TITRATABLE ACID OR BASE OF BODY FLUIDS Ole Siggaard-Andersen

Department o f Clinical Chemistry, Rigshospitalet, Copenhagen, Denmark

Actual Acidity and Titratable Acidity

S.P.L. Sgrensen not only introduced the p H concept, but in his classical paper from 1909 “Enzymstudien 11” he also described in great detail the colorimetric as well as the potentiometric methods for p H measurement, and he pointed out that two kinds of acidity must be distinguished: Actual acidity and titratable acidity.

Previous to Sgrensen the acidity of a solution was usually determined by titration. But with the dissociation theory of Arrhenius (1887) it became clear that “true” acidity was related to the concentration of free hydrogen ions. Thusan aqueous solution of 0.1 molar HC1 and an aqueous solution of 0.1 molar lactic acid show the same acidity when titrated with base using phenolphthalein as indicator (100 mmol./l), but the “true” acidity, i.e. the hydrogen ion concentration is widely different, being 100 mmol./l and 3.7 mmol./ 1, respectively.

In the newer terminology the word acidity is reserved for the hydrogen ion concentration or pH, while titratable acidity is simply called titratable acid. Decreasing acidity may also be designated increasing basicity (alkalinity). In the following the acid-base definitions of Brgnsted (1923) are used.

Titratable acid (T .A . ) can be defined as the excess of acid in a solution determined by the amount of strong base it is necessary to add to the solution in order to produce a given final state. Titratable base (T.B.) can be defined as the excess of base in a solution determined by the amount of strong acid it is necessary to add to the solution in order to produce a given final state. The f ina l state is defined by the hydrogen ion concentration (activity), e.g. a p H of 7.40 in the case of urine, and as the hydrogen ion concentration varies with temperature the latter must also be specified. For a given final state a positive value for T.A. corresponds exactly to a negative value for T.B. Or in other words, an addition of strong acid is completely equivalent to the removal of strong base with respect to the acid-base conditions. The water concentration and the concentration of electrolytes will of course be different .

The relationship between p H and T.A. is indicated by means of the titration curve (PH as abscissa and added strong base as ordinate, see FIGURE 1). The slope of the titration curve gives the buffer capacity defined by Van Slyke (1922) : p = d Base/d p H .

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42 Annals New York Academy of Sciences

3 4 5 6 7 8 9 1 0 1 1 1 1 1 1 I l l I . . I 1

90 -

- -

-

1 I I I I l l 1 1 1

3 4 5 6 7 8 9 1 0 1 1

90

70

50

30

10

10

30

50

70

90

FIGURE 1. Titration curve for human urine. The curve was obtained as follows: 200 rl. urine, + 200 rl. 0.1 N HCl, vacuum for

elimination of COZ, + 800 ~ 1 . H20, titrating with 0.1 N NaOH using glass electrode and Titrigraph (Radiometer, Copenhagen). T.A. (end point pH = 7.40 at Pco2 = 0) is 15 meq./l.

The dotted curve represents the result by stopping the titration at pH 7.40, adding 200 pl. 30 per cent formaldehyde solution and titrating once again with 0.1 N NaOH. The extra T.A. (end point 7.40) represents the ammonia. The result is 59 meq./l. minus the blank value for the formaldehyde solution which was 24 meq./l, i.e. NH; = 35 meq./l. T.A. + NH; = net acid = 50 meq./l.

The concept “total acid” of a complex buffer solution has no meaning. This postulate is explained as follows: An acid is defined as a molecule which will give off a hydrogen ion under the given conditions. Thus, in order to determine the amount of acid as such i t is necessary to add strong base to determine the amount of hydrogen ions given off. In other words, it is in- herent in the definition of an acid that it must be determined by titration. The problem is that no single end point can be selected which is satisfactory for each of several weak acids. In the case of urine, e.g., it is impossible to decide whether the titration should be continued to pH 9, 10, or 11.

Similarly the concept “total base” of a solution has no meaning. According to the old “medical” acid-base definition “total base” indicated the sum of positive charges of cations. Using BrBnsted’s acid-base definition “total base” is often erroneously taken to indicate the sum of negative charges of the anions. However, not all bases are anions, as e.g. the amino groups of proteins.

Siggaard-Andersen: Titratable Acid or Base 43

The concept buffer base means a base which acts as a buffer, i.e. a weak base. However, it depends on the actual pH value which bases act as buffers. Thus a t a p H of 7.4 lactate is not a buffer base, but as pH falls and approaches the pK of lactic acid the lactate begins to function as a buffer base, and if the p H falls even further chloride may become a buffer base. One might define the buffer bases a t one given pH value as those bases for which the pKa lies within + or - 3 p H units. Bases with a lower pKn practically do not act as bases a t this pH value, whereas bases with a higher pKa are practically completely converted to the corre- sponding acid. However, defined in this way the concept loses its interest as indicator of the added strong acid or base, because as the pH changes new buffer bases appear.

Description of the Acid-Base Status of Body Fluids

The characteristic of the body fluids as buffer solutions is that they are open systems in equilibrium with a gas phase, namely alveolar air, with a certain carbon dioxide tension.

This state of affairs is represented schematically in FIGURE 2. The pH of the solution can be altered by two essentially independent means: by increasing or reducing the carbon dioxide tension of the gas phase, or by adding acid or base (acids other than carbonic acid, i.e. non-volatile acid). The actual change in pH depends on the buffer capacity of the non-bicarbonate buffers.

We call the added nonvolatile acid or base the T.A. because we are able to titrate back to the original pH and Pco2. In other words p H is a function of the Pcoz and the T.A. together with the buffer capacity: pH = F(Pco2, T.A.), where F depends on the nonbicarbonate buffer capacity.

FIGURE 2. Acid-base characteristics of a biological system. It is an open system in equilibrium with a gas phase. Changes in Pco2 alter the pH without affecting the content of non-volatile acid or

base. Addition of nonvolatile acid or base alters the pH without affecting the PCOZ.

44 Annals N e w York Academy of Sciences

Added nonvolatile acid is buffered by bicarbonate and the carbon dioxide tension tends to rise. However, because we are talking of a system in equilibrium with a constant Pco2 the liberated CO, disappears in the gas phase.

Determination of the excess acid or base when titrating to the original p H in a closed system where the total-C02 is constant has no physiological interest.

T.A. is a parameter which has been extensively used for the description of urine or gastric juice. However, it has not been much employed for descrip- tion of plasma or whole blood.

Titratable Acid or Base of Urine

Different endpoints of titration may be used for the determination of this quantity:

(1) T.A. p H determined by titrating to p H = 7.40 a t Pcor = O . This is easily achieved by adding a known excess of strong acid, eliminating the evolved C 0 2 and back titrating to p H = 7.40 with strong base.

p c o 2 - 4 n determined by titrating to p H = 7.40 with strong base a t Pco2 = 40. This should be performed by titrating while bubbling with a gas mixture with a Pco2 of 40. This value is approximately 25 meq./l greater than the first value due to the bicarbonate at p H = 7.40 and

(3) T.A. p ~ i ; 4 fol.i C O ~ determined by titrating to p H = 7.40 without loss of C02. This is performed by titrating in a closed vessel or often by carrying through the titration in an open vessel so rapidly that little CO, loss occurs. The value is greater than the first value mentioned by an amount equal to the bicarbonate concentration of the urine calculated from the constant total- COZ and pH = 7.40.

Dilution o f the urine also influences the result by decreasing the T.A. (Nutbourne, 1961).

The temperature of titration is seldom specified. Most often the titration is carried out a t room temperature but the result of course will be different when titrating a t 38°C.

All these conditions can be standardized (J@rgensen, 1957). However, difficulties arise when one attempts to asign a physiological significance to the result. Two questions may be asked:

p l o z -

(2) T.A. p H

P C O ~ = 40.

(1) How much H ' is excreted by the kidney tubules into the urine? (2) How much acid or base is lost from the organism? Ad (1): Theoretically the most correct performance must be to titrate

the urine to the p H of the ultrafiltrate a t the Pco2, temperature, and ionic strength of the ultrafiltrate. I t is evident that this can not be performed as a routine method.

Ad (2): Several substances behave as acids or bases in the orgar8i:m but are not determined by titrating the urine to p H 7.40, e.g. NPT. lwtate,

Siggaard-Andersen: Titratable Acid or Base 45

acetoacetate, etc. Excretion of ammonia represents a loss of acid. I t can be determined titrimetrically by a modified SQrensen formal titration (FIGURE 1). JQrgensen (1957) calls the sum of T.A.,740, P e o 2 = 0 , and NHZthe net acid of the urine. Lactate and several other organic anions, which do not act as bases in uitm a t a p H of 7.4, nevertheless behave as bases in the organism because they accept a H + during oxygenation to COs and H20. The only theoretically possible method for determination of such substances which act as acids or bases in uiuo but not in uitro must be a biological method, measur- ing the effect of administration to an organism. The same method must be employed for determination of the in uiuo acid-base properties of the diet and the stools. This is of course very complicated and exact methods for determining the hydrogen ion balance of the organism under all conditions are not available.

Titratable Acid of Gastric Juice

This parameter can be determined following different types of stimulation of the gastric secretion (standard meals, caffeine, insulin, histamine or histamine analogs). Previously the titration was carried out using different indicators, e.g. the Giintzburg reaction, Congo, Litmus, Phenolphthalein, corresponding to p H end points of approximately 3.3, 4, 6.5, and 9.5 respectively. Nowadays it is easier to use a glass electrode and record the whole titration curve, or simply the T.A. upon titration to p H = 7.40 a t Pco? = 0.

Titratable Acid or Base of P l a s m a or Whole Blood

The most natural end point for a titration of the excess of non-volatile acid or base of plasma or whole blood is a normal pH a t normal PCOZ and temperature. The endpoint has been chosen to be pH = 7.40 a t PCO~ = 40 mm. Hg a t 38°C. (Siggaard-Andersen, 1962). Originally in our laboratory the end point was defined on the basis of normal values measured on NaF stabilized whole blood a t p H = 7.38 and Pco2 = 40 mm. Hg a t 38%. (Siggaard-Andersen & Engel, 1960). However it was shown that NaF alters the acid-base values of the blood (Siggaard-Andersen, 1961).

We have chosen to call this quantity the base excess (BE), a negative value indicating a base deficit. Singer & Hastings (1948) call the quantity A buffer base, while Davenport (1958) calls it extra fixed acid or base.

Other quantities which have been used as indicators of the excess of nonvolatile acid or base of plasma or whole blood:

(1) Actual bicarbonate or total-C02 (Henderson, 1909; Peters & Van Slyke, 1932; Schwartz & Relman, 1963) falls when nonvolatile acid accumulates and rises when base accumulates. However, this quantity also varies with the Pco2. In many cases where the PCOZ can be estimated clinically the bicarbonate

46 Annals New York Academy of Sciences

or total-COz will be sufficient as indicators of the excess of non-volatile acid or base. In other cases the quantity can be used as indicator of hypercapnia or hypocapnia.

I t could be argued that a knowledge of the bicarbonate ion concentration per se is important for biochemical or physiological reasons. Thus a specific effect of bicarbonate on lipid metabolism has been demonstrated (Hastings, 1964). Further a knowledge of the bicarbonate concentration is necessary in cases where the fraction of anions called “undetermined anions” is to be calculated in order to estimate accumulation of different organic anions.

Thus the bicarbonate concentration and the BE are not equivalent para- meters, but different from a biochemical as well as a physiological point of view.

(2) Standard bicarbonate at a standard Pcor of 40 mm. Hg. In order to avoid the variation of the bicarbonate concentration with Pco2 Hasselbalch (Hassel- balch, 1916; Van Slyke & Cullen, 1917; Henderson & Morris, 1917; Astrup, 1957; Peirce, 1960) choose to standardize the PcoZ a t 40 mm. Hg. before measuring the pH, total COz or bicarbonate. The quantity obtained falls and rises with accumulation of nonvolatile acid and base, respectively, but the change in the standard bicarbonate also depends on the buffer capacity of nonbicarbonate buffers and does not directly indicate the accumulated amount of acid or base.

(3) Bicarbonate at a standard p H of 7.40 (Van Slyke standard bicarbonate) represents the bicarbonate concentration of the sample when the PCOZ is changed until the pH is 7.40. This quantity was proposed by Van Slyke (1921) because the change from the normal value equals the accumulated nonvolatile acid or base. Because when the pH is 7.40 all buffers have their normal distribution between buffer acid and buffer base except the carbonic acid/bicarbonate system. Thus all excess of nonvolatile acid or base have been buffered by the latter system and changes in this system reflects the base excess.

A small detail is worth mentioning. Not all buffers are restored to normal distribution by changing the Pco2 until1 the p H is 7.40. Theoretically the pK of some aminogroups is lowered by the formation of carbamino com- pounds. As a result the change in bicarbonate (actually bicarbonate plus carbaminate) a t constant p H should be slightly lower than the BE (Warburg, 1922). However, this effect could not be demonstrated in plasma (Siggaard- Andersen, 1963).

(4) Buffer base (BB) introduced by Singer & Hastings (1948) is defined as the biological buffer salts capable of neutralizing strong acid, for blood the sum of bicarbonate and proteinate.

This quantity rises with accumulation of bases and falls with the ac- cumulation of acids except when the corresponding base of the acid is itself bicarbonate or proteinate. Thus BB is independent of accumulation of carbonic acid, and independent of changes in oxygen saturation of the sample.

Siggaard-Andersen: Ti t ra table Acid or Base 47

I have previously mentioned that the BB concept is not well defined. A more exact definition of BB is given later.

Direct determination of the BE of plasma or whole blood: (1) Titration with strong acid or base to a p H of 7.40 a t a Pcoz of 40 mm.

Hg a t 3 8 T . is difficult as a routine method due to the difficulties in main- taining the P C O ~ = 40 during titration.

(2) Titration with strong base to a pH of 7.40 a t Pcoz = 0 a t 38°C. after elimination of COn with a known excess of strong acid (Van Slyke, Stillman & Cullen, 1919). The titratable base determined by this method gives the Van Slyke standard bicarbonate and the BE can be calculated from this value by subtracting the bicarbonate concentration of the sample when p H = 7.40 and Pco:! = 40 (with S = 0.0300 and p K = 6.10, bicarbonate = 24.0 meq./ I).

(3) The determination of "the titratable alkali of the blood" according to Greenwald 8z Lewman (1922) is mainly of historical interest but deserves to

FIGURE 3. CO, equilibration curve of plasma. Plotted in the pH, HCO, diagram the equilibration curve is practically a straight

line (actually slightly convex downwards). The vertical displacement of the equilibra- tion line gives the base excess. This curve is identical with the titration curve for the nonbicarbonate buffers when titrating with strong acid.

48 Annals New Y ork Academy of Sciences

be mentioned because the value obtained by their method actually is the buffer base of Singer & Hastings. They add an excess of picric acid which precipitates the proteins and eliminates COZ. The total amount of picric acid of the filtrate is measured together with the titratable acid of the filtrate, using a convenient color indicator. When titrating to the original pH of the sample the result will be the sum of bicarbonate and proteinate of the sample, i.e. BB.

(4) Titrating with CO,. The Van Slyke standard bicarbonate can be determined by equilibrating the sample with different gas mixtures of known carbon dioxide tension and measuring or calculating the bicarbonate con- centration. By interpolation or extrapolation to pH = 7.40 the Van Slyke standard bicarbonate and the BE can be calculated. Plotted in the pH, HCOi diagram (FIGURE 3) the equilibration curve is practically a straight line, actually slightly convex downwards for plasma and oxygenated whole blood (Van Slyke, 1921; Warburg, 1922; Siggaard-Andersen, 1963). Plotted in the pH, log Pcoz diagram (FIGURE 4) a good approximation to a straight line is also obtained (Astrup, 1954; Siggaard-Andersen 1963).

- 20

FIGURE 4. CO, equilibration curve of plasma. Plotted in the p H , log PCO~ diagram the equilibration curve is also practically a

straight line (actually slightly convex upwards). The vertical displacement of the equilibration line is a measure of the base excess.

Siggaard-Andersen: Titratable Acid or Base 49

Calculation of the base excess:

(1) Calculation of BE from change in bicarbonate at pH = 7.40. The base excess of the plasma (BE,) can be calculated as: BE, = a Van Slyke standard bicarbonate = Van Slyke st. bic.-24. For whole blood the change in Van Slyke st. bic. also equals the base excess of the whole blood (BEb) provided the whole blood bicarbonate is used. The change in plasma bicarbonate a t pH 7.40, how- ever, is greater than the change in whole blood bicarbonate (FIGURE 5), and the formula for BEb according to own experimental data becomes:

BEb = (1 - 0.0143 Hb) . A Van Slyke st. bic.,

where Hb is hemoglobin concentration in g./100 ml. A slightly different formula is obtained when using the data of Van Slyke & Sendroy (1928) for the relationship between whole blood bicarbonate and plasma bicarbonate.

The Van Slyke st. bic. can be calculated from the actual bicarbonate and pH, when the slope of the C 0 2 equilibration line is known. The latter depends

69 7.0 7.1 7.2 7.3 7A 7.5 7.6 pH I , l , l , l , l r l , l , l , l ,

FIGURE 5. CO, equilibration curves for plasma and whole blood, plotted in a diagram with p H as abscissa and plasma bicarbonate as ordinate. Note that the vertical displace- ment a t constant p H is greater for whole blood than for plasma.

50 Annals New York Academy of Sciences

on the buffer capacity of nonbicarbonate buffers, i.e. mainly on the hemoglobin concentration. On the assumption that all the equilibration lines are parallel for a given Hb. conc. the slope can be expressed as

p = = (9.5 + 1.63 Hb) meq./lperpH unit

The formula derived from Peters & Van Slyke (1932) is given for comparison:

The complete formula for the calculation of BEh from the actual bi- @ = - (8.2 + 1.39Hb).

carbonate of the plasma and the p H thus is:

BEh = (1 - 0.0143 Hb) (act. bicarb. - (9.5+1.63 Hb) (7.40-pH) - 24.0)

( 2 ) Calculation of BE from base excess and buffer base curves. (A) pH, HCO; diagram: The assumption made in the preceeding section,

that the pH, HCO; equilibration lines for a given Hb. conc. are all parallel, is theoretically exact only for plasma not for whole blood. The equilibration lines for whole blood converge slightly towards the acid side due to the change in bicarbonate distribution between plasma and red cells with pH.

It is possible, however, to calculate the degree of displacement of the equilibration line by addition of nonvolatile acid or base, and the change in slope of the equilibration line, by means of the base excess curve and the buffer base curve.

The base excess curve is the locus for corresponding pH, Pco?, and HCO, values of plasma, where the excess of nonvolatile acid or base is equal in plasma and red cells. For a given blood sample with pH, Pco~, and bicarbonate coordinates fitting the base excess curve, we can remove red cells anaerobically and thus change the hemoglobin conc. and the slope of the equilibration line without changing the whole blood base excess. In other words with the addition of nonvolatile acid or base the equilibration line is displaced meq, for meq. along the base excess curve irrespective of the hemoglobin conc. of the sample.

In order to calculate the change in slope of the equilibration line during this displacement another “strong acid or base displacement curve” situated at some distance from the base excess curve is necessary. The exact position is unimportant. However, it is conveniently chosen so that it represents the locus for corresponding pH, Pco?, and HCO, values of plasma, where the buffer base is equal in plasma and red cells (Siggaard-Andersen & Engel, 1960). The nomogram shown in FIGURE 6 represents the function

pH = F(pH, HCOB, BEb, Hb)

When the pH, the plasma bicarbonate conc., and the Hb. conc. of the sample are known the BEh is calculated by drawing the equilibration line for BE = 0 corresponding to the Hb. conc, and displacing the line equally much along the BE and the BB curves until the line passes the point repre-

Siggaard-Andersen: Titratable Acid or Base 51

FIGURE 6. pH-bicarbonate nomogram for blood.

sented by the actual pH and bicarbonate concentration. The BEb is then read off a t the intersection point between the equilibration line and the BE curve.

It was previously mentioned that the BB, representing the sum of buffer bases, is an ill-defined concept, and hence the buffer base curve should be an ill defined curve. I t is possible, however, to define a chemical quantity the numerical value of which approaches the sum of bicarbonate and proteinate of arterial blood, namely the titratable base when titrating with strong acid to a pH of 7.093 a t a Pcoz of 92.0 a t 38%. adding an arbitrary constant of 41.7.

(B) pH, log Pcoz diagram: The relationship between pH, Pcoz, BEb, and Hb. conc. can also be expressed by means of the pH, log Pcoz diagram (FIGURE 7). This diagram is especially useful1 when the equilibration technique for determining the P C O ~ is used.

This type of representation of the functions using a Cartesian coordinate system may be termed a Cartesian nomogram or a curue nomogram. The same functions may be represented by a d’0cagne nomogram or alignment nomo- gram (FIGURE 8) (Van Slyke & Sendroy 1928; Singer & Hastings, 1948; Siggaard-Andersen, 1963). The latter represents the simplest possible illustration of the relationship between pH, Pcoz, (HCO,),, (total-COd,,,

52 Annals New York Academy of Sciences

50 i 1 almdard btcarbonslr

FIGURE 7. pH-log PcoL nomogram for blood (Siggaard-Andersen, 1962)

BEb, Hb. conc. The alignment nomogram of Singer & Hastings gives the BBb and includes whole blood total-C02 as well.

Base release on reduction of hemoglobin: It is well known that oxyhemo- globin in the p H range from 6.1 to 9.0 is a stronger acid than reduced hemoglobin. However, the amount of base formed on reduction of hemoglobin varies with the pH and also with the oxygen saturation (German & Wyman, 1937; Margaria, 1957). In the pH range 7.0 - 7.8, however, it may be reckoned without any great error that there is a constant value of 0.3 meq./lO g . Hb for human blood (Siggaard-Andersen, 1963). Thus, the BE of venous blood is higher than the BE of the completely oxygenated blood by an amount of 0.3 x H b Y oxygen saturation/lOO meq./l.

As an indicator of the accumulation of nonvolatile acid or base in the organism (the “metabolic” acid-base factor) some authors prefer the whole blood BE of the in uitra oxygenated blood (Crawford & Holaday, 1964), while others, including myself, prefer the BEb of the actual oxygen saturation. (Siggaard-Andersen, 1964).

Total-C02 mMal/ I plasma 6%

KO; mEq/ I plasma.

Bare Excess

mEq/l blood or ~ I r r m a .

Y o - 30

Pco, mm Hg

I0 1 15

130

FIGURE 8. Acid-base alignment nomogram for blood (Siggaard-Andersen, 1963).

54

6 10

680

6 90

7 00

7 10

7 20

7 30

7 40.

DH

- - 220

200

at teC

- - - - -

-

-

-

-

Annals New York Academy of Sciences

-170 -160

-150

-140

-130

-120

-110

-100

90

pH-Ro, -temperature nomogram for whole Mood

wordmp to Sq@ard-And*rsm,l963 Pco2 rnrn Hg at t°C at 38.C

- - - - - - -

t e c

140 130

-120 110 100

90

80

. I 0

60

50

LO

30

20

FIGURE 9. p H - Pco? - temperature nomogram.

80

70

60

50

Lo

30

20 - 770-

Base excess at varying temperatures: The base excess is defined at 38°C. If the temperature of plasma or whole blood in a closed system is changed both pH and Pco2 changes, but as long as nonvolatile acid or base is not added, the BE remains constant.

- - - - - - - -

10

Siggaard-Andersen: Titratable Acid or Base 55

The methods mentioned in the previous sections for calculating the BE frompH and Pcoz requires that the pH and PCOZ refer to 38°C. Conversion of values for p H and Pcoz valid a t other temperatures to values valid a t 38°C. can be performed by means of the nomogram shown in FIGURE 9. The formulas underlying the nomogram are (Rosenthal, 1948, Siggaard-Andersen, 1963) :

pHt = pH3a0 + 0.0147(38 - to)

log PCOz, t = log Pco2.3a0 - 0.21(38 - to)

The change in whole blood base excess by changes in PCOZ in vivo: I t is well known that changes in Pco2 in whole blood in uitro causes a redistribution of water, chloride, and hydrogen ions between red cells and plasma (Van Slyke, Wu & McLean, 1923). As the Pco2 rises the hydrogen ion concentration falls more in the poorly buffered plasma than in the well-buffered red cells, and thus H + passes from the plasma into the red cells. As a consequence the BE of the plasma rises, while the BE of the cells falls.

FIGURE 10. CO, equilibration curve for whole organism (Siggaard-Andersen, 1962). Line C: I n uiw relation between pH and Pcol in plasma in acute CO, inhalation or hyperventilation. Line B: CO, equilibration curve for whole blood in uitro (Hb =

15 g./100 ml). Line A: CO, equilibration curve for separated plasma in uitro.

56 Annals New York Academy of Sciences

A similar change in distribution of H+ is seen when the PCOZ is altered in vivo. I t has generally been supposed that the whole blood BE would be constant during acute changes of the Pco2 in vivo (Henderson & Haggard, 1918; Davies, Haldane & Kennaway, 1920; Shock & Hastings, 1935). Actually, however, as the Pcoz rises hydrogen ions pass from the less buffered extravascular tissues into the better buffered blood. As a consequence the BE of the blood falls. This state of affairs was originally described by Shaw & Messer (1932) and recently by Siggaard-Andersen (1962) and Cohen, Brackett & Schwartz (1964). The relationship between the equilibration curve for plasma and whole blood in uitro, and the equilibration curve for the whole organism in vivo is shown in FIGURE 10.

Conclusions

(1) Titratable acid, or base excess, is a valuable parameter for the description of biological fluids.

(2) The concepts “total acid” or “total base” are rather poorly defined. In contrast to the electrolytes (Na+ , K + , HCOR, etc.) the amount of acid or base must be given as the titratable acid or base, i.e. as an excess or deficit.

(3) The actual bicarbonate concentration is not an independent acid-base variable as i t changes with the Pco2 as well as with the excess of nonvolatile acid or base.

(4) The base excess of the blood changes with accumulation of nonvolatile acid or base in the blood, whether this accumulation is due to an accumula- tion within the whole organism or to a redistribution of hydrogen ions between the body compartments. The changes due to a redistribution as seen in hypercapnia or hypocapnia (or potassium depletion) are small and usually without clinical significance.

References

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ASTRUP, P. 1954. Om erkendelsen af forstyrrelser i organismens syre og base stofskifte. Ugeskr. Laeg. 1 1 6 758-776.

ASTRUP, P. 1957. Erkennung der Storungen des Saure/Base-Stoffwechsels und ihre klinische Bedeutung. Klin. Wochenschr. 35: 749-753.

BRBNSTED, J. N. 1923. Einige Bemerkungen iiber den Begriff der Sauren und Basen. Rec. Trav. Chim. Pays-Bas 42: 718-728.

COHEN, J. cJ., N. C. BRACKETT, JR . & W. B. SCHWARTZ. 1964. The nature of the carbon dioxide titration curve in the normal dog. J. Clin. Invest. 43: 777-786.

CRAWFORD, S. & D. HOLADAY. 1964. Acidibase disturbances (letter to the Editor). Lancet 1.

DAVENPORT, H. W. 1958. The ABC of acid-base chemistry. 4th Ed. Univ. Chicago Press. Chicago, Ill.

GERMAN, B. & J. WYMAN, JR. 1937. The titration curves of oxygenated and reduced hemoglobin. J. Biol. Chem. 117: 533-550.

DAVIES, H. W., J. B. S. HALDANE & E. KENNAWAY. 1920. Experiments on the regulation of the blood’s alkalinity. J. Physiol. 54: 32-45.

Siggaard-Andersen: Titratable Acid or Base 57

GREENWALD, I. & G. LEWMAN. 1922. The determination of the titratable alkali of the

HASSELBALCH, K. A. 1916. Die “reduzierte” und die “regulierte.” Wasserstoffzahl

HASTINGS, A. B. 1964. Personal communication. HENDERSON, L. J. 1909. Das Gleichgewicht zwischen Basen und Sauren im tierischen

Organismus. Ergebn. Physiol. 8: 254-325. HENDERSON, Y. & H. W. HAGGARD. 1918. Respiratory regulation of the COZ capacity

of the blood. J. Biol. Chem. 33: 333-371. HENDERSON, Y. & W. H. MORRISS. 1917. Applications of gas analysis. I. The determina-

tion of CO, in alveolar air and blood, and the C 0 2 combining power of plasma, and of whole blood. J. Biol. Chem. 31: 217-227.

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