. Chemical Stability
`,
`of Pharmaceuticals ‘
`.
`A Handbook for Pharmacists ,
`
`Second Edition
`
`Kenneth A. Connors
`
`.
`
`School of Pharmacy, The University of Wisconsin" _
`
`a A“ '
`Gordon L. Amidon’
`College of Pharmacy, The University of Michigan
`'35
`
`Valentino J. Stella 4
`School of Pharmacy, The University of Kansas
`
`A Wiley-Interscience Pnhlieation ‘
`
`JOHN WILEY & soNs
`New York 0 Chichester I Brisbane I Toronto 0‘ Singapore
`Noven EX. 1015
`
`Page 1 of 35
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`Page 1 of 35
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`Noven Ex. 1015
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`

`

`Copyright © 1986 by John Wiley & Sons, Inc.
`
`All rights reserved. Published simultaneously in Canada.
`
`Reproduction or translation of any part of this work
`beyond that permitted by Section 107 or 108 of the
`1976 United States Copyright Act without the permission
`·of the copyright owner is unlawful. Requests for
`permission or further information should be addressed to
`the Permissions Department, John Wiley & Sons, Inc.
`
`Library of Congress Cataloging in Publication Data:
`Connors, Kenneth A. (Kenneth Antonio), 1932-
`Chemical stability of pharmaceuticals.
`
`"A Wiley-lnterscience publication."
`Includes bibliographies find index.
`I. Drug stability.
`I. Amidon, Gordon L. II. Stella,
`I. Drug
`Valentino J., 1946-
`.
`III. Title.
`[DNLM:
`Stability-handbooks. 2. Kinetics-handbooks.
`QV 735 C752c]
`
`· RS424.C66 1986
`ISBN 0-471-87955-X
`
`615'.18
`
`85-31455
`
`Printed in the United States of America
`
`10 9 8 7 6 5 4 3 2 I
`
`Page 2 of 35
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`Noven Ex. 1015
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`

`

`CHAPTER 5
`Oxidation and Photolysis
`
`Oxidative and photochemical reactions are, for the
`most part, one-electron reactions as opposed to reac(cid:173)
`tions discussed in Chapter 4, which are two-electron
`reactiona. For the hydrolytic reactions in Chapter 4,
`a free pair of electrons on a heteroatom in one mole(cid:173)
`cule, a nucleophilic center, attacked an electrophilic
`center on a second molecule, whereas oxidative and
`photochem-ical reactions proceed through free radical
`or free-radical-like reaction pathways.
`Most drugs exist in a reduced form, so the presence
`of 20% oxygen in the atmosphere creates obvious poten(cid:173)
`tial stability problems for ~hese molecules. That is,
`many molecules tend to be converted to a more oxidized
`state. Kinetically, however; there is a sufficient
`energy barrier to ma,n y such reactions (the energy 0 f
`activation) that not all molecules are subject to
`measurable rates of spontaneous oxidation or autox(cid:173)
`idation. The radiation from
`the_ sun and artificial
`light, particularly visible and ultraviolet light, is
`also ubiquitous, so that molecules capable of rear(cid:173)
`ranging upon absorption of radiation ener-gy must be
`protected.
`Our overall mechanistic understanding of oxidative
`and photochemical reactions is poor.
`The reason for
`this will be understandable as this chapter proceeds.
`Simply stated, many oxidative and photochemical r•ac(cid:173)
`tions involve very complex reaction pathways with mul(cid:173)
`tiple intermediates so that even though the stoichi(cid:173)
`ometry of a reaction might be given by Eq. (5.1) the
`kinetic law is not as simple as Eq. (5.2).
`> ROOH
`
`( 5. 1)
`
`82
`
`I-
`
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`

`

`Oxidation V
`
`83
`
`dIROOHI
`dt
`.
`
`-
`? k[RH][02]
`
`'
`(5.2)
`
`rAlso, unlike two—electron reactions where catalysis is
`often limited to acid/base or nucleophilic catalysis,
`trace quantities of environmental agents can powerful—
`ly catalyze one—electron reactions.
`For example,
`trace contamination of metal
`ions can catalyze oxida—
`tive reactions by many orders of magnitude,
`and the
`presence of a photosensitizing agent can cause a-mole—
`cule that
`in the absence of
`the photosensitizing agent
`is not photolabile to undergo an apparent photochemi~
`cal reaction.
`I
`'
`‘
`introduce,
`In this chapter we
`from a basic view—
`point,
`the kinetics and other factors affecting oxida—
`tive and photochemical
`reactions and describe how
`these reactions can be prevented or at
`least
`inhibited.
`‘
`-
`
`A.
`
`OXIDATION
`
`1.
`
`'Nature of Oxidation
`
`,
`
`to real—
`When one considers oxidation, it is important
`ize that
`this reaction is a complementary one;
`its
`partner
`is reduction.
`.One cannot happen without
`the
`other. Oxidation/reduction (redox)
`reactions involve
`the transfer‘of one or more oxygen or hydrogen atoms
`or
`the transfer of electrons.
`The classical,
`and
`familiar,
`inorganic redox system can be described by
`Eq.
`(5.3), where e‘ represents an electron and n the
`
`reduced form -———4> oxidized + ne~
`.
`.
`<——"—
`‘
`
`,
`
`'
`
`(5.3)
`
`redox reactions are elec—
`Thus
`number of electrons.
`tron—transfer processes, and this aspect must be con—
`sidered if the basic process is to be understood.
`In the case of organic compounds and especially the
`oxidation state of carbon,
`the oxidation state is de—
`termined by the number of bonds from carbon to oxygen.
`For example,
`the state of oxidation of one—carbon com—
`pounds
`increases as shown in Eq.
`(5.4).
`As stated
`earlier,
`the mechanism of
`this process is not as sim—
`ple as suggested by the stoichiometry of
`the reaction.
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`

`84
`
`Oxidation and Photolysis
`
`oxidation
`
`H
`H
`I
`I
`H,
`~
`H-C-H ~ H-C-OH ~ C=O ~ H-C ~ O=C=O
`I
`I
`"'-oH
`W""'
`H
`H
`
`reduction
`Also the simple redox system illustrated by Eq. (5.3)
`is made more complex by the medium in which the reac(cid:173)
`tion occurs.
`For example,
`the oxidation of hydroqui(cid:173)
`none
`( 1, 4_-dihydroxybenzene) to its quinone (p-benzo(cid:173)
`quinone) is .often illustrated in the textbooks by
`
`(5.4)
`
`Eq. (5.5). HOVH
`
`quinone
`hydroquinone
`Yet in aqueous solution, free electrons, e-, do not
`exist and the state of ionization of the hydroquinone
`is affected by the solution pH. Therefore in aqueous
`solution the oxidation of hydroquinones is more accu(cid:173)
`rately described by Eq. (5.6).
`
`(5.5)
`
`As will be discussed later, the oxidation of hydro(cid:173)
`quinone and other phenols is even more complex than
`shown by Eq. (5.6)
`in that the product of the immedi(cid:173)
`ate oxidation, the quinone, can catalyze the oxidation
`
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`

`Oxidation
`
`85
`
`This process is called autocatalysis or
`further.
`product catalysis.
`
`2. Kinetics of Oxidation
`
`take place spontaneously under mild
`Oxidations that
`conditions are often called "autoxidation";
`the major—
`ity of
`these are free—radical
`reactions.
`Free radi—
`cals are chemical species that possess an unpaired
`electron.
`0xygen,,in its ground state,
`is a diradical
`with the electronic configuration
`
`0.0.:ch
`O. .0
`
`therefore, would "like" to fill its outer
`Oxygen,
`electron shell
`to produce 022“,
`the peroxy dianion.
`To do so, oxygen must accept
`two electrons from a
`donor molecule(s)
`and in so doing could in theory
`generate other free—radical molecules.
`In most autoxidation reactions, even though oxygen
`is often involved,
`the initiation of
`the oxidation re—
`action does not
`involve molecular oxygen,
`that
`is,
`oxygen itself in its ground state does not really ini—
`tiate oxidation reactions. Let us consider the gener—
`al kinetic behavior of olefin autoxidation described
`by Eqs.
`(5.7)—(5.9)
`M
`
`RI-
`
`;
`'
`k
`+ R"—CH2—CH=CH—R"' —l€>R"—CH—CH=CH—R"'I+ R'H (5.7)
`‘
`0-0”
`
`.
`R"—CH—CH=CH—R"'
`
`I
`k2
`+ 02 -—€>R“-CH—CH=CH-R"'
`
`.
`
`(5.8)
`
`0-0'
`
`'
`I
`R"—CH-CH=CH—R"'
`
`.
`.
`k3 .
`+ R"—CH2—CH=CH-R"' ——€>R"—CH—CH=CH-R"'
`
`+
`
`(5.9)
`
`o—oH
`l
`R"-CH—CH=CH~R"'
`
`a hydroperoxide
`
`(5.7) represents the initia-’
`The reaction given by Eq.
`tion reaction. Generally the species R" is not
`
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`

`86
`
`Oxidation and Photolysis
`
`oxygen but some other peroxy radical present in the
`solution, trace quantities of metal
`ions such as
`ferrous or cupric ions, or radicals formed
`in the
`solution from the absorption of light (visible or
`ultraviolet).
`The reactions given by Eqs. (5.8) and (5.9) repre(cid:173)
`sent chain propagation reactions, i.e., one radical
`produces one radical plus a hydroperoxide molecule.
`Radical species in solution, apart from reacting with
`oxygen or another unreacted molecule to produce anoth(cid:173)
`er free radical, can also react with each other to
`produce sta_ble or metastable products. This step is
`called a chain termination step and three examples 'are
`given in Eqs. (5.10)-(5.12).
`
`Ro 2 · + Ro 2 ·
`
`k4
`~
`
`k
`Ro 2 · + R" ~-
`k6
`~
`
`R.· + R·
`
`stable products
`
`(5.10)
`
`(5.11)
`
`(5.12)
`
`Considering this reaction mechan-ism, and simplifying
`Eqs. (5.7)-(5.9) to Eqs. (5.13)-(5.15),
`
`initiator
`
`k1
`----=--7 R·
`
`R· + 02
`
`k2
`---=-7 Ro 2 ·
`
`Ro 2 · + RH
`
`ks
`~ ROOH + R·
`
`(5.13)
`
`(5.14)
`
`( 5.15)
`
`assuming normal levels of oxygen, applying a steady(cid:173)
`state assumption to the radical species R02 · and R ·,
`letting k 2 [R"][0 2 ] = ks[Ro 2 ·][RH], and k4 = k 5 = k6 =
`kt, it can be shown that the rate of hydroperoxide
`formation is given by Eq. (5.16) where ri is the ini(cid:173)
`tiation rate.
`
`It is obvious.from this expression that the rate of
`hydroperoxide formation is proportional to the square
`root of the initiation rate,. r1. Also, if oxygen
`
`(5.16)
`
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`

`Oxidation
`
`87
`
`concentration is very high, k2[02]
`Eq.
`(5.16) collapses to Eq.
`(5.17):
`
`>> k3[RH];
`
`. Z
`
`'
`
`dt
`
`kt
`
`The reaction will apparently be first order in start~
`ing material, RH.= On
`the other hand,
`if k3[RH]
`>>
`k2[02],
`then Eq.
`(5.18) is realized.
`
`Wdflkzwz]
`
`1.x
`
`'
`
`-
`
`(5.1s)
`
`"pseudo
`these conditions the reaction will be
`Under
`zero order" with respect
`to RH,
`that is,
`[RH] does not
`appear
`in Eq.
`(5.18),
`and the reaction will be first
`order in [02].‘
`If\the term ri actually involves RH or
`02,
`the order of
`the reaction with respect
`to RH and
`02 could be as high as 1.5.
`(5.7)—(5.15) assumes
`The mechanism defined by Eqs.
`that
`the species ROOH is stable and that
`the termina—
`tion products are stable.
`In the case of olefins,
`the
`hydroperoxide can break down to produce volatile and
`nonvolatile products as well as multiple radicals.
`
`0-0H
`
`.
`I
`R"—CH—CH=CH—R""-%R"-CHO + R"'-CH=CH'
`
`+ :0H_
`
`(5,19)‘
`
`The rancidity of unsaturated cooking oils and oil—
`based paints is the result of
`this fragmentation of
`olefinic bonds
`to produce aldehydes, acids, and alco—
`hols as well as multiple radicals.
`As can be seen in
`Eq.
`(5.19); one hydroperoxide molecule produces two
`radicals.
`'If this reaction is favorable, not only do
`we have a chain propagation reaction, but chain
`branching reactions will be observed.
`If branching
`does occur,
`the oxidation kinetics and products become
`even more complex than those defined by the initia—
`tion, propagation, and termination sequence.
`Qualitatively,
`the kinetics of oxidative free—radi-
`cal reactions follow the pattern illustrated in Figure
`5.1.
`A characteristic of many such reactions is a lag
`time or
`lag phase corresponding to the gradual buildup»
`of radicals via the initiation step..
`If the radicals
`
`
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`

`88
`
`Oxidation and Photolysis
`
`100
`
`80
`
`co
`
`c :s 60
`"' E
`~
`co
`::>
`0 40
`*
`
`20
`
`0
`
`Time
`
`Illustration of percentage of drug re(cid:173)
`FIGURE 5.1.
`maining vs .. time for an oxidative free-radical reac(cid:173)
`tion:
`curve a initiation step only; curve b initia(cid:173)
`tion plus propagation; and curve c init~ation, propa(cid:173)
`gation, and chain branching. Arrows indicate the lag
`times.
`
`produced from the initiation step are stable then as
`soon as the catalytic species is consumed the reaction
`stops (curve a in Fig. 5.1).
`If the radicals produced
`from the initiator go into a propag_ation cycle, curve
`b results.
`The overall loss of drug will then often
`follow a first-order decay curve with respect to drug,
`depending on the oxygen dependency of the reaction.
`If chain branching occurs the overall loss of drug
`shows an acceleration phase (see curve c) with maximum
`acceleration occurring at -50% drug remaining.
`The reaction kinetics defined by Eqs. (5.7)-(5.19)
`were for a reaction in .which the reactant RH was not
`capable of ionization. For the oxidation of drugs or
`pharmaceutical additives, the kinetics are further
`complicated wh~n the state of ionizatibn of the mole(cid:173)
`cule is affected by solution pH.
`In the oxidation of
`sodium sulfite (or bisulfite) a first approximation of
`the oxidation mechanism is given by Eqs. (5.20)(cid:173)
`(5.27),
`
`Page 9 of 35
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`

`Oxidation
`
`89
`
`3037+ M+ ————4> 305'
`
`+ M'
`
`so;-
`
`+ 02
`
`:———:>‘sog'
`
`’
`
`305-
`_
`SO5
`
`-
`
`+ H803
`2-
`o
`S 3
`
`+
`
`2_
`o
`s 3
`
`_
`+ H805
`
`————e> Hso5
`2_‘
`05
`
`-———:> s
`<T____
`
`+ 303'
`_
`+ $03
`
`-
`
`(pH 5 7)
`
`<
`(pH g 7)
`
`————4> HSO
`
`_
`4 + 304
`
`_
`
`(pH 5 7)_
`
`303‘
`
`+ So§_
`
`-—-—4> 230%— (pH < 7)
`W“
`—
`
`_
`
`'
`
`_.
`$03
`
`'
`
`_
`+ so ' ——~—4>
`5
`<______
`
`2;
`s
`.206
`
`+ 02
`
`~
`
`(5.20)
`
`(5.21)
`
`(5.22)
`
`.
`(5 23)
`
`,
`‘
`(5 24)
`
`-(5 25)
`
`.
`(5 26)
`
`SOE'
`
`w—
`,
`+ inhibitor ——~—4> nonreactive products
`
`(5.27)
`
`'
`
`(5.21)—
`(5.20) is the initiation step, Eqs.
`where Eq.
`(5.23) are propagation steps, Eqs.
`(5.24) and (5.25)
`_oxidation steps leading to the ultimate oxidation
`product, 3042*, and Eqs.
`(5.26) and (5.27)
`termination
`steps.
`M+
`is a metal
`ion catalyst.
`The overall ki—
`netics of sulfate formation or sulfite loss is very
`complicated, although it has been shown that at pHs
`<
`8
`the proportionality given by Eq.
`(5.28)
`is observed
`(1)_
`.
`.
`
`.
`
`d[SO42‘]
`‘ dt
`
`u [M+][SO3
`
`2‘.
`
`][Hso§]-
`
`.
`
`»
`
`(5.28)
`
`in this expression there is no oxygen con—
`Note that
`centration dependency,
`that
`the reaction proceeds
`faster at higher pHs, and that
`the reaction is sensi—
`tive to metal
`ion catalysis, especially by Fe3+, Mn2+,
`and Cu2+.
`The pH dependency arises because the frac-
`tions of sulfite_and bisulfite ions are pH dependent.
`The kinetics and mechanism of oxidation of phenols
`and substituted 0— and p—dihydroxybenzenes
`in aqueous
`solution are also very complex and pH dependent.
`In
`general
`terms the oxidation of such molecules is very
`sensitive to the presence of metal
`ions, oxygen con—,
`centration, and pH, with increasing rates of oxidation
`J
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`

`90
`
`Oxidation and Photolysis
`
`at higher pHs. This suggests that it. is the anionic
`form of the phenol that is most sensitive to oxida(cid:173)
`tion.
`In the case of dihydroxybenzenes, such as epi(cid:173)
`nephrine and hydroquinones,
`the rate of oxidation
`often exhibits apparent second-order dependency on
`
`HO
`
`epinephrine
`hydroquinone
`hydroxide ion c~ncentration. The oxidation rates of
`hydroquinone and alkyl-substituted hydroqtiinones have
`been extensively studied because of their use in the
`photographic industry (2). Their mechanism of oxida(cid:173)
`tion appears .to be
`
`¢~¢·
`
`OH
`
`o-
`HQ2-
`
`However, it was found that the kinetics of oxidation
`of hydroquinone seemed to be dependent on quinone con(cid:173)
`centration; that is, the quinone; the immediate prod(cid:173)
`uct of the reaction ,
`c'a t a 1 y zed the oxidation of the
`hydroquinone. This was explained by the following re(cid:173)
`action,
`in which the quinone reacts with· the dianion
`of the hydroquinone to form a very unstable semiqui(cid:173)
`none radical, which very rapidly and spontaneously re(cid:173)
`acts with oxygen to form two molecules of the quinone
`and hydrogen peroxide (H202), or dimerizes to a stable
`product (3).
`·
`
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`

`Oxidation
`
`91
`
` semiquinone
`
`radical
`
`‘dimer
`
`2
`7_
`
`+
`
`H202
`
`O
`In both of the above reaction schemes it is the di—
`anion of the hydroquinone that appears to be the reac—
`tive species.
`The
`two pKa's of hydroquinone are )9
`and under pH conditions found in most formulations the
`fraction of hydroquinone present as
`its dianion is
`given by Eq.
`(5.29), where Kal and Kaz are the first
`and second dissociation constants of hydroqninone.
`
`[HQZ—]
`_
`fHQz_'_ [HQ]T0TAL
`
`=
`
`[H+]2 +
`
`.
`KalKag
`[H+]Ka1 + KalKag
`
`>
`
`5 2
`' 9)
`
`(
`
`With the conditions [H+]
`becomes Eq.
`(5.30).
`
`$> Kal and Kaz, Eq.
`-
`
`(5.29)
`
`Ra
`Ra
`Ra Ra
`fHQZ_ _ -_iL___Ei= __l___ii.[onj]2
`+
`2
`2
`[H ]
`Kw
`
`.
`
`'
`
`,
`
`-
`
`/
`
`'
`'
`(5.30)
`
`the hydroqui-
`if the rate of oxidation of
`Therefore,
`none is proportional
`to the hydroquinone dianion con—
`centration,-it can be seen that
`the rate of oxidation
`‘will be apparent second—order
`in hydroxide ion concen—
`tration.
`This
`is confirmed by the data in Fig. 5.2
`for the oxidation of m—dimethylhydroquinone and hydro—
`quinone in aqueous buffer solutions.
`Similar pH de—
`pendencies on the rate of oxidation of ascorbic acid
`(see curves 3 and 4 of Fig.
`1 of
`the L—ascorbic acid
`monograph) and captopril
`(see Fig.
`1 of the captopril
`monograph) have been observed.
`In each case the rate
`of oxidative decomposition under aerobic conditions is,
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`

`92
`
`Oxidation and Photolysis
`
`7.0
`
`pH
`FIGURE 5.2. pH-rate profiles for the oxidation of
`( 0> at
`hydroquinune (G) and m-dimethylhydroquinone
`The slopes of the lines are 1.96 and 1.98, re(cid:173)
`25°C.
`spectively. Both systems were studied in phosphate
`buffer.
`
`proportional to the fraction of the drug in the anion(cid:173)
`ic form, namely the ascorbate anion or the thiolate
`(RS-) in the case of ~aptopril.
`Interpretation of temperature effects on oxidative
`reactions is made difficult by the multiple steps in
`many of the reactions and because oxygen solubility in
`water (and other solvents) is temperature dependent.
`Since each reaction in a complex scheme will have its
`own activation energy, it is possible that as. the tem(cid:173)
`perature is changed a different reaction will become
`rate determining.
`Theoretically~ under such circum(cid:173)
`stances, the adherence of the reaction-rate/tempera(cid:173)
`ture relationship to the Arrhenius equation will b~eak
`down. Practically, however, over a
`limited tempera(cid:173)
`ture range, Arrhenius behavior may be observed, but
`the activation energy is very much an "apparent" acti(cid:173)
`vation energy for which the reaction conditions must
`be clearly stated.
`Included in this "apparent" acti(cid:173)
`vation energy is the temperature dependence of the
`oxygen solubility. Table 5.1 gives the 02 content of
`water at various temperatures if fhe water is satu(cid:173)
`rated by air or by pure oxygen. As can be seen for
`the air data, a 20°C change in temperature (5 ----725°C)
`results in a 40% decrease in oxygen concentration.
`If
`
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`
`
`Oxidation
`
`93
`
`TABLE 5.1. Oxygen Content of Water Under Air and Pure
`Oxygen at Atmospheric Pressure and Various
`Temperatures.
`
`
`
`Millimoles of 02 Millimoles of 02
`from air per mL
`from pure OE per
`Temperature (°C)
`of Hzoarb
`mL of H20 ,c
`
`
`-—
`o
`2:18'x 10-3
`0.386 x 10-3
`5
`e—
`0.34
`x 10—3
`10
`——
`0.304 x 10“3
`15
`4—
`0.267 x 10—3
`20
`——
`0.232 x 10—3
`25
`1.29 x 10-3
`——
`50
`9.28 x 10-4
`—-‘
`100
`.51 x 10-4
`'q
`
`
`'
`
`'
`
`aFrom Reference 4.
`bCalculated from cc of 02 in H20 and the expression
`PV = nRT.
`CFrom Reference 5.‘
`
`
`
`
`
`in
`the rate of reaction under study is first order
`oxygen concentration,
`then a ninefold increase in rate
`(Q10 = 3) due to the direct effect of
`temperature on
`the rate—controlling step will
`show up experimentally
`as only a 5.5—fold increase im mate owing to the con—
`/
`comitant change.in oxygen-concentration.
`
`3. Oxidative Pathways @f Phanmaceutical Interest
`
`A few selected oxidative reactions of pharmaceutical
`interest are illustrated here;
`the stability mono~
`graphs
`include other examples. Many drug Compounds
`have been reported to be subject
`to autoxidation,
`in—
`cluding adriamycin hydrochloride, amphotericin B,
`apo—
`morphine, ascorbic acid, Capropril, chlorpromazine and
`other phenothiazine derivatives, cyanocobalamin,
`cys~
`teine, epinephrine, ergometrine, hydrocortisone,
`iso—
`amyl nitrite,
`isoproterenol, kanamycin, 6—mercaptopur—
`ine, morphine, neomycin, norepinephrine, novobiocin,
`p—aminobenzoic acid, paraldehyde, penicillin, pheny1—'
`ephrine, physostigmine, prednisolone, prednisone,
`
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`94
`
`Oxidation and Photolysis
`
`procaine, resorcinol, riboflavin, streptomycin and
`dihydrostreptomycin, sulfadiazine, terpenes,
`the
`tetracyclines, __ thiami_I!_EL,_ and vitamins A, D, and E.
`The unwanted conversion of fats, oils, flavors, and
`perfumes to a rancid state is due to oxidation of
`these unsaturated molecules.
`The double bonds are
`oxidized to form hydroperoxides, as demonstrated ear(cid:173)
`lier, which then produce aldehydes;
`the latter cause
`the offensive odors and unpleasant flavors. Vitamin A
`(see monograph) and amphotericin B (see monograph) are
`two drug molecules with extended conjugated double
`bonds that are very susceptible to oxidative break-
`down.
`Epinephrine forms colored products on oxidation
`[Eq. (5.31)].
`
`epinephrine quinone
`
`melanin
`pigments
`
`-H 0 2
`
`adrenochrome
`
`< 5. 31)
`
`The activity of riboflavin is contingent on its
`ability to take part in this redox equilibrium:
`
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`

`Oxidation
`
`95
`
`6—..—
`
` ————>
`
`I
`(fr‘HOH) 3
`CHEOH
`
`.
`
`I
`
`H
`‘gHa
`(EHOH) 3
`CHZOH
`
`
`
`
`
`riboflavin
`"
`'
`
`‘
`
`dihydroriboflavin'
`‘
`. (5 . 32)
`
`Morphine dimerizes when oxidized.> Equation (5.33)
`shows the first step.
`
`
`
`morphine
`
`_
`
`,
`
`" '
`
`(5.33)
`
`formed morphine free radical couples with a mor—
`The
`phine molecule (at
`the free position ortho to the phe—
`nolic oxygen)
`to give the dimer
`(bimorphine or pseudo—
`morphine).
`”Hydrogen peroxide is also produced and can
`cause additional oxidation to the N—oxide.
`
`Vitamin E can form an epoxide, which then produces
`a quinone:
`vitamin E («Ftocopherol):
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`96
`
`Oxidation and Photolysis
`
`HO
`
`;FH3
`TH3
`H2-(cH2-cH2-CH-CH2J 2-cH2-CH2~H
`CH3
`CH3
`
`vitamin E epoxide
`
`~-tocopherylquinone
`
`(5.34)
`
`The phen.othiaziqes readily oxidize, producing a
`multitude of products.
`The degradation mechanism of
`promethazine and the influence of pH, metals, chelat(cid:173)
`ing agents, and antioxidants, have been extensively
`studied (6,7). The following r~actions have been pro(cid:173)
`posed and their products isolated.
`
`Sulfide and sulfhydryl-containing molecules are par(cid:173)
`ticularly vulnerable to oxidation. Captopril (see
`
`(5.35)
`
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`

`Oxidation
`
`97
`
`is an example of a recently developed sulf—
`monograph)
`hydrylfldrug-with oxidative stability problems.
`.Sulf—
`”hydrylrcontaining molecules generally oxidiZe to their
`corresponding disulfides, Eq.
`(5.36).
`
`R—su
`
`§92—€>
`
`R-S—S~R + H20-
`
`.
`
`,
`
`(5.36)
`
`4.
`
`Inhibition of Oxidation
`
`Exclusion of Oxygen
`
`If a molecule requires the presence of oxygen to de—
`grade why not exclude oxygen from the formulation?
`For parenteral drugs this can be achieved by packaging
`the drug in glass ampuls that are heat sealed under an
`inert atmosphere.
`For
`tablets, capsules,
`and so on,'
`packaging of
`the formulation in a hermetic strip may
`be useful
`in preventing the oxidation.
`Quite often manufacturers would like to formulate a
`drug for parenteral administration in,a multidose
`vial, which is sealed with a
`rubber stopper.
`It
`is
`very difficult
`to formulate very oxygen~sensitive
`drugs
`in such multidose vials because rubber
`is rea—
`sonably permeable to oxygen and both synthetic and
`natural
`rubber
`tend to release additives capable of
`catalyzing oxidative reactions. Oxidatively unstable
`drugs formulated in multidose vials require more than
`
`just oxygen exclusion during sealing to prevent oxida—
`tion from occurring.
`‘
`The
`importance of deoxygenating the water and the
`headspace atmosphere in an ampul prior to sealing can
`be seen in the following calculation using/captopril
`as an example.
`The stoichiometry of
`the oxidative
`breakdown of captopril
`is given by Eq.
`(5.37), showing
`that 2 mol-of captopril are lost for each half—mole of
`oxygen consumed.
`
`1/202
`
`.
`captopril disulfide
`
`COO
`
`(5.37)
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`98
`
`Oxidation and Photolysis
`
`injection of captopril, in a 2-mL ampul,
`A 2 mg/2 mL
`has a headspace of -1 mL.
`The 2 mg of captopril is
`equivalent to 9.2 x 1o-8 mmol of captopril. At 25°C
`there are 0.23 x Lo-3 mmol of dissolved oxygen in each
`mL of water .for a total of 0.46 x 10-3 mmol of oxygen.
`The 1 mL of headspace, assuming that it is air, con(cid:173)
`tains 8.6 x 10-3 mmol of oxygen. Based on the stoi(cid:173)
`chiometry of Eq. (5.37) it can be seen that the 9.06 x
`1o-3 mmol of oxygen in the system would be capable of
`degrading 3.62 x 10-2 mmol of captopril, that is,
`there is more ·than enough oxygen present in this for(cid:173)
`mulation to completely d~grade the captopril.
`If the
`oxygen in the headspace is removed by flushing with an
`inert gas, the 0.46 x 1o-3 mmol of oxygen in the water
`are still capable of degrading '1.84 x 10-3 mmol of
`captopril, or 20% of the formulation.
`By· performing
`such calculations it is possible to predict how thor(cid:173)
`ough the exclusion of oxygen from the system must be
`in order to prevent the oxidative breakdown of sensi(cid:173)
`tive drugs.
`
`Alteration of Solution pH
`
`As has already been discussed, the oxidation of many
`drugs is pH sensitive. Acidic drugs such as ascorbic
`a c i d , ph e n o 1 s ,
`a n d
`s u 1 f h,Y d r y 1 c om pound s a 11 d e g r ad e
`more rapidly in neutral to alkaline pH conditions.
`For such drugs the pH range 3 to 4 is generally forind
`most useful in minimizing oxidation. Obviously, this
`pH range would not be useful for acidic drugs that
`have limited aqueous solubility at low pH values.
`Amine drugs such as the phenothiazines appear to be
`most stable in their protonated forms,
`that is, also
`at low pH values.
`
`Protection from Light
`
`Oxidative breakdown of drugs gerrerally proceeds
`through the sequence of initiation, propagation (and
`maybe chain branching), and termination. As mentioned
`earlier, a triggering force that may promote oxidation
`is "light" (namely, certain components of the electro(cid:173)
`magnetic spectrum). Not all photolytic re~ctions are
`oxidative in nature and not all oxidative reactions
`require light as"either an initiator or as an integral
`component of the propagation steps.
`If, however,
`
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`

`Oxidation
`
`99
`
`light does initiate or promote an oxidative breakdown
`of.a drug,
`the exclusion of
`the particular wavelength
`“range of
`light responsible for
`the catalysis will'
`often suppress the oxidation. This can be achieved by
`the total exclusion offlight using an opaque container
`or the use of pigmented glass capable of excluding the
`damaging wavelengths.
`This will be discussed further
`in the section of this chapter on photolysis.
`
`Use of Chelating Agents and Antioxidants
`
`Oxidation reactions can be inhibited by agents that
`are
`
`(a)
`
`(b)
`
`ion initiators of
`chelating agents for metal
`free— radical oxidation reactions;
`reducing agents,
`that
`is,
`substances that can
`reduce an oxidized drug;
`is,
`that
`(c) preferentially oxidized compounds,
`ag-ents
`that are more readily oxidized than the
`agents they are to protect; or
`re—
`chain terminators,
`that is, agents capable of
`V acting with radicals in solutions to produce a
`new species,
`a chain terminator radical, which
`does not
`reenter the radical propagation cycle.
`The new radical may be
`intrinsically stable or
`may dimerize to form an inert molecule.
`
`(d)
`
`four categories are often classified
`in all
`Compounds
`as antioxidants (or antoxidants).
`It is probably more
`accurate to classify compounds
`in categories (b)—(d)
`as antioxidants and chelating-agents as synergists.
`These categories are now treated in more detail.)
`
`(a) Chelating Agents. Oxidative reactions are
`often initiated by metal
`ions such as Fe3+, Cu2+,
`003+, N12+, Mn3+.
`These metal
`ions act as initiators
`because in their oxidation states they are capable of
`acting as radicals.
`For example, Cu2+ has 27
`elec—
`trons and requires one more electron to complete the
`electron pair.
`Metal
`ions catalyze oxidative reactions in a number
`of ways.
`They can react directly with oxygen to pro-
`duce an oxygen radical, which can then initiate an
`autoxidation.
`The metal
`ion can form a cbmplex with’
`oxygen and subsequently form a peroxy radical.
`The
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`

`100
`
`Oxidation and Photolysis
`
`ion can react with the drug itself to r'orm a
`metal
`radical, as illustrated by Eq. (5.38), which is then
`able to enter into a propagation cycle [see Eqs. (5.7)
`and (5.13)].
`
`Mn+ + RH.
`
`(5.38)
`
`ion can also react with a hydroperoxide in
`The metal
`the formulation to catalyze the breakdown, given by
`Eq. (5.39).
`Mn+ + R'OOH ~ M(n- 1 )+ + H+ + R'o
`
`(5.39)
`
`·
`
`2
`
`R'OOH co~ld be a hydroperoxide of the drug itself or
`of some other component of the formulation.
`Chelating agents act in an antioxidant capacity by
`binding metal ions, thus removing them,
`thermodynami(cid:173)
`cally speaking, from solution. The most effective
`chelating agents used pharmaceutically are ethylenedi(cid:173)
`aminetetraacetic acid (EDTA), citric acid, many of the
`amino ac·ids, phosphoric acid (weak), and tartaric
`acid.
`EDTA and citric acid are the two most useful
`agents. Their metal-binding capacity is dependent on
`their state of ionization, both being most effective
`when their carboxyli~ acid groups are fully ionized.
`Thus, they lose their chelating capacity at low pH.
`
`EDTA tetraanion (edetate)
`
`citric acid trianion
`
`Just because an agent is able to chelate metal ions
`does not mean that it will reduce the effectiveness of
`the metal ion to act as a catalyst.~ There are circum(cid:173)
`stances in which the metal ion may bind to some func(cid:173)
`tional groups and in this bound capacity actually may
`be a better catalyst than in the unbound state. The
`chelating agents mentioned above, how~ver, generally
`act to lower the catalytic activity of the metal ions
`towards radical chain reactions.
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`
`
`Oxidation
`
`101
`
`Reducing Agents.‘ This approach is generally
`(b)
`not used as-a means of preventing oxidation.
`Sodium
`,thiosulfate and ascorbic acid are two reduCing agents.~..
`that have been used in this capacity.
`
`'These are
`(a) Preferentially Oxidized Compounds.
`compounds
`that are moreJreadily oxidized than the
`agents they are to protect. Essentially these agents
`act as oxygen scavengers.
`Two good pharmaceutical
`examples of oxygen scavengers are sodium bisulfite
`(and sulfite) and ascorbic acid.
`Sodium sulfite re—
`acts with oxygen according to Eq.
`(5.40).
`
`5032‘
`
`+ 3502 ——-—> 3042‘
`
`((5-40)
`
`is possible to calculate the approximate amount of
`It
`antioxidant needed to use up all the oxygen in an am—
`pul,
`for example, by calculating the amount of oxygen
`dissolved in the water and the headspace of the ampul.
`In this example,
`2 mol of
`sodium sulfite equivalents
`(as sodium bisulfite) would be needed to consume each
`mole of oxygen.
`The sulfites are very commonly used, and a word of
`caution is in order.
`In the process of acting as an—
`tioxidants, sulfites yield acid sulfates, which cause
`a drop in pH. Also, sulfites can readily form inac~
`tive addition compounds, as with epinephrine (see
`monograph).
`They react with compounds
`such as
`a1—
`kenes, alkyl halides,
`and aromatic nitro and_carbonyl
`compounds.
`Sometimes, as with thiamine,
`they may
`cleave molecules.
`_
`,
`the sulfite ions
`Although the nucleophilicity of
`can be a disadvantage,
`there are circumstances where
`this nucleophilicity is an advantage.
`For example,
`in
`the oxidatidn of hydroquinone,
`it appears that the im—
`mediate prOduct of
`the oxidation,
`its corresponding
`quinone,
`can act as a catalyst for further oxidation.
`Sodium sulfite (or bisulfite) can react with the qui—
`none and hydrogen peroxide to form hydroquinone mono—
`sulfonate, and sulfate as described by Eq.
`(5.41).
`
`
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`102
`
`Oxidation and Photolysis
`
`0
`
`OH
`
`(5.41)
`
`It appears that. thiols can also act as antioxidants in
`a similar manner. All the major water-soluble antiox(cid:173)
`idants, namely acetone sodium bisulfite, ascorbic
`acid, cysteine hydrochloride, isoascorbic acid, sodium
`bisulfite, sodium formaldehyde sulfoxylate; sodium
`metabisulfite, sodium sulfite, thioglycerol, thiogly(cid:173)
`colic acid, and thiosorbitol, act as oxy·gen scaven(cid:173)
`gers, although the thiol (sulfhydryl) group-containing
`antioxidan~s can also act as chain terminators (cate(cid:173)
`gory d).
`Few if any of the lipid-soluble antioxidants
`act as true ·oxygen scavengers except for, perhaps,
`ascorbityl palmitate.
`
`(d) Chain Terminators. Referring to the earlier
`discussions on the oxidation of olefins
`[Eqs. (5.7)(cid:173)
`(5.15)], bisulfite [Eqs. (5.20)-(5.27)], and phenols,
`all of these reaction·s proceed' through a radical mech(cid:173)
`anism. Therefore any substance that can donate a hy(