds the earth from flying off into space. If the temperature goes on rising, the orbits of the molecules become greater, the liquid expands, and finally molecules fly off at a tangent out of reach of the attractions of the mass of liquid, and are only diverted from their course by colliding with solids or with other flying molecules, from which they rebound. This constitutes the state of vapour or gas.
The molecules usually consist of groups of atoms. Thus in the vapour of water, each molecule contains one atom of oxygen combined with two of hydrogen, and it is only at immense temperatures that this inner grouping is broken up. Naturally, the more complicated and heavier the molecular group, the more easily it is broken up by outside causes into simpler groupings, and molecules may exist in liquids or solids, which break up before they reach the gaseous form. Of such substances the chemist says that they “cannot be volatilised without decomposition.” In very rare instances does the gaseous molecule consist of a single atom; even those of the most perfect gases, such as hydrogen, oxygen and nitrogen consist of pairs which are not broken up at any known temperature. The pressure of a gas, and its tendency to expand is due simply to the motion and impact of its flying molecules, and it may be noted that at the same temperature and pressure equal volumes of all gases have the same number of molecules, the lighter molecules making up for their want of weight by their greater velocity. The average velocity of a molecule of oxygen (O2) at freezing point is 461 meters per second or about that of a rifle-bullet. It must not be taken however, that in any given solid, liquid, or gas, all the molecules at any temperature move at a uniform velocity, but that each individual molecule may vary from moment to moment from rest up to a very high speed, while the temperature of the mass only represents the average. Thus it happens that in all liquids, and even in solids, a certain proportion of the molecules at any temperature will have a speed sufficient to enable them to leave the surface, and take the form of vapour, while a certain proportion will fall back and be caught and retained. Thus every liquid, and theoretically every solid, has a “vapour-pressure,” rising with the temperature, and depending on it only, and at the boiling temperature of the liquid equal to that of the atmosphere, or about 15 lb. per square inch, and therefore able to form bubbles in the interior of the liquid. If a little of a liquid is confined in a flask, the flask will become filled with its vapour, and so long as any of the liquid is present, the pressure of the vapour will depend only on the temperature and not at all on the respective quantities of liquid or vapour. Neither will it be affected by the pressure of other vapours or gases present in the flask, the total pressure in which will be the sum of the “partial” pressures of all the gases and vapours present.[52]
The behaviour of gases and vapours has been described in some detail because it possesses very close analogy to that of substances in solution. The molecules of liquids are held together by attractions which are very powerful over the short distances which separate them, amounting in most cases to many tons per square centimeter of sectional area, but the range over which they act is very small. In the interior of the liquid the attractions on one side of a molecule are of course exactly balanced by those on the opposite side, so that it is free to move within the liquid without hindrance, but at the surface a very small part of the force due to the attractions of the surface-layer is unbalanced and acts as a sort of elastic skin holding the liquid together, and is called “surface-tension.” Familiar examples of this are found in the force which supports a drop on the end of a tube, the possibility of laying a slightly oily needle on the surface of water without sinking, and the ability of some flies to walk on water as if it were covered with a sheet of india-rubber. Many liquids will mix or dissolve in each other in any proportions, e.g. water and alcohol; the attraction of the alcohol for the water-molecule being as great or greater than that of alcohol for alcohol, or water for water. In other cases, such as water and oil, or water and petroleum spirit, practically no mixture takes place, their mutual attraction being small; and each retains a considerable surface-tension at the points of contact, though less than that of the free surfaces, since each exerts an attraction on the other. There are also many intermediate cases, such as water with chloroform, carbolic acid, or ether, in which each solvent dissolves a portion of the other, but the two solutions do not mix, but form separate layers. In these cases an equilibrium is attained, in which there is just as much tendency for either of the liquids to pass into as out of the other layer. In this there is an extraordinary resemblance to what has been said of vapour-pressures; and the tendency to pass into solution is often called solution-pressure; and it may be noted that when equilibrium has been reached, not only is the solution-pressure, but the vapour-pressure of each constituent equal in both solutions. Like vapour-pressures, the solution-pressures usually increase with rise of temperature, more of each constituent passing into the other, till at last the composition of the two layers becomes identical, their surface-tensions disappear, and complete mixture takes place. With phenol (carbolic acid) and water this takes place at about 70° C.
Most of what has been said of the mutual solution of liquids is also true of the solution of solids, but the latter may be divided into two very distinct classes, colloids and crystalloids (which, however, shade off into each other). The colloid or gluey bodies are mostly miscible in any proportion with liquids in which they dissolve, and there is no such thing as a definite point of saturation. There are however some which form jellies which have great analogy to the partially miscible liquids; there is a mutual solubility, a portion of the solid dissolving to a liquid solution, while the remainder of the liquid dissolves in the solid, increasing its volume, but still retaining the characteristics of the solid state. As the temperature is raised, this mutual solubility generally increases, till at a given point the jelly melts, and complete solution takes place, as in the case of partially miscible liquids. These phenomena are of prime importance in the theory of tanning, but their further consideration must be deferred till a few words have been said about the crystalloids. These are characterised by regular crystalline form, indicating that the attractive forces of their molecules are exerted in definite directions, giving them a tendency to attach themselves together in definite geometrical arrangements. They dissolve in themselves no part of the solvent, but are dissolved by it till an equilibrium is reached in which the tendency of further particles of the solid to pass into the solvent is balanced by that of those already dissolved to attach themselves to the remaining solid, or “crystallise out.” Such a solution is “saturated” with respect to the solid residue, but the word has no meaning unless solid crystals are present, and where a body has, as sometimes happens, more than one crystalline form, a solution may be saturated with regard to one of them, and more or less than saturated with regard to another. In “supersaturated” solutions, crystallisation is at once started by the addition of a “seed” crystal of the proper form.
If a crystalloid substance, such, for instance, as copper sulphate, be placed in a solvent (e. g. water), the dissolved salt will gradually spread itself through the whole body of the solvent, though in the complete absence of currents in the liquid, the motion is extremely slow, and years may be taken for the diffusion to rise through a few feet. In many cases salts diffuse through aqueous jellies at the same speed as they would through still water. Colloid substances on the other hand have little or no power of diffusion and mostly cannot pass through jellies at all. This is the reason why tannage with mineral salts is so much more rapid than with vegetable tannins which are of colloidal character, and diffuse through the gelatinous fibres of the hide with extreme slowness.
All dissolved crystalloids do not pass through gelatinous membranes with equal ease, and substances are known, mostly gelatinous precipitates, which do not permit the diffusion of dissolved salts, though they allow water to pass freely. Thin layers of such precipitates form what are called “semipermeable membranes.” The existence of such membranes affords us the possibility of direct measurement of the tendency to diffusion, or as it is generally called the “osmotic”[53] pressure of dissolved bodies. Thus a porous earthenware battery-cell may be immersed in a solution of copper sulphate, and filled with one of potassium ferrocyanide. In this way its pores will be filled with a gelatinous precipitate of copper ferrocyanide, which is pervious to water, but impervious to most dissolved substances. If now the cell be filled with a dilute solution of some crystalloid, say sugar, and its top closed by a perforated cork fitted with a vertical tube, and the cell be plunged in water, the latter will pass into the cell, and the dilute solution will rise in the tube to a height of many feet above the water outside. By substituting a mercury pressure gauge for the vertical tube, exact measures of the pressure in the cell can be made, which is the osmotic pressure of the dissolved substance. At first sight it is paradoxical that the water should flow into the solution, apparently against a heavy pressure, but the explanation is simple. Mention has already been made of the enormous internal pressures of liquids produced by the attractions of their molecules. In the solution a portion of this is borne by the dissolved substance, and the water flows in from the outside till an internal mechanical pressure is produced, equal in amount to the osmotic pressure of the dissolved substance. The resemblance of the phenomena of solution to those of vapour-pressure has already been mentioned, and it is found to be even quantitative, since the measured osmotic pressures are exactly equal in amount to those which the dissolved body would produce if it were in the state of vapour at the same temperature and occupying the same volume as the solution. It acts, in fact, precisely as the “partial pressure” of a vapour. There are several indirect ways of measuring the osmotic pressure of dissolved bodies, as for instance, from the lowering of the freezing point, or the raising of the boiling point of the solution as compared to those of the pure solvent, all of which confirm the direct measurements, and show that in a given volume at the same temperature, the same number of molecules will produce the same osmotic pressure whatever their nature, or conversely, that at the same osmotic pressure and temperature equal volumes of any solution must contain the same number of molecules. The use of these facts in determining molecular weight is obvious.
[53] Solution-pressure and osmotic pressure are really two names for the same force; the former being employed to signify the tendency of a solid to dissolve, and the latter the pressure produced by the dissolved body which tends to prevent further solution. Thus, in a saturated solution in contact with its solid, the two pressures are always equal, but exerted in opposite directions.
A curious apparent deviation from this law is however noticed in solutions of salts, acids, and alkalies, and indeed of electrolytes generally; thus a dilute solution of sodium chloride produces an osmotic pressure nearly double that corresponding to the number of molecules of NaCl present; and in fact behaves as if it were a solution of Na and Cl existing separately. Such a solution conducts a current of electricity very readily, while at the same time the chlorine is carried to the positive, and the sodium to the negative pole, where they separate as Na2 and Cl2 (the Na decomposing the water present and forming NaOH). In fact, the modern theory of electrolysis asserts that these dissociated atoms are not separated from each other by electricity, but that they exist already separated in the solution of the electrolyte, and merely act as carriers for the electricity, and that the work done by the latter is not that of breaking up the salt-molecule, but of giving its dissociated atoms fresh charges of electricity which enable them to combine as new molecules, and escape from the electrolyte. Complex salts do not always break up into single atoms, thus calcium sulphate dissociates into Ca and SO4, hydrogen sulphate (sulphuric acid) into 2H and SO4, and so on. These dissociated atoms and atom-groups are called “ions,” and may be monovalent, divalent, and so on; the divalent ion carrying double the electrical quantity or charge of the monovalent. Without discussing the ultimate nature of electricity itself, the matter is most easily pictured by assuming that the molecule of the undissolved salt is made up of an ion with a + charge (“kation,” e.g. Na), and an ion with a - charge (“anion,” e.g. Cl), by the electrical attraction of which charges they are held together. In the solution these attractions are balanced by those of other ions, so that they can wander freely within the liquid, but in order to take the molecular form of free elements and escape, say as Na2 and Cl2, the pair of kations must go to the - pole and give up one + charge, and at the same time a pair of anions must go to the + pole and receive a + charge. Thus the Na and all other kations separate at the - pole, and the Cl and all other anions at the + pole.