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The kinetics of the complexes of peroxidase formed in the presence of chlorite or hypochlorite
The Kinetics of the Complexes of Peroxidase Formed in the Presence of Chlorite or Hypochlorite~ Britton Chance From the Johnson Research Foundation, University of Pennsylvania, Philadelphia, Pennsylvania
Received June 5, 1952
George (1) has found that a variety of oxidizing agents other than hydrogen peroxide (2) or alkyl hydrogen peroxide (3) cause the formation of the secondary complex of horse-radish peroxidase. In this paper we have studied whether the formation of complex I precedes complex II when chlorite or hypochlorite is used as an oxidizing agent andwhether the complex II that is formed has the same activity as that formed from peroxides. Some kinetic and equilibrium data on the reaction of chlorite and hypochlorite with peroxidase are included. PREPARATIONS The peroxidase preparation of paper I of this series is used (4). Pure sodium chlorite was available and the hypochlorite solutions were assayed by iodine titrations.*
METHODS The experiments were carried out in the open cuvette ing spectrophotometer described previously (4, 5).
EXPERIMENTAL
of a double-beam
record-
RESULTS
The Existence of Complex I Figures 1A and B show the kinetics of formation of the peroxidase complexes in the presence of hypochlorite and chlorite; the formation of 1 These researches were supported in part by a grant from the Division of Research Grants and Fellowships, United States Public Health Service, and from the Office of Naval Research. This is paper III of a series on peroxidase. 2 The solutions were the same as those standardized by Dr. P. George and Mr. T. M. Devlin as will be described elsewhere. 425
426
BRITTON
CHANCE
complex I is registered as an upward deflection of the traces at 412 rnM (corresponding to a decrease of optical density) and complex II is registered as an upward deflection of the traces at 427 rnp (corresponding to
B 1. The existence of complex I in the reactions of hypochlorite (A) or chlorite (B) with horse-radish peroxidase. The formation of the complexes gives an upward deflection of the traces, 412 rnp for complex I, 427 rnp for complex II; 0.53 PM horse-radish peroxidase, S”, pH = 5.4, 0.01 M phosphate (Expt. 8593, c). FIG.
ComplexlI m..f+427mu 20JJM NoC102
I 25sec. I
B
2A. The effect of added chlorite concentration upon the rate of formation of complex II. The half-time of the reaction (ti ..) is measured from records such as Fig. 1B. The slope of the line is approximately lo6 M-l X sec.+. Horse-radish peroxidase, 0.53 PM, 8”, pH = 5.4, 0.01 M phosphate (Expt. 859b). FIG. 2B. The reaction of chlorite with complex I. The sense of the deflections is as in Fig. 1. Horse-radish peroxidase, 0.9pM, 25”, pH = 6.0, 0.01 M phosphate buffer (Expt. 962). FIG.
an increase of optical density). Upon addition of hypochlorite, the sequence of reactions is very nearly identical to those that occur when peroxide is used as an oxidant; the rapid formation of complex I is followed by the slow formation of II, and the latter reaction is completed upon the addition of a donor such as nitrite [compare with Fig. 1 of Ref. (4)].
PEROXIDASE
IN
CHLORITE
OR
HYPOCHLORITE
427
If, however, chlorite is added, a rather different seauenceof reactions occurs. In this case a scarcely measurable deflection is recorded at 412 rncc;nevertheless, the slow formation of complex II proceeds at about the same rate as in Fig. IA. In order to eliminate the possibility that chlorite forms a primary complex of a different spectrum that does not show at 412 rnp, we added methyl hydrogen peroxide next and verified the presence of free peroxidase by the formation of complex I. Nitrite addition completes the formation of complex II. The Kinetics of Formation of Complex II from Chlorite
Whereas the added peroxide concentration does not affect the speed of formation of complex II (3, B), we observe an increase in the speed of formation of complex II with increasing chlorite concentration, as illustrated by Fig. 2A. There is a roughly linear relationship between the rate of formation of complex II and the added chlorite concentration. Inasmuch as no complex I was observed under these conditions we may assume that the I - II transition reaction proceeds more rapidly than does the formation of I. On this basis the second-order velocity constant for the formation of complex I is the slope of the line in Fig. 2A, ~10~ M-l X sec.-’ at pH = 5.4 and 4”. A slower reaction is observed at pH = 7 (-1500 M-l X sec.-‘) and indicates that the free acid is the reactive species. On this basis, the latter value would correspond to ~4 X lo5 M-l X sec.? at pH 5.4. But before such a figure can be considered seriously, we must explain how the velocity constant for the formation of complex II could reach a value of 0.2 sec.-’ in the absence of added donor; with hydrogen peroxide a value of only ~0.02 sec.-l is obtained with this peroxidase preparation. The explanation is provided by Fig. 2B. In this casewe have added methyl hydrogen peroxide to form complex I and to initiate the slow formation of complex II. Then chlorite is added with rather striking results: the concentration of complex I immediately falls to zero and that of complex II rises to 80% of the maximum value: chlorite acts both as a substrate and a donor toward peroxidase. Proof that an oxidation of chlorite to chlorine dioxide is occurring is afforded later by Fig. 4B. Based on this discovery of the dual possibilities of chlorite as a substrate and donor for peroxidase, we can readily explain the lack of complex I in the traces of Fig. lB-considerable donor was indeed present. And our preliminary calculation of kl from the kinetics of formation of complex II would appear to be justifiable.
428
BRITTON
CHANCE
Studies of the effect of pH on complex II of horse-radish peroxidase show that only complex I appears at about pH 3.5 (7). We have repeated this experiment using chlorite and find a similar result. We may safely
FIG. 3. The “dissociation constants” of complex II formed from hypochlorite (A) and chlorite (B). The abscissas represent added reagent. The ordinates represent the percentage of the maximum deflection obtained at 427 rnp in experiments such as those of Fig. 1A for hypochlorite (i.e., nitrite is added to convert complex I to II) and the first portion of Fig. 1B for chlorite (no electron donor need be added to convert complex I to II). Peroxidase, 0.53 &Z, 8”, pH = 5.4, 0.01 M phosphate (Expts. 8596,860).
conclude that a small amount of complex I is formed from chlorite at higher values of pH. The “Dissociation
Constants” for Chlorite and H ypochlorite
Figure 3 gives an indication of the amounts of added chlorite and hypochlorite that are required to form complex II at 8” and pH 5.4. The free chlorite and hypochlorite concentrations giving half saturation of complex II are approximately 0.5 X 10ds and 1 X 10M6M, respectively. In the case of chlorite, no donor need be added, while in the case of hypochlorite, nitrite was added to secure complete conversion of complex I to II. The amount of these reagents required to form complex II is considerably greater at pH 7.0. In the case of chlorite this effect is caused by a decrease in the velocity of the combination reaction. These “dissociation constants” are then over 100 times larger than the value estimated for hydrogen peroxide (< 5 X lo-* M). This dis-
PEROXIDASE
IN
CHLORITE
OR
HYPOCHLORITE
429
crepancy leaves open the possibility that a portion of the hypochlorite or chlorite is converted into peroxide upon reaction with peroxidase. The Reactivity of Complex II Toward Donor Molecules By using a large excess of nitrite (so that k4ao- 0.2 sec.-l) and a low concentration of chlorite (/&a0 for chlorite is <0.03 sec.-‘), the donor
Substrata
N&tO*
A
cn,oc+t
B
4A. A comparison of the “cycles” of appearance and disappearance of complex II formed from 3.3 PM chlorite or from 3.3 PM methyl hydrogen peroxide. Peroxidase 0.53 PM, 25”, pH = 5.4,O.Ol M phosphate (Expt. 859a). FIG. 4B. The formation of chlorine dioxide in the reaction of chlorite and peroxidaae. The experimental conditions are identical to those of Fig. 2B except that optical density changes at 350 w following the addition of chlorite are recorded on a slow time scale. The steady-state concentration of complex II is 0.75 pM (Expt. 962). FIG.
action of the latter is rendered negligible and the activity of complex II formed from chlorite may be compared with that formed from methyl hydrogen peroxide, as shown by Fig. 4A. The smaller velocity constant for the combination of chlorite and peroxidase gives a smaller steadystate concentration of Complex II t,han is obtained with methyl hydrogen peroxide. xevertheless, the values of k4 calculated by the usual formula (8) are 270 and 320 M-* X sec.-‘, respectively, on the basis of total nitrit,e concentration, pH = 5.4, 25”; 290 and 280 M-l X sec.-’ being obtained with 1.33 PM chlorite and methyl hydrogen peroxide, respectively, in another experiment. It is clear that there is no appreciable difference in the reactivity of complex II formed from the two types of oxidizing agents. Under the conditions of Fig. 4A where a large excess of donor is present, there is no question that the transition from complex I to II exceeds the rate of formation of I. Thus the steady-state concentration of complex
430
BRITTON
CHANCE
II depends upon the values of kl and lc, [see Ref. (9), Eqs. 13 and 14, remembering that ka = k4a]. Since kq for nitrite is the same for both methyl hydrogen peroxide and chlorite, and k, for methyl hydrogen peroxide is 1.5 X lo6 M-l X sec.-’ (9),
k 1NaC102
= 1.5 x lo6
CH300H M-1
x
sec*-l
From Fig. 4A and other experiments an average value of 4 X lo6 M-l X sec.-’ at 25” and pH 5.4 is obtained. If chlorite is used as a donor molecule, instead of as a substrate molecule, the formation of chlorine dioxide can be measured by its absorption at 350 rnpCca as shown in Fig. 4B,which corresponds to the experimental conditions of Fig. 2B. The addition of chlorite, following the addition of methyl hydrogen peroxide, causes the downward deflection of the recorder trace that corresponds to an increase of optical density caused by the formation of chlorine dioxide. About half the initial chlorite is converted to chlorine dioxide. The initial slope of this trace corresponds to a value of k4 of lo4 M-l X sec.-l at pH 6.0 and 25”. Chlorine dioxide is formed at about the same value of k4 when methyl hydrogen peroxide is omitted and chlorite acts both as substrate and electron donor. A similar formation of chlorine dioxide has been observed in the presence of catalase (Expt. 857~) and will be reported in detail later. In numerous attempts to measure k4 for the complex II formed in the presence of hypochlorite, we have never obtained satisfactory results: complex II disappears in a much shorter time than one would expect from the value of k, obtained when peroxide is used as a substrate. We attribute this to the uncatalyzed reaction of hypochlorite with the various donor molecules that were used. Such reactions proceed at rates that are over 50-fold more rapid than the peroxidase-catalyzed reactions. Thus we advance no proof of the reaction of electron donors with complex II formed from hypochlorite. However, the data afforded by Fig. 1A prove that the complexes formed from hypochlorite behave qualitatively as do those formed from peroxide. Our studies do not attempt to prove that complex I formed from S This method of assaying who, with Mr. T. M. Devlin, tion we used cab0 = 0.99 cm.-’
chlorine dioxide was suggested by Dr. P. George, found cS6* = 0.995 cm.+ X mW1. In this calculaX mM-I.
PEROXIDASE
IN CHLORITE
OR HYPOCHLORITE
431
peroxides is identical to that formed from the various oxidizing agents; we only show that the reaction kinetics of this complex are similar to those of complex I. Nor do our studies show whether a peroxide produced by the interaction of peroxidase and t,hese oxidizing agents causes the effects observed [but see also George (lo)]. SUMMARY
The addition of chlorite to horse-radish peroxidase appears to initiate the same sequence of reactions as peroxide: the formation of a complex I followed by its transition to complex II. However, the amount. of free chlorite or hypochlorite required to give half-maximal formation of the peroxidase complexes is over 100 times t,he amount of free hydrogen peroxide that is required. Complex I is not observed at an appreciable concentration because chlorite also acts as an electron donor (from which chlorine dioxide is formed) and accelerates the transition from complex I to II. The complex II formed from chlorite has the same react.ivity toward nitrite as does that formed from peroxide, and the complexes may be considered to be identical. With hypochlorite, the formation of a complex I is readily observed, as is its transition to complex II. An accurate evaluation of the reactivity of this complex II toward donors has not yet been obtained. REFERENCES 1. GEORGE, P. Advances in Catalysis k, 367 (1952). 2. KEILIN, D., AND HARTREE, E. F., hoc. Roy Sot. (London) 3. CHANCE, B., Arch. Biochem. 21, 416 (1949). 4. CHANCE, B., Arch. Biochem. Biophys. 41, 404 (1952). 5. CHANCE, B., Rev. Sci. Instruments 22, 619 (1951). 6. CHANCE, B., Arch. Biochem. Biophys. 41, 416 (1952). 7. CHANCE, B., Arch. Biochem. Biophys. 40,153 (1952). 8. CHANCE, B., J. Biol. Chem. 161, 553 (1943). 9. CHANCE, B., Arch. Biochem. 22, 224 (1949). 10. GEORQE, P., Nature 169, 612 (1952).
Bl22,
119 (1937).
Received June 5, 1952
George (1) has found that a variety of oxidizing agents other than hydrogen peroxide (2) or alkyl hydrogen peroxide (3) cause the formation of the secondary complex of horse-radish peroxidase. In this paper we have studied whether the formation of complex I precedes complex II when chlorite or hypochlorite is used as an oxidizing agent andwhether the complex II that is formed has the same activity as that formed from peroxides. Some kinetic and equilibrium data on the reaction of chlorite and hypochlorite with peroxidase are included. PREPARATIONS The peroxidase preparation of paper I of this series is used (4). Pure sodium chlorite was available and the hypochlorite solutions were assayed by iodine titrations.*
METHODS The experiments were carried out in the open cuvette ing spectrophotometer described previously (4, 5).
EXPERIMENTAL
of a double-beam
record-
RESULTS
The Existence of Complex I Figures 1A and B show the kinetics of formation of the peroxidase complexes in the presence of hypochlorite and chlorite; the formation of 1 These researches were supported in part by a grant from the Division of Research Grants and Fellowships, United States Public Health Service, and from the Office of Naval Research. This is paper III of a series on peroxidase. 2 The solutions were the same as those standardized by Dr. P. George and Mr. T. M. Devlin as will be described elsewhere. 425
426
BRITTON
CHANCE
complex I is registered as an upward deflection of the traces at 412 rnM (corresponding to a decrease of optical density) and complex II is registered as an upward deflection of the traces at 427 rnp (corresponding to
B 1. The existence of complex I in the reactions of hypochlorite (A) or chlorite (B) with horse-radish peroxidase. The formation of the complexes gives an upward deflection of the traces, 412 rnp for complex I, 427 rnp for complex II; 0.53 PM horse-radish peroxidase, S”, pH = 5.4, 0.01 M phosphate (Expt. 8593, c). FIG.
ComplexlI m..f+427mu 20JJM NoC102
I 25sec. I
B
2A. The effect of added chlorite concentration upon the rate of formation of complex II. The half-time of the reaction (ti ..) is measured from records such as Fig. 1B. The slope of the line is approximately lo6 M-l X sec.+. Horse-radish peroxidase, 0.53 PM, 8”, pH = 5.4, 0.01 M phosphate (Expt. 859b). FIG. 2B. The reaction of chlorite with complex I. The sense of the deflections is as in Fig. 1. Horse-radish peroxidase, 0.9pM, 25”, pH = 6.0, 0.01 M phosphate buffer (Expt. 962). FIG.
an increase of optical density). Upon addition of hypochlorite, the sequence of reactions is very nearly identical to those that occur when peroxide is used as an oxidant; the rapid formation of complex I is followed by the slow formation of II, and the latter reaction is completed upon the addition of a donor such as nitrite [compare with Fig. 1 of Ref. (4)].
PEROXIDASE
IN
CHLORITE
OR
HYPOCHLORITE
427
If, however, chlorite is added, a rather different seauenceof reactions occurs. In this case a scarcely measurable deflection is recorded at 412 rncc;nevertheless, the slow formation of complex II proceeds at about the same rate as in Fig. IA. In order to eliminate the possibility that chlorite forms a primary complex of a different spectrum that does not show at 412 rnp, we added methyl hydrogen peroxide next and verified the presence of free peroxidase by the formation of complex I. Nitrite addition completes the formation of complex II. The Kinetics of Formation of Complex II from Chlorite
Whereas the added peroxide concentration does not affect the speed of formation of complex II (3, B), we observe an increase in the speed of formation of complex II with increasing chlorite concentration, as illustrated by Fig. 2A. There is a roughly linear relationship between the rate of formation of complex II and the added chlorite concentration. Inasmuch as no complex I was observed under these conditions we may assume that the I - II transition reaction proceeds more rapidly than does the formation of I. On this basis the second-order velocity constant for the formation of complex I is the slope of the line in Fig. 2A, ~10~ M-l X sec.-’ at pH = 5.4 and 4”. A slower reaction is observed at pH = 7 (-1500 M-l X sec.-‘) and indicates that the free acid is the reactive species. On this basis, the latter value would correspond to ~4 X lo5 M-l X sec.? at pH 5.4. But before such a figure can be considered seriously, we must explain how the velocity constant for the formation of complex II could reach a value of 0.2 sec.-’ in the absence of added donor; with hydrogen peroxide a value of only ~0.02 sec.-l is obtained with this peroxidase preparation. The explanation is provided by Fig. 2B. In this casewe have added methyl hydrogen peroxide to form complex I and to initiate the slow formation of complex II. Then chlorite is added with rather striking results: the concentration of complex I immediately falls to zero and that of complex II rises to 80% of the maximum value: chlorite acts both as a substrate and a donor toward peroxidase. Proof that an oxidation of chlorite to chlorine dioxide is occurring is afforded later by Fig. 4B. Based on this discovery of the dual possibilities of chlorite as a substrate and donor for peroxidase, we can readily explain the lack of complex I in the traces of Fig. lB-considerable donor was indeed present. And our preliminary calculation of kl from the kinetics of formation of complex II would appear to be justifiable.
428
BRITTON
CHANCE
Studies of the effect of pH on complex II of horse-radish peroxidase show that only complex I appears at about pH 3.5 (7). We have repeated this experiment using chlorite and find a similar result. We may safely
FIG. 3. The “dissociation constants” of complex II formed from hypochlorite (A) and chlorite (B). The abscissas represent added reagent. The ordinates represent the percentage of the maximum deflection obtained at 427 rnp in experiments such as those of Fig. 1A for hypochlorite (i.e., nitrite is added to convert complex I to II) and the first portion of Fig. 1B for chlorite (no electron donor need be added to convert complex I to II). Peroxidase, 0.53 &Z, 8”, pH = 5.4, 0.01 M phosphate (Expts. 8596,860).
conclude that a small amount of complex I is formed from chlorite at higher values of pH. The “Dissociation
Constants” for Chlorite and H ypochlorite
Figure 3 gives an indication of the amounts of added chlorite and hypochlorite that are required to form complex II at 8” and pH 5.4. The free chlorite and hypochlorite concentrations giving half saturation of complex II are approximately 0.5 X 10ds and 1 X 10M6M, respectively. In the case of chlorite, no donor need be added, while in the case of hypochlorite, nitrite was added to secure complete conversion of complex I to II. The amount of these reagents required to form complex II is considerably greater at pH 7.0. In the case of chlorite this effect is caused by a decrease in the velocity of the combination reaction. These “dissociation constants” are then over 100 times larger than the value estimated for hydrogen peroxide (< 5 X lo-* M). This dis-
PEROXIDASE
IN
CHLORITE
OR
HYPOCHLORITE
429
crepancy leaves open the possibility that a portion of the hypochlorite or chlorite is converted into peroxide upon reaction with peroxidase. The Reactivity of Complex II Toward Donor Molecules By using a large excess of nitrite (so that k4ao- 0.2 sec.-l) and a low concentration of chlorite (/&a0 for chlorite is <0.03 sec.-‘), the donor
Substrata
N&tO*
A
cn,oc+t
B
4A. A comparison of the “cycles” of appearance and disappearance of complex II formed from 3.3 PM chlorite or from 3.3 PM methyl hydrogen peroxide. Peroxidase 0.53 PM, 25”, pH = 5.4,O.Ol M phosphate (Expt. 859a). FIG. 4B. The formation of chlorine dioxide in the reaction of chlorite and peroxidaae. The experimental conditions are identical to those of Fig. 2B except that optical density changes at 350 w following the addition of chlorite are recorded on a slow time scale. The steady-state concentration of complex II is 0.75 pM (Expt. 962). FIG.
action of the latter is rendered negligible and the activity of complex II formed from chlorite may be compared with that formed from methyl hydrogen peroxide, as shown by Fig. 4A. The smaller velocity constant for the combination of chlorite and peroxidase gives a smaller steadystate concentration of Complex II t,han is obtained with methyl hydrogen peroxide. xevertheless, the values of k4 calculated by the usual formula (8) are 270 and 320 M-* X sec.-‘, respectively, on the basis of total nitrit,e concentration, pH = 5.4, 25”; 290 and 280 M-l X sec.-’ being obtained with 1.33 PM chlorite and methyl hydrogen peroxide, respectively, in another experiment. It is clear that there is no appreciable difference in the reactivity of complex II formed from the two types of oxidizing agents. Under the conditions of Fig. 4A where a large excess of donor is present, there is no question that the transition from complex I to II exceeds the rate of formation of I. Thus the steady-state concentration of complex
430
BRITTON
CHANCE
II depends upon the values of kl and lc, [see Ref. (9), Eqs. 13 and 14, remembering that ka = k4a]. Since kq for nitrite is the same for both methyl hydrogen peroxide and chlorite, and k, for methyl hydrogen peroxide is 1.5 X lo6 M-l X sec.-’ (9),
k 1NaC102
= 1.5 x lo6
CH300H M-1
x
sec*-l
From Fig. 4A and other experiments an average value of 4 X lo6 M-l X sec.-’ at 25” and pH 5.4 is obtained. If chlorite is used as a donor molecule, instead of as a substrate molecule, the formation of chlorine dioxide can be measured by its absorption at 350 rnpCca as shown in Fig. 4B,which corresponds to the experimental conditions of Fig. 2B. The addition of chlorite, following the addition of methyl hydrogen peroxide, causes the downward deflection of the recorder trace that corresponds to an increase of optical density caused by the formation of chlorine dioxide. About half the initial chlorite is converted to chlorine dioxide. The initial slope of this trace corresponds to a value of k4 of lo4 M-l X sec.-l at pH 6.0 and 25”. Chlorine dioxide is formed at about the same value of k4 when methyl hydrogen peroxide is omitted and chlorite acts both as substrate and electron donor. A similar formation of chlorine dioxide has been observed in the presence of catalase (Expt. 857~) and will be reported in detail later. In numerous attempts to measure k4 for the complex II formed in the presence of hypochlorite, we have never obtained satisfactory results: complex II disappears in a much shorter time than one would expect from the value of k, obtained when peroxide is used as a substrate. We attribute this to the uncatalyzed reaction of hypochlorite with the various donor molecules that were used. Such reactions proceed at rates that are over 50-fold more rapid than the peroxidase-catalyzed reactions. Thus we advance no proof of the reaction of electron donors with complex II formed from hypochlorite. However, the data afforded by Fig. 1A prove that the complexes formed from hypochlorite behave qualitatively as do those formed from peroxide. Our studies do not attempt to prove that complex I formed from S This method of assaying who, with Mr. T. M. Devlin, tion we used cab0 = 0.99 cm.-’
chlorine dioxide was suggested by Dr. P. George, found cS6* = 0.995 cm.+ X mW1. In this calculaX mM-I.
PEROXIDASE
IN CHLORITE
OR HYPOCHLORITE
431
peroxides is identical to that formed from the various oxidizing agents; we only show that the reaction kinetics of this complex are similar to those of complex I. Nor do our studies show whether a peroxide produced by the interaction of peroxidase and t,hese oxidizing agents causes the effects observed [but see also George (lo)]. SUMMARY
The addition of chlorite to horse-radish peroxidase appears to initiate the same sequence of reactions as peroxide: the formation of a complex I followed by its transition to complex II. However, the amount. of free chlorite or hypochlorite required to give half-maximal formation of the peroxidase complexes is over 100 times t,he amount of free hydrogen peroxide that is required. Complex I is not observed at an appreciable concentration because chlorite also acts as an electron donor (from which chlorine dioxide is formed) and accelerates the transition from complex I to II. The complex II formed from chlorite has the same react.ivity toward nitrite as does that formed from peroxide, and the complexes may be considered to be identical. With hypochlorite, the formation of a complex I is readily observed, as is its transition to complex II. An accurate evaluation of the reactivity of this complex II toward donors has not yet been obtained. REFERENCES 1. GEORGE, P. Advances in Catalysis k, 367 (1952). 2. KEILIN, D., AND HARTREE, E. F., hoc. Roy Sot. (London) 3. CHANCE, B., Arch. Biochem. 21, 416 (1949). 4. CHANCE, B., Arch. Biochem. Biophys. 41, 404 (1952). 5. CHANCE, B., Rev. Sci. Instruments 22, 619 (1951). 6. CHANCE, B., Arch. Biochem. Biophys. 41, 416 (1952). 7. CHANCE, B., Arch. Biochem. Biophys. 40,153 (1952). 8. CHANCE, B., J. Biol. Chem. 161, 553 (1943). 9. CHANCE, B., Arch. Biochem. 22, 224 (1949). 10. GEORQE, P., Nature 169, 612 (1952).
Bl22,
119 (1937).