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The effects of temperature and ionic strength on the formation of the hydroxide complex of ferric horseradish peroxidase
ARCHIVES
OF
BIOCHEMISTRY
The Effects of the
AND
of Temperature
Hydroxide
and
Complex
WII,LIAi/l Department
133, 313-317 (1969)
BIOPHYSICS
of Ferric
D. ELLIS2
of Chemistry, Received
ionic
AND
University July
Strength
Horseradish
H. BRIAK
of Alberta,
1, 1968; accepted
on the
Formation
Peroxidase’
DUNFORD
Edmonton,
dlberta,
Canada
June 9, 1969
The equilibrium constants for the formation of the hydroxide complex of ferric horseradish peroxidase have been determined spectrophotometrically at 25.0” for seven values of ionic strength over the range 0.02-0.20 and also at 18.0” and 35.0” with p = 0.11. The results were analyzed on the assumption that the hydroxide complex is formed by the ionization of a water molecule in the sixth coordination position of the heme ferric iron. The effective charge on the aquo species is indicated to be in the range +2 f 0.5; its dissociat,ion constant is 2.0 X lo-i1 M and the thermodynamic parameters for its ionizat,ion are indicated t,o be AG” = 15 kcal/mole; AH” = 6 kcal/ mole; AS” = -30 entropy unit.s at 25” and zero ionic strength.
It is widely accepted that HRP,4 like hemoglobin and myoglobin, contains water in the sixth coordination position of the heme ferric iron (24), although the suggestion has been made that the heme of HRP lies in a crevice with both the fifth and sixth positions occupied by protein residues (5). When a solution of HRP is made alkaline, there are accompanying changes in the spectrum of HRP (2), and its magnetic moment changes from 5.45 to 2.66 Bohr magnetons (6), indicating the formation of a new species, HRPOH. The pK for hydroxide formation has been measured to be 10.9 by spectrophotometry and 11.3 by use of a magnetic balance (6), and 10.6 by a study of the reduction i Supported financially by the National Research Council of Canada. 2 Holder of an Izaak Walton Killam Memorial Fellowship. 3 To whom reprint, requests should be sent. 4 Abbreviations used: HRP, ferric horseradish peroxidase; P, refers specifically to a form of HRP in which it is assumed that water is in the sixth coordinatiou position; HRP-OH or POH, hydroxide form of HRP; p, ionic strength; RZ, purity number, after Theorell and Maehly (1); e.u., entropy units. 313
potentials of HRP (3), but these studies were not carried out at constant ionic strength. A study of the kinetics of cyanide binding to HRP at an ionic strength of 0.11 yielded a value of 10.8 for this pK (7). We report here on a spectrophotometric study of the formation of HRP-OH at 25.0” for seven values of k over the range 0.02-0.20 and also at 18.0’ and 35.0” with ~1= 0.11. MATEKIALS
AND
METHODS
Experimental. Two grades of HRP were obtained from Boehringer-Mannheim Corp. N.Y.: “analytical reagent grade” (RZ = 2.9 lot 648206) and “purified lyophilized powder” (RZ = 0.8). All HRP was exhaustively dialyzed against conductivity water and centrifuged before use, Analytical reagent grade potassium nitrate (Mallinckrodt) and potassium hydroxide (B.D.H.) were used without further purification. Some of the “purified” HRP was treated by passage through a 1.1 X 50-cm column of C-50 CM-Sephadex cation exchange gel. The HRP was eluted using pH 6.2 phosphate buffer, with ionic strength increasing from 0.1 to 0.2 during the course of the elution. A single passage through the column resulted in the recovery of a fraction of RZ = 2.7 which represented 65% of the initial “purified” HRP as determined by absorbance measurements at 497 mN. Absorbance measure-
314
ELLIS
AND
ments were made with a Beckman DU spectrophotometer which had the cell compartment thermostated to f0.1”. A Beckman Expandomatic pH meter was used for all pH measurements. For a typical experiment, an appropriate amount of 1 M KNO, and enough stock HRP to make a final HRP concentration of about 1.7 X lo-& M were pipetted into a IO-ml volmetric flask, which was then filled to the mark with water. The 10 ml of solution was transferred to a 30-ml beaker which was covered with Parafilm and placed in a constant temperature bath. A Beckman 39183 combination pH electrode, a Teflon syringe needle, and a platinum stirring wire were inserted through slits in the Parafilm. The pH was measured and was found typically to have a value of 5.5. A 3-ml aliquot of the solution was transferred with a syringe to a cuvette and the HRP concentration was determined by measuring the absorbance of t,he solution at 497 rnp where the molar absorptivity coefficient is 1.00 X lo4 M-lcrn-l (2). The absorbance of the solution was t.hen measured at 416 rnp, the wavelength corresponding to the maximum of the Soret peak of HRP-OH (2). B second HRP solution of about pH 5.5 was prepared to use as a blank in the spectrophotometer. Its concentration was such that the difference in absorbances between the blank and the experimental solution was about 0.1 at 416 rnp, t.hus keeping t,he absorbance readings during the course of the titration in the range 0.14.6, which includes the region of minimum error due to instrument,al uncertainty. The experimental HRP solution was returned to the beaker and the pH was raised by the addition from a lo-p1 Hamilton syringe of a few microliters of approximately 1 M KOH, an amount small enough that no volume correction was necessary. The solution was stirred, the pH was remeasured, the absorbance of an aliquot was measured at 416 w in a capped cuvette, and the aliquot was returned to the beaker. Nine to eleven points were obtained this way, with the titration being carried out up to about. pH 11.35. Above this pH, the heme splits from the protein of HRP at an appreciable rate (8). After each increase in pH, the beaker was flushed gently with nitrogen so that there would be no change in pH or ionic strength due to absorption of carbon dioxide a problem encountered from the atmosphere, particularly above pH 10. This method was effective as proved by the pH stability of the solution at any given pH over a period of 15 min. While it is true that the addition of KOH to the solution contribut’es to the ionic strength, this contribution is negligible for most cases. At pH 11.3, the KOH added results in an increase in ionic strength of 0.002. This means that the great-
DUNFORD est effect due to the added base would be a 107, increase in the ionic strength for the last point taken in t.he experiment performed at p = 0.02. If it is assumed that a water molecule is in the sixth coordination position of t,he heme iron of HRP; its ionization can be represented by: P @ POH + Hf.
(1)
Charges on the protein species have been omitted since they are not known. The reversibility of the process represented in Eq. (1) was tested in the following experiment. The absorbance curve of the Soret peak of an HRP solution at pH 5.5 was taken, the pH was raised to 11, then readjusted to 5.5 where the spectrum was remeasured and found to be invariant wit,hin instrumental error (<0.5%). RESULTS
The total absorbance of the solution at any pH is the sum of the absorbances due to the P and POH species, which for a l-cm cell is given by: A = +[Pl + +o,WHl,
(2)
where + and +oH are the molar absorptivities of P and POH at 416 rnp. The value of tpoH cannot be determined directly because it is not possible to raise the pH of the solution high enough to form 100% HRP-OH without causing irreversible changes in the constant correenzyme. The equilibrium sponding to Eq. (1) is given by:
K 0 = POHl[H+l IPI *
(3)
Equations (2) and (3) can be combined to give the expression : [H+]AA AA = [P]oAe - K , 0
(4)
where AA = A - A0 ; A0 is the absorbance of the solution when all HRP is in the P form; [PI0 is the total concentration of HRP; and Ae = cPOH - +. The value of [H+] is obtained from the equation: pH = -log
[H+].
(5)
The validity of Eq. (5) will be discussed later. It follows that a plot of AA vs. [H+]AA should yield a straight line, with a slope of
HYDROXIDE
COMPLEX
315
OF HRP
value of cpq03= 9.1 X lo4 al-‘crnM1 (2). From the intercept of the plot of Eq. (4) we obtain Ae = EpoH416- Ep416= 3.74 x lo4 d cm-l, from which EpoH416= 9.82 X lo4 M-I cm-‘. The effect of ionic strength on the value of K. can be treated on the basis of the theory of the secondary salt effect. For a reaction of the type
0.5
04
0.3
HA” $ H+ + ,4 (‘-l),
2l
(6)
where Z is the charge on the undissociated acid, the equilibrium constant, Kc , expressed in terms of concentrations (Eq. 3), is related to the thermodynamic equilibrium constant, Kt , by the expression
0.2
0.1
0 0
0.5
1.0
_
1.5
AA [H+]x 10” FIG. 1. Plot of AA vs. AA[H+] for a spectrophotometric titration of pure horseradish peroxidase at 25.0’ and 0.14 p. The pH at which each experimental point was determined is indicated on the plot.
-l/K, and an intercept of [PgAe. Deviations from linearity at lower pH values were always found to occur that were largest for “purified” HRP and smallest for “analytical reagent grade” HRP. A plot testing Eq. (4) is shown in Fig. 1, the data being from an experiment using ‘lanalytical reagent grade” HRP at 25.0’ and p = 0.14. Table I summarizes the results of the linear plots of Eq. (4) obtained at high pH for “analytical reagent grade” HRP. DISCUSSION
The molar absorptivity of HRP-OH, determined at pH 11.4, has been reported as +oH416 = 8.85 X lo4 11-l cm-’ (2). However, from the results of the present study, it is known that hydroxide formation is not complete at pH 11.4. We measured the ratio &,3/A a16for an “analytical reagent” HRP solution at pH 7, and obtained the result 0.668, which yields a value of ~416= 6.08 X lo4 1z1-l cm-’ from Keilin and Hartree’s
log Kc = log K, + 2(1 - Z)D. (7) The symbol D in the above equation stands for a portion of the extended Debye-Htickel expression due to Davies (9), which at 25’ and with water as the solvent is given by:
& D = 0.509 __-- 0.20/~ . (8) ( 1 + d/J > The experimental values of Kc, listed in Table I, were obtained by assuming that Eq. (5) is valid. However, the definition of pH is an operational one. The term “pH” does not have a simple meaning; it is not a true measure of either the hydrogen ion concentration or hydrogen ion activity, but probably of a quantity between the two (10). If TABLE
I
VALUES OF THF: EQUILIBRIUM CONSTANT, Kc; = (POH) (H+)/(P) OWT.II~VEDas .i FUNCTION OF IONIC STREKGTH AND TEMPEK.ITURE L1 0.02 0.05” 0.08 0.11” 0.14 0.17” 0.20” 0.11 0.11 B Duplicate conditions.
t, "C
Kc (M)x loli ~~
25.0 25.0 25.0 25.0 25.0 25.0 25.0 18.0 35.0
1.44 1.13 1.00 1.12 0.91 0.94 0.90 0.87 1.51
experiments
performed
PKC .~ 10.84 10.95 11.00 10.95 11.04 11.03 11.05 11.06 10.82 for
these
316
ELLIS r
1
AND
parameter, D, with temperature is assumed to be negligible. The logs of these values vs. l/T are plotted in Fig. 3; from linear leastsquares analysis AH” is found to equal 5.7 kcal/mole. Values for AG” and AS” at 25O are readily obtained :
I
I
DUNFORD
AG” = -RT
In Kt = 14.7 kcal/mole
Aso = AH” - AG”
T
= -30
e.u.
These values may be compared to the corresponding thermodynamic quantities for other ionizations of water in different environments, shown in Table II. As can be seen in that table, the thermodynamic con-
FIG.
(-\/;/(l
2. Plot of log Ko vs. D where D = 0.509 +v$) - 0.20 p) and p is the ionic strength.
it were a measure of hydrogen ion activity, then for this study, Eq. (7) would take the form : log Kc’ = log K, +
(1 -
22)D,
(9)
where Kc ’ is the same as Kc except that hydrogen ion concentration is replaced by hydrogen ion activity. Figure 2 is a linear plot of log Kc vs. D with a slope of -2.5 & 0.4 as determined by least-s&ares analysis. Therefore, values of 2, determined from Eq. (7 and 9) are 2.3 and 1.8. The linear plot obtained in Fig. 2 is in contrast to result; obtained on studies of the formation of ferrimyoglobin hydroxide (11) and ferrihemoglobin hydroxide (12). The heat of ionization can be obtained from a linear plot of log K, vs. l/T; the slope of such a plot is AH0/2.303R. From the intercept in Fig. 2 a value of K, = 2.00 X 10-l’ M at 2.5’ is obtained. Values of Kt at 18.0” and 35.0”, calculated from either of Eq. [7 or 91, using a value of -2.5 for the slope in Fig. 2, are 1.77 X lo-l1 and 3.00 X 10-l’ M. The variation of the Debye-Htickel
Z 52
E 3 (3 0.3 9 0.2 -
I 3.3
0.1 3.2
FIG.
I 3.4
ho3 I 3. Plot of log KT vs. l/l’
thermodynamic
equilibrium TABLE
where constant.
KT is the
II
THERMODYNAMIC OF WATER, VARIOUS
CONSTANTS FOR THE IONIZATION AND WATER COORDINATED TO Fe(II1) COMPLEXES, AT 25.0"
--__ Hz0 Fe(OH#+ Ferrihemoglobin Ferrimyoglobin Ferri-HRP4
19.1 13.5 2.96 10.4 11.90 3.91 12.12 5.75 14.1 5.1
-18.7 25.0
-27.3 -21.7 -30
15 16
12 11 This study
HYDI~OSIT)E
COMPLEX
stank for the ionization of water in the sixth position of the heme proteins are not too different from one another, but as a group they differ from either the ionization of water alone or 011 Fe (OH:!)G3+, showing the considerable effect of the protein-heme structure on the nature of the hydrogen-oxygen bond in the bound water molecule. The deviations from linearity in Fig. 1 and similar plots for other data obtained with “analytical reagent grade” HRP appear within experimental error. However, the consistency of the deviation and the trend toward smaller deviations with increased purity of HRP suggest that the deviations may be caused by the presence of impurity or of an isozyme of HRP (13, 14) the concentration of which appears negligible in the “analytical reagent grade” HRP. REFERENCES 1. THEORELL, H., AND MAEHLY, A. C., Acta Chem. &and. 4, 422 (1950). 2. KEILIN, L)., .\ND HARTREE, E. F., Biochem. J. 49, 88 (1951).
OF HRP
:317
3. HARHURY, H. il., J. Biol. Chem. 225, 1009 (1957). 4. BLUMTIERG, W. E., PISUCH, J., WITTI,:NHI~:IG, B. A., .\ND WITTESBRRO, J. B., .I. Viol. Chem. 243, 1854 (1968). 5. GEORGE, P., AND LYSTER, R. L. J., T’roc.. .Vatl. Acad. Sci. U.S. 44, 1013 (1958). 6. THEORELL, T-I. =Irkic ICo?li Minrwl. GPO/. 16A, No. 3 (1942). 7. ELLIS, W. D., AND DUSFORD, H. B., f~ioch~tt~istry 7, 2054 (1968). 8. MAEHLY, A. c., Riochivl. Riophys. .-Icln 8, 1 (1952). 9. D.4v11c3, c. w., J. Chem. sot. 1938, 209:3. lo. GOLD, v., “pH Measuremertts,” pp. 34-45. Wiley, New York (1950). 11. GEORGE, P., AND HINANIA, G., Eioch~n~. J. 62, 517 (1952). 12. GEORGZ, P., AND H.1N.1~1.1, G., Bioche,,l. J. 56, 236 (1953). 13. SHANNON, L. ill., Kal-, E., AND L+:H., J. W., J. Biol. Chem. 241, 2166 (1966). 14. Kay, E., SHANNON, L. M., AND LEN-, .J. W., J. Biol. Chem. 242, 2470 (1967). 15. CLEVIX, H. L., J. Chem. Ed. 46, 231 (1968). 1G. MILBURX, 11. M., J. --Lnl. Chem. Sot. 79, 537 (1957).
OF
BIOCHEMISTRY
The Effects of the
AND
of Temperature
Hydroxide
and
Complex
WII,LIAi/l Department
133, 313-317 (1969)
BIOPHYSICS
of Ferric
D. ELLIS2
of Chemistry, Received
ionic
AND
University July
Strength
Horseradish
H. BRIAK
of Alberta,
1, 1968; accepted
on the
Formation
Peroxidase’
DUNFORD
Edmonton,
dlberta,
Canada
June 9, 1969
The equilibrium constants for the formation of the hydroxide complex of ferric horseradish peroxidase have been determined spectrophotometrically at 25.0” for seven values of ionic strength over the range 0.02-0.20 and also at 18.0” and 35.0” with p = 0.11. The results were analyzed on the assumption that the hydroxide complex is formed by the ionization of a water molecule in the sixth coordination position of the heme ferric iron. The effective charge on the aquo species is indicated to be in the range +2 f 0.5; its dissociat,ion constant is 2.0 X lo-i1 M and the thermodynamic parameters for its ionizat,ion are indicated t,o be AG” = 15 kcal/mole; AH” = 6 kcal/ mole; AS” = -30 entropy unit.s at 25” and zero ionic strength.
It is widely accepted that HRP,4 like hemoglobin and myoglobin, contains water in the sixth coordination position of the heme ferric iron (24), although the suggestion has been made that the heme of HRP lies in a crevice with both the fifth and sixth positions occupied by protein residues (5). When a solution of HRP is made alkaline, there are accompanying changes in the spectrum of HRP (2), and its magnetic moment changes from 5.45 to 2.66 Bohr magnetons (6), indicating the formation of a new species, HRPOH. The pK for hydroxide formation has been measured to be 10.9 by spectrophotometry and 11.3 by use of a magnetic balance (6), and 10.6 by a study of the reduction i Supported financially by the National Research Council of Canada. 2 Holder of an Izaak Walton Killam Memorial Fellowship. 3 To whom reprint, requests should be sent. 4 Abbreviations used: HRP, ferric horseradish peroxidase; P, refers specifically to a form of HRP in which it is assumed that water is in the sixth coordinatiou position; HRP-OH or POH, hydroxide form of HRP; p, ionic strength; RZ, purity number, after Theorell and Maehly (1); e.u., entropy units. 313
potentials of HRP (3), but these studies were not carried out at constant ionic strength. A study of the kinetics of cyanide binding to HRP at an ionic strength of 0.11 yielded a value of 10.8 for this pK (7). We report here on a spectrophotometric study of the formation of HRP-OH at 25.0” for seven values of k over the range 0.02-0.20 and also at 18.0’ and 35.0” with ~1= 0.11. MATEKIALS
AND
METHODS
Experimental. Two grades of HRP were obtained from Boehringer-Mannheim Corp. N.Y.: “analytical reagent grade” (RZ = 2.9 lot 648206) and “purified lyophilized powder” (RZ = 0.8). All HRP was exhaustively dialyzed against conductivity water and centrifuged before use, Analytical reagent grade potassium nitrate (Mallinckrodt) and potassium hydroxide (B.D.H.) were used without further purification. Some of the “purified” HRP was treated by passage through a 1.1 X 50-cm column of C-50 CM-Sephadex cation exchange gel. The HRP was eluted using pH 6.2 phosphate buffer, with ionic strength increasing from 0.1 to 0.2 during the course of the elution. A single passage through the column resulted in the recovery of a fraction of RZ = 2.7 which represented 65% of the initial “purified” HRP as determined by absorbance measurements at 497 mN. Absorbance measure-
314
ELLIS
AND
ments were made with a Beckman DU spectrophotometer which had the cell compartment thermostated to f0.1”. A Beckman Expandomatic pH meter was used for all pH measurements. For a typical experiment, an appropriate amount of 1 M KNO, and enough stock HRP to make a final HRP concentration of about 1.7 X lo-& M were pipetted into a IO-ml volmetric flask, which was then filled to the mark with water. The 10 ml of solution was transferred to a 30-ml beaker which was covered with Parafilm and placed in a constant temperature bath. A Beckman 39183 combination pH electrode, a Teflon syringe needle, and a platinum stirring wire were inserted through slits in the Parafilm. The pH was measured and was found typically to have a value of 5.5. A 3-ml aliquot of the solution was transferred with a syringe to a cuvette and the HRP concentration was determined by measuring the absorbance of t,he solution at 497 rnp where the molar absorptivity coefficient is 1.00 X lo4 M-lcrn-l (2). The absorbance of the solution was t.hen measured at 416 rnp, the wavelength corresponding to the maximum of the Soret peak of HRP-OH (2). B second HRP solution of about pH 5.5 was prepared to use as a blank in the spectrophotometer. Its concentration was such that the difference in absorbances between the blank and the experimental solution was about 0.1 at 416 rnp, t.hus keeping t,he absorbance readings during the course of the titration in the range 0.14.6, which includes the region of minimum error due to instrument,al uncertainty. The experimental HRP solution was returned to the beaker and the pH was raised by the addition from a lo-p1 Hamilton syringe of a few microliters of approximately 1 M KOH, an amount small enough that no volume correction was necessary. The solution was stirred, the pH was remeasured, the absorbance of an aliquot was measured at 416 w in a capped cuvette, and the aliquot was returned to the beaker. Nine to eleven points were obtained this way, with the titration being carried out up to about. pH 11.35. Above this pH, the heme splits from the protein of HRP at an appreciable rate (8). After each increase in pH, the beaker was flushed gently with nitrogen so that there would be no change in pH or ionic strength due to absorption of carbon dioxide a problem encountered from the atmosphere, particularly above pH 10. This method was effective as proved by the pH stability of the solution at any given pH over a period of 15 min. While it is true that the addition of KOH to the solution contribut’es to the ionic strength, this contribution is negligible for most cases. At pH 11.3, the KOH added results in an increase in ionic strength of 0.002. This means that the great-
DUNFORD est effect due to the added base would be a 107, increase in the ionic strength for the last point taken in t.he experiment performed at p = 0.02. If it is assumed that a water molecule is in the sixth coordination position of t,he heme iron of HRP; its ionization can be represented by: P @ POH + Hf.
(1)
Charges on the protein species have been omitted since they are not known. The reversibility of the process represented in Eq. (1) was tested in the following experiment. The absorbance curve of the Soret peak of an HRP solution at pH 5.5 was taken, the pH was raised to 11, then readjusted to 5.5 where the spectrum was remeasured and found to be invariant wit,hin instrumental error (<0.5%). RESULTS
The total absorbance of the solution at any pH is the sum of the absorbances due to the P and POH species, which for a l-cm cell is given by: A = +[Pl + +o,WHl,
(2)
where + and +oH are the molar absorptivities of P and POH at 416 rnp. The value of tpoH cannot be determined directly because it is not possible to raise the pH of the solution high enough to form 100% HRP-OH without causing irreversible changes in the constant correenzyme. The equilibrium sponding to Eq. (1) is given by:
K 0 = POHl[H+l IPI *
(3)
Equations (2) and (3) can be combined to give the expression : [H+]AA AA = [P]oAe - K , 0
(4)
where AA = A - A0 ; A0 is the absorbance of the solution when all HRP is in the P form; [PI0 is the total concentration of HRP; and Ae = cPOH - +. The value of [H+] is obtained from the equation: pH = -log
[H+].
(5)
The validity of Eq. (5) will be discussed later. It follows that a plot of AA vs. [H+]AA should yield a straight line, with a slope of
HYDROXIDE
COMPLEX
315
OF HRP
value of cpq03= 9.1 X lo4 al-‘crnM1 (2). From the intercept of the plot of Eq. (4) we obtain Ae = EpoH416- Ep416= 3.74 x lo4 d cm-l, from which EpoH416= 9.82 X lo4 M-I cm-‘. The effect of ionic strength on the value of K. can be treated on the basis of the theory of the secondary salt effect. For a reaction of the type
0.5
04
0.3
HA” $ H+ + ,4 (‘-l),
2l
(6)
where Z is the charge on the undissociated acid, the equilibrium constant, Kc , expressed in terms of concentrations (Eq. 3), is related to the thermodynamic equilibrium constant, Kt , by the expression
0.2
0.1
0 0
0.5
1.0
_
1.5
AA [H+]x 10” FIG. 1. Plot of AA vs. AA[H+] for a spectrophotometric titration of pure horseradish peroxidase at 25.0’ and 0.14 p. The pH at which each experimental point was determined is indicated on the plot.
-l/K, and an intercept of [PgAe. Deviations from linearity at lower pH values were always found to occur that were largest for “purified” HRP and smallest for “analytical reagent grade” HRP. A plot testing Eq. (4) is shown in Fig. 1, the data being from an experiment using ‘lanalytical reagent grade” HRP at 25.0’ and p = 0.14. Table I summarizes the results of the linear plots of Eq. (4) obtained at high pH for “analytical reagent grade” HRP. DISCUSSION
The molar absorptivity of HRP-OH, determined at pH 11.4, has been reported as +oH416 = 8.85 X lo4 11-l cm-’ (2). However, from the results of the present study, it is known that hydroxide formation is not complete at pH 11.4. We measured the ratio &,3/A a16for an “analytical reagent” HRP solution at pH 7, and obtained the result 0.668, which yields a value of ~416= 6.08 X lo4 1z1-l cm-’ from Keilin and Hartree’s
log Kc = log K, + 2(1 - Z)D. (7) The symbol D in the above equation stands for a portion of the extended Debye-Htickel expression due to Davies (9), which at 25’ and with water as the solvent is given by:
& D = 0.509 __-- 0.20/~ . (8) ( 1 + d/J > The experimental values of Kc, listed in Table I, were obtained by assuming that Eq. (5) is valid. However, the definition of pH is an operational one. The term “pH” does not have a simple meaning; it is not a true measure of either the hydrogen ion concentration or hydrogen ion activity, but probably of a quantity between the two (10). If TABLE
I
VALUES OF THF: EQUILIBRIUM CONSTANT, Kc; = (POH) (H+)/(P) OWT.II~VEDas .i FUNCTION OF IONIC STREKGTH AND TEMPEK.ITURE L1 0.02 0.05” 0.08 0.11” 0.14 0.17” 0.20” 0.11 0.11 B Duplicate conditions.
t, "C
Kc (M)x loli ~~
25.0 25.0 25.0 25.0 25.0 25.0 25.0 18.0 35.0
1.44 1.13 1.00 1.12 0.91 0.94 0.90 0.87 1.51
experiments
performed
PKC .~ 10.84 10.95 11.00 10.95 11.04 11.03 11.05 11.06 10.82 for
these
316
ELLIS r
1
AND
parameter, D, with temperature is assumed to be negligible. The logs of these values vs. l/T are plotted in Fig. 3; from linear leastsquares analysis AH” is found to equal 5.7 kcal/mole. Values for AG” and AS” at 25O are readily obtained :
I
I
DUNFORD
AG” = -RT
In Kt = 14.7 kcal/mole
Aso = AH” - AG”
T
= -30
e.u.
These values may be compared to the corresponding thermodynamic quantities for other ionizations of water in different environments, shown in Table II. As can be seen in that table, the thermodynamic con-
FIG.
(-\/;/(l
2. Plot of log Ko vs. D where D = 0.509 +v$) - 0.20 p) and p is the ionic strength.
it were a measure of hydrogen ion activity, then for this study, Eq. (7) would take the form : log Kc’ = log K, +
(1 -
22)D,
(9)
where Kc ’ is the same as Kc except that hydrogen ion concentration is replaced by hydrogen ion activity. Figure 2 is a linear plot of log Kc vs. D with a slope of -2.5 & 0.4 as determined by least-s&ares analysis. Therefore, values of 2, determined from Eq. (7 and 9) are 2.3 and 1.8. The linear plot obtained in Fig. 2 is in contrast to result; obtained on studies of the formation of ferrimyoglobin hydroxide (11) and ferrihemoglobin hydroxide (12). The heat of ionization can be obtained from a linear plot of log K, vs. l/T; the slope of such a plot is AH0/2.303R. From the intercept in Fig. 2 a value of K, = 2.00 X 10-l’ M at 2.5’ is obtained. Values of Kt at 18.0” and 35.0”, calculated from either of Eq. [7 or 91, using a value of -2.5 for the slope in Fig. 2, are 1.77 X lo-l1 and 3.00 X 10-l’ M. The variation of the Debye-Htickel
Z 52
E 3 (3 0.3 9 0.2 -
I 3.3
0.1 3.2
FIG.
I 3.4
ho3 I 3. Plot of log KT vs. l/l’
thermodynamic
equilibrium TABLE
where constant.
KT is the
II
THERMODYNAMIC OF WATER, VARIOUS
CONSTANTS FOR THE IONIZATION AND WATER COORDINATED TO Fe(II1) COMPLEXES, AT 25.0"
--__ Hz0 Fe(OH#+ Ferrihemoglobin Ferrimyoglobin Ferri-HRP4
19.1 13.5 2.96 10.4 11.90 3.91 12.12 5.75 14.1 5.1
-18.7 25.0
-27.3 -21.7 -30
15 16
12 11 This study
HYDI~OSIT)E
COMPLEX
stank for the ionization of water in the sixth position of the heme proteins are not too different from one another, but as a group they differ from either the ionization of water alone or 011 Fe (OH:!)G3+, showing the considerable effect of the protein-heme structure on the nature of the hydrogen-oxygen bond in the bound water molecule. The deviations from linearity in Fig. 1 and similar plots for other data obtained with “analytical reagent grade” HRP appear within experimental error. However, the consistency of the deviation and the trend toward smaller deviations with increased purity of HRP suggest that the deviations may be caused by the presence of impurity or of an isozyme of HRP (13, 14) the concentration of which appears negligible in the “analytical reagent grade” HRP. REFERENCES 1. THEORELL, H., AND MAEHLY, A. C., Acta Chem. &and. 4, 422 (1950). 2. KEILIN, L)., .\ND HARTREE, E. F., Biochem. J. 49, 88 (1951).
OF HRP
:317
3. HARHURY, H. il., J. Biol. Chem. 225, 1009 (1957). 4. BLUMTIERG, W. E., PISUCH, J., WITTI,:NHI~:IG, B. A., .\ND WITTESBRRO, J. B., .I. Viol. Chem. 243, 1854 (1968). 5. GEORGE, P., AND LYSTER, R. L. J., T’roc.. .Vatl. Acad. Sci. U.S. 44, 1013 (1958). 6. THEORELL, T-I. =Irkic ICo?li Minrwl. GPO/. 16A, No. 3 (1942). 7. ELLIS, W. D., AND DUSFORD, H. B., f~ioch~tt~istry 7, 2054 (1968). 8. MAEHLY, A. c., Riochivl. Riophys. .-Icln 8, 1 (1952). 9. D.4v11c3, c. w., J. Chem. sot. 1938, 209:3. lo. GOLD, v., “pH Measuremertts,” pp. 34-45. Wiley, New York (1950). 11. GEORGE, P., AND HINANIA, G., Eioch~n~. J. 62, 517 (1952). 12. GEORGZ, P., AND H.1N.1~1.1, G., Bioche,,l. J. 56, 236 (1953). 13. SHANNON, L. ill., Kal-, E., AND L+:H., J. W., J. Biol. Chem. 241, 2166 (1966). 14. Kay, E., SHANNON, L. M., AND LEN-, .J. W., J. Biol. Chem. 242, 2470 (1967). 15. CLEVIX, H. L., J. Chem. Ed. 46, 231 (1968). 1G. MILBURX, 11. M., J. --Lnl. Chem. Sot. 79, 537 (1957).