`
`Kinetics
`
`of Morphine in Aqueous Solution I11
`of Morphine Degradation in Aqueous Solution
`
`By SHU-YUAN YEH and JOHN I,. LACH
`
`The degradation of morphine in aqueous solution is dependent on the pH of the
`solution and on the presence of atmospheric oxygen in the system. The overall
`reaction rate, in systems containing excess oxygen, was found to be equal to
`Ka
`
`H + )] (Morphine)
`
`Degradation mechanisms for morphine, based on kinetic data and previously re-
`ported data on naphthol oxidation, are presented.
`
`M sedative since its isolation in 1805. Due
`
`ORPHINE has been used as an analgesic and
`
`to the limited solubility of morphine base, the
`acid salts, chiefly as the sulfate and hydrochlo-
`ride, have been used extensively in various
`pharmaceutical preparations.
`The stability of morphine in aqueous solution
`has been studied by many investigators since
`such solutions, after prolonged storage, undergo
`decomposition, as evidenced by discoloration.
`This decomposition of morphine is believed to be
`due to an oxidation reaction resulting in the for-
`mation of pseudomorphine (oxymorphine) and
`morphine N-oxide in the ratio of 9 : 1, together
`with a trace of a base believed to be methylamine
`(1). The oxidation of morphine and subsequent
`condensation
`to
`the dimer pseudomorphine
`is assumed to involve the phenolic group, as in
`the oxidation of naphthols to dimolecular com-
`pounds. Morphine derivatives not possessing
`the free phenolic group, as in the case of codeine
`and diacetyl morphine, do not undergo this type
`reaction (2). The oxidation of morphine is
`catalyzed by oxygen of air (3), sunlight (4),
`ultraviolet irradiation
`( 5 ) , iron and organic
`impurities (6), rat liver slices ( 7 ) , tissue ho-
`mogenates (8), and cytochrome (9). Ionescu-
`Matin, et ul. (lo), claimed that the deterioration
`of morphine in presence of oxygen takes place
`through condensation of morphine a t the phe-
`nolic hydroxy group with the formation of the
`dimer, pseudomorphine. However, in the ab-
`sence of oxygen, and in the presence of light,
`they also claimed that the deterioration is due to
`peroxidation or dimerization which can take place
`through the oxygen of the hydroxy group which
`was thought to be activated by ultraviolet light,
`
`resulting in the formation of bimorphine. Abood
`and Kun (8) reported that in the course of oxi-
`dation of morphine by tissue, one mole of mor-
`phine utilizes one-half mole of oxygen and that
`the phenolic hydroxy group is oxidized to a
`quinone. Thorn and Agren (11) reported that
`the pseudomorphine formed in this oxidation
`was extremely stable and did not undergo further
`decomposition. However, Balls (12) pointed
`out that pseudomorphine was quite unstable,
`that it decomposes on either oxidation or re-
`duction, and that in alkaline solution pseudomor-
`phine gradually decomposed to higher oxidized
`products.
`The stability of morphine in aqueous solution
`is largely dependent on the hydrogen ion con-
`centration.
`In alkaline or neutral solution,
`morphine deteriorates rapidly at room tempera-
`ture, whereas acidic solutions are relatively
`stable (11, 13, 14). The effect of temperature
`on the stability of morphine has been reported
`to be less important than the hydrogen ion con-
`centration (14).
`Although the stability of morphine in aqueous
`solution has been extensively investigated, no
`quantitative studies have been conducted This
`report deals with a kinetic study of the degrada-
`tion of morphine in aqueous solution, the degrada-
`tion of these solutions carried out in a light proof
`oven. No study was made to investigate the
`effect of light on morphine degradation.
`
`Received November 13, 1959, from the State University of
`Iowa. College of Pharmacy. Iowa City.
`Abstracted in part from a dissertation presented to the
`Graduate School of the State University of Iowa by Shu-
`Yuan Yeh in partial fulfillment of the requirements for the
`degree of Doctor of Philosophy.
`
`EXPERIMENTAL
`Reagents and Apparatus
`Morphine sulfate U. S. P. recrystallized from
`alcoholic aqueous solution and dried under vacuum
`for six hours, m. p. 250"; Beckman spectrophotom-
`eter, model DU, equipped with photomultiplier, 1-
`cm. silica cells. The buffers used in this study are:
`acetate buffer: 0.2 M and 0.4 M at pH 3.0, 4.5.
`5.0, and 5.5; phosphate buffer: 0.2 M and 0.4 M
`at pH 4.0, 5.0, 5.5, 6.0, 6.5, and 7.0. Phosphate
`35
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`36
`buffer, 0.2 M , at pH 2.0 and 2.5 were madc by
`adjusting 0.1933 M phosphoric acid with mono-
`potassium phosphate. All buffer solutions were
`adjusted using a Reckman pH meter model H2.
`The phosphate buffer solutions used here are inde-
`pendent of the temperature (151, and the change of
`pH of acetate buffer a t higher temperature is
`tiegligible (16).
`Degradation of Morphine
`Sulfate in Sealed Ampuls
`Effect of pH on the Degradation of Morphine
`Sulfate Solution Sealed Under Atmosphere at 95".-
`Five milliliters of 0.3y0 morphine sulfate solution in
`0.2 M phosphate buffer; pH 2.0, 6.0, 6.5, and 7.0,
`was introduced into 5-ml. ampuls, sealed under
`atmosphere, and stored in an oven a t 95". These
`solutions were assayed for morphine content a t
`various time intervals by a previously described
`chromatographic procedure (17).
`The results of the study are shown in Fig. 1.
`The
`data obtained indicate that the rate and extent of
`decomposition of morphine is dependent on the pH
`of the solution. I t is also apparent that after a
`time interval,
`this decomposition is halted, as
`evidenced by the plateaus. This behavior was
`probably due to the lack of oxygen in the systems.
`Calculation of the atmospheric oxygen present in
`the solutions contained in the ampuls (including the
`MIL.
`void space) was approximately 3.2 X
`Since a t pH 7.0 approximately 3.1 X lo-' M/L. of
`morphine had undergone decomposition, it appeared
`that morphine and oxygen react on a mole to mole
`basis.
`
`Journal of Pharmaceutical Sciences
`
`Y C
`
`,
`
`,
`
`,
`
`,
`
`,
`
`,
`
`,
`
`,
`
`,
`
`,
`
`,
`
`F
`
`Fig. 2.-Effect of inner gas on degradation of mor-
`phine sulfate in deionized water, a t 95".
`
`Time 10 +our;
`
`0 42
`
`0 3G
`
`\
`
`3
`E 0 3 0
`I0
`a, fu
`
`.2 0 0 24
`"
`c 0 n
`2 006
`a
`
`12
`
`0 04
`
`0 02
`
`T- ----
`
`I
`
`; I
`
`h
`
`v
`
`0 oc
`
`Time in Hours
`of sodium bisulfite, sulfur dioxide
`Fig. 3.-Effect
`on degradation of morphine sulfate a t 95": 1 and 2
`contained 1% NaHSO3, sealed under atmosphere;
`3, MzSOa in deionized HzO, sealed under SO2; 1,
`assayed by direct spectrophotometric method: 2
`and 3, by the chromatographic method.
`
`I
`0
`
`.
`
`.
`4 0
`
`,
`
`.
`80
`
`,
`
`
`i20
`160
`Tsme 8n H o u r i
`of pH on degradation of mor-
`Fig. 1.-Effect
`phine sulfate in 0.2 M phosphate buffer, sealed under
`atmosphere, at 95".
`
`200
`
`740
`
`28"
`
`I
`
`In order to verify the oxygen dependency of the
`reaction,
`the following study was undertaken.
`Five milliliters of 0.3yo morphine sulfate solution
`was introduced into 5-ml. ampuls and sealed under
`sulfur dioxide, atmosphere, commercial nitrogen,
`and absolute nitrogen. An additional solution of
`morphine sulfate containing l(yb of sodium bisulfite
`was prepared and sealed under atmosphere. Using
`sulfur dioxide and commercial nitrogen, the ampuls
`were bubbled for one minute, three minutes for
`absolute nitrogen, prior to sealing. These ampuls
`were stored in a 95" oven and assayed for morphine
`at periodic time intervals. Results of this study are
`shown in Figs. 2 and 3. The data indicated that
`the rate of degradation of morphine is oxpgen-
`dependent since solutions sealed under absolute
`
`nitrogen showed no evidence of decomposition.
`The slight degradation noticed in systems saturated
`with commercial nitrogen was probably due to im-
`purities in the nitrogen. Absolute nitrogen was
`prepared by passing commercial nitrogen through
`three wash bottles containing Fieser's solution (18)
`and through a saturated solution of lead acetate.
`I t is interesting to point out here that morphine
`solutions underwent a color change to yellow when
`sulfur dioxide was introduced. This color intensity
`was dependent on the length of time used to saturate
`the solution with sulfur dioxide. Chromatographic
`analysis of
`these freshly prepared sulfur dioxide-
`saturated solutions gave low recoveries for morphine.
`Direct spectrophotometric analysis, however, of
`these same samples, using deionized water as the
`diluent, gave only slight absorbance difference when
`compared to the original solutions indicating that
`some reaction had probably taken place These
`sulfur dioxide solutions become highly colored after
`three to four days' storage a t 95' and developed a
`yellow precipitate after one month.
`.4n investiga-
`tion of this phenomenon is, a t present, continuing.
`I t was also noted
`that the chromatographic
`
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`
`Vol. 50, No. 1, January 1961
`
`analysis of the morphine sulfate solution containing
`1% sodium bisulfite prior to sealing and storage a t
`95" resulted in low recovery of morphine (707,).
`Direct spectrophotometric analysis of the freshly
`prepared solution, using deionized water as the
`diluent gave, on the other hand, higher values for
`morphine. Further study dealing with this aspect
`showed there was no appreciable change in pH of the
`solution, although the peak of the spectrogram was
`shifted slightly to the lower wavelength from 286 to
`284 mp and a stronger absorbance was noted, indi-
`cating that some reaction had taken place between
`morphine and sodium bisulfite. This phenomenon
`will be discussed in a future comniunication (19).
`This initial
`investigation confirmed previous
`studies and showed that the decomposition of
`morphine sulfate in aqueous solution was dependent
`on the hydrogen ion and oxygen concentration of the
`system. Attempts to employ a manometric pro-
`cedure using a Warburg manometer in the study of
`the oxygen dependency regarding the rate of degra-
`dation of morphine were unsatisfactory
`in that
`morphine sulfate, in 0.2 M phosphate buffer pH 7.0
`at 60°, underwent only slight color change after
`twelve hours, the amount of oxygen uptake being
`negligible. However, it was felt that valuable in-
`fcrmation with respect to the kinetics of this reac-
`tion could be obtained by maintaining a relatively
`constant oxygen concentration in the system. To
`this end, the rate of degradation of morphine was
`studied as a function of hydrogen ion concentration,
`molarity of buffers used, ionic strength, and concen-
`tration of morphine in oxygen-saturated systems.
`
`Degradation of Morphine Sulfate in the
`Presence of Excess Oxygen
`Effect of pH.-Approximately
`0.15-Gm. portions
`of anhydrous morphine sulfate were accurately
`weighed, transferred to 100-ml. vaccine bottles, and
`dissolved in 50 nil. of the following buffers: 0.2 M
`phosphate buffer, pH 2.5, 6.0, 6.5, and 7.0; 0.2 A!!
`acetate buffer, pH 4 0, 4.5, 5.0, and 5.5. Three
`milliliters of each solution was withdrawn and
`assayed chromatographically for the original mor-
`phine concentration. These solutions were bubbled
`with oxygen for two minutes, rubber stoppered,
`sealed with aluminum caps, and placed in a 95"
`oven. At various time intervals the bottles were
`removed, chilled, and 3-ml. aliquots of solution
`were withdrawn. After each sample removal, the
`remaining solution was again saturated with oxygen
`for two minutes, and stored in the oven The proc-
`ess was repeated for a total of ninety-six hours.
`The data obtained, as shown in Fig. 4, indicated
`that the rate of degradation of morphine was
`hydrogen ion-dependent. A plot of the log of the
`concentration of undecomposed morphine against
`time gives a straight line, indicating that the reac-
`tion is pseudo first order with respect to morphine a t
`constant hydrogen ion and oxygen concentration.
`A plot of the log of the specific rate constant of these
`reactions as a function of pH resulted in an "S"-
`shaped curve (Fig. 5). This plot resembles a dis-
`sociation curve and indicates that the rate of degra-
`dation of morphine is dependent on the type of
`morphine species present in the solution. Data
`obtained from studies conducted in 0.2 M phosphate
`buffer a t pH 3.0, 4.0, 4.5, 5.0, 5.5, and 0.4 M phos-
`
`37
`
`-2 0 5
`
`-2 15
`
`-2 25.
`
`-8
`
`3
`
`0 p - 2 3 5 .
`aJ " c s
`a s
`4 -2 55
`
`Iu
`5 - 2 4 5 -
`r
`
`0
`
`-2 6 5 .
`
`\
`
`
`
`80
`
`100
`
`0
`
`2 0
`
`,
`60
`40
`Time m Hours
`of pHs on degradation of morphine
`Fig. 4.-Effect
`sulfate in excess oxygen, at 95".
`
`phate buffei a t pH 5.0 were unsatisfactory in that
`the capacity of these buffers was inadequate.
`Effect of Buffer Molarity and Ionic Strength.-
`Data obtained in this study using 0.2 M and 0.4 M
`acetate buffer at pH 5.0 and 0.2 M acetate buffer at
`pH 5 containing 1 and 370 sodium sulfate indicate
`that the rate of degradation of morphine is inde-
`pendent of the molarity of buffer and of the ionic
`strength present, as shown in Figs 6 and 7.
`Effect of Morphine Concentration.-Although
`data already showed that the overall mte of degrada-
`tion of morphine a t c o n s t a t hydrogen ion and
`oxygen concentration is pseudo first order with
`respect to the concentration of morphine present,
`this first-order reaction was further verified by a
`study of 0.3 and 0.15% morphine sulfate in 0.2 M
`acetate buffer, pH 5 0, as shown in Fig, 8.
`Effect of Temperature.So1utions of morphine
`sulfate in 0.2 M phosphate buffer, pH 6.0, 6.5,
`0 2 M acetate buffer, pH 5.0, were subjected to
`degradation a t 85, 90, and 95'. Results of this in-
`vestigation are shown in Figs. 9, 10, and 11. By
`plotting the log of specific rate constant, k, against
`the reciprocal of temperature, T, as shown in Fig. 12,
`the apparent energy of activation, E,, under the
`condition employed in this study, was calculated
`from the slopes of these lines and was found to be
`of the order of 22 8 Kcal.
`
`DISCUSSION
`Oxidation of Morphine.-As
`has already been
`pointed out, morphine undergoes decomposition
`resulting in the discoloration of the solution and the
`formation of precipitates. Our present study indi-
`cates that this degradation of morphine is chiefly
`dependent on the pH of the solution and the pres-
`
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`38
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`Journal of Pharmaceutical Sciences
`
`- 2 50 i
`o
`
`)O
`
`LO
`
`~c
`
`in
`
`60
`
`do
`T,me ,n Hours
`
`70
`
`BC
`
`90
`
`I
`100
`
`Fig. 8.-Effect of morphine concentration on deg-
`radation of morphine sulfate, in excess oxygen, a t
`95".
`
`-2 05
`
`- 2 15
`
`>
`\
`%
`.? - 2 2 5
`c
`
`c s -
`0 - e
`
`-235
`
`a) "
`-2 45
`
`U
`m
`0 i
`
`-2 55
`
`-2 65
`v
`
`10
`
`20
`
`50 60 70 80
`30 40
`Time ~n Hours
`Fig. 9.-Effect of temperature on degradation of
`morphine sulfate in 0.2 M phosphate buffer pH 6.0.
`in excess oxygen.
`
`90 100
`
`that the deterioration of morphine takes place
`through a condensation process a t the phenolic
`hydroxyl group with the formation of pseudomor-
`phine (10). Derivatives of morphine in which this
`phenolic group is alkylated, as in codeine, do not
`undergo this type of oxidation (2). The fact that
`deterioration of morphine increases rapidly in the
`presence of ultraviolet light (5) and decreases in a
`more acidic solution (11) gives additional support to
`the oxidation reaction. This oxidation of morphine
`to pseudomorphine may be somewhat analogous to
`the oxidation of phenols to dimolecular compounds
`and is presented here in some detail.
`The ferricyanide oxidation method commonly
`used to prepare pseudomorphine involves one elec-
`tron transfer which is common in free radical sys-
`This reaction can be illustrated as follows:
`tems.
`
`[Fe(CN)sJ a- -I- e - [Fe(CN)6l4-
`
`Pummerer, et al. (20-23), in their studies of the
`oxidation of naphthols with alkaline potassium
`ferricyanide reported that the oxidizing agent attacks
`
`2
`
`3
`
`5
`
`6
`
`PH
`Fig. 5.-Plot of the log of specific rate constants as a
`function of pH.
`
`1 , m e I" HOU.5
`Fig. &--Effect of molarity of acetate buffer a t pH
`5 on degradation of morphine sulfate, in excess oxy-
`gen, a t 95".
`
`0" - 2 20
`
`- 7 2 2
`
`0
`
`10
`
`20
`
`30
`
`4 0
`
`6 0
`
`5c
`I H ~ , , . ~
`
`7 0
`
`RC,
`
`9C.
`
`l@C
`
`effect of ionic strength on degrada-
`Fig. '?.-The
`tion of morphine sulfate in 0.2 M acetate buffer, a t
`pH 5.0 and in excess oxygen, a t 95".
`
`ence of atmospheric oxygen. It has been reported
`(1) that the degradation products of morphine are
`pseudomorphine and morphine N-oxide together
`with traces of the base said to be methylamine, and
`
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`Vol. 50, No. 1, January 1961
`-
`
`.
`
`'
`
`- 2 0 5 b
`
`" .
`
`'
`
`'
`
`'
`
`'
`
`39
`
`I
`
`\
`
`t
`
`07 O6
`
`Time In Hours
`Fig. 10.-Effect of temperature on degradation of
`morphine sulfate in 0.2 M phosphate buffer solution,
`pH 6.5, in excess oxygen.
`
`-3 0
`2 7 G 1 2
`
`3 4 5 6
`f x 1.000
`of the log specific rate constants
`Fig. l2.-Plot
`as a function of the reciprocal of the absolute tem-
`perature.
`
`7 8
`
`9 2 8 0 1 2
`
`-E -2 I 5
`
`5 - 2 2 0
`
`Time I" H o u r i
`Fig. 11.-Effect of temperature on degradation of
`morphine sulfate in 0.2 acetate buffer solution a t pH
`5.0, in excess oxygen.
`
`the OH group and has a direct dehydrogenating
`action and that the primary oxidation product is a
`resonant aroxyl radical. They confirmed this by
`isolating and identifying the possible dimeric forms.
`The oxidation of naphthol gives binaphthol which is
`further oxidized to oxy-binaphthylene-oxide and
`binaphthylene dioxide as illustrated in Diagram I.
`Oxidation of p-cresol was found to give dicresol
`and an ether (24, 25). If the resonant forms of p-
`cresol are examined, it becomes evident that these
`are all cogent in the transition of the resonant aroxyl
`radical as illustrated in Diagram 11.
`Considering such a reaction and by direct analogy,
`morphine could undergo a similar type of process in
`its degradation and subsequent production of pseu-
`domorphine as shown in Diagram 111.
`Mechanism of Degradation of Morphine.-Since
`data obtained from this investigation showed that
`the rate of decomposition of morphine in solution
`was dependent on the presence of oxygen and that
`no decomposition occurred in systems void of oxy-
`
`gen, it was concluded that a free radical reaction was
`involved in this process. It was also found that
`the rate of degradation was dependent on tbe
`hydrogen ion concentration of the solution since the
`rate was considerably greater at the higher pH. It
`is interesting to note here that an '3'"-shaped curve,
`similar to the typical dissociation curve, was ob-
`tained when the log of the specsc rate constant of
`these reactions was plotted as a function of pH.
`This indicated that the degradation was dependent
`on the type of morphine species present in solution.
`Based on this information, it appeared that the un-
`dissociated morphine molecules undergo oxidation
`more easily. At the pH values employed in this
`study and considering the pKa and pKb of morphine,
`only the protonated and undissociated morphine
`species are involved. The amount of anionic
`species of morphine present at these hydrogen ion
`concentrations is negligible and has been included in
`the concentration of
`the undissociated species.
`Since the rate of decomposition of morphine at
`lower pH's, i. e., a t pH 2.5 and pH 4.0, is nearly the
`same, this indicated that the protonated morphine
`species also undergoes oxidation, but at a different
`rate to that of the free undissociated morphine base.
`The undissociated (free base form) and protonated
`morphine are both oxidized by atmospheric oxygen
`to give a semiquinone (MO) and a free radical per-
`oxide (Hot. ). This semiquinone is further trans-
`formed to a free radical quinone (MO.), which can
`( b ) the un-
`undergo coupling with:
`( a ) itself,
`dissociated morphine, and (c) the protonated mor-
`phine. Since the amount of activated or free radical
`morphine species present in the system is small com-
`pared to the protonated or free base forms, inter-
`action or union of two such activated species is un-
`
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`40
`
`Interaction of this activated species with
`likely.
`the protonated or free base form of morphine is
`more probable resulting in the formation of the
`dimer pseudomorphine (PSM) with the simultaneous
`elimination of a hydrogen-free radical H . . This
`hydrogen-free radical can then react with the per-
`oxide-free radical (HO:!. ) to form hydrogen peroxide.
`The hydrogen peroxide which formed in such a
`process can react with morphine to form morphine
`N-oxide (MNO), or may decompose to give water
`and a free radical oxygen which can also react with
`morphine base to give the N-oxide. The postulated
`mechanism is given in Diagram IV.
`Derivation of the Rate Equation.-From
`the
`postulated mechanism and kinetic data obtained in
`this investigation the following rate equation may
`be derived
`M t = M + +HM = total concentration of morphine.
`M and +HM represent
`the undissociated and
`protonated form of morphine, respectively.
`MO and +HMO represent the activated forms of
`the undissociated and protonated morphine
`species, respectively.
`ki
`
`Journal of Pharmaceutical Sciences
`DIAGRAM I
`Oxidation of Naphthol (20-23)
`
`'
`
`e
`
`-
`
`g O H % \ OHHO '
`
`\
`
`/
`
`
`
`\
`
`/
`
`
`
`-
`
`Oxy-binaphthylene Oxide Binaphthylene Dioxide
`
`DIAGRAM I1
`Oxidation of p-Cresol (24, 25)
`
`M + 0 2 -+ MO. + HOz.
`+HM + 0 2 -* +HMO. + HO2.
`+HMO. + M -+ PSM + H .
`'HMO. + +HM -P PSM + H .
`MO. + M -+ PSM + H .
`MO- + 'HM -+ PSM + H
`H - + HOn.
`H202 + M -+ MNO + H20
`
`kz
`
`k3
`
`k4
`
`k5
`
`k6
`
`k g
`
`-+
`
`H202
`
`ka
`
`Application of the usual steady-state treatment
`for elimination of the unstable free radical and inter-
`mediates gives the rate.
`
`-d(Mt)/dt = 3ki(Oz)(M) + 3kz(Oz)(+HM)
`
`The dissociation equation of an ampholyte may
`be represented as follows
`
`+HM -). M + H f
`
`Mt = M + +HM
`
`Therefore, the overall rate can be written as
`
`+ "' (Ka
`
`[k" (K*+)
`
`Since oxygen was maintained in excess, this rate
`equation becomes pseudo first order.
`- d ( M i ) / d t =
`
`+ H + )] (Md
`H+
`When a I<a value of 1.7 X
`was chosen for
`morphine, the values of kl' and k2' were calculated
`per hour, and 1.77 X
`and found to be 8.03 X
`lop3 per hour. The success of this derived equation
`is shown in Table I. I t is interesting to note the
`close agreement of the values obtained using this
`overall rate equation as compared to the reported
`experimental data.
`
`CONCLUSION
`
`The degradation of morphine in aqueous solu-
`tion is dependent on the p H of the solution and
`on the presence of atmospheric oxygen in the
`
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`L'd
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`50, No. I , January 1961
`
`41
`
`DIAGRAM I11
`Formation of Pseudomorphine
`
`fj)
`
`- eH3
`eH3
`-
`&y*q5JL$H~
`*eH3
`e4
`
`5554
`
`0
`
`K3Fe(cN)ep
`
`0
`
`DIAGRAM IV
`Degradation of Morphine
`
`+ 0 2
`
`'
`
`0
`
`
`
`OH
`
`OH
`
`OH
`
`OH
`
`00
`
`OH
`
`0
`
`-I
`
`&H3- I 0
`
`HO;+
`
`'
`
`0
`
`
`
`0
`
`0
`
`OH
`
`0
`
`OH
`
`' 0
`
`OH
`
`0
`
`OH
`
`eH'j+
`
`0 1
`
`0
`
`0
`
`& +
`
`OH
`
`*&+
`
`OH
`
`OH
`
`0
`
`0 "
`
`0
`
`H* + HO, -
`
`OH
`
`H3
`
`+ Ha
`
`0
`
`0
`
`OH
`
`OH
`
`OH
`
`0
`
`OH
`
`OH
`
`OH
`
`OH
`
`H202
`
`OH
`
`OH
`
`I
`
`OH
`
`OH
`
`OH
`
`OH
`
`OH
`
`OH
`
`I1
`
`0
`1.
`
`OH
`
`OH
`
`AND CALCULATED SPECIFIC
`TABLE I.-oBSERVED
`RATE CONSTANTS AT VARIOUS pH's
`k X 102 hr.-',
`Calcd. by the
`Derived Eq.
`1.77
`1.90
`2.19
`3.08
`5.77
`13.17
`29.23
`51.20
`Mean dev.
`
`PH
`2.5
`4.0
`4.5
`5.0
`5.5
`6 . 0
`6.5
`7.0
`
`k X loa hr.?,
`Observed
`1.76
`1.95
`2.16
`2.99
`5.61
`13.78
`29.36
`48.70
`
`Deviation,
`% +0,5
`-2.6
`f 1 . 4
`+3.0
`+2.8
`-4.6
`-0.5
`+ 5 . 1
`0.62
`
`system. The overall reaction rate, in systems
`containing excess oxygen, was found to be
`equal to
`
`) 1
`(Ka + H +
`+
`[klf
`The apparent energy of activation, E,, was
`calculated to be 22.8 Kcal. for this reaction.
`
`(9, Ka + H + ) $-
`
`(Morphine)
`
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`
`
`
`42
`
`Journal of Pharmaceutical Sciences
`
`REFERENCES
`
`9 , \ I Y t L I .
`(6) Balan J. and Csere E., Chem. eoesti, 1, 407(1953).
`(7) Bernheim, H., and Bernheim, M. L. C., J . Phavmacol.
`Erptl. Therap., 81, 374(1944).
`(8) Abood, Is. G., and Kun, E . , Federation Pvoc., 18,
`207(19491.
`(91 Hosoya E. and Brody, T. M . , 3. Pharmacol. Erptl.
`Therap., 120, 5b5(i957).
`(10) Ionescu-Matin A. Papescu, A,, and Moncium, L.,
`Ann. Pharm. f r a n ~ . , 6 , 13’7(1948).
`(11) Thorn. N., and Agren, A., Svensk Farm. Tidskr., 55,
`61(1951).
`(12) Balls A. R. J . Biol. Chem. 71, 537(1927).
`(13) Zocccha, F., ’Giorn. farm. Lhirn. 67 60 90(1918).
`W.?‘ A;ch.‘ Pharnc., 266,
`(14) Dietzel, R., and Huss,
`641(1929).
`
`(15) Bower, V., and Bates, R. C., J . Research Natl. B w .
`Slandards 59 261(1957).,,
`(16) K&thh€, I. M., Acid-Base Indicators,” The Mac-
`millan Co. New York N. Y. p. 265.
`(17) Ye8, S. Y . , a n i Lach’ J. I,. to be published.
`(18) Fieser, L. F., Expe;imen<s in Organic Chemistry,”
`3rd rev. ed., D. C. Heath & Co., Boston, Mass., 1955. p. 299.
`(19) Yeh, S. Y., and Lach, J. L., to be published.
`(20) Pummerer, R., and Frankfurter, F., Ber., 47, 1472
`114141
`l-”--,. (21) Pummerer, R., and Cherbuliez, E., ibid., 47, 2957
`(1914).
`(22) Pummerer, R., and Cherbuliez, E., ibid., 52, 1392,
`1402, 1403, 1414(1919).
`(23) Pummerer, R . , and Fiankfurter, F., ibid., 52, 1416
`(1919).
`(24) Pummerer R . , Melamed, D., and Puttfarcken, H.,
`ibid., 55, 3116(19;2).
`(25) Pummerer, R.. Puttfarcken, H., and Schopdecher, P.,
`ibid.. 58, 1808(1925).
`
`Antifungal Activity of Some
`Dichloroacetaldehyde and
`
`Amides
`Bromal
`
`of
`
`By WILLIAM D. EASTERLY, Jr., and JAMES E. DUSENBERRY
`
`The fungus-inhibiting properties of a series of amides of dichloroacetaldehyde and
`bromal were studied. Three of the compounds showed antifungal activity when
`subjected to tests with Aspergihs niger and Tricboderma viride. Terminal chlorine
`atoms appeared to be a factor in enhancing the antifungal activity.
`
`R dichloroacetaldehyde (1) and bromal ( 2 )
`
`ECENTLY the syntheses of some amides of
`
`In view of the fact that studies
`were reported.
`have shown the potentialities of certain sub-
`stituted amides as antifungal agents and have
`indicated an increase in antifungal activity in
`certain compounds following halogenation (3, 4),
`it was decided to subject these compounds to
`preliminary tests.
`
`EXPERIMENTAL
`
`the fungus-inhibiting
`study of
`Procedure.-A
`properties of these compounds was carried out by a
`modification of a method used by Bateman ( 5 ) ,
`Vincent (B), and also by Leonard and Blackford ( 3 ) .
`This procedure consists in comparing the growth
`rates of the test fungus upon nutrient agar contain-
`ing known concentrations of the compound to be
`tested with that of a control, identically treated but
`containing none of the test compound. All com-
`pounds were tested in triplicate for each concentra-
`tion.
`Preparation of Culture Medium.-The
`culture
`medium, Sabouraud’s dextrose agar, was prepared
`
`Received July 18, 1960, from the University of Arkansas,
`School of Pharmacy, Little Rock.
`Accepted for publication August 4, 1960.
`Supported b y grants from the Upjohn Co. and the Sigma Xi
`Society.
`
`in the usual manner; and prior to being steam
`sterilized for twenty minutes a t 15 pounds pressure,
`a glass enclosed magnetic stirring bar was added to
`the flask containing it. When the flask was removed
`from the autoclave it was placed in a hemispherical
`heating mantle, which in turn was placed on a mag-
`netic stirring motor. A weighed amount of the test
`compound was added to the agitated culture medium
`which was maintained in the liquid state.
`Most of the amides of dichloroacetaldehyde, in the
`indicated concentrations, were readily soluble in the
`culture medium, while the amides of bromal were
`difficult to dissolve. The latter could be uniformly
`dispersed, however, by keeping the magnetic stir-
`ring bar moving rapidly, and gradually lowering the
`temperature of the heating mantle until the con-
`sistency of the culture medium was still suitable for
`pouring, yet keeping the test compound in suspen-
`sion. The culture medium was then poured into
`previously chilled, sterile 9-cm. Petri dishes.
`Inoculation.-The
`test organisnis to be used had
`been grown on Sabouraud’s agar slants for seven
`days at 30” Separate spore suspensions of each
`of the two test organisms were prepared by washing
`the agar slants with a 5-ml. portion of normal saline.
`As an inoculum, 0.03 ml. of this spore suspension
`w d S added to the center of the Petri dishes contain-
`ing test substance, and a like amount was added to
`the control.
`plates, incubated a t
`Growth Measurement.-The
`3 7 O , were observed daily and measurements of
`
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