[To be continued.]
IRON IN THE LIVING BODY.
By M. A. DASTRE.
Iron occurs, in small and almost infinitesimal proportions, in numerous organic structures, in which its presence may usually be detected by the high color it imparts; and in the animal tissues is an important ingredient, though far from being a large one. It is essential, however, that the animal tissues, and particularly the liquids that circulate through them, should be of nearly even weight, else the equilibrium of the body would be too easily disturbed, and disaster arising therefrom would be always imminent. Hence the iron is always found combined and associated with a large accompaniment of other lighter elements which, reducing or neutralizing its superior specific gravity, hold it up and keep it afloat. Thus the molecule of the red matter of the blood contains, for each atom of iron, 712 atoms of carbon, 1,130 of hydrogen, 214 of nitrogen, 245 of oxygen, and 2 of sulphur, or 2,303 atoms in all. Existing in compounds of so complex composition, iron can be present only in very small proportions to the whole. Though an essential element, there is comparatively but little of it. The whole body of man does not contain more than one part in twenty thousand of it. The blood contains only five ten-thousandths; and an organ is rich in it if, like the liver, it contains one and a half ten-thousandths. When, then, we seek to represent to ourselves the changes undergone by organic iron, we shall have to modify materially the ideas we have formed respecting the largeness and the littleness of units of measure and as to the meaning of the words abundant and rare. We must get rid of the notion that a thousandth or even a ten-thousandth is a proportion that may be neglected. The humble ten-thousandth, which is usually supposed not to be of much consequence, becomes here a matter of value. Chemists working with iron in its ordinary compounds may consider that they are doing fairly well if they do not lose sight of more than a thousandth of it; but such looseness would be fatal in a biological investigation, where accuracy is necessary down to the infinitesimal fraction. The balances of the biologists must weigh the thousandth of a milligramme, as their microscopes measure the thousandth of a millimetre.
The great part performed by iron in organisms, what we may call its biological function, appertains to the chemical property it possesses of favoring combustion, of being an agent for promoting the oxidation of organic matters.
The chemistry of living bodies differs from that of the laboratory in a feature that is peculiar to it—that instead of performing its reactions directly it uses special agents. It employs intermediaries which, while they are not entirely unknown to mineral chemistry, yet rarely intervene in it. If it is desired, for example, to add a molecule of water to starch to form sugar, the chemist would do it by heating the starch with acidulated water. The organism, which is performing this process all the time, or after every meal, does it in a different way, without special heating and without the acid. A soluble ferment, a diastase or enzyme, serves as the oxidizing agent to produce the same result. Looking at the beginning and the end, the two operations are the same. The special agent gives up none of its substance. It withdraws after having accomplished its work, and not a trace of it is left. Here, in the mechanism of the action of these soluble ferments, resides the mystery, still complete, of vital chemistry. It may be conceived that these agents, which leave none of their substance behind their operations, which suffer no loss, do not have to be represented in considerable quantities, however great the need of them may be. They only require time to do their work. The most remarkable characteristic of the soluble ferments lies, in fact, here, in the magnitude of the action as contrasted with infinitesimal proportion of the agent, and the necessity of having time for the accomplishment of the operation.
Iron behaves in precisely the same way in the combustion of organic substances. These substances are incapable at ordinary temperatures of fixing oxygen directly, and will not burn till they are raised to a high temperature; but in the presence of iron they are capable of burning without extreme heat, and undergo slow combustion. And as iron gives up none of its substance in the operation, and acts, as a simple intermediary, only to draw oxygen from the inexhaustible atmosphere and present it to the organic substance, we see that it need not be abundant to perform its office, provided it have time enough. This action resembles that of the soluble ferments in that there is no mystery about it, and its innermost mechanism is perfectly known.
Iron readily combines with oxygen—too readily, we might say, if we regarded only the uses we make of it. It exists as an oxide in Nature; and the metallurgy of it has no other object than to revivify burned iron, remove the oxygen from it, and extract the metal. Of the two oxides of iron, the ferrous, or lower one, is an energetic base, readily combining with even the weakest acids, and forming with them ferrous or protosalts. Ferric oxide, on the other hand, is a feeble base, which combines only slowly with even strong acids to form ferric salts or persalts, and not at all with weak acids like carbonic acid and those of the tissues of living beings. It is these last, more highly oxidized ferric compounds that provide organic substances with the oxygen that consumes them, when, as a result of the operation, they themselves return to the ferrous state.
Facts of this sort are too nearly universal not to have been observed very long ago, but they were not fully understood till about the middle of this century. The chemists of the time—Liebig, Dumas, and especially Schönbein, Wöhler, Stenhouse, and many others—established the fact that ferric oxide provokes at ordinary temperatures a rapid action of combustion on a large number of substances: grass, sawdust, peat, charcoal, humus, arable land, and animal matter. A very common example is the destruction of linen by rust spots; the substance of the fiber is slowly burned up by the oxygen yielded by the oxide. About the same time, Claude Bernard inquired whether the process took place within the tissues, in contact with living matter in the same way as we have just seen it did with dead matter—the remains of organisms that had long since submitted to the action of physical laws—and received an affirmative answer. Injecting a ferric salt into the jugular vein of an animal, he found it excreted, deprived of a part of its oxygen, as a ferrous salt.