Biochimica et Biophysica Acta, 393 (1975) 389-395 © Elsevier Scientific Publishing Company, Amsterdam - - Printed in The Netherlands

BBA 37066

RECOMBINATION OF CARBON MONOXIDE WITH HEMOGLOBIN AFTER FLASH PHOTOLYSIS OF THE CARBOXYDERIVATIVE IN MIXED SOL- VENTS AT SUBZERO TEMPERATURE

MICHEL BERNARD, CLAUDE BALNY, RAMAPRASAD BANERJEE and PIERRE DOUZOU

Ecole Pratique des Hautes Etudes, Institut de Biologie Physico-chimique, 13 Rue Pierre & Marie Curie, 75005 Paris and U-128 1NSERM, BP 5051, 34033 Montpellier (France)

(Received December 16th, 1974)

S U M M A R Y

The kinetics of recombination of carbon monoxide to hemoglobin produced by total flash photolysis of its carboxyderivative are studied at low temperatures (down to --55 °C) in mixed hydroalcoholic solvents. The rates are found to be differ- ent in two solvents used, namely ethylene glycol/buffer and methanol/buffer; for the former, the rates at subzero temperatures are simply explained by cooling and are consistent with the activation energy as measured in aqueous solution, while those in methanol/buffer show evidence of a specific solvent effect.

Values are reported for the rate constants and activation energies in the two solvents.

INTRODUCTION

It is well known that the recombination of carbon monoxide with hemoglobin (Hb) generated in situ by flash photolysis of its liganded form carboxyhemoglobin (HbCO) follows a time course different from that measured in flow experiments [1 ]. Very limited photolysis (5-10 ~) gives a much higher rate, while for total photolysis, the time course of CO-recombination is often biphasic [2]. For the latter, Gibson has postulated the scheme:

h~ HbCO - - ~ Hb* q- Hb ÷ CO

1" Hb* + CO ---~ HbCO

1' H b + C O - - ~ H b C O

k~ Hb* - - Hb,

where hemoglobin* (Hb*) represents a "quickly reacting form" of hemoglobin which combines with CO at a rate (l*) much higher than that (l') of ordinary hemoglobin;

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hemoglobin* itself is unstable and reverts to the ordinary form with a rate constant kl. In view of subsequent work [3-5], the above scheme should be considered only in a formal sense since the identity of the quickly reacting form as well as the process of its reconversion to the slow species may depend on experimental conditions, particu- larly dilution. While in concentrated solution, hemoglobin* may possibly, as initially postulated, be a deoxy tetramer which conserves its oxy quaternary structure though having lost its ligand rapidly through photolysis; in dilute solution the fast phase arises principally from the presence of fast-reacting dimers. However, the Gibson scheme still remains a convenient formal representation of the kinetic steps involved.

A previous work [6] from this laboratory was concerned with the above reaction conducted at --55 °C in methanol/water (60/40, v/v) mixture. It was shown that the main kinetic expressions of cooperative ligand binding were basically pre- served at --55 °C and that the biphasic ligand recombination course after total photolysis could be represented by Gibson's scheme after suitable adjustments of activation parameters. Specific solvent effects were considered with a view to assess the influence of factors other than temperature (i.e. chemical composition, organiza- tion and viscosity of the solvent) which may influence the rates. For these reasons, we put more emphasis on experiments which compared the carbon monoxide recom- bination rate after partial photolysis with that after total photolysis, in the same solvent and at the same temperature.

Since then, such mixed solvents have been widely utilized for the study of proteins and nucleic acids at subzero temperatures [7-10]. In order to rationalize their use, detailed studies have been performed in this laboratory, with a view to define the optimum conditions, with respect to protonic activity, pH* [11], dielectric constant [12, 13], viscosity, etc. The physical properties of several alcohol/water mixed solvents have been found to depend, as expected, on the nature of the alcohol; in addition, one has to take into account any specific chemical action exercised on the protein by the solvent. Being in possession of some such data, we have considered it worthwhile to compare carbon monoxide binding to hemoglobin obtained by flash photolysis of carboxyhemoglobin, at several temperatures below 0 °C and in two different mixed solvents, namely methanol/water (60/40) and ethylene glycol/water (50/50).

EXPERIMENTAL

1. Oxyhemoglobin was prepared according to the method of Gibson from fresh sheep blood treated with heparin (20 U/ml) [2]. The diluted blood was stored under nitrogen at + 4 °C and used the day after preparation. The protein concen- tration was determined by spectrophotometry [2]; CO was purchased from Air Liquide; solutions of known concentration were prepared by dilution of a saturated solution of the gas in buffer (20 °C), taken to be 1.0 mM in CO.

2. Both the flash photolysis and the optical detection systems were designed and built in this laboratory, particularly with a view to work on liquid supercooled solutions.

The main characteristics of this apparatus are: duration of the flash at half height of 50 #s for a maximum energy 1 k J; measurements from 250 to 700 nm (xenon source and monochromator, Jobin and Yvon); detection by an EMI 9558

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photomultiplier. An electronic device has been developed which switches off the high voltage applied to the photomultiplier during the flash, to protect this detector. The measuring cell (internal diameter 5 mm, optical length 70 mm) is thermostated from +50 to --65 °C with a liquid thermostat already described [14]. The temperature measured in the sample with a chromel-alumel thermocouple (0 0.3 mm) indicated a temperature gradient of less than 0.05 °C.

3. For preparation of protein solutions, two mixed solvent systems have been used in these experiments: ethyleneglycol/Tris buffer in a volume ratio 50/50, freezing point --45 °C and methanol/Tris buffer in a volume ratio 60/40, freezing point --60 °C. The final buffer concentration was 0.1 M. From the protonic activity (pH*) measurements carried out on Tris buffer in mixed solvent systems [11], we find the following: an ethyleneglycol/Tris buffer mixture (50/50) using Tris pH 7.8 had, at 20 °C, a protonic activity pH*20 of 7.8 and at --40 °C, a pH*-40 of 10.3. Similarly, for the 60/40 methanol buffer mixture using the same aqueous buffer, pH*20 was 8.0 and pH*-50 was 10.8. The choice of these mixed solutions was dictated by their relatively high basic character which gave a high concentration of the quick reacting form during the recombination of the photodissociated carboxyhemoglobin [15]. Both buffer and solvents were degassed with pure nitrogen and reduced by traces of dithionite.

According to the earlier observations [6], when using methanol/buffer mixtures the solution must be prepared at low temperatures to avoid denaturation of the hemoglobin. We have adopted the experimental procedure already described.

A few/~1 of a concentrated solution of the hemoglobin (sufficient to provide the desired final concentration from 1 to 0.4/~M on heme basis) was mixed at 0 ° C with 4 ml of Tris buffer pH 7.8 containing 0.5 mM of CO. 3 ml of methanol at 0 °C were added and the mixture immediately cooled to --20 °C. A second addition of 3 ml of methanol is followed by a cooling to --40 °C. At this temperature, the mixture was finally diluted with a precooled (at --40 °C) mixture of methanol/buffer (6 ml: 4 ml). The solution is then quickly transferred in the measuring cell of the flash apparatus precooled also at --40 °C. In these conditions, the medium is nearly iso- dielectric with water throughout the dilution with MeOH. At --50 °C, the media has a dielectric constant D = 82 [11]. In this example, the CO concentration is 100/~M, but can be differently adjusted.

Similar precautions were taken while using ethyleneglycol/buffer mixtures, though many experiments have shown the deoxy- and carboxyhemoglobin to be quite stable in this solvent even at ordinary temperature.

RESULTS AND DISCUSSION

The maximum energy of the flash (1 kJ with a duration of 50/~s at half height) resulted in complete photolysis as seen from the initial absorbance change followed at 421 nm or at 435 nm. With the help of the switch-off device, it was possible to record kinetics 0.5 ms after the maximum of the flash, a value very close to the charac- teristics given for Gibson's apparatus [2]. This dead time does not in any way restrict the observation of recombination kinetics, particularly since they are slowed down due to low temperature.

Fig. 1 (a, b) represents typical recombination kinetic curves after total photo-

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AA+ AA ,~ I + , ~ 42 In m . [ ~ _ . ~ 42O,3 0.3

II~ (b)

0.1 "ft ~ 0.1

o.orf/+oo\ o-o I 0 0.04 0.(38 0.12 0 0.1 0+2 0.3 0.4 0.5 0.6 0.7 0.8

T i m e in s T i m e in s

Fig. 1. Semi logarithmic plots of the variation of the absorbance at 421 nm after flash photolysis as a function of time at different temperatures. Hemoglobin -- 0.43/~M, CO -- 0.l raM. (a) in ethyl- eneglycol/Tris buffer 50/50; (b) in methanol/Tris buffer 60/40.

lysis respectively in ethyleneglycol/buffer and methanol/buffer systems. In the ethyl- eneglycol/buffer solvent, the recombination kinetics was found to be biphasic between + 4 °C and --7 °C (not shown in the figure) in agreement with the observation of Gibson in aqueous buffer at temperature above 0 °C. The slow phase represented only about 10~ of the total reaction. Between -7 and --43 °C, the slow phase could not be recorded, for it was severely slowed down under our experimental conditions, due to lowered temperature. The fast phase could be represented by second-order kinetics with a single rate constant which must be closely related to l*. The time constant for the fast phase can be represented by ~ = 1/(kl ÷ 1 *[CO]). Experiments with varying CO concentration showed that kl at low temperature must be negligible compared to 1" [CO]. We have therefore assimilated the second-order rate constant of the fast phase with 1" and measured the corresponding activation energy. This 1" at 0 °C in ethyleneglycol/buffer solvent is 1.8.106 M - ~ . s - l ; the activation energy measured over a relatively wide range of temperature is 5.4 ~k 0.5 kcal .M -1, both values being very close to those of Gibson (Table 1).

Significant differences are observed when the reaction is carried out in methanol/buffer system (Fig. 2). While the activation energy calculated from experi- ments carried out from --55 to --11 °C are about the same as in ethyleneglycol/buffer system, the absolute value of the rate constants for a given temperature are quite different, lower by about one order of magnitude (Fig. 2). 1 * at --55 °C is 1.4.104 M-I" s-1 similar to the one reported by Banerjee et al. [6]. The low rate constants obtained in this solvent explain why these authors reported a higher activation energy of about 11.1 kcal. M -1 for this reaction, since the latter was calculated on the basis of their value at --55 °C and of Gibson's 1 * at 1 °C (1.8.106 M -~.s -1) in aqueous buffer. For reasons of protein stability, we did not make measurements at 0 °C in methanol/buffer solvent, but extrapolation (Fig. 2) gives a value, for this temperature, of 1.8. l0 s M - l - s - ; , one order of magnitude lower than that in aqueous solution.

In brief, the two solvents methanol/buffer and ethyleneglycol/buffer exhibit

i -M- l ,s -1

393

106

105

10 4 i i i i

3,5 4 4,5 103 . -1 T

. . . . . . , 'o ' ' 7: - *20 0 - 2 0 - - 5 5 *

Fig. 2. Arrhenius plots of the rate constant l *. (a) in ethyleneglycol/Tris buffer 50/50, ( O - ' - O G) values reported by Gibson [2] in pure aqueous buffer; (b) in methanol/Tris buffer 60/40, ( . . . . . . . . . ) extrapolation to 0 °C. ( . . . . . . ) the abrupt fall in the rate constant between 0 °C in aqueous buffer and --11 °C in methanol/Tris buffer.

ve ry d i f ferent p rope r t i e s wi th respec t to the reactioia s tud ied in this work . W e t end to

cons ide r the e thy l eneg lyco l /bu f f e r m i x t u r e as the " n o r m a l " solvent , s ince the k ine t i c

ra te c o n s t a n t a t 0 °C in this so lven t is the same as tha t at 0 °C in water . T h e A r r h e n i u s

p lo t d r a w n f r o m 20 °C in wa te r d o w n to - - 4 0 °C in the so lven t is a s t ra igh t l ine wi th-

TABLE I

RATE CONSTANT l* AND ACTIVATION ENERGY OF THE QUICK-REACTING FORM IN DIFFERENT CONDITIONS

Aqueous Ethyleneglycol/buffer Methanol/buffer buffer

1.8.106a 1.8.106 1.8-105c (at 1 °C)

--40 °C = 3.2 -104 --40 °C = 3.105 --55 °C 1.4 .10 ~

[--55 °C = 9.75.103b

5.6 a 5.4 ± 0.5 5.4 d_ 0.5 11.1 b

1" at 0 °C (M- l . s - l ± 0.2)

1 * at low temperature (M-1"s -1 i 0.2)

E (kcal' M- 1)

a Values reported by Gibson [2]. Values reported by Banerjee et al. [6].

c Value obtained by extrapolation at 0 °C in Fig. 3.

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out break (Fig. 2) signifying no abrupt changes of reaction parameters, including pH. On the other hand, a very significant break in this plot from 20 and 1 °C at water and 11 °C and --55 °C in methanol/HzO (Fig. 2) reveals the "abnormal" nature of this solvent which cannot be simply a pH effect. In fact, as reported earlier, the pH* of methanol/Tris buffer mixture is substantially different from the pH of the aqueous Tris buffer at the same temperature. An increase of pH would riormally enhance the kinetic constant [15] instead of diminishing it as actually happens.

We want to stress that the activation energy of 1" is the same in the two solvents, the only abnormal behaviour of the methanol/H20 mixture consists in slowing down the reaction, at all temperatures. As stated in the experimental section ; the reactions were carried out at low carboxyhemoglobin concentrations (from 1 to 0.4 #M heme), when the carboxyhemoglobin is largely present in dimers form. The solvents may have enhanced the proportion of dimers, since the fast phase represented up to 9 0 ~ of the total reaction. Even if this were so, the difference of solvent behav- iour could not arise from any difference in their dissociating properties, since the proportion of the fast phase as obtained in the two solvents was about the same. Again, the relaxation of the quick-reacting material to the slow form (formally represented by kl) was very severely slowed down in both the solvents. At this stage we cannot satisfactorily interpret the abnormal behaviour of the methanol/H20 sol- vent and can only say that the action of methanol on the reactions studied originates from a specific solvent effect, probably related to its microscopic structure which nevertheless does not show any spectral alterations of the protein. Such effects may indeed be specific for the system studied, although differences between solvent effects had been found for a peroxidasic reaction [16, 17] whereas none were reported for hydrolysis of synthetic substrates by chymotrypsin [18]. We suggest that the continuity of Arrhenius plots from the aqueous to the mixed solvent should be a good criterion for the choice of optimum reaction medium at low temperature.

ACKNOWLEDGEMENTS

This work was supported by grants from CNRS (E.R.A. No. 262 and E.R. No. 157), the DGRST (No. 71. 73139101), INSERM and the Fondation pour la Recherche M6dicale Franqaise.

REFERENCES

1 Gibson, Q. H. (1956) J. Physiol. 134, 123-134 2 Gibson, Q. H. (1959) Biochem. J. 71, 293-303 3 Antonini, E., Brunori, M. and Anderson, S, (1968) J. Biol. Chem. 243, 1816-1822 4 Gibson, Q. H. and Antonini, E. (1967) J. Biol. Chem. 242, 4678-4681 5 Antonini, E., Anderson, N. M. and Brunori, M. (1972) J. Biol. Chem. 247, 319-321 6 Banerjee, R., Douzou, P. and Lombard, A. (1968) Nature 217, 23-25 7 Douzou, P. (1973) Mol. Cell Biochem. 1, 15-27 8 Douzou, P., Sireix, R. and Travers, F. (1970) Proc. Natl. Acad. Sci. U.S, 66, 787-792 9 Debey, P., Balny, C. and Douzou, P. (1973) Proc. Natl. Acad. Sci. U.S. 70, 2633-2636

10 Travers, F., Michelson, A. M. and Douzou, P. (1970) Biochim. Biophys, Acta 217, 1-6 11 Hui Bon Hoa, G. and Douzou, P. (1973) J. Biol. Chem. 248, 4649-4654 12 Travers, F. and Douzou, P. (1970) J. Phys. Chem. 74, 2243-2244 13 Travers, F. and Douzou, P. (1974)Biochimie 56, 509-514

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14 Hui Bon Hoa, G. and Douzou, P. (1973) Anal. Biochem. 51, 127-136 15 Gibson, Q. H. (1959) Progress in Biophysics, Vol. 9, pp. 1-53, Pergamon Press, London 16 Maurel, P. and Travers, F. (1973) C.R. Acad. Sci. Paris 276D, 3057-3060 17 Maurel, P. and Travers, F. (1973) C.R. Acad. Sci. Paris 276D, 3383-3386 18 Bielski, B. H. J. and Freed, S. (1964) Biochim. Biophys. Acta 89, 314-323