(c) Treat moist spot with 2 drops of glacial acetic acid, a few crystals of benzidine (preferably white), and finally 2 drops of hydrogen peroxide. A greenish-blue color is “positive.”
Acidosis
Everyone is familiar with that form of respiratory disturbance associated with diabetic coma that is known as Kussmaul’s air hunger. Here we have hyperpnoea, a form of dyspnoea typically without cyanosis, and furnishing the best clinical evidence of acidosis. Acidosis, however, is now recognized to be but a particular phase of disturbance of the acid-base equilibrium of the body, and recent work has radically changed our conceptions of its features and its intricate relationships.
Van Slyke restricts the use of the term “acidosis” to describe a condition caused by acid retention sufficient to lower either the bicarbonate or the pH of the blood below normal limits. The pH of the blood may be considered the danger sentinel; as long as it is normal, the acid-base equilibrium is normal or compensated; otherwise, it is uncompensated, and life is seriously threatened. The normal pH of the blood may be given as 7.3 to 7.5 (a slightly alkaline reaction), each individual, however, probably having normally narrower limits of variation. That of the blood serum is about 0.2 pH higher, and that of the other body fluids (not the excretions) probably closely approximates and promptly follows any change in that of the blood plasma. Variations to the acid side may, for a short time at least, be as low as 7.0, although not much lower without fatal results; 7.0 is considered the point where coma occurs. Variations to the alkaline side (Alkalosis) beyond 7.8 are accompanied by symptoms of tetany, although one is not at present justified in assuming that all tetany is either caused, or accompanied, by alkalosis. So, the extreme range of reaction compatible with life probably lies approximately between pH of 7.0 and 7.8.
Recent, but as yet unconfirmed, work suggests that the severe reactions following intravenous medication or infusions may be due, at least in part, to the fact that the pH of the fluid introduced is decidedly more acid or alkaline than that of the blood. This applies to solutions of glucose, the salines, and possibly also to sodium citrate, arsphenamine, sera, antitoxins, etc. The question of suitably buffering such solutions, e.g., with suitable phosphate mixtures, in order to avoid disturbance of the acid-base equilibrium, is being studied, and the preliminary results are promising.
The hydrogen-ion concentration (or its derivative, pH) of the blood varies as the ratio between the concentrations of dissolved carbonic acid and bicarbonate (generally indicated by (H2CO3)/(NaHCO3)), i.e., a relative increase in the H2CO3 increases the hydrogen-ion concentration and lowers the pH, and vice versa. The stability of this ratio is preserved by body mechanisms operative in controlling its two factors,—the H2CO3 being under respiratory control, and the NaHCO3, considered as representing the alkali reserve, being normally maintained by food.
The erythrocytes control the concentration of bicarbonate by virtue of their haemoglobin and the reversible reaction.
H2CO3 + NaCl ⮂ HCl + NaHCO3
The HCl passes into the cell, and is probably held by the haemoglobin. In the lungs, the CO2 is excreted and NaCl reformed. This ability of haemoglobin to form bicarbonate is important inasmuch as the corpuscles can conceal 5 to 10 times as much acid as the plasma bicarbonate can ordinarily neutralize. A full appreciation of the significance of this ratio being the basis of an intelligent comprehension of acid-base equilibrium, a detailed analysis of factors that tend to influence the ratio is given.