CLIN. CHEM. 32/2, 314-319 (1986)
Stabilityof NADPH: Effectof VariousFactorson the Kineticsof Degradation
James
T. Wu, Lily H. Wu, and Joseph A. Knight
Seeking to minimizedegradationof NADPH duringstorage,
reagentpreparation,and assays, we investigatedthe effects
of pH, temperature,and ionicstrengthas well as the effects
of phosphate and acetate. NADH was also included for
comparison. Our resultsindicatethat the rate of degradation
of NADPH is proportionalmost importantlyto temperature
and concentrationsof hydroniumion, but also to concentrations of phosphate and acetate. The degradation rate decreased with increasing ionic strength at neutral pH, but
increased slightly at lower pH. NADPH generallyis less
stable than NADH under the same conditions.The reaction
orders with respect to hydroniumionand anionswere near 1
for NADH degradationreactions,about0.5 for NADPH. Rate
constants for NADH and NADPH differedmore at higher pH
and lower phosphateand acetate concentrations.
AddftlonalK.yphrases: variation, source of
N4DH compared
Reduced forms of both nicotinamide adenine dinucleotide
(NADH) and nicotinamide adenine dinucleotide phosphate
(NADPH) are coenzymes of many dehydrogenases.
Although there have been extensive studies concerning stability of NADH, relatively few have involved NADPH, perhaps
because no difference between the two was expected. However, it is interesting that they are fundamentally
different in
most biochemical reactions: NADH is oxidized by the respiratory chain to generate ATP, whereas NADPH serves as a
hydrogen and electron donor in reductive biosynthesis.
Both NADH and NADPH are known to be unstable in
acidic medium (1-4). However, almost all studies were
related to NADH and its model compounds such as 1,3substituted
1,4-dihydropyridine
(1, 2). As a result, little is
known about the degradation kinetics, the products, or the
spectral changes associated with NADPH degradation,
although Lowry et al. (3) found NADPH degradation
to be
about 80% faster than that of NADH at 30#{176}C
and at pH 2 to
4.5.
The importance of studying the stability of NADPH can
also be appreciated because of the wide clinical applications
involving NADPH. Information
on NADPH stability is
important to clinical chemists. Information of this kind may
also help in minimizing
degradation during storage and
assays.
Here we have investigated
the effects of some factors on
the stability and the kinetics of degradation of NADPH. In
most cases, NADH was included for the purpose of comparison.
Materials and Methods
Chemicals
NADP
(free
of EDTA),
(G6P), glucose-6-phosphate
NADH, glucose-6-phosphate
dehydrogenase
(G6PDH; EC
Department of Pathology, University of Utah School of Medicine,
Salt Lake City, UT 84132.
Received July 5, 1985; accepted November 14, 1985.
314 CLINICALCHEMISTRY,Vol.32, No.2, 1986
1.1.1.49), and 6-phosphogluconate
dehydrogenase
(EC
1.1.1.44) were purchased
from Sigma Chemical Co., St.
Louis, MO 63178. Diethylaminoethyl
(DEAE)-cellulose
WE
52) was obtained from Whatinan, Inc., Clifton, NJ 07014.
Sephadex 0-10 was from Pharmacia
Fine Chemicals, Inc.,
Piscataway, NJ 08854. All other chemical reagents were of
analytical grade. The water used for buffer preparation was
doubly distilled, de-ionized, and degassed by boiling.
Preparation of Salt-Free, Pure NADPH
To ensure purity (5), and because various anions(6) have
been shown to affect the kinetics of NADH degradation
reaction, we prepared an essentially salt-free NADPH for
the present study. Batches of NADPH were produced by
enzymatic reduction of 100 mg of NADH at 30#{176}C
in the
presence of 1000 U of G6PDH and excess glucose 6-phosphate in 10 mL of water. The pH was maintained at pH 8 by
continuously adding a 10 mmol/L solution of NaOH, dropwise. The reaction took about 30 min. At the end of the
reaction, the mixture was chilled toO #{176}C
and applied immediately to a 1.5 x 30 cm column containing about 10 g (dry
weight) of diethylaminoethyl-cellulose
(7). Ammonium bicarbonate (pH 7.8, 1 mmol/L) at 4#{176}C
was used for column
equilibration.
The column was eluted at 20 mLfh at 4#{176}C
with a linear gradient of 1 to 500 mmol/L ainmonium
bicarbonate, pH 7.8. Figure 1 shows a typical elution profile.
The chromatographic
procedure was similar to that described by Margolis et al. (8). The peak corresponding
to
NADPH was identified by the absorbance at 340 nm. Pooled
NADPH solution was immediately adjusted to pH 9.5 with 5
moL/L NH4OH and lyophilized.
Most of the buffer was
removed
during lyophilization. Any residual salt or buffer in
the NADPH fraction could be further removed by passing
aqueous NADPH solution through a column containing
Sephadex
G-10 (Figure 1) (9). The eluates containing
NADPH were immediately
lyophilized. Table 1 presents
data on the purity of a typical NADPH preparation.
The
ratio A]A
was equal to 2.33, which indicates that the
preparation is about 99% pure and free of any potent
enzyme inhibitors (5). The absorbance of our pure NADPH
at 340 nm was also completely removed by oxidation with
glutamic dehydrogenase
in the presence of excess a-ketoglutarate.
Buffer Preparations
To study the effect of pH and temperature
on NADPH
stability, we prepared buffers of various compositions, covering the pH range from 2 to 10 (Table 2). The pH was
carefully adjusted at the temperature
specified. To adjust
the ionic strength, the concentration of the various buffer
species was estimated
from the Henderson-Hasselbalch
equation: pH = pK + log (A)/(HA). The pH and plC values
are known and (A) + (HA) = 0.05 mol/L, so the concentrations of the ionized (A) and protonated
forms (HA) of the
weak acid can be calculated. Similar calculations may also
be applied to the weak bases. Knowing the concentrations of
the ionic species, we calculated the ionic strength (&) of the
solution from the equation L = 1/2 ICIZ2, where C is the
1NAOPII
concentration
and Z the charge. KNO3 was added to adjust
0.05. KNO3
dissociates completely in the solution, so the concentration
of KNO3 will be numerically
equal to the ionic strength. In
Table 2, the compounds listed second were the acids or bases
used to adjust the pH. Accuracy of the pH adjustment was
within ± 0.02 unit.
the final ionic strength if it was less than
CELLULOSE
DEAE4
3
z
2
I--
0
Monitoring Absorbance Change at 340 nm
Lu z
0-.
C,
I-J
4
U)
C,
z
150
4
200
250
Vt
V0
SALT
0
U)
G-io
<3
NADPM
j\ i
k1 = 2.303/t log (A)o/(A)
In equation 1, k1 is the first-order rate constant
A
2
40
60
80
100
ELUTION VOLUME (mL)
Fig. 1. Isolationof salt-free NADPH
NADPHwas Isolated,by chromatographyon diethylamlnoethyl.cellulose
fromthe
enz,me mention mbtiure and desaltedon a column of G-1O Sephadex. After
ctwamatography, NADPH and NADP were identifiedby their ateodences
at
340 nm and260 nm, repectWely
andglucose 6-phosphatewere
measured by using glucoee-6.phosphste dehydrogenase In the presence of
excess amount of glucose 6-phosphate (G6P) and NADP,
6Phoephogkconate (6PG) was quantitated by 6-phosphoglucnnste dehydrogensee Inthe presence of excess of NADP. , salt concentration Inmolt; A, A
nm; x, ft nm alter reactingwithgkioose-6-phcephatedehydrogenasein the
presence of excess of NADP*; 0, A, nm
To remove
sail, we used aof
1.5NADPH
x 30cm columnof
Saphadex G-1O.
Tableresidual
1. PreparatIon
by Enzymatic
Samplevolume was 2 mL The salt peak (ammonlumacetate or ammonlum
blcssbcnate)was Identifiedby mbdng2 mLof eluatefrom each lube wIth02 mL
of a 20 g/L r*ihyddn solution in acetone. The A
nm was read In a
epectrophofometer after the mbdure was Incubated in a boilingwater bath for 2
mm.t, A; 0, A
Reduction of NADP
-
-
Amount, gimol
(NADP)
Before reduction
108
Afterreduction
AfterDEAE-
103 (NADPH)
100
(NADPH)
Ratio
Recovery,
S’
(260/340 nm)
-
2.56
2.33
-
95
97
Yd,
S
100
95
92.6
cellulose
Refers to analytical recovery from the previous step.
Table 2.
ComposItIon of the Buffers Used for the pH
and Temperature StudIes
Buffer composition
Glycine HCI
Acetate-imidazole
Imidazole-acetic acid
Imidazole-acetic
Tris-acetic acid
acid
C3Iyclne(NaOH-adjusted)
pH
Buffer compositIon
Glyclne HCI
Acetate-Imidazole
Imldazole-acetic acid
Tris-acetlc acid
8.06 Tris-acetlc acid
9.11 Glycine (NaOH-adjusted)
2.33
3.91
6.26
7.99
To study the stability of the reduced pyridine nucleotides,
we monitored the absorbance change at 340 nm (2,6), using
either a Gilford or a Cary 14 spectrophotometer.
Cuvet
compartments
in both were temperature-controlled.
Buffer
was first pipetted into the cuvet and, after temperature
equilibration,
a small aliquot of NADPH stock solution was
added and measurement
immediately
started. The initial
concentration of NADPH was usually about 1 x iO mol/L
(absorbance =0.6). No pH change was detected within the
time frame of our experiments. The loss of absorbance at 340
nm followed first-order kinetics; therefore
pH
2.96
5.08
6.9
7.54
9.17
10.03
(1)
(min’), A is
the absorbance at 340 nm either at time zero, (A)0, or at
time t, (A). A plot of log (A) vs time was linear, the slope
being equal to -k1/2.303. On the other hand, at low temperatures and at neutral and alkaline pH, the NADPH and
NADH solutions were sealed in test tubes and incubated in
a temperature-controlled
water bath, because the absorbance changes were slow. Aliquots were withdrawn
at
appropriate time intervals for A
measurement.
At alkaline pH, weeks of incubation were needed to produce any
appreciable absorbance change at the 340 nm; therefore,
enzymatic assay with glutamate dehydrogenase
(EC 1.4.1.2)
had to be used to monitor the absorbance change. Turbidity
frequently developed after prolonged
incubation at increased temperature.
Effects of Phosphate and Acetate
In these experiments,
all reaction mixtures
contained the
same concentration
of reduced pyridine nucleotide and the
same 0.05 mol/L imidazole-acetate
buffer and were adjusted
to the same ionic strength by addition of either KC1 or NaCl.
For those at the same pH, the only variable was the
phosphate or acetate concentration.
Inasmuch as the concentrations
of phosphate and acetate
were in excess as compared with the concentration of
NADPH or NADH, and because plotting log (A) vs time
resulted in a straight line, the reaction fit pseudo-first-order
kinetics (2, 6):
k21’ = k11/(d)
=
2.303/(d)t
log (A)0/(A)
(2)
In this equation, k1’ is the pseudo-first-order
rate constant
(min_1), which is the rate constant actually obtained from
the plot; k2” is the second-order rate constant (min’
L’
mol), which can be calculated
from k1’; and (d) is the
concentration
of the active species of either phosphate or
acetate.
Monitoring the Spectral Change of NADPH
To determine
whether NADPH degrades similarly to
NADH (2,6) we decidedto monitor the difference-spectrom
change from 250 to 390 nm as well as absorbance changes
specifically at 340 nm and 280 nm. Knowledge
of the
spectral change during the course of NADPH degradation
CLINICALCHEMISTRY,Vol.32, No.2, 1986 315
allows one to assess the purity of the NADPH preparation
by monitoring the absorbance at proper wavelengths. For
example, the AIA290
ratio was used by Margolis et al. (8)
to assess the purity of NADH. During the measurement
of
the difference spectrum, NADPH was dissolved in pH 8
buffer in the blank cuvet, to avoid any change in the
absorption spectrum of NADPH. Fresh solutions of NADPH
were prepared frequently when we used extended incubation times.
Determination of Reaction Orders
According
rate
follows the following expression:
k1 (NADPH)
(H)’
=
Since the absorbance change of NADPH and NADH at 340
nm follows first-order kinetics, na = 1. Therefore
rate = k1’ (NADPH)
where k1’ = k1 (Hb
and
log k1’ = log k1 + nb log (He)
or
log k1’ = log k1
nb pH
-
Here k1’ is the rate constant obtained by plotting log (A) vs
time. By plotting log k1’ vs pH, the reaction order with
respect to hydronium
ion (nb) can be estimated from the
slope (Figure 2).
To study the effect of phosphate or acetate on the stability
1.0
expression
can be used at
fixed pH condition:
rate
=
k1 (NADPHYIa (phosphate
or acetate)nc
When the molar concentration of phosphate or acetate
much higher than the concentration of NADPH, then
or
and
is
rate = k1” (NADPH)
k1” = k1 (phosphate or acetate)nc
log k1” = log k1 + nc log (phosphate or acetate)
In this
to the rate law, the rate of NADPH degrada-
tion or the loss of A
of NADPH and NADH, a similar
case
k1” is the rate constant
for the NADPH
obtained jn equation 2 and was used in Figure
3. By plotting log k1” vs log (phosphate or acetate) the
reaction order (nc) with respect to either phosphate or
acetate can be obtained (Figure 3).
degradation
Results
Effects of pH and Temperature on the Stability of
NADPH
The stability
of NADPH was studied over the pH range 219, 30, and 41 ‘C. The logarithms
of the apparent first-order rate constants were plotted vs pH
(Figure 2). The degradation
of NADH at 41#{176}C
was studied
for purposes of comparison. All the plots were linear between pH 3 and 7.5. Above pH 7.5, the rate of degradation
became much slower.
In harmony
with previous reports for NADH (2, 6),
degradation
of NADPH
was also pH- and temperaturedependent. Both reduced forms of pyridine nucleotides were
less stable at low pH and high temperatures.
However, the
NADPH and NADH reaction orders (slopes, nb) were different for NADPH and NADH. That for NADH was close to 1,
while that for NADPH was 0.59.
10 at three temperatures:
ll0-’
10
-. -
p144.13
C
!l110-2
U)
Ox
1.10
_-
I-
z
U)
5.2
z
0
(_)
p145.2
.10-
-
#{149}pH6.27
II
LLI
1.10
I6 l0
l.IO
1,10-’
001
TOTAL
0.1
PHOSPHATE
(130 I/L
I
CONC,
0.15
TOTAL
ACETATE CONC
(1301 /L)
FIg. 3. Effect of phosphate and acetate on the rate
NADPH at 41 ‘C
pH
Ionic strength
Slope
Atate
2
4
6
8
10
Symbol
4.13
1.5
0.5
A
5.2
2
0.5
0.5
I
6.27
Phosplate
pH
of degradation of
3
FIg.2. Effect ofpH andtemperature on thestabIlity of NADPH
The degradation of NADH (-0-)
at 41 ‘C is Induded, for comparison.
NADPH at 41 ‘C; -0--, NADPH at 30 ‘C; -A-, NADPH at 19#{176}C
316 CLINICALCHEMISTRY,Vol.32, No.2, 1986
-I-,
The Ionic strength for each pH was adjusted to be the same by adding eIther KCI
(darksolidsymbols)or NaCI(open symbols)
Effect of Phosphate and Acetate
We have confirmed the report by Winer and Schwert (10)
that phosphate
accelerates
the degradation
of NADH. In
addition, we have shown the same effect occurs with
NADPH. Figure 3 also shows that both phosphate and
acetate increased
the degradation rate of NADPH and that
these ratios were concentration dependent. The reaction
orders with respectto phosphate and acetate did not change,
even though the concentration of the different ionized forms
of the weak acids changed so much at different pH condi-
tions.
The effectsof acetate, in general, were very similar to that
shown in Figure 3 for phosphate. The only differenceis at
pH 4.13. An increased concentrationof CH3COOH evidently
inhibited the degradation of reduced pyridine nucleotides.
The scatter of rate constants for NADPH at pH 7.5 was
found to be due to the different salts used to adjust the ionic
strength. Buffer containing KC1 increased the degradation
rate as compared
with buffer containing NaC1. This difference could only be observed at pH 7.5, where the rate was
relatively slow.
Effect of Ionic Strength
For most of our studies, we adjusted all solutions to the
same ionic strength.
Addition of salt (NaCl or KC1) to
buffered NADPH solution reduced the rate of NADPH
degradation (Table 3), in contrast to the accelerating effect
of acetate
and phosphate (weak acid). This salt effect was
also pH dependent. At pH 4.13, salt actually increased
NADPH degradation.
Effect of NADPH Concentration
A possible explanation of the reaction orders for NADPH
being close to 0.5 would be dinierization of NADPH. To test
this possibility, we measured
the rate of NADPH degradation at four different initial NADPH concentrations,
ranging from 1 x io to9 x 10 moJJL. The dirner concentration would be expected to vary with the concentration of
NADPH, and so the rate constants should also be different.
However, our results show no evidence of NADPH dimerization, because the rate constants did not change with the
change of initial NADPH concentrations.
Half-lives of NADH and NADPH Stabilities
A quantitative
appreciation of the effects of various conditions on the rate of NADPH degradation can be gained from
the rate constants and half4ives listed in Table 4. It would,
for example, take more than 8 h for half of the NADPH to
disappear
at 19#{176}C
but only about 1 h at 41 #{176}C.
Therefore,
one should be aware of the lability of NADPH at the
temperature
(37 #{176}C)
at which most clinical enzyme assays
are now performed.
As Table 4 shows, the pH has the greatest effect on
NADPH stability. This rapid loss of NADPH in acid media
suggests that NADPH should never be exposed to a pH
below 7.4. If a stock solution of NADPH is prepared, the
solution should be maintained
at a slightly alkaline pH (pH
8). Preparation of NADPH solution with pure distilled
water is not recommended, because the pH of distilled water
is usually 5 to 6.
Phosphate and acetate both accelerate the degradation of
NADPH. However, in practice, one seldom uses buffer
containing
phosphate or acetate at concentrations higher
than 0.1 molfL. In Table 5 we compare the degradation rates
of NADPH and NADH under similar conditions. NADPH
has a shorter half-life (i.e., faster degradation
rates) under
most conditions.
Difference Spectrophotometric Measurement
Loss of the 340-nm absorption peak with concomitant
formation of a new peak at 290 nm has been reported by
Burton and Kaplan (1) for NADH and its model compounds.
Alivasatos
et al. (2), on the other hand, found that absorbance at 280 nm increased at pH 6.62 as NADH was
gradually degraded.
In our study, the absorbance
change at both 340 am and
280 am of NADPH were monitored over the pH range 2 to9;
the results are shown in Figure 4. Similar to NADH, we also
observed a concomitant
increase
at 280 nm as the abeorbance of NADPH at 340 am declined. The 280-nm abeorbance was also unstable.
As shown in Figure 4, when
incubation was continued, the 280-nm absorption also decreased, to a degree depending on the pH of the solution.
These results are very similar to those reported for NADH
(6).
Table 4. Summary of the Rates of NADPH Degradation
under Various Conditions
Expsdmsntal cenditlons
AnIon
concn
Ionic
strength
Temp.
pH
Temperature
6.0
0
mel/L
‘C
Half-Ufs,
I mln
mm
525
141
55
0.05
19
1.32 x iO
6.0
0
0.05
30
4.9
6.0
0
0.05
41
3.0
0
0.05
30
5.0
7.0
0
0.05
30
0.05
30
0.05
30
2.53 x lO1.82 x 10_2
1.34 x iO
2.75 x i0
x iO
1.26 x 10_2
pH
9.0
o
0
2.7
38
517
25 200
Phosphate
Table 3. Effect of Salt on the Apparent First-Order Rate
Constant of NADPH Degradation
ploufta
A
pH
4.13
5.2
6.27
Salt added5
B
min’
1.4 x 10
3.1 x 10_2
7.9 x 1o
2
x10
2.1 x 10_2
lenlc
strength
1.8 x i0
1.08
0.4
3
41
4.97 x 1O
3
41
2.3
7.5
7.5
7.5
0.16
0
0
3
3
41
0.05
41
1.59 x i03
4.06 x 10
1.8 x iO
2.34
1.34
3
3
41
41
0.335
3
41
41
x i0
140
301
435
1700
385
Acetate
1.5
0.7
6.27
2
1.5
6.27
4.2 x io3
1.9
4.4 x iO3
4.1
‘NADPH solutioncontained only 0.05 mol of ammonlum acetate per liter.
or NaClwas added to solution A to adjust the Ionic strength.
7.5
7.5
7.5
6.27
6.27
1.45 x 102
1.19 x 10_2
8.02 x iO
5.39 x
7.9 x
41
0
3
6.27
0
0.05
41
a k Is the measuredapparent first-orderrate constant.
47
58
86
129
88
CUNICAL CHEMISTRY, Vol. 32, No.2, 1986
317
A
0.S
0.S
0.25
0.4
0.1
0.I
0.05
tOo
too
0.2
-
00
300
Thw)
-
-0.2
E
C
0
N)
7.54
pH
J
pH 9.11
as
0.4
0.8
329 1w.
/
0.4
151w.
02
0.2
5.5 1w.
“Th2%1w.
Tins (iurlnut.)
1533 Iw\
00
0.2
0.0
1t7
r.
-0.2
350
250
2
350
Wavelength (nn
a
7.12
TIme(hour)
Flg4.AandBChabscthancosat340nm(-)and28onm(--)
of NADPH during Incubation at 30 ‘C under various pH condItions. C:
Disappearance of absotbance at 280 nm after NADPH was incubated for
a prolonged period at 30 ‘C at three dilterent pHs
We also measured
the difference spectra of NADPH
degradation,
to determine why some investigators
have reported the increase at 290 am instead of the absorbance at 280 am for NADH at low pH, and whether a
similar sequence of reactions occur with NADPH. Figure 5
shows that as the abeorbance at 340 am declined there was
an initial increase
in absorbance at 280 am. However, soon
thereafter this 280 am peak also decreased, with a shift in
the absorption peak to 285 am and ultimately to 300 am at
pH 2.96. At acid pH, the changes in absorbance are very
rapid, so that early investigators
might have missed the
appearance of that at 280 am and observed only the increase
at 290 nm. At pH 9.11, the change in the spectrumwas
so
slow that we observed an additional decrease at 260 am
before the increase at 280 am.
during
Discussion
Of the various factors and variables we evaluated, temperature and pH had the greatest impact on the rate of
degradation
of both NADPH and NADH. As shown in
Figure 2, for each 10#{176}C
increase in temperature
the rate
constant for NADPH increased about 2.7 times. For every
pH unit increase between pH 3 to 7.5, there was a corre-
318 CLINICALCHEMISTRY,Vol.32, No.2, 1986
FIg. 5. lime-dependent change of differencespectra of NADPH
incubated at 30 ‘C at four different pHs
sponding fivefold decrease in the rate constants. Below pH 3
and above pH 8, the rate of NADPH degradation
is either
too rapid or too slow to allow accurate measurement.
The
deviation of the rate constants from a straight line above pH
8, and the observation of a different time-dependent spectral
change at alkaline pH, support the recent report by Bernofsky and Wanda (11) that different degradation mechanisms are involved with NADH and NADPH at slightly
alkaline pH from that at acid pH.
Our study not only confirms the finding by Lowry et al. (3)
that the rate of degradation is much faster in NADPH than
in NADH (Figure 2, and Table 5), we also undertook a more
systematic study and included the effects of other parame-
Table 5. Rate of Degradation (Hail-Life) of NADH and
NADPH Compared
IonIc
pH
strength
4
5
0.05
0.05
6
0.05
5.2
5.2
5.2
2
2
2
5.2
5.2
5.2
5.2
2
2
2
-
Acstts or
cencn,
mmel/L
Hall-Ills, mln
NADH
5.5
0
0
0
43
400
NADPH
4
15
56
Phosphate
1.6
0.8
0.2
Acetate
2.66
1.35
0.9
0.04
0
Data extracted from results presented
in Figure
3.
6.6
5.3
9.6
30
7.4
17
7
11
17
6
8.5
11
77
21
ters. The degradation rates of NADH and NADPH at the
same temperature
(e.g. 41#{176}C)
(Figure 3) differ by a factor of
five at pH 6, but are practically the same at pH 3.7. This pHdependent change results from the different reaction order
with respect to hydronium
ion between
NADH and
NADPH. The effect of phosph4e
and acetate on NADPH
degradation
rate was also investigated
at three different
pHs (Figure 3). The effect of acetate was very similar to that
of phosphate. The acceleration of degradation of NADH by
other anions, such as H2AsO4,
and HSO,
has been
reported previously (7). We found that the degradationrate
of NADPH is only minimsilly affected by phosphate
or
acetate at concentrations
of <100 mmolJL, the increased
rate of degradation by phosphate and acetate being offset by
the protective effect of increasing the ionic strength.
Burton and Kaplan (1,12) have studied the acid degradation of NADH and many of its analogs. From a plot of the
logarithm of the initial degradation
rate vs pH, they obtained a slope of 0.94 for NADH and 0.81 for the reduced 3acetylpyridine-NAD,
but only about 0.5 for NAD-dihydroxyacetone and NAD-acetone.
Their interpretation
was
that only one proton was required per molecule of either
NAD or 3-acetylpyridine-NAD
for the degradation reaction, whereas two protons were required per molecule of
NAD-dihydroxyacetone
and NAD4-acetone. Our results
not only confirm the earlier report made by Burton and
Kaplan (12) that the reaction order with respect to hydromam for NADH degradation
is close to 1, they also suggest
that it may require two hydronium ions, two acetate, and
two phosphate species to participate
in the degradation of
NADPH.
The results of our studies on the disappearance
of the 340
am absorbance peak and the increased absorbance at 280
am suggest that a serial reaction similar to that reported for
NADH (2,6) also occurs with NADPH. Therefore,the purity
of NADPH should also be assessed by measuring
both the
abeorbance ratio, Asso/A
and AssoIAo,
as suggested by
Margolis
et al. for NADH (8). These ratios will reflect any
degradation products formed from NADPH degradation. For
a pure NADPH preparation
these ratios should be near 2.3
and greater than 10, respectively.
We have also found that
products may be removed or isolated by chromatography on diethylaminoethyl-cellulose.
These degradation products may be better isolated by high-pressure
liquid
chromatography
(9, 13).
degradation
References
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