Buffer Properties

BUFFER PROPERTIES, the ability of many substances to mitigate changes in the active reaction (see) of a solution that would otherwise occur upon the addition of acids or alkalis. This stabilizing effect on the reaction of a solution is called a buffer action. Buffer action. If a solution of caustic soda of the same concentration is gradually added to ten cubic centimeters of a decinormal solution of acetic acid, the acidity of the solution, determined by the concentration of free hydrogen ions contained in it (see), will decrease. At pH 8 7 6 / cubic cm of NaOH with the addition of 10 cubic cm of NaOH, the process of binding the acid by the alkali, the neutralization process, will be completed, all the acetic acid will be converted into the corresponding salt—sodium acetate, and the combined H and OH ions will yield water molecules. Further addition of NaOH will result in the predominance of free hydroxyl ions—an alkaline reaction. The curve presented here (see figure, solid line) illustrates the changes in the reaction expressed in terms of pH (hydrogen ion exponent, —see Hydrogen ions) observed during the neutralization of acetic acid. The dashed line in the same figure depicts the corresponding change in reaction (pH) upon the addition of NaOH to decinormal hydrochloric acid. If one compares both curves and observes how much alkali was required for an identical change in reaction, for example, for a change in pH from 4 to 5, the results will be very different: in the first case—about 5 cubic cm of NaOH, in the second—barely perceptible traces of the latter. The amount of alkali (or, accordingly, acid) required for a definite change in reaction is a measure of the stability of the solution's reaction, the magnitude of its buffer action. In the first case it is very significant, in the second—completely negligible. If the number of gram-equivalents of alkali (or, accordingly, acid) added to a liter of the test solution is denoted by the symbol \Delta B, and the resulting change in reaction by \Delta pH, then, according to Van Slyke, the buffer action will be equal to the ratio of these quantities: Buffer action = \Delta B / \Delta pH. The difference in the course of the curves for both solutions considered above is due to the properties of both acids. Hydrochloric acid belongs to strong acids, completely dissociated into its ions. In contrast, acetic acid is relatively weakly dissociated: only a small part of its molecules (about 1.3% in a decinormal solution) dissociates and yields hydrogen ions, which determine the acid reaction of the solution. Therefore, acetic acid has a significantly less acidic reaction (larger pH) than hydrochloric acid at the same molar concentration. Upon the addition of NaOH, the hydroxyl ions of the alkali bind the hydrogen ions. But by virtue of the general conditions of chemical equilibrium, the removal of dissociation products causes the dissociation of new, previously undissociated molecules, releasing ever-new amounts of H-ions to replace those bound by the alkali. Thus, acetic acid (unlike fully dissociated hydrochloric acid), in addition to free, active H-ions determining the active reaction of the solution, possesses reserve, backup hydrogen ions in its undissociated molecules—a reserve acidity capable of rapidly replenishing the depletion of free ions. These acid reserves (or alkaline ones, if the solution can release reserve OH-ions and bind added acids) determine its buffer action; it is all the more significant, the more reserve ions are mobilized for a given change in reaction. The very name (buffer action) was given by analogy with railway buffers, which soften the severity of mechanical shocks. A more correct comparison would be with vessels of various capacities, in which the addition of an equal amount of liquid causes a different change in level. The greater the capacity of the vessel, the more liquid is required for a definite rise in level; similarly, the amount of alkali (or acid) necessary for a given change in the "level" of the reaction depends on the amount of reserve H- or OH-ions ("buffer capacity"). Buffer solutions. Electrolytic dissociation of weak acids and alkalis drops sharply in the presence of salts having an ion in common with them. For example, acetic acid is significantly less dissociated in the presence of its sodium salt (sodium acetate, which, like acetic acid, yields the acetate ion) and yields significantly fewer hydrogen ions than in a pure solution. The concentration of hydrogen ions is directly proportional to the concentration of acetic acid molecules and inversely proportional to the concentration of acetate ions. Since neutral salts belong to strong electrolytes almost completely dissociated into their ions, one can take, with sufficient approximation, simply the concentration of the corresponding salt instead of the concentration of acetate ions. The concentration of hydrogen ions in such a solution containing a weak acid and its salt is then expressed by a simple formula (in which square brackets denote the concentration of the substances enclosed in them): [H'] = K ([acid] / [salt]) (1). Similarly, in a mixture of a weak alkali and its salt, the concentration of hydroxyl ions (from which it is likewise easy to calculate the closely related H-ion concentration and solution reaction) is determined by an analogous expression: [OH] = K ([alkali] / [salt]) (2). For a more precise calculation, it would be necessary in both formulas to slightly decrease the denominator by multiplying it by the degree of dissociation of the salt (a value less than one). Such mixtures possess particularly large amounts of reserve, easily mobilized H- and OH-ions and, correspondingly, a particularly large buffer action. At the same time, they make the reaction of the solution stable simultaneously with respect to both alkalis and acids. Thus, for example, a mixture of acetic acid with sodium acetate (obtained during the partial neutralization of acetic acid with caustic soda, see figure), as we saw, changes its reaction relatively little upon alkalization. Likewise, upon the addition of a strong acid, e.g., hydrochloric, its effect is weakened due to the fact that it combines with sodium, displacing an equivalent amount of weak acetic acid from its salt. Solutions of such mixtures of a weak acid or alkali with the corresponding salt, the so-called buffer solutions, have acquired special importance due to the ease with which their reaction can be calculated according to formulas (1) and (2) given above. The constant K in these formulas represents a constant characteristic of each acid or alkali—the so-called dissociation constant. If the acid and its salt are present in equal (equivalent) concentration, then, obviously, the concentration of hydrogen ions becomes numerically equal to the dissociation constant ([H'] = K). Thus, the dissociation constant of the acid (or, accordingly, alkali) directly indicates the average reaction in the region of which the buffer action of the given mixture is manifested. At this point, the buffer action is mostly particularly great. The following table (see table on page 274) gives the hydrogen exponent of several buffer solutions: mixtures of acetic acid and sodium acetate (acetate mixture), primary (monometallic) and secondary (dimetallic) sodium phosphate (NaH₂PO₄ and Na₂HPO₄), and ammonia with ammonium chloride. From formulas (1) and (2) one can directly derive one very property of buffer solutions: the reaction yielded by the buffer mixture depends (in the first approximation) exclusively on the ratio of its components, and not on their absolute concentration. Therefore, in the table given, one could also, without citing the concentration of the acid (or alkali) and salt, limit oneself to indicating their ratio. Dilution of a buffer solution does not affect its reaction. Of course, the same cannot be said about the buffer action. For a given reaction, it is the more significant, the higher the concentration of the buffers. The considered properties of buffer solutions determine their most important practical applications: 1. Very many biochemical and biological processes are highly sensitive even to insignificant changes in reaction (see Active reaction and Hydrogen ions). In the very course of these processes, large amounts of acidic or alkaline products are often generated, which could alter or even completely halt their further course. For the precise study of such processes, it is necessary to carry them out under conditions that exclude the possibility of any significant fluctuations in reaction. Buffer solutions serve this purpose, used here as reaction regulators. This method was applied by Sörensen (1909) to study the effect of the active reaction on enzyme activity. Depending on the amount of generated acidic or alkaline products, on the one hand, and the desired degree of reaction constancy, on the other hand, one has to use solutions with a greater or lesser significant buffer action. 2.

In other cases, the magnitude of buffer action is of no particular significance, and the use of buffer solutions is based on the possibility they provide of preparing stable solutions of any desired reaction (see table). By means of indicators (see), which are substances that change their color depending on the active reaction of the solution, one can compare the test solution with a series of buffer solutions of known reaction. By establishing in which of these solutions a given indicator assumes the same color as in the test solution, the reaction of the latter can be determined. Thus, buffers are used here as standard solutions by comparison with which the reaction is measured. The use of such standard buffer solutions lies at the basis of the indicator or colorimetric method of measuring reactions. Other buffer systems. Other chemical systems can also exert a greater or lesser significant buffer action. This may depend, for example, on the precipitation of the added alkali or acid. Thus, if caustic soda is added to sea water, the solution will become alkaline until its pH becomes approximately equal to 8.6. At this reaction, Mg(OH)₂ begins to precipitate, formed from magnesium salts and the added NaOH; a further increase in alkalinity will stop until all the magnesium has precipitated from the solution. Furthermore, even insoluble substances (e.g., animal charcoal) can capture added acids or alkalis by adsorption. Finally, proteins and other amphoteric substances (see Ampholytes) are distinguished by a very strong buffer action. Due to their dual ("amphoteric") nature, they can bind both acids and alkalis. The amphoteric character of cellular colloids is of great importance for the constancy of the intracellular reaction. - Sea water buffers. Changes in reaction exert a tremendous influence on vital phenomena; life is possible only within a certain, for most organisms relatively narrow, interval of concentrations of hydrogen and hydroxyl ions. Therefore, in nature, buffers play a major role in maintaining the constancy of reaction necessary for life. Sea water, which represents the natural external environment of most aquatic organisms, possesses a very significant buffer action, which depends on the bicarbonate mixture contained within it—a combination of carbonic acid and sodium bicarbonate (sodium bicarbonate). Due to the presence of this buffer, the usual weakly alkaline reaction of sea water is preserved, and fluctuations in the reaction produced by aquatic organisms that absorb CO₂ during photosynthesis or release acidic metabolic products are moderated. Buffer properties of the blood. Of particular interest are the buffer properties of the internal environment of the organism, in particular, the blood. Blood has a weakly alkaline reaction characterized by great constancy. Even in vitro, blood stably retains its reaction and possesses a very large buffer action. It is necessary to add several tens of times more caustic soda to it than to distilled water to cause the same alkalinization of the solution, and several hundred times more HCl for the same acidification. Just as in sea water, the main buffer of blood serum is the bicarbonate mixture—a combination of CO₂ and NaHCO₃. The hydrogen ion concentration it yields is approximately determined as follows: where K is approximately equal to 3.10^-7. Phosphates are also contained in the serum, however, compared to bicarbonates, their amount and role are small. With respect to buffer action, the bicarbonate solution is quite similar to blood serum. Thus, for example, both liquids dissolve the same amount of CO₂ proportional to its partial pressure in the surrounding air. When this pressure changes, as formula (3) shows, the concentration of hydrogen ions in them changes by the same factor. Whole blood with its formed elements exhibits under the same conditions a noticeably greater constancy of reaction. This additional buffer action, compared to serum, depends on the amphoteric protein substances of the blood, in particular on the Hb located in erythrocytes. The latter represents a very weak acid, so weak that its acidic character cannot manifest itself in the presence of an excess of CO₂. But when the pressure of the latter is lowered, e.g., in arterial blood, oxyhemoglobin, as an acid, decomposes a certain amount of bicarbonate, displacing CO₂ from it. As a result, the denominator in formula (3) decreases and the influence of the reduced CO₂ content is partly compensated. Thus, Hb exerts a significant influence on the carbon dioxide binding curve, and thereby on the blood reaction. In particular, it moderates the differences associated with the different CO₂ pressures in arterial and venous blood. In any case, ultimately the blood reaction is entirely determined by the ratio of carbonic acid and bicarbonate, i.e., the ratio of free (dissolved) CO₂ and chemically bound CO₂. The former is easily liberated from the blood, the latter can be displaced by decomposing bicarbonates with acids. Both of these quantities—the amount of free and bound CO₂—jointly characterize the buffer properties and reaction of the blood. Their measurement has recently become widespread and important. Regarding its reaction, blood possesses the same properties as other buffer solutions. We have seen that the reaction of a buffer mixture is determined by the ratio of the acid and its salt, and not by their absolute concentration. Correspondingly, the blood reaction remains practically unchanged even upon multiple dilution with an isotonic NaCl solution (or any other unbuffered solution). This property of blood is frequently utilized when measuring its reaction, employing for this purpose a small amount of blood diluted with a NaCl solution. It also makes harmless the intravenous infusion of various so-called "physiological solutions," which frequently have an abnormal reaction that would prove fatal to the organism if even a small admixture of blood did not bring it close to the physiological norm. Upon the addition of alkali to blood in vitro, the latter is neutralized by carbonic acid; conversely, any acid reacts with bicarbonate and, forming a neutral salt, is replaced by an equivalent amount of CO₂ displaced by it from the bicarbonate. This explains the remarkable fact that has already repeatedly attracted the attention of researchers: by introducing various acids into the blood (in vivo)—from the weakest to the strongest—it turns out to be completely impossible to achieve different (according to the strength of the acid used) changes in the blood reaction. As long as some amount of bicarbonate buffer remains in the blood, the changes in reaction turn out to be equally negligible in all cases. Then, simultaneously with a sharp disruption of the reaction, death ensues. These crude experimental exposures give a graphic representation of what happens in the organism under natural conditions. The vast majority of metabolic products are acidic in nature (phosphoric, carbonic, lactic, butyric, and other acids). The blood buffers must protect its normal reaction from these acids continuously entering from the tissues. The latter is weakly alkaline, i.e., characterized by a small excess of active hydroxyl ions. The hydrogen ion exponent (pH) of blood averages 7.4, the hydrogen ion concentration is 0.44.10^-7, and the hydroxyl ion concentration is about 7.10^-7 (at 37°). Compared to this insignificant concentration of free hydroxyl ions, the amount of reserve ions that can be liberated for binding added acids is very large (about 2.10^-2). Their amount, however, is far from possessing such constancy as the active reaction of the blood, and can undergo severe changes, especially under pathological conditions. Alkaline solutions represent only the first barrier against acidic products introduced from the outside or generated in the organism. The disruption of the reaction produced by the latter is weakened many times over by the blood buffers, but cannot be completely eliminated by them: the binding of part of the bicarbonate molecules and the liberation of CO₂ shifts the initial ratio of this primary buffer mixture. A finer regulation of the reaction is accomplished by the lungs. Any increase in the hydrogen ion concentration serves as a stimulus to the respiratory center and immediately increases pulmonary ventilation (see Respiration). Due to the high sensitivity of the respiratory center to hydrogen ions, the pulmonary regulation apparatus operates unusually precisely: removing larger or smaller amounts of CO₂ from the blood, depending on the active reaction existing in it, it automatically restores the normal ratio between it and the bicarbonate. The blood buffers protect the organism from sharp fluctuations in reaction that would be fatal to it; the respiratory apparatus ensures a constant ratio of the components of the buffer mixture (even with sharp changes in their absolute concentration) and thereby the exact constancy of the active reaction. A particularly significant pathological accumulation of non-volatile acids and a corresponding decrease in alkaline reserve are observed in acidosis (see).

However, even this usually does not lead to a change in the active reaction of the blood: by means of enhanced pulmonary ventilation, a decrease in the CO₂ content is achieved, which in most cases compensates for the drop in bicarbonate concentration ("compensated acidosis"). The opposite phenomenon is represented by compensated alkalosis, in which an increase in alkaline reserves is compensated by a proportional rise in CO₂ pressure. Changes in the CO₂ content in the alveolar air of the lungs can serve in both cases as a direct indicator of changes in blood bicarbonate concentration. The total amount of buffers in the blood decreases in the first case and increases in the second, but the active reaction remains practically constant.

D. Rubinshtein. Bufo, toads, tailless amphibians, family Bufonidae. The common species are Bufo vulgaris, the common toad, and Bufo viridis, the green toad. They live in forests, shrubs, gardens, cellars, old walls, under tree trunks, and in other places. They are nocturnal animals. They are very useful to humans by destroying harmful insects. In their skin, they have sac-like poison glands, especially powerfully developed behind the eyes (so-called parotids); this secretion is not absorbed by human skin, which is why toads can be handled with undamaged hands without any fear, but it is highly toxic when introduced directly into the blood. The secretion of the skin glands of some toads from tropical countries is used for making "arrow poison" (see Amphibia, Poisonous animals).