ARCHIVES
OF
BIOCHEMISTRY
AND
A Potentiometri~
BIOPHYSICS
(1968)
Study of the Fhvin
ROY D. DRAPER2
Department
802808
I!%,
of Biochemistry
LLOYD
AND
and Biophysics,
Received September
Semjqui~one
University
Equilibrium’
L. INGRAHAM
of California,
11, 1967; accepted November
Davis,
California
21, 1967
Potentiometric
studies of the riboflavin and flavin mononucleotide
(FMN) equilibria with reduced and semiquinone species have been performed by the classical
method of Michaelis. A value has been determined for the pk’ of the semiquinone of
FMN and s corrected value for the semiquinone pK of riboflavin is reported. The
equilibrium
constants between reduced, oxidized, and semiquinone forms of riboflavin in baeic solution are less than one; this is in contradiction to the second paper of
Michaelis but in agreement with the first paper.
Michaelis et al. (1) measured the equilibria
between reduced, oxidized, and the semiquinone forms of riboflavin at various pH
values and found values for the forma~on
constant below one in all cases. Michaelis
and Schwarzenbach (2) later reported a
different dependence of the equilibria on
pH with values of the formation constant
greater than one in solutions of high pH.
These latter values are those accepted in
the literature today. However, this dependence of the equilibria on pH requires a
value for the pK of the semiquinone that is
not consistent with modern spectral studies
of riboflavin. We have therefore reinvestigated the dependence of the equilibria on
pH by the classicalpotentiome~,ric method of
Michaelis (3, 4) in order to reevaluate the
pK of the semiquinone and to determine the
formation constant in basic solution.
EXPERIMENTAL
PROCEDURE
~a~er~aZ~. R,iboflavin from the California Corporation for Biochemical Research gave only one
fluorescent spot when chromatographed
in three
different solvent systems, namely butanol-acetic
acid-water
(4: 1: 5) ; pyridine-propanol-water
1 This work was supported by U.S. Public
Health Service grant GM-#~.
2 Recipient of a predoctoral fellowship from the
National Institutes of Health. Present address:
Department of Chemistry, Sacramento State College, Sacramento, California.
802
(3:5:2); and 5’% Na2HP04.12 H20. The RF of the
material in all three solvents agreed well with the
value reported by Whitby (5) and Yagi (6). The
ratios of the extinction coefficients at 260 mp/375
m&, 260 m&50 rnp, and 260 mp/510 rnp were 2.564,
2.229, and 68.03, respectively, which also agree well
with the values of Whitby (7): 2.5F6, 2.230, and
68.00. The riboflavin
was used without further
purification.
Saturated
stock solutions of riboflavin were freshly prepared each week in deionized water and stored in a refrigerator. Aliquots
of the stock solution were mixed in a loo-ml Berzelius beaker with known quantities of buffer. The
solution was then deoxygenated for 15 minutes
with a stream of purging gas. Care was taken
during the entire operation to prevent light from
entering the solution.
Commercial flavin moIlonucleatide
(FMN) from
the California Corporation
for Biochemical
Research was reported to be only 70% pure and
showed five fluorescent, yellow spots in butanolacetic acid-water
(4:1:5) on Whatman No. 3 filter
paper. The dominant spot was eluted with water
and the eluate was lyophilized
to dryness after
altering. Re~hromatographing
in the three solvents used to test the purity of riboflavin gave
only one fluorescent spot. The ratios of the extinction coefficients of the lyophilized material at 260
mr(/375 w, 260 mp/450 rnp, and 260 me/510 mp were
2.604, 2.212, and 67.14, respectively, which again
agree well with the values of Whitby (7): 2.606,
2.221, and 67.75. Stock solutions of the lyophilized
material were made as with riboflavin.
Methods.
All potentiometric
measurements were
made using a Leeds and Northrup Type K-3 Uni-
FLAVIN
SEMIQUINONE
versa1 potentiometer,
which is capable
of measuring potentials
to the nearest
0.01 mV with a 0.01%
error.
The reference
electrode
was a 3.5 M KClcalomel
electrode
constructed
as described
by
Daniels
and Alberty
(8). The voltage
of the 6-V
working
lead
storage
battery
was
calbrated
against
an Eppley
laboratory
certified
unsaturated
cadmium
cell.
A lOO-ml Berzelius
beaker
fitted with a number
10 rubber
stopper
was used as the titration
vessel.
A number
of holes were drilled
in the stopper
to
provide
access
for two platinum
electrodes,
a
burette,
a salt bridge,
a thermometer,
and an inlet
and outlet
for a purging
gas. Other
holes were
filled with rubber
plugs that could be removed
to
add material
during
the course of a titration
or to
insert
a glass electrode
to measure
the pH. The
titration
vessel was covered
with
a black
paper
jacket
during
a titration
of either
riboflavin
or
FMN
to minimize
the possibility
of photodecomposition
of the flavin.
The temperature
was recorded
at the beginning
and end of the titration,
and the average
was used in the analysis
of the
data.
The temperature
was maintained
within
a
degree of 20”.
In reductive
titrations
it is necessary
to remove
as much oxygen
as possible
from the solution
before starting
the titration
and to perform
the titration in an oxygen-free
atmosphere.
Either
helium
or deoxygenated
nitrogen
was used to purge
the
oxygen
before
titration.
The
nitrogen
was deoxygenated
by the method
of Fieser
(9). The
stream
of purging
gas bubbling
through
the titration solution
also served
as a mixer.
A 5-ml microburette
accurate
to 0.001 ml was
fitted
into the top of the titration
vessel such that
the tip was in contact
with the solution.
A small
arm sealed to the uncalibrated
portion
allowed
the
burette
to be flushed with purging
gas before it was
filled with sodium
hydrosulfite
solution.
The filling
of the burette
was accomplished
by gravity
flow
from
a reservoir
that
contained
deoxygenated
sodium
hydrosulfite
solution.
Two glass-sealed
platinum
electrodes
supplied
by E.H.
Sargent
and Company
were used in the
flavin
half-cell.
The second electrode
was used to
check the potential
measured
by the first electrode
to within
0.05 mV.
DATA
ANALYSIS
Early in the investigation
it became
apparent
that
the Michaelis
graphical
method of treating the data introduced large
errors in the value of the semiquinone
formation constant. This becomes evident if
we realize that an error of only 0.08 mV in
the midpoint potential (Em) will cause an
503
EQUILIBRIUM
8 % error in a semiquinone formation constant of 0.05. The slopes in the center portion of the curve are extremely difficult to
estimate accurately by graphical methods.
When the semiquinone constant is 0.1, an
error of only 3% in the slope will cause an
error of 35% in the semiquinone formation
constant.
Errors in data analysis were reduced by
using an iterative method for analysis of the
potentiometric
curves. Iterative
methods,
not feasible in the time of Michaelis, are
now possible because of the availability
of
computers.
The dependence
of the experimental
potentials, E, on the parameter P(P = -1
for zero equivalents oxidant, P = 0 for one
equivalent oxidant and p = +l
for two
equivalents oxidant) were fitted to Eq. (1)
(Refs. 3, 4) by a computer program developed in this laboratory.
E=
E
+
+ftTh(l-p)
m
2%’
(1
+
3n
2F
P
+
r(l
TABLE
OF FMN
-
P2>
constant
K = 4/o
COIIC.
(X 1OSP)
+
K is related
to
+ 7).
(2)
I
AND
RIBOFLAVIN
DIMERIZATION
FMN
PH
(1)
[l + r(l - /.L2)- /.P2
The equilibrium
Y by Eq. (3:
STUDY
PU>
SEMIQUINONE
Riboflavin
k”
PH
(X
COIIC.
105 M)
ka
4.00
1.87
3.66
8.53
12.8
15.4
0.100
0.099
0.102
0.125
0.162
4.50
1.59
3.55
8.00
12.0
0.095
0.097
0.100
0.108
8.00
1.87
3.66
5.90
8.53
12.8
15.4
0.103
0.095
0.104
0.128
0.166
0.202
8.00
1.59
3.55
8.00
12.0
16.0
0.128
0.125
0.131
0.174
0.215
tion
a The abbreviation
constant.
k is the semiquinone
forma-
804
DRAPER
AND
INGRAHAM
When the program was checked with calculated data, the error in the E, was usually
less than 1 part in 10,060, while the error in
the formation
constant of 0.05 was 2%,
0.4% when the K was 0.5, and 0.4% when
K was 10,000. In a typical titration curve
consisting of 43 points the average fit of experimental to calculated voltage was 0.05%
error.
RESULTS
AND
12.8 X 1O-5 M, while at pH 8.06 the formation constant
increases between 5.90 X
10m5and 8.53 X 1O-5 M. Lowe and Clark (10)
showed the occurrence of dimerization above
10 X 10-5 M FMN
at pH 7.01 and at
1 X 10V2 M FMN at pH 6.22, but only used
two FMN concentrations
at any pH value.
At pH 4.50 riboflavin shows little tendency
for
dimerization.
Michaelis
and
Schwarzenbach
(2) have reported the same
results at pH 4.62. At pH 8.00 dimerization
of the semiquinone begins between 8.00 X
1O-5 and 12.0 X 1CP5 M riboflavin.
At
PH 6.92 Michaelis and Schwarzenbach
(2)
have shown that above 1 X 10”’ M riboflavin, dimerization becomes noticeable and
DISCUSSION
Table I summarizes the results of titrations with varying starting concentrations
of
riboflavin and FMN at two different pH
values. In the case of FMN at pH 4.06, the
semiquinone
formation
constant increases
between concentrations
of 8.53 X UP5 and
40.10
0.00
-0.10
Y
2
-0.P
e
-0.30
-0.40
-0.50
2
3
4
5
6
7
8
9
10
11
12
13
PH
1. The three normal
potentials
of riboflavin
plotted
against
normal
potential
of the oxidized-semiquinone
system;
E2 is the
reduced
semiquinone
system;
and the EM is the normal
potential
system.
The ~KR, pKs, and PKT values
are the dissociation
constants
quinone,
and totally
oxidized
forms, respectively.
The individual
Ei
S-10% error in the formation
constant
from which they are calculated.
FIG.
pH at 20 f 1”. El is the
normal
potential
of the
of the oxidized-reduced
of the reduced,
semiand Et points reflect the
FLAVIN
SEMIQUINONE
rapidly increases with increasing concentration.
Although the concentration
at which dimerization begins to occur has not been
determined for either riboflavin
or FMN
through the entire pH range, it is felt that a
starting
concentration
of FMN and riboflavin less than 6 X W5 M can be used
lvithout fear of dimerization.
Most of the
titrations
were run at starting concentrations of FMN
and riboflavin
between
3 X 10e5 and 5 X 1O-5 M.
A discrepancy exists between the value of
the formation constant for riboflavin at pH
8.00 (Table I) and the value indicated by
Michaelis and Schwarzenbach
(2) at this
pH. When Michaelis and Schwarzenbach
plotted
the values of riboflavin’s
three
potentials against pH over a wide range, the
El, E,, and Ez curves intersected at pH 8.2,
and at pH values greater than 8.2 the positions of the El and ES curves were reversed
with respect to the E, curve. The point of
TABLE
RIBOFLAVIN
PHb
2.52
3.01
3.01
3.50
4.00
4.49
5.00
5.50
6.01
6.50
6.95
7.44
8.00
8.52
9.52
10.00
10.50
10.95
11.55
11.96
12.50
12.95
K
0.091
0.174
0.179
0.043
0.111
0.097
0.052
0.038
0.072
0.766
0.079
0.130
0.125
0.093
0.130
0.103
0.037
0.037
0.016
0.016
0.001
0.048
ELll
t-O.0412
+0.0148
$0.0121
-0.0180
-0.0446
-0.0748
-0.1089
-0.1329
-0.1614
-0.1898
-0.1990
-0.2155
-0.2328
-0.2480
-0.2864
-0.3013
-0.3257
-0.3504
-0.3818
-0.4101
-0.4395
-0.4695
805
EQUILIBRIUM
intersection indicates a formation constant
value of one, and the reversal of the positions of the El and Ez curves above pH 8.2
indicates a value of the formation constant
greater
than one. Thus,
according
to
Michaelis and Schwarzenbach,
the value of
the formation
constant should be in the
neighborhood of one at pH 8.00. The values
reported in Table I are much lower than
one. Furthermore,
titrations
of riboflavin,
at pH 8.5 and above, showed the formation
constant not to exceed 0.130 in any case.
This suggested that the earlier results of
Michaelis and Schwarzenbach
are in error,
and it was therefore necessary to completely
redetermine
the dependence of riboflavin
reduction upon pH. Even though the values
in Table I for the formation constants of
FMN are in agreement with those reported
by Lowe and Clark (lo), it was also decided
to check their work as regards the dependence of FMN reduction upon pH. Table
II shows the results of the riboflavin titraII
TITRATIONB
El
+0.0111
-0.0073
-0.0096
-0.0576
-0.0723
-0.1042
-0.1464
-0.1742
-0.1946
-0.1932
-0.2309
-0.2412
-0.2590
-0.2778
-0.2130
-0.3249
-0.3673
-0.3922
-0.4338
-0.4618
-0.5187
-0.5079
E2
$0.0714
+0.0369
+o .0338
+0.0217
-0.0168
-0.0454
-0.0714
-0.0915
-0.1281
-0.1864
-0.1671
-0.1848
-0.2066
-0.2182
-0.2607
-0.2728
-0.2841
-0.3087
-0.3298
-0.3584
-0.3603
-0.4310
Buffer systemC
Phosphate
Phosphate
Phosphate
Acetate
Acetate
Acetate
Acetate
Acetate
Phosphate
Phosphate
Phosphate
Phosphate
Tris-HCl
Tris-HCl
Carbonate
Carbonate
Carbonate
Carbonate
Phosphate
Phosphate
Phosphate
Phosphate
a Riboflavin
concentration
in the first two titrations
was 2.91 X lo-& M. All the rest were 3.55 X 10-5 M.
* This was the starting
pH; the final pH was within
0.01 pH unit of this value.
c All buffer
concentrations
were 0.1 M; their pH was checked
against
a standard
commercial
buffer
at pH 8.00 and also a standard
within
two pH units of the unknown
buffer.
806
DRAPER
AND
INGRAHAM
tions at various pH values. This data is
plotted in Fig. 1 in order to determine the
pK of the semiquinone.
The EI and E2 from each titration curve
were evaluated using the error limits of the
formation
constant from which they are
calculated, thus giving a range of values for
the El and Ez as indicated in Fig. 1. Since
the E, possesses a much lower error than
either the El and Ez, the E, curve dictated
a great deal of the relationships
within the
rest of Fig. 1. Three linear sections appear
on the E, curve, the first from pH 2.5 to
about pH 5.5, the second between pH 7.0
and 9.0, and the third between pH 10.0 and
13.0. The theoretical slope of the first linear
section was 0.0581; the second, 0.0290; and
the third, 0.0581. Straight lines with these
slopes were drawn through the points in the
appropriate
sections and were extended
(dotted lines) until they intersected.
According to the rules defined by Clark (1 l),
the intersecting
points represent the dissociation constants of the R (Reduced) and
T (Totally oxidized) forms, 6.25 and 10.00,
respectively.
These values agree quite well
with the values report’ed by Michaelis (1,
2) (~KR = 6.1 and pKT = 9.5) and Eiuhn
and Boulanger (12) (~KR = 6.3 and ~KT
= 10.0). Both of these groups used potentiometric methods to obtain the values reported. Acid-base titrations
yielded a ~KT
of 9.93 (13), while studies of fluorescence
intensity at different pH values has yielded a
pKT of 10.2 (14). The Ez curve was drawn
so that pKR equaled 6.25, and the El curve
to.1
0.0
-0.1
Y
-0.2
2
e
-0.3
-0.4
-0.5
I
2.0
I
3.0
I
4.0
I
5.0
I
6.0
I
7.0
I
8.0
II
9.0
10.0
II
11.0
12.0
I
13.0
PH
FIG. 2. The three normal
potentials
is the same as that found in Fig. 1.
of FMN
plotted
against
pH at 20 f
1”. The
notation
FLAVIN
SEMIQUINONE
was drawn such that ~KT equaled 10.00.
The slopes of the El and Ez curves were
equal to the slope of the E, curve in the
latter’s first and third regions of linearity.
Graphically,
lines of best fit were drawn
through the points t’hroughout
the entire
region for both and R and Ez curves, and
following the method used by Michaelis et
al. (1, 2) to find the pKa from the El and Ez
curves, the rest of Fig. 1 was constructed.
The value of pKs (8.27) agrees with the
value of 8.00 reported by Michaelis et al. (l),
but varies considerably from the presently
accepted value of 6.5 reported by Michaelis
and Schwarzenbach
(2). The latter value has
been a point of confusion with regards to
experimental
work by other investigators
whose results have not been entirely consistent with a pKs of 6.5. Using spectrophot’ometric methods, Beinert (15) was unable to find any difference in the spectrum
of the semiquinone form at pH 2.4 and 6.2.
If the pKs is 6.5, it is felt that at pH 6.2,
n-here one-third of the semiquinone would be
present as the anion, certain aberrations in
TABLE
FMN
pHb
2.52
3.01
3.50
4.00
4.50
5.00
5.49
6.01
6.50
6.95
7.45
8.00
8.52
9.04
9.50
10.00
10.50
10.95
11.55
11.95
12.50
12.95
a The concentration
h~c See footnotes
the spectrum
of the semiquinone
should
exist. At pH 8.9, Beinert (15) reports a large
change in the spectrum of the riboflavin
semiquinone. If the pKs is 8.27, the spectrum of pH 2.4 and 6.2 should be very nearly
the same, but at pH 8.9, where four-fifths
of the radical would be present as the anion,
it is expected that the semiquinone would
have a different spectrum. Thus the spectrophotometric
evidence reported by Beinert
(15) indicates a pK greater than 6.50 for the
semiquinone.
In another investigation,
Ehrenberg (16)
finds that the ESR spectra of riboflavin
shows changes above the reported pK of 6.5
for t’he semiquinone and in the pH range 7-9.
Table II is a summary of the reduction of
FMN studied over a wide pH range. The
three normal potentials plotted versus pH
are shown in Fig. 2, which was constructed
in the same manner as Fig. 1. The three
linear sections of the E, curve were drawn
with slopes of 0.0571, 0.0290, and 0.0580.
The error in the individual El and Ez points
are, in general, not as great as those shown
III
TITRATIONS”
K
0.115
0.115
0.091
0.099
0.087
0.094
0.100
0.109
0.085
0.073
0.071
0.120
0.112
0.098
0.127
0.277
0.094
0.076
0.057
0.068
0.065
0.063
807
EQUILIBRIUM
E2
+0.0419
-0.0131
-0.0159
-0.0449
-0.0718
-0.1022
-0.1320
-0.1610
-0.1830
-0.2053
-0.2274
-0.2423
-0.2568
-0.2716
-0.2876
-0.3031
-0.3260
-0.3518
-0.3815
-0.4050
-0.4361
-0.4630
of FMN
in all titrations
b and c in Table
II.
f0.0146
-0.0140
-0.0461
-0.0741
-0.1024
-0.1322
-0.1611
-0.1889
-0.2142
-0.2384
-0.2607
-0.2692
-0.2844
-0.3008
-0.3136
-0.3192
-0.3558
-0.3844
-0.4175
-0.4390
-0.4705
-0.4979
was 3.80 X 10-S M.
+O. 0693
+0 .0402
+0.0143
-0.0158
-0.0412
-0.0723
-0.1029
-0.1329
-0.1520
-0.1721
-0.1941
-0.2155
-0.2243
-0.2423
-0.2616
-0.2869
-0.2962
-0.3193
-0.3455
-0.3710
-0.4018
-0.4280
Buffer systed
Ph0sphat.e
Phosphate
Acetate
Acetate
Acetate
Acetate
Acetate
Phosphate
Phosphate
Phosphate
Phosphate
Tris-HCl
Tris-HCl
Tris-HCl
Carbonate
Carbonate
Carbonate
Carbonate
Phosphate
Phosphate
Phosphate
Phosphate
808
DRAPER
AND INGRAHAM
in the case of riboflavin, especially above
pH 11. In general, the potentials of the
FMN system are more stable and lead to
better analyses. The dissociation constants
of the R, S, and T forms are 6.72, 8.55, and
10.35, respectively.
The three dissociation constants of FMN
are higher than their corresponding values in
riboflavin. The formation
of the anionic
forms of FMN is delayed due to the negative charge effect of the phosphate group of
FMN.
The pH values of the reduced and oxidized
forms are in excellent agreement with those
reported by Lowe and Clark (10): 6.70 and
10.40, respectively. The pK of the semiquinone has not been previously reported.
The observation of a formation constant
for the semiquinone of less than one above
pH 8.2 has been reported by other investigators (10, 12) who used potentiometric
methods. Spectral investigation
also indicates a formation constant, for the semiquinone of less than one. For example,
Swinehart (17) estimates that ~(570 mp)
for the semiquinone is 500. If this value is
used in the data of Holmstrom
(18), the
formation constant for FMN-semiquinone
becomes -0.1, which is in agreement with
our potentiometric
data.
REFERENCES
1. MICHAELIS,
2.
3.
4.
5.
6.
L., SCHUBERT, M. P., AND SMYTHE,
C. V., J. Biol. Chem. 116,587
(1936).
MICHAELIS,
L., AND SCHWARZENBACH,
G., J.
Biol. Chem. 123,527
(1938).
MICHAELIS,
L., Chem. Rev. 16,243 (1935).
MICHAELIS,
L., AND SCHUBERT,
M. P., Chem.
Rev. 22, 437 (1938).
WHITBY,
L. G., Biochem. J. 60,433 (1952).
YAGI,
K., Methods
Biochem. Anal.
10, 319
(1962).
7. WHITBY, L. G., Biochem. J. 64,437
F., AND ALBERTY,
R.
8. DANIELS,
Chemistry.” Wiley, New York
9. FIESER,
L. F., J. Am. Chem.
(1924).
10. LOWE, H. J., AND CLARK, W. M.,
221, 983 (1956).
11. CLARK,
W. M.,
(1953).
A., “Physical
(1955).
Sot.
J. Biol.
46,
2639
Chem.
“Oxidation-Reduction
Potentials of Some Organic Systems.” Williams &
Wilkins, Baltimore, Maryland (1960).
12. KUHN,
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