mains to be understood about the interplay of chemical activity among the various parts of these structures, particularly the substituents on the ring. There are indications that some of the nitriles originally formed may react with free radicals to increase their chainlength- for example, the apparent presence of pentanenitrile, hexanenitrile, and 2-methylpropenenitrile among the products of sodium barbital. If this is so, i t is likely that some of the unidentified peaks may also be nitriles. The loss of a side chain, or more than one, may be expected to result in the formation of unsaturated products; hence further studies should give consideration to unsaturated nitri!es, in addition to the propenenitrile which is reported here. Hydrocarbons, the group of products reported by Janak (4, 5 ) , were not explored in this study as they did not appear to be major products. The probable presence of benzene among the products of several barbiturates was noted, and in fact the data reported here are not inconsistent with the presence of a considerable quantity of hydrocarbons. It appears probable that the conditions of chromatography utilized by Janak may have led to retention of nitriles in the column.
Ethanol was the only alcohol indicated, and acetone, the only ketone. The absence of other ketones would suggest that the acetone is formed from the barbituric acid ring itself. Water, which was found among the products of several barbiturates, would probably have been found in all of them if they had all been chromatographed a t 40' C. Gases such as CO, CO,, Sz, 02, and NzO were not investigated in this study. One or more of them might reasonably be expected among the pyrolysis products. A quantitative study of their formation might usefully complement the data presented here. Two interesting sidelights appeared in this study. (a) The matching of retention times revealed the likely identity of an early negative peak as being that of the long-postulated hydrogen isocyanide, HNC ( 7 ) . (b) Contrary to the usual laboratory methods for preparing methane and acetone by the pyrolysis of sodium and calcium acetates, respectively, acetone was found to be the main pyrolytic product of both of these salts under the conditions of this investigation. This may indicate that the hydrocarbons produced by the pyrolysis of the sodium salts of the fatty acids result from secondary reactions among the primary products.
ACKNOWLEDGMENT
We acknowledge the technical assistance of Charles R. Fontan in the preparation of nitriles and isonitriles used as reference standards, and the advice of James Cason of the College of Chemistry. Chemical standards were supplied by the Chemistry Department. LITERATURE CITED
1,LaboratoryJames, Rapoport, Text in Organic
(1 Cason,
Henry, Chemistry," pp. 72, 84, 89, Prentice-Hall, K.J. ( 2 ) Friedman, L., Schechter, H., J . Org. Chem. 25, 877 (1960). (31 Jackson, H. L., McKiusick, B. C., Org. Syn. 35, 62 (19;5). ( 4 ) Janak, J., C'ollectzon of Czech. Chem. Commun. 25, 1780 (1960). ( 5 ) Janak, J., Suture 185, 684 (1960). (6) Selson, D. F., Kirk, P. L., AXAL. CHEW34,899 (1962). ( 7 ) Selson, D. F., Kirk, P. L., J. Chromatog. 12, 167 (1963). (8) Strassburger, J., Brauer, G. M., Tryon, M., Forziati, A. E., Ax.4~. CHEM.32,454 (1960). RECEIVED for revie-- September 5 , 1963. Accepted January 22, 1954. Fall Meeting, 1962, The California Association of Criminalists, Concord, Calif. This work was supported by grants from the Kational Institutes of Health, U. S. Public Health Service (RG-4372 and RG-5802) and from the Research Committee, University of California.
Correlation between Apparent pH and Acid or Base Concentration in ASTM Medium OREST POPOVYCH' Analytical Research Division, Esso Research & Engineering Co., linden, N. 1.
b
Equations were derived which correlate the apparent p H of acid and base solutions in amphiprotic nonaqueous solvents with their concentrations. Within moderate concentration ranges, these relationships are closely approximated b y straight-line equations. The slopes of these lines are a function of the nature of the ionic dissociation, while the intercepts depend on the magnitude of the ionic dissociation constant, on changes in the liquid-junction potentials and on the primary medium effect. The above relationships were verified for solutions of perchloric, hydrochloric, SUIfuric, nitric, benzoic, acetic, a mixture of benzoic and acetic acids, potassium hydroxide, piperidine, and lutidine in the ASTM medium. The latter consists of 50.0% toluene, 49.5% isopropyl alcohol, 0.5% water, by volume, and i s used in the ASTM titrations of acids and bases ( I ) . Elec-
878
ANALYTICAL CHEMISTRY
trolytic conductance of HC104, HzS04, HN03, and HCI in the ASTM medium was also studied. A combination of apparent p H and conductance data showed that the sum of liquid-junction and primary medium effects was roughly a constant characteristic of the medium only.
T
definition of pH (8)makes it possible to convert to apparent p H any e.m.f., E , developed between a conventional glasscalomel electrode pair previously standardized against aqueous standard p H buffers : pH
HE OPERATIONAL
=
+
pH, (E
- E8)/(2.3026 RT/F) (1)
where the subscript s designates ap aqueous standard buffer solution. The above procedure coupled with the availability of meters which read pH
directly has extended the measurement of apparent p H to partially and totally nonaqueous solvents. The potential misinterpretation of such p H readings has prompted extensive reviews of the concept and the limitations of p H (2-5, 10, 17, 25). Although it is generally recognized that pH readings do not represent hydrogen ion activity outside of dilute aqueous solutions, the apparent p H in nonaqueous solveiits is by now well established in industrial use as an index of relative acidity. .Is a result, need does exist for obtaining a better insight into the meaning of apparent p H in nonaqueous solvents. I n the petroleum industry, those acidbase measurements which involve ASTM specifications are made in a medium consisting of 50.0yo toluene, 49.5% 1 Present addrese, Department of Chemistry, Brooklyn College of the City University of New York, Brooklyn 10, N. Y.
isopropyl alcohol, and 0.5% water (1); we label this mixture i,he ASThl medium (dielectric constant 2% 8). The results of such measurements are often used as criteria for product quality and performance. Because of its importance in the petroleum industry, this threesolvent medium was chosen for testing the derived pH relationships. EXPERlMliNTAL
Measurements of pH. All p H measurements were made on a Beckman expanded scale p H meter, hlodel 76. The same genwal-purpose glass electrode and slec ve-type calomel electrode, standardi aed against aqueous Beckman buffers, were used throughout. The measurements were made in a constant-temperature bath at 25" C., with nitro:en blanketing for bases and weak acids. Equilibrium (stable p H reading) in the ASTM medium required as little as 1 to 3 minutes for strong acids and bases. However, for solutions of weal.: acids and bases, the observed p H drifted with time. Therefore, the readings were followed with the aid of a recorder, until a reasonably stable reading was obtained after 30 to 60 minutes. I n all cases, the electrodes were repeatedly equilibrated with a t least three fresh portions of a given solution. Conductance Measurements. The purification of solve i t s was described elsewhere (29). Measurements were made on an Industrial Instruments RCl6R2 bridge a t 1000 cycles. A platinum cell having a constant of 0.0100 was used. -111 measurements were made a t 35' C. with several equilibrations before a final reading. The concentrations were determined by potentiometric titration. It was observed that the conductance of strong acids in the dSTR4 solvent increases appreciably upon standing. Therefore, the solutions were prepared immediately before measurements were made and they were used when fresh. THEOI!Y
A p H determination is made by measuring the potentials of two solutions: an aqueous standard buffer to which a known p H kind been assigned by the National Bur1:au of Standards, and an unknown solution of interest, a nonaqueous solution iJi this case. Both are measured by the same glass calomel electrode system, whose potential is given by the equation below:
where E , is the sum of liquid-junction potentials a t both ends of the salt bridge and (Cl-), the activity of chloride ions in the calomel reference cell. When the above cell potential is combined with the operational definition of pH, the terms EF and log (C1-)
cancel out, because they are the same for the standard and for the unknown. Thus, the (apparent) pH reading ( ~ H R ) is always in error compared to true p H (hydrogen ion activity with respect to aqueous standard state), first of all by a factor involving the difference between the liquid-junction potentials in the aqueous buffer and in nonaqueous unknown: pH
-~
H = R
-AEj X F 2.3 RT
(3)
(This error arises also in aqueous solutions, but is, in general, assumed to be negligible.) Since the above true p H is expressed retaining dilute aqueous solution as the standard state, it is related to the nonaqueous hydrogen ion concentration through two activity coefficients of the hydrogen ion, fc and fm. The first is the familiar concentration activity coefficient, which approaches unity a t infinite dilution in any solvent. The second is the activity coefficient reflecting energy changes of the proton in the transition from the aqueous to the nonaqueous medium, known as the primary medium effect: it has also been referred to as the "degenerate activity coefficient" (20). The final expression for the negative logarithm of hydrogen ion concentration in a nonaqueous solvent in terms of the above parameters is: --log /H+]N= ~ H R
Thus, the major sources of discrepancy between the apparent p H reading and the negative logarithm of hydrogen ion activity are the values of AEj and of the primary medium effect, fm. Liquid-junction potentials have been estimated to introduce an error of as many as three pH units (f7,27), while fm has been reported to attain values corresponding up to five p H units in common solvents and higher values in strongly acidic or basic media (17, 22). Neither of these parameters can be evaluated strictly by thermodynamics. As a result, it may be of greater practical interest a t the present time to correlate the apparent p H of a nonaqueous solution with its formal acid (base) concentration, and also, with the p H exhibited by an aqueous solution containing the same formal concentration of the same acid or base. The latter is actually a superimposition of the first relationship upon that between the p H and the formal concentration of the same acid (base) in water. I n the present paper, experimental data are analyzed in terms of the above two correlations. Equation 4 gives the relationship between the p H reading ( ~ H Rin ) a nonaqueous solvent and the hydrogen ion concentration and activity. To obtain a similar relationship for the PHR us. acid concentration, we must know the nature of its ionic dissociation in the given solvent. Assuming the simplest case, where, omitting solvation,
the net dissociation of an acid HA can be expressed by the constant:
(5) we can rewrite the equation so as to express the p H in terms of total acid concentration present a t equilibrium, [HA] :
-
PHR = -'/zlogK
1% [HA] - A (6a)
'/?
The activity of the undissociated acid (HA) is assumed to equal concentration. Although Equation 5 is strictly valid only if [HA] is the concentration of acid a t equilibrium and not the total acid concentration added to the solution, the two may be equated in lowdielectric media as a first approximation. Thus, for the simplest case, a plot of apparent p H us. the negative logarithm of acid concentration in a nonaqueous solvent should yield theoretically 0.5 as the slope and (--I/? log K - A) as the intercept. The hopeful assumption involved here is that A is not a function of hydrogen ion concentration. Such an assumption is not without foundation. Because log fm is a lunction of the medium only, the question reduces to whether the unknown liquidjunction potential a t the boundary between an aqueous salt bridge and a nonaqueous solution is independent of concentration. Previous experience in nonaqueous potentiometry (9, 11-13) indicated that a constancy of the liquidjunction potential could be assumed, as a good approximation. Furthermore, soon after the first communication of the present study appeared in preprint form (30), Bates, Paabo, and Robinson (9) published their extensive investigations on the p H in alcohol-water solvents, in which the liquid-junction potentials were found to be constant for buffered solutions. Finally, to obtain a similar relationship for the apparent p H in a nonaqueous solvent us. the p H of a formally equivalent acid solution in water, we have to combine the above relationships with those between the pH and the [HA] in water. For strong acids, pHw = -log [H.4]
- logf,,
(7)
For weak acids,
pH, = '/z log K ,
-
'/2
log [HA]
(8)
(Subscript w stands for aqueous solutions.) Substituting the correspcnding expressions for log [HA] into Equation 6, we obtain: For strong acids,
PER= '/z pHw '/z logfcw
+
-
'/2
log K
-A
VOL. 36, NO, 4, APRIL 1964
(9)
879
and for a weak base,
For weak acids,
PHR = pHw
+
'/z
log '/z
Kw log K -
A
(10)
For concentration ranges where log can be assumed to be approximately constant, Equations 9 and 10 show that a plot of apparent p H us. the p H in water should yield straight lines with slopes of 0.5 for strong acids and of unity for weak acids. The intercept in the case of strong acids should be about the same as for the plot of apparent p H us. -log [HA]. For weak acids, the intercepts differ from the above by the term '/?log K,. Similar relationships hold between the apparent p H of basic solutions in nonaqueous amphiprotic solvents and the p H in water. For any base in a nonaqueous solvent of low dielectric constant, we have:
fc
- log K , - A
(11)
where K , is the autoprobolysis constant of the solvent. In terms of aqueous pH, Equation 10 becomes, for a strong base,
~ H= R
+
'/z '/2
Consequently, for strong and weak bases, as in the case of acids, plots of apparent p H us. aqueous p H should yield straight lines with slopes of 0.5 and 1, respectively. Unfortunately, the simple dissociation H-4 = H+ A- is not the only common process in nonaqueous solvents. At moderate acid concentrations, the reaction
+
H+
2HA
+ HA2
was found to be predominant in many s y s t e m (14, 15, 18, 21, 24, 28, 81-84), Similar processes are known for bases (16, 26):
+
PHR = '/z log K '/zlog PI
Table 111. Limiting Equivalent Conductance and Ionic Dissociation Constants of Acids in the ASTM Medium
pH, - 7 log K - log K ,
-A
(12)
Table 1. Apparent ptf in ASTM Medium vs. Formal Acid (Base) Concentration M
2B
+ HS e B H + + BS-
It is easy to see that for systems in which processes 14 and 15 are governing, plots of apparent p H us. the negative logarithm of equilibrium concentration would yield slopes of unity, and not of 0.5. Other dissociation processes are also possible, and a generalization can be made that if 1: and y are the powers to which ( H t ) and (HA) are raised in the equilibrium constant, the theoretical slope for the plot of p H us. -log [HA] is y/x.
Eauation: RESULTS
Acid or base Perchloric Hydrochloric Sulfuric Xitric Benzoic Acetic Benzoic acetic (equimolar) Potassium hydroxide Piperidine 2,4-Lutidine
0.776 0.702 0.774
+
-1.18
0.588
0.635 0.906
-0.20 4.71 4.40
0.968
4.15
-0.573 -0.956 -0.549
17.1 12.8 10.5
Measurement of pH. The apparent pH's of several acids and bases (listed in Tables I and 11) in the to lO-4M concentration range were measured both in the ASTM medium and in water, using the same glasscalomel electrode system. To obtain the straight-line equations, the results were analyzed by the method of least squares, assuming either the p H ~ , o ,or the negative logarithm of the known
Apparent p H in ASTM Medium vs. pH of Corresponding Solutions in Water
Table It.
Acid or base Perchloric Hydrochloric Nitric Benzoic Acetic Benzoic acetic equimolar) Potassium hydroxide Piperidine 2,4-Lutidine
+
880
B -1.85 -1.04
m
+
PHASTM= m' pHmo B' in ' B' 0.742 -1.78 0.706 -1.04 0.588 -2.200 1.105 2.52 1.57 0.872
ANALYTICAL CHEMISTRY
1.57 0.578 1.32 0.944
1.01 9.00
-4.24 0.749
Compared to pH in water, apparent pH is: Lower by 2-3 units Higher by 3-4 units Higher by 4-5 units Same, within 1 unit
HClOn H?S04' HC1 HNOI
44. I 33 0 37 4 34 1
3 . 9 2 x 10-5 2 57 x 10-8 3 86 x 6 86 x lo-'
molar concentration Jf to be the errorfree variable, 2 , and taking pHAsTM as the dependent variable, y. The straight-line constants are shown in Tables I and 11. Conductance Measurements. Electrolytic conductance of perchloric, sulfuric, hydrochloric, and nitric acid was measured as a function of con- 1 X 10-3kf centration (1 x range) in the ASTM solvent a t 25' C. Results are shown in Table 111. Limiting equivalent conductance .io and ionic dissociation constants K were calculated by the Fuoss-Shedlovsky method (19), programmed on an IBhI 7090 computer. DISCUSSION
Data in Table I show that most of the slopes observed experimentally deviate from the predicted value of It is possible that the more complex behavior predicted by Equations 14 and 15 is exhibited by acetic acid, its mixture with benzoic acid, and by piperidine, where the observed slopes are close to unity. However, for weak acids and bases there is no additional key available for testing the validity of the equations and, moreover, the measurements in such systems are comparatively uncertain. For the so-called strong acids, on the other hand, conductometric data made it possible to elucidate some of the apparent discrepancies between theory and experiment. From the ionic dissociation constants (Table 111) of strong acids, it was possible to estimate their mean ionic concentration, artivity, and the nonaqueous pH(-log of mean ionic activity). Equations 4 and Ab show that the difference between the p H reading and the p H nonaqueous obtained independently from conductance makes i t possible to calculate the overall medium effect A for all systems where the measurements mere available. The A was found to be independent of the nature of the acid and a comparatively insensitive function of hydrogen ion concentration (Figure 1). Indeed, if one considers that the proton is by far the most important
1.8-
3.40
CORRECTION FOR CONCENTRATION DEPENDENCE OF A .
1.6
1.4 1.2
SLOPE
1.0
A 3.10
I a
.
0.6
tic104 HCI
2.90 5
15 [H+] x l o 5 -L 10
20
25
Figure 1. Overall medium effect, A, as function of hydrogenion concentration
species in determining both the solvation energy and the imic mobility of all acids, this finding coines as no surprise. Since the primary mt.dium effect f m is a constant, the slight lrariation of A with concentration reflects primarily changes in the liquid-junction potentials. These in turn, should follow variations in hydrogen ion concentration in a consistent manner, which is indeed observed for all the acids studied. The A decreases with decreasing ionic concentration. Howevei., it is likely that changes in the asymmetry potential are also included in the observed variation of A. The approximate constancy of A for a given nonaqueous medium facilitates interpretation of apparent-pH data. It means, for example, that the differences in the apparent pH’s of acids in the ASTM medium are essentially due to differences in their cverall dissociation constants. Consequently, by knowing A, we can estimate from a reading of apparent p H both the true nonaqueous pH and the ionic dissociation constant of a n acid, without resorting to conductance measurements. This is very useful for those acsids whose conductances are too low for reliable direct determinations. The slight inconstsncy of A as a function of hydrogen ion concentration is the cause of the slight curvature in the relationships which were written in the form of straight-line equations and it is part of the cause for the slopes being higher than the predi:ted l/*. This is part of the reason why the derived relationships can be ,approximated by straight lines only within moderate concentration ranges. Initial neglecting of the ionic dissociation is another cause for the deviation of the slope from l/z. On the example of HC1 (Figure 2), it can be seen that a plot of the pHR vs. the -log acid molarity uncorrected for ionic dissociation (formal molarity) has a slope of 0.70; correction for ionic dissociation-Le., plotting -log molarity of undissociated acid-reduces the slope
DISSOCIATION ONLY A----CORRECTED F O R IONIC DISSOCIATION AND CONCENTRATION DEPENDENCE OF A.
0.4 0.2
2.80 0
0.702
0.8
2.401
0.0 2.0 Figure 2. tration
2.4
2.8
3.2
- log M
I
4.2
apparent pH vs. concen-
Hydrochloric acid:
I
3.6
0.50
I
I
PHN
Figure 3.
Apparent
v5.
true pH of acids in ASTM medium
to 0.65; and when A is normalized to a constant value (when each point is corrected for its concentration dependence), the slope becomes exactly 0.50, as predicted by theory. Such corrections were found to bring the slopes to 0.50 for the remaining strong acids as well. Finally, conductance measurements enabled us to express p H readings as a function of p H nonaqueous, which again is almost a straight line (Figure 3). Other relationships and additional information obtainable from the apparent p H equations are easily understood from our discussions of the theory and are left up to the reader. The findings in the present study indicate that the glass-calomel system standardized us. aqueous buffers does respond to hydrogen ion activity in the ASTM solvent to a reasonably predictable degree. Earlier, the glass electrode was reported to behave as a hydrogen electrode in a similar systembenzene-ethanol (23). In conclusion, considering all the simplifying assumptions and approximations involved, the apparent p H relationships derived here from theory do meet the test of experiment to a reasonable degree of approximation. Apparent p H in the ASTM solvent is a meaningful quantity, provided it is interpreted within the framework of the limitations discussed here.
ACKNOWLEDGMENT
The author thanks Professor Stanley Bruckenstein for stimulating discussions of t,his work. He acknowledges the capable assistances of E. S. hfcBride and G. B. Treacy in carrying out many of the experiments. LITERATURE CITED
(1) “ASTM
Standards on Petroleum Products and Lubricants,” 38th ed., Vol. 1, pp. 298-306, Am. SOC.Testing Materials, Philadelphia, Pa., 1961. (2) Bates, R. G., Am. SOC.Testing Materials, Spec. Tech. Publ. 190, l (1956). (3) Bates, R. G., ANAL.CHEM.29, 15A (#5) (1957). (4) Bates, R. G., Analyst 77, 653 (1952). (5) Bates, R. G., Chem. Rev. 42, 1 (1948). (6) Bates, R. G., Chimia 14, 111 (1960). (7) Bates, R. G., J . Electroanal. Chem. 2, 93 (1961). (8) Bates, R. G., “Treatise on Analytical Chemistry,” I. M. Kolthoff and P. J. Elving, eds., Part I, Vol. 1, Chap. 10, Interscience, New York, 1959. (9) Bates, R. G., Paabo, hl., Robinson, R. A., J . Phys. Chem. 67, 1833 (1963). (10) Bates, R. G., Schwarzenbach, G., Helv. Chzm. Acta 38, 699 (1955). (11) Bruckenstein, S., Kolthoff, I. M., J . Am. Chem. Soc. 78, 2974 (1956). (12) Bruckenstein, S., Mukherjee, L. M., J. Phys. Chem. 64,1601 (1960). (13) Zbzd., 66, 2228 (1962). (14) Bruss, D. B., Harlow, G. A., ANAL. CHEM.30, 1833, 1836 (1958). (15) Bryant, P. R., Wardrup, A. H., Chem. SOC.1957, p. 895. VOL. 36,
NO. 4, APRIL 1964
0
881
(16) Coetzee, J. F., Padmanabhan, G. R., J . Phys. Chem. 66, 1708 (1962). (17) Feldman, I., ANAL.CHEM.28, 1895 (1956). (18) French, C. LI., Roe, I. G., Trans. Faradall SOC.49, 314 (1953). (19) F u o s ~ ,R. AT,, Sheldlovsky, T., J . A m . Chem. SOC.71, 1496 (1949). (20) Grunwald, E , , Ibid., 73,4939 (1951). (21) Hummelstedt, L. E. I., Hume, D. Y . , AUAL.CHEM.32, 1792 (1960). (22) Izmaylov, N. A , , Dokl. Akad. Sauk SSSR 127, 104 (1959). (23) Izmaylov, S . A., Aleksandrova, A. A I . , Uch. Zap. Khar’kovsk. Gos. Univ.
Tr. Khim. Fak. i Nauchn. Issled. Inst. Khirn. 17, 121 (1961). (24) Kolthoff, I. M., Bruckenstein, S., Chantooni. &K.. I. Jr.. J . Am. Chem. SOC.83, 3927 (1961). (25) McKinney, D. S., Fugassi, P., Warner, J. C., Am. Soc. Testzng Materials, Spec. Tech. Publ. 73, 19 (1946). (26) Muney, W. S.,Coetzee, J. F., J . Phys. Chem. 66, 89 (1962). (25) Nelson, I. V., Iwamoto, R. T., ANAL. CHEM.33, 1795 (1961); 35, 867 (1963). (28) Pocker, Y., J . Chem. SOC.1958, p.
240. (29) Popovych, O., J . Phys. Chem. 66, 915 (1962).
(30) Popovych, O., preprints, General Papers, Division of Petroleum Chemistry, ACS, Vol. 8, No. 8, August 1963, p. 63. (31) Romberg, E., Cruse, K., Z. Elektrochenr. 63,404 (1959). (32) Van Heiide. H. R.. Anal. Chim. Acta 16, 392”(1955). (33) Van Looy, H., Hammett, L. P., J . Am. C h e m . Sor. 81, 3972 (1959). (34) \‘an ?VIeurs, N., Dahmen, E., Anal. Chim. Acta 19, 64 (1958). RECEIVED for review October 23, 1963. Accepted January 14, 1964. Division of Petroleum Chemistry, 145th Meeting, ACS, Sew York, September 1963.
Estimation of Pyridine Nucleotides in a Crude Tissue Extract by Color Reaction of Nucleotide Ribose with Orcinol and Ferric Chloride ANIMA DEVl and UMA SRIVASTAVA Department of Biochemistry, Northwestern University, Medical School, Chicago, Ill., and Department of Biochemistry, Faculty of Medicine, Lava1 University, Quebec, Canada
b On the basis of the color reaction of ribose with orcinol and ferric chloride, a simple and sensitive method of estimating total pyridine nucleotide in a crude system has been developed. It is easily workable and reproducible. It provides a sensitivity of micrograms of diphosphopyridine nucleotide (DPN) in tissue. The concentration of DPN in various tissues of rats has been estimated b y this method and then compared with the results obtained by the cyanide methoc’
R
of increased DPNase activity with age ( 9 ) in guinea pig lung tissue suggests changes, due to aging, of D P N concentration in the same tissue. The procedure for determination of total pyridine nucleotide as the cyano derivative (IO) seems less sensitive to minute changes of pyridine nucleotide concentration, especially when the change is transient. X sensitive method suitable for the determole in 1 ml. of mination of 0.3 X pyridine nucleotide was developed by Glock and McLean (6). Later Bassham ( 1 ) introduced a modified method based on the fluorometric procedure of Lowry, Roberts, and Kapphahan ( 7 ) . These methods, although sensitive for the determination of a very small amount of pyridine nucleotide, are not applicable to a routine study of the change of diphosphopyridine nucleotide (DPN) and triphosphopyridine nucleotide (TPN) in the presence of a high concentration of adenosine triphosphate (ATP) in a crude system. The enzyECENT STUDY
882
ANALYTICAL CHEMISTRY
matic procedure of the former becomes unsuitable due to the degrading enzymes that are present to a considerable amount in tissues which destroy endogenous D P K and T P K during the assaying procedure. Also there is difficulty in purifying the enzyme, D P N linked-cytochrome -C reductase ( 6 ) , as required for the assay, and the fluorometric technique needs purified D P N or T P N for their estimation ( 7 ) . Hence, for the detection of a minute change of pyridine nucleotide concentration in a crude system, a simple but sensitive chemical method was needed. -4method, based on the rate of formation of chromogens with orcinol and ferric chloride of D P N or ATP ribose, has been developed. This report deals with the workability, reproducibility, and sensitivity of the method. THEORY
The method is based on Mejbaum’s orcinol color reaction of ribose of ATP and D P N (8). The ratios of absorbance of the green chromogens, obtained when definite concentrations of D P S and ATP are boiled with orcinol and ferric chloride in the presence of concentrated hydrochloric acid for 3 minutes and 20 minutes, have been calculated. Since D P N is composed of adenine nucleotide as one part and pyridine nucleoside as the other, i t needs a longer hydrolysis period to release the ribose which then reacts with the orcinol reagent. The solution of D P N therefore produces less chromogen than does an equal amount of ATP in solution during the 3-minute period. Both ATP and D P N solutions
develop maximum color after 20 minutes of hydrolysis ( 2 ) . Using a mathematical relationship between the ratios of the absorbance by the chromogens the concentrations of D P N and ATP in a crude sample can easily be derived. If the total amount of nucleotide in a sample is composed of X parts of ATP and Y parts of D P N then the algebraic relation is represented by the equation, X + Y = 1
(1)
where 1 represents total nucleotide. If P and Q are the ratios of the absorbance by the chromogens of known concentrations of pure ATP and pure D P N after hydrolysis for 3 minutes and 20 minutes, and R is the ratio of absorbance by the chromogens after hydrolysis of the unknown sample (mixture of ATP and DPY) for the same period of time, then the ratios P and Q must not change with variation of concentration of pure ATP and pure D P N under the same experimental conditions. The ratios are devoid of units; they are simple constant numbers. If these numbers are multiplied by the respective concentrations of nucleotides in the crude system, then the total nucleotide must be equivalent to the sum of the two real concentrations. Similarly R, being the ratio of the absorbance of total chromogen at 3- and 20-minute hydrolysis, has no unit, and changes from sample to sample depending on the amount of these two nucleotides in the tissue extract. Therefore, total nucleotide must be equal to R (X Y). From the rule of identity, one can derive a n equation,
+
