THE BRONSTED THEORY APPLIED TO ACID-BASE BALANCE AND RESPIRATION' ARTHUR W. DEVOR The Ohio State University College of Medicine, Columbus, Ohio
INPREVIOUS
report^,^,^ the Briinsted concept of acids and bases was commended as an advance in chemistry that offers an excellent "tool" for explaining certain biological processes. The purpose of this paper is to point out that advantages can be gained by using this theory when explaining acid-base balance and respiration. The theory shows that there is a direct correlation between pH and acid-base balance. I t clearly illustrates the equilibrium existing between acids, bases, and H ions of the various buffer pairs as indicat.ed in the following equations:
=
+
= =
+ H+
acid base H.HCOa HCOSH.HPO&-SHP0,--+ H.Pr4 ProrgH.Org'
+ HHproton + +
(hydrogen ion)
+ H+
These acids and bases are important buffer pairs of the animal body. Excess hydrogen ions combine with the bases, thus preventing a rapid decrease in the pH of the solution, while the acids furnish hydrogen ions to prevent a rise in pH. It should be understood that Presented hefore the Division of Biological Chemistry at the 124th Meeting of the A m ~ r ~ c aChemical n Society, Chicago, September, 1953. 'DEYOR, A. W., Science, 104, 492 (1946). 3 D~vo, aA. . W.. Proe. 5. Dakota Aead. Sci.. 28. 39-42 (1949). 4 H .Pr = protein in acid form. H .Org = organic acid.
H ions are combined with water and actually exist as H 3 0 + chydronium) ions. One could illustrate t,he equilibrium between the acid and the base as: hydron!um acid base ,on H.A+H?O*A-+HaO+
This is not necessary because it is logical to assume that all ions are hydrated in 1-120 solutions. To further illust,rate the direct correlation between pH and the acid-base equilibrium one can use the Henderson-Hasselbalch equation: pH = pK.
+ log [base]/[aeidl
This mathematical expression indicates that the pH of a solution is affected by the ratio of base concentration t,o the acid concentration. As one would expect, an increase in base-to-acid ratio increases the pH and a decrease in base-to-acid ratio decreases the pH of the solution. What then is thc function of the cations (Xaf, K+, etc.) in acid-base balance? It is not difficult to answer this question, because where anions occur, there must, be an equivalent amount of cations. Therefore the cations must be present as "companions" for the hasic anions. With a loss of such cations, there d l be a decrease in pH if there is a decrease of the base conren-
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tration (HCOa-, for example) to such an extent that the baseto-acid ratio is decreased. The COrcombining capacity of the blood is sometimes called the "alkali reserve." When COz is taken up by the blood, it forms H2C03, which is neutralized by the basic anions (HP04--, protein-) yielding HCOa-, while the cations are present as "companions" to the bases. The basic behavior of nitrogen bases (purines, pyrimidines, etc.) and the amphoteric behavior of amino acids and proteins fits into the picture because they function as proton (H ion) donors or acceptors, or both. It is unnecessary to think in terms of the hydrated molecules as is the case when one uses the Arrhenius theory. Likewise, the basic behavior of substances snch as NaHCOa and NazHPOa can be easily understood without thinking of the hydrolyzed salts. A Teaching Aid. Students of chemistry are now being taught to think of an acid as a proton donor and of a base as a proton acceptor. Therefore it seems that advantages would be gained if we used these ideas for the explanation of the acidic and basic behavior of biological materials. We should realize that chemical theories have changed considerably since the ideas of acids and bases were first introduced into our biochemistry texts. The Bronsted theory helps to avoid the confusion that is sometimes encountered when one tries to explain the basic properties of body fluids on the cation content. Instead, i t makes possible a simple explanation as to why the cations have been called bases. Students find the theory very helpful in explaining the principles of respiration and the chloride shift. When the Henderson-Hasselbalch equation is applied, it is not difficult for students to understand that the basetwacid ratio must not change if the pH is t,o remain constant. This approach to acid-base balance is very simple because as one adds base or removes acid, the solution tends to become more alkaline, while the addition of acid, or removal of base, tends to cause a more acidic solution. The "Alkaline Tide." When HCl is produced by the parietal cells and secreted into the stomach, hydrogen ions are removed from the always available H.HCOa (COz and H,O are always present) as illustrated in the following equation:
The concentration of HCOa- of the blood is temporarily increased, and the urine may become more alkaline. However, most of this "alkaline tide" bicarbonate is soon secreted into the intestine where its function is to neutralize the HCI coming from the stomach, resulting in an over-all neutral action. In other words, the HC03- formed during the production of gastric HCI is soon eliminated via the intestine. Utilization of Acids and Bases. One of the first steps in utilization of amino acids for energy is the removal of the basic amino group to produce NHa and an organic acid. The liver prevents an accumulation of the strong base, NH,, by converting i t into the nearly
neutral amides, urea and glutamine. Since these amides are weakly acid and weakly basic, they have very little effect upon the pH of the body fluids. Organic acids such as a-keto glutaric acid, acetic acid, oxalw acetic acid, citric acid, and pyruvic acid do not decrease the pH of the blood because they are utilized by the body through the Krebs cycle or converted to neutral compounds such as glycogen and neutral fats. When the Na or K salts of these organic acids are taken into the body, there will be a basic reaction because the negative ions of such salts accept H ions from H.HC03 (see equation above) to form more HCOa(alkali reserve) and an organic acid. Organic acids thus formed have a neutral action as described above. Therefore, many salts of organic acids will cause an alkaline nrine while the acids themselves have a neutral action. This explains why fruits and vegetables cause an alkaline urine. Function of the Kidneys in Acid-Base Balance. Acids like sulfuric and phosphoric acids are not utilized as are the organic acids mentioned above, and when they are produced in the body, the kidneys produce NHa, which combines with the H ions of these strong acids, forming NH,+, which is immediately excreted. The negative "partners" for NH4+ are SOr--, &Po&-, HPOI--, and CI-. The HCOa- is thus spared along with its "partner," Na+. If the kidney fails to produce NH, as needed, the H + concentration of the blood increases, and these excess H ions are in part taken up by the HCOa-, (see equation above) to form H . HC03. The ratio of base to acid (HCOI- to H . HCO3 is then decreased and the pH of the blood falls as is illustrated by the Henderson-Hasselbalch equation. Although the sodium ion is potentially a weak acid, one remember that it is verv in acid- ~ - shonld -~ - - ~ " im~ortant base balance, and i t is often considered to be a b a s e forming element. When too much sodium is lost, HCOa- cannot be formed unless this Na+ is replaced, because the Na+ is the common or natural "companion" for HC03- in the extracellular fluid. When the kidneys function normally, they can produce NHa, which accepts H ions from H.HC08, and the NH4+ may be excreted as NH4C1. Thus NaCl may act as a "base-former." However, it is the HC03- formed which is responsible for the basic action, and it should be recognized as the base. Acidosis. If organic acids enter the blood stream a t a very rapid rate, uncompensated acidosis may occur. Such is probably the case of acidosis in a diabetic individual, where an excessive amount of acetoacetic and p-hydroxybutyric acids are formed. Upon examination of the blood during snch a condition of acidosis, we find that the HCOa- concentration is comparatively low and that the base-twacid ratio is below normal. This happens as the base (HC03-, for example) accepts H ions from the stronger acids to form other acids (H.HCOa, for example). In some cases of diarrhea. there is considerable loss of HC03- via the intestine. In addition, there may ~~
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be a starvation effect and the metabolic process will use body tissue, resulting in a release of acids which must be neutralized. The "alkali reserve" may be drawn upon when such acids are produced a t a very rapid rate, causing a decrease in the base-to-acid ratio (HCOa- to H.HC03, for example). These conditions can be better understood when illustrated by the Henderson-Hasselbalch equation. Alkalosis. During excessive vomiting there will be a loss of H ions from the body. The body can produce more hydrogen ions from the "always available" COz and H,O as illustrated in the following equation: COz
+ HsO = H.HCO3 + H + + HCOa-
However, excessive loss of H ions causes an increase in the base-to-acid ratio and the pH of the blood will then rise above 7.4. In such cases, uncompensated alkalosis results. When individuals with peptic ulcers take large amounts of alkalies, the hydrogen ions may be used up as they are secreted into the stomach. When this happens, the HC03- concentration of the blood is then increased as in the cases of excessive vomiting. During hyperventilation a comparatively large amount of COzis lost and reserve H.HCOa is then used to produce more CO,. Although there is some decrease in HCOa-, the base-to-acid ratio increases, resulting in a rise in pH. This condition of alkalosis can be better understood by the aid of the Henderson-Hasselbalch equation. Respiration and the Chloride Shift. The Bronsted concept is very helpful in explaining the COrcarrying capacity of the blood. As the blood is pumped through the lungs, 0 2 diffuses into the erythrocytes and combines with the HHb (hemoglobin) to form HHbOz (oxyhemoglobin). Oxyhemoglobin is a stronger acid than HHb. Therefore, HHb02 yields a hydrogen ion to the HC03- ion:
the red blood cells. In the presence of carbonic anhydrase, this COzrapidly combines with HzOto form more H.HCOa, which in turn yields more H ions to the Hbas fast as it is formed. One can see that the major portion of Copis carried as HCOa-, while the hemoglobin of the cells is responsible for the COTcanying capacity of the blood. From this mechanism it is not difficult to understand why the pH of the blood remains so constant as it picks up and yields COz. Hemoglobin is continually going through the following cycle:
- -b2 doaO2fmm lungs
C
H W +HHbOz H+ to bicarbonate weaker acid stronger acid Hbstronger base weaker base to tissues H+ from H.HCO1
Oxygen is continually leaving the lungs while Cop is leaving the tissues. Therefore one need only to use reason to conclude that H.HbOe is a stronger acid than HHb. To illustrate this point, let us consider COa, which enters the erythrocytes and combines with HZO to form H.HC03. Hydrogen ions from this acid do not go to H20 or there would be a change of pH. Instead, nature must provide a stronger base to accept these H ions. This base must be provided as the O2 leaves the oxyhemoglobin of the blood: Hb0,--HIweaker stronger base base
+ 0,
Therefore it is logical to conclude that hemoglobin is the strager base and oxyhemoglobin is the stronger acid. In the charts on the following page, ionic equations are used to illustrate the respiration cycle. This cycle does not show Na+ and K+ because they do not shift across the membrane. The K ions are "companions" to the anions within the cells and the Na As the HC03- in the cells is used to produce HHCOa, ions are "companions" to the anions of the plasma. more of these ions diffuse in from the plasma and the They themselves do not have a basic action. It should C1- must then shift out of the cells because Na+ re- also be remembered that there is an equilibrium existr mains in the plasma. The HHCOa rapidly decomposes ing between ionized and unionized forms of HHb and into H20and C02in the presence of carbonic anhydrase HHb02. and the COz is thus excreted. Therefore, the blood leaves the lungs with a lower HC03- concentration. SUMMARY There is only a very slight change in the pH of the blood The Bronsted theory shows that there is a direct because the hydrogen ions are simply transferred from correlation between acids, bases, and pH. The a p one base to another. plication of the theory clearly explains the amphoteric As the blood passes through the tissues, Oa diffuses behavior of proteins and amino acids and one need no out of the erythrocytes as it is released from the oxy- longer think of the hydrated molecules as in the case of hemoglobin. The Hb- formed is a stronger base than the Arrhenius theory. The basic behavior of HCOa- is Hb02- and it readily accepts H ions from H.HC03: emphasized, while Na+ and K+ are shown to be necessary as "companions" to the basic anions. It is not difWon- 0 1 Hbficult to understand why such cations were called bases weaker base stronger base before the Bronsted theory was introduced. The theH.HCOs HHb HCOaHbory helps to explain the COrcarryiug capacity of the As the bicarbonate ions are formed in the red blood blood by showing how H.Hb02 furnishes H ions to the cells, they diffuse out and C1 ions shift in because the HC03- and how H.HCOJ furnishes H ions for basic K+ remains in the cells. Carbon dioxide diffuses into Hb-. When the H ions are transferred from acid to
+
-+ -
+
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JOURNAL OF CHEMICAL EDUCATION
m,
t o tissues
Plasma
u CO, from tissues
CO, to lungs
base without being taken up by the water, the pH remains constant. By aid of the Henderson-Hasselbalch equation, the theory emphasizes the importance of the baseto-acid ratio. In addition, this mathematical expression is easier to understand when the theory is applied. As one would expect, more base tends toward a higher pH and more acid tends toward a lower pH. In a few words, one can say that the Bronsted theory "puts the action where it is," which helps to avoid confusion, and therefore it should be the approach to
the explanation of acid-base balance. Theories have changed and we now have a better understanding of acidic and basic behavior of matter. The basic action of fruits and vegetables is easily explained by the fact that the organic acids in such foods are utilized by the body, and the organic anions in these foods accept protons from H.HCOB to produce more HC03- base, while the K+ and Na+ act as "companions" to the base. Protein foods produce an acid reaction because they yield H2SOnand H,POa, which are not utilized and must be neutralized and excreted.
