Pharmaceutical Research, Vol. 4, No. I , [987
`
`Report
`
`Stability of Lidocaine in Aqueous Solution: Effect of
`Temperature, pH, Buffer, and Metal Ions on Amide Hydrolysis
`
`Michael F. Powell1
`
`Received May 5, I986; accepted August 26, I986
`
`The degradation of lidocaine in aqueous solution obeys the expression
`k0,,s = (kH,[H’] + k°)[H+]/([H+] + IQ + kgK,/([H*] + K“)
`
`where k“. is the rate constant for hydronium ion catalysis, and [to and k", are the rate constants for the
`spontaneous (or water-catalyzed) reactions of protonated and free-base lidocaine. At 80°C, the rate
`constants for these processes are 1.31 x 10411!“1 sec“, 1.37 x 10'9 sec", and 7.02 x 10’9sec”';
`the corresponding activation energies are 30.5, 33.8, and 26.3 kcal mol", respectively. It was found
`that the room temperature pH of maximum stability is ~3-6 and that lidocaine is more reactive in the
`presence of metal ions such as Fe2+ and Cu“. The dissbciation constant, K,, for lidocaine at 25~80“C
`was also measured at 0.1 M ionic strength and a plot opr, versus l/Tgave a slope of(l .88 i- 0.05) X
`10’ K" and intercept 1.56 t 0.16.M
`KEY WORDS: kinetics of lidocaine degradation; aqueous stability; pK_; amide hydroylsis.
`
`water~catalyzed reaction), as observed for N—(l-amino-
`alkyl)amides (11,12) and certain hydroxy anilides (13). In
`order to determine the pH of maximum stability for lido-
`caine, and to delineate the factors that affect the rate of lido-
`caine degradation, the effect of pH, temperature, buffer,
`propylene glycol, and added metal ions on the hydrolysis of
`lidocaine in aqueous solution was investigated.
`
`EXPERIMENTAL
`
`Apparatus and Reagents. Reverse-phase chromatog-
`raphy of lidocaine was carried out using a high-performance
`liquid chromatography (HPLC) system consisting of a
`Model 725 Micromeritics autoinjector, a Model 110A Altex
`pump, a Model 770 Spectra Physics spectrophotometric de—
`tector, and an SP 4000 computing integrator. A 250 X 4.6-
`mm Partisil S ODS 3 column (Whatman) was used for anal-
`ysis. Lidocaine hydrochloride USP was used for all experi-
`ments herein. HPLC-grade methanol,
`tetrahydrofuran
`(Burdick and Jackson), and distilled, deionized water were
`used for the preparation of mobile phase. Potassium hy~
`drogen phosphate, sodium hydroxide, hydrochloric acid,
`iron (Ii) chloride, copper (II) chloride, heptane sulfonic acid
`sodium salt, propylene glycol, and 2,6-dimethy1aniline were
`of the highest grade commercially available (Aldrich or Mal—
`linckrodt) and were used without further purification.
`HPLC Conditions. A linear response (21%) through
`out the range of 0.1—- 10 pg lidocaine injected and separation
`of lidocaine from its degradation products, 2,6-dimethylani-
`line and N,N—diethylaminoacetic acid, were achieved using
`the following conditions: mobile phase, waterzmethanolztet-
`rahydrofuran (48.2:48.2:l.6, v/v) containing 0.02 M heptane
`sulfonic acid sodium salt and 0.02 M KHZPO4; flow rate,
`
`INTRODUCTION
`
`Lidocaine exhibits exceptional stability in most paren-
`teral solutions (1—3) and is extremely resistant toward hy—
`drolysis at room temperature. even in strongly acidic or
`basic media (4—7). At higher temperatures, however, lido-
`caine degrades slowly, to give primarily 2,6-dimethylaniline
`and N,N-diethylaminoacetic acid, as shown in Scheme I (4).
`
`n 9c»,
`can
`5 w
`N’c’ w:
`can,
`"3
`
`+
`
`lien; £sz
`HO—C’
`‘N\
`cz"s
`
`NHg
`CH3
`
`Scheme 1
`
`Despite the importance of this drug as a powerful local aneSv
`thetic, little is known about the temperature and pH depen-
`dence on lidocaine degradation in aqueous solution. Con-
`flicting accounts regarding the stability of lidocaine have
`been reported; for example. an early study found only 0.05%
`drug loss after 3 hr at 116°C (4), whereas another study
`showed more rapid drug 1085 (2.5%) after reaction for 12 hr
`at 96°C (5). Furthermore, the pH of maximum stability for
`lidocaine has not been accurately determined. Previous re-
`ports of lidocaine hydrolysis in aqueous solution (4—7) have
`not discerned conclusively whether lidocaine exhibits a
`sharply defined pH of maximum stability, as do formamide
`(8), acetamide (9), and ot-propylamino-Z—methylpropionani-
`_ lide (10), or whether it shows a broad pH range of maximum
`stability (characteristic of a predominant spontaneous or
`
`W '
`
`institute of Pharmaceutical Sciences. Syntex Research. Palo Alto,
`California 94304,
`
`0724»8741/87/0100~0042$DS.00I0 D 1987 Plenum Publishing Corporation
`
`42
`
`
`
`P3991
`
`Teoxane S .A.
`Exhibit 1031
`
`
`Page 1
`
`

`

`43
`
`RESULTS AND DISCUSSION
`
`life“ ofpH and Temperature. For all experiments, ap-
`parent first-order kinetics were observed, and the pH de-
`pendence on lidocaine degradation is given by the data in
`Table 1. The log (rate)-p1-1 profiles are shown in Fig. 1; the
`best-fit lines in this figure were generated using Eq. (1).
`
`kob: = (ka+[H*l + k9) [HQ/(1H3! + K.)
`+ ksKa/(lHW + K.)
`
`(1)
`
`The rate constants of Eq. (1) are for catalysis by hydronium
`ion (km) and for the spontaneous (or water-catalyzed) reac-
`tions of protonated (kc) and free-base forms of lidocaine (kg).
`The factors [H+]/([H*] + K.) and [ta/([11+] + K.) represent
`the fractions of protonated and unprotonated lidocaine, re-
`spectively, and are an integral part of Eq. (1) necessary for
`fitting the pronounced curved region near pH 7 in the
`log(rate)-pH profile. The K. values used for the fit of Eq. (1)
`were determined by least-squares analysis of pKa versus l/T
`(Table 11), as shown in Fig. 2. The secondary rate constants
`derived from the data in Table 1 are given in Table III; the
`corresponding activation parameters are summarized in
`Table IV.
`'
`1 show that lidocaine
`The log (rate)—pH profiles in Fig.
`hydrolysis is only weakly catalyzed by hydronium ion and
`even less so by hydroxide ion; in fact, the hydroxide term
`kHo_[HO‘] was not required in Eq. (1) for fitting the data up
`to pH ~12. Hydroxide ion-catalyzed amide hydrolysis for
`other amides has been observed in strongly alkaline solu-
`tions, although it went undetected for lidocaine, even in 0.1
`M KOH at 100°C (4,5). In the midva region, the dominant
`rate constants of Eq. (1) are kD and k1,, denoting the sponta-
`neous or water-catalyzed reaction of protonated and frce—
`base lidocaine, respectively. At 80°C. the spontaneous reac-
`tion for free-base lidocaine was found to be approximately
`five times faster than for protonated lidocaine. This rate
`ratio becomes even more pronounced at lower tempera-
`tures, and at 25°C; k; is calculated to be approximately 36
`
`109It(see-1)
`
`1
`
`3
`
`5
`pH
`
`7
`
`9
`
`11
`
`Fig. I. Log(rate)~pH profile for the deg»
`radation of lidocaine in aqueous solution
`at 80 and 100°C. The dashed line is the
`calculated log(rate)—pH profile for lido-
`caine at 25°C and is included to show the
`pH range of maximum stability (pH 3—6)
`at room temperature.
`
`Stability of Lidocaine in Aqueous Solution
`
`1.0 ml/min; detection, 230 nm; injection volume, 10~50 pl;
`and typical retention times, 12—125 min for lidocaine and
`9.5—10 min for 2,6-dimethylaniline (capacity factors, 5.66
`and 4.33, respectively).
`Kinetics. For all experiments, 0.01 M buffer solutions
`(p. = 0.10) containing lidocaine (35 ug/ml) were prepared
`shortly before use and the pH’s were determined at 21, 60,
`and 80°C; pH's at 100°C were extrapolated from linear plots
`of pH (at 21 —80°C) versus the reciprocal of absolute temper-
`ature. A summary of the buffer solutions used and pH’s is
`given in Table I. For strongly acidic or alkaline solutions at
`80 and 100°C, the hydronium and hydroxide ion activities
`were calculated from published activity coefficients (14,15)
`or HD values (16,17). In a typical experiment, lO-ml aliquots
`of reaction solution containing lidocaine were transferred to
`pretreated ampoules, flame—sealed, and stored at 80 or
`100°C. Several of these samples were also refrigerated im~
`mediately after flame-sealing and were later used as controls
`for the initial time points. At known time intervals, up to '
`~300 days, ampoules were removed and refrigerated until
`6—10 samples for each kinetic run were taken. Upon re-
`moval of the last samples, the stored solutions were allowed
`to warm to room temperature and then all samples were an-
`alyzed by HPLC on the same day. Rate measurements were
`carried out using varying acetate buffer concentrations up to
`0.15 M in order to determine the effect of buffers on this
`reaction. Reaction kinetics were also carried out at 100°C
`with added iron ([1) chloride and copper (II) chloride (10 and
`50 ppm) at pH 7.6 or with added 20 and 80% (v/v) aqueous
`propylene glycol. Peak area integration values were used di-
`rectly. in first-order fits of the data and most reactions were
`followed to less than 90% remaining. Nonlinear least-
`squares analysis (18) was used to obtain the best-fit rate
`constants from the log(rate)—pH profiles.
`pKa Determinations. The pK. of lidocaine at 0.1 M
`ionic strength (NaCl) was determined by preparing a solu-
`tion of lidocaine (0.0103 M) and lidocaine hydrochloride
`(0.0103 M) and measuring the pH‘s at the temperature of
`study using a Radiometer PHM 64 pH meter and Model
`GK2401C combination electrode. Buffer dilutions were not
`carried out during the pH measurements because the pK. of
`lidocaine is insensitive to ionic strength (19).
`
`Table 1. Summary of Rate Constants for the Degradation of
`Lidocaine in Aqueous Solution”M
` 80°C 100°C
`
`
`Bufi'er
`108 k(sec"‘)
`pub
`109k (sec-I)
`pH“
`(n = 0.1)
`15.5 i 0.10
`1.0
`15.8 2 1.31
`1.0
`0.1MHCl
`2.40 t 0.79
`2.3
`1.58 i 0.31
`2.3
`0.005 M HCl
`1.60 i 0.19
`4.6
`1.32 z 0.07
`4.5
`0.01 M phosphate
`2.76 I 0.10
`5.8
`2.38 I 0.15
`5.6
`0.01 M phosphate
`6.54 t 0.33
`7.2
`4.91 1 0.35
`7.2
`0.01 M phosphate
`5.76 t 0.24
`8.3
`5.97 z 0.66
`8.3
`0.01 M phOSphate
`5.23 i 0.37
`8.5
`7.33 2 0.42
`8.8
`0.01 Mphosphate
`4.05 t 0.06
`10.1
`6.32 1' 0.17
`10.6
`0.01 MKOH
`
`
`
`
`7.88 z 0.18 11.111.70.1MKOH 4.10 '1: 0.08M
`
`_
`a pH determined at 80°C.
`4’ pH extrapolated from a linear plot of pH (21—80°C) versus the
`reciprocal of absolute temperature.
`‘
`
`Page 2
`
`Page 2
`
`

`

`44
`
`Powell
`
`Table II. Effect of Temperature on the pK. of Lidocaine”W
`
`Reference
`pK.
`Temperature (°C)
`20
`8.24
`10
`This paper
`7.97
`2|
`20
`7.92
`25
`20
`7.57
`38
`This paper
`7.41
`50
`This paper
`7.14
`65
`This paper
`6.91
`80
`
`
`6.62“100 This paperm
`" lonic strength = 0.1 M.
`0 Calculated from the least-squares line of pK. versus l/T having a
`slope of (1.88 t 0.05) X 103 K‘1 and intercept 1.56 1 0.16.
`
`Table III. Summary of Rate Constants for the Degradation of
`Lidocaine in Aqueous SolutionW
`Rate constant‘1
`TemperatureW
`
`k0 two")km W" sec”)(’0 kl, (sec‘l)W
`
`
`100
`1.35 X 10"
`l.82 X lo"
`5.26 X 10—3
`80
`l.31 X 10‘7
`L37 X 10“9
`7.02 X 10‘9
`
`L90 X 10'”4.32 X l0‘"25" 6.86 X IO“2M
`
`
`“ Obtained from a nonlinear least-squares fit (Ref. 18) of the rate
`data to Eq. (1).
`1’ Rate constants at 25°C were calculated from Ea and log A given in
`Table IV.
`
`times larger than k0. This difference in relative rates is not
`critical for drug formulation, however, because even the
`faster reaction (k1,) has a calculated room»temperature shelf
`life (190) of >400 years. At 25°C. the pH of maximum sta—
`bility for lidocaine is ~3—6, as shown by the calculated log
`(rate)~—pH profile (dashed line in Fig. 1). At temperatures
`above 100°C the rate constants k(, and k; are much closer in
`magnitude, and at 145°C they are predicted to be equal at 2.4
`x 10" sec"; above this temperature k0 > k’. Thus, at
`145°C, lidocaine stability is independent of pH from pH 2 to
`pH 12, where the predicted shelf life is approximately 12 hr.
`Effect of Additives. A study was carried out to deter-
`mine if the anomalous rate constants in the literature for li-
`docaine degradation (4-7) could be caused by buffers, by
`the adventitious presence of metal ions, or by the addition of
`added alcohols such as propylene glycol. Inspection of Table
`V shows that acetate buffer up to 0.15 M did not noticeably
`accelerate the rate of drug degradation, and little, if any.
`buffer catalysis is expected in the 0.01 M phosphate solu-
`tions used herein. This lack of buffer catalysis for amide hy«
`drolysis has been reported previously for N-(lvamino-
`a1kyl)amides (12,13), compounds that show log(rate)-—pH
`
`profiles remarkably similar to that of lidocaine. Added pro—
`pylene glycol increased the rate of .lidocaine degradation
`only slightly; such rate enhancements with added alcohols
`have been observed previously (21) and may be due to a
`change in the reaction mechanism from hydrolysis to esteri-
`fication or, perhaps, to a solvent effect.
`The largest increase in the rate of lidocaine degradation
`was observed when FeClz and CuClz were present. Even at
`10 ppm each these metal salts provided a 14-fold increase in
`degradation rate, possibly by increasing the rate of nucleo-
`phllic attack on the carbonyl carbon by complexation of the
`metal with the carbonyl oxygen and amine nitrogen. Other
`explanations involving breakdown of the tetrahedral inter‘
`mediate formed in base-catalyzed amide hydrolysis are less
`satisfactory but may still account for the small rate aceelera~
`tions observed herein (22). This sensitivity to added metal
`ions shows that the disparity in reaction rates reported ear-
`lier (4—7) could be due to the presence of adventitious metal
`ions, especially since untreated glass ampoules can leach
`metal ions or other impurities which may effect the reaction
`rate (23). Thus, to ensure minimal drug loss upon auto—
`claving aqueous lidocaine solutions, care should be taken to
`avoid contamination by metal ions such as Fe“ and Cu“.
`Reaction Mechanism. It is interesting to speculate on
`the reason(s) why k1, is greater than kg. The following kinet-
`ically indistinguishable mechanisms predict k; > k0: intra-
`molecular nucleophilic catalysis (Scheme 11), intramolecular
`general-acid— specific—base catalysis (Scheme 11]), and intra-
`molecular general-base catalysis (Scheme IV).
`
`91:». tea“: W“
`
`— [c.mm‘h"s ——
`
`Schemell
`
`en;
`H est-W" c H
`me
`*u’ 3 5 _.
`CH; '0“
`
`@ m...
`
`Scheme [II
`
`“a
`N
`H3
`
`.
`
`+
`
`mo,
`
`too-g
`
`.wz. £2":
`
`N‘q‘l
`
`c ..
`t
`+ "Wyatt"; 2 5
`
`Table IV. Summary of Activation Parameters for the Degradation
`of Lidocaine in Aqueous Solution
`Rate
`E.
`AH:
`ASt
`constant
`(keal mol' |)
`log A
`(kcal mol ’ ')
`(cal moi“ K ‘ ')
`
`
`~5.96
`29.8
`12.0
`30.5
`k...
`—5.65
`33.1
`12.l
`33.8
`k0
`
`
`
`
`26.3 8.16 25.5k1, —23.5M
`
`Page 3
`
` 6.4
`
`33
`31
`29
`1/!“ x 1040(4)
`Fig. 2. Relationship between lidocaine
`pK. and the reciprocal of absolute tem~
`perature. The dashed circle is the pK_
`(6.62) at 100°C calculated from linear
`least-squares analysis. The data points
`at 10, 25. and 38°C are from Ref. 20.
`
`27
`
`35
`
`Page 3
`
`

`

`4s
`
`ried out by others (4—7) showed nonufirst-order kinetics
`(where pseudo-first-order kinetics are expected for a simple
`amide hydrolysis) or were followed‘for short reaction times.
`such that reliable rate constants could not be obtained. It is
`concluded that lidocaine is extremely stable across the en-
`tire pH range, however, less so in the presence of metal ions
`as Fe2+ and Cu“.
`
`ACKNOWLEDGMENTS
`
`The author expresses thanks to Drs. L. Gu, D.
`Johnson, and W. Lee for helpful discussions during the prep-
`aration of the manuscript.
`
`REFERENCES
`
`00an9wN
`
`10.
`
`l. T. E. Lackner, D. Baldus. C. D. Butler, C. Amyx, and G.‘
`Kessler. Am. J. Hosp. Pharm. 40:97~101 (1933).
`. H. L. Kirschenbaum, W. Aronoff. G. P. Perentesis, G. W. Plitz,
`and A. J. Cutie. Am. J. Hosp. Pharm. 39:1013-1015 (1982).
`. F. M. Smith and N. O. Nuessle. Am. J. Hosp. Pharm. 38:1745~
`1747 (1981).
`K. Bullock and J. Grundy. J. Pharm. Pharmacol. 7:755—773
`(1955).
`. S. Goto and T. Itano. Yakugaku Zasshi 99:146—154 (1979).
`. J. Katz. Anesthesiology 27:835-837 (1966).
`. E. Zollner and G. Vastagh. Pharm. Zentral. 105:369—372 (1965).
`. J. Hine. S.-M. King. R. Midden, and A. Sinah. J. Org. Chem.
`46:3186—3189(198l).
`1°
`T. C. Bruice and F. H. Marquardt. J. Am. Chem. Soc. 84:365-
`370 (1962).
`S. O. Eriksson and .130. Omdal. Acta. Pharm. Suec. 1:77-90
`(1964).
`.
`11. G. M. London and J. .1. Jacob. Chem. Soc. Chem. Commun.
`377—378 (1980).
`12. G. M. Loudon, M. R. Almond, and J. N. Jacob. J. Am. Chem.
`Soc. 103:4508—4515 (1981).
`13. T. Yamana, A. Tsuh, and Y. Mizukami. Chem. Pharm. Bull.
`21:721-728 (1973).
`H. W. Hamed. Physical Chemistry of Electrolyte Solutions,
`Reingold. New York, 1958.
`15. R. S. Greelct. Anal. Chem. 32:1717 (1960).
`16. A. J. Kresge, H. J. Chen, G. L. Capen. and M. F. Powell. Can.
`J. Chem. 61:249-256 (1983).
`17. C. H. Rochester. Acidity Functions, Academic Press, New
`York, 1971.
`18. P. R. Bevington. Data Reduction and Error Analysis for the
`Physical Sciences, McGraw-Hill. New York, 1969.
`19. R. H. Levy and M. J. Rowland. J. Pharm. Pharmacol. 24:841—-
`847 (1972).
`20.
`H. Kamaya. J. J. Hayes. Jr., and [. Ueda. Anesth. Analg.
`62:1025~ 1030 (1983).
`K. A. Wyatt and I. H. Ditman. Aust. J. Pharm. Sci. 8:77—85
`(1979).
`T. H. Fife and T. J. Przystas. J. Am. Chem. Soc. 108:4631—-
`4636 (1986).
`23. 1... Gu. Personal communication.
`24. H. Bundgaard and E. Falch. Int. J. Pharm. 23:223—237 (1985).
`25. T. C. Bruice and S. Benkovic. Bioorgam'c Mechanisms, W. A.
`Benjamin, New York, 1966.
`
`I4.
`
`21.
`
`22.
`
`
`
`CH3
`iii/”'3‘ (Cal‘s -
`(‘O'H *er clfli
`CH3
`.4
`
`9
`”3
`NH; + Hoe’ch’CZH-l
`‘csz
`CH3
`Scheme 1V
`
`It has been argued that intramolecular nucleophilic ca-
`talysis may be the favored mechanism for the reaction of
`some amides such as allopurinol derivatives, since the allo-
`purinol moiety is a good leaving group (24,25). This mecha-
`nism probably does not occur for lidocaine hydrolysis, how-
`ever, because the ensuing lactam product would be strongly
`sterically strained and so its formation would be unlikely.
`For the mechanism of Scheme III to be kinetically compe-
`tent, the protonated form of lidocaine must react with hy-
`droxide ion at pH ~7. resulting in an unusually fast reaction
`between a protonated amide and a hydroxide ion. This is
`unlikely inasmuch as specific-base catalysis does not readily
`occur with free-base lidocaine, even in 1 M KOH at 80°C
`(Fig. 1), and it is improbable that N-protonation of lidocaine
`should promote the specific-base catalysis reaction by a
`factor of >105. It is more probable that lidocaine reacts by a
`mechanism involving general-base catalysis by an intramo—
`lecular reaction with the tertiary amine group through a
`water molecule (Scheme IV). Intramolecular base catalysis
`by the amine group is plausible because the five-membered
`cyclic transition state is entropically favored, and because
`the base strength (and sensibly the nucleophilicity) of 1ido~
`caine is relatively high (pK. 2 7.92 at 25°C). In this case the
`rate enhancement by intramolecular general-base catalysis
`is expected to be less than an order of magnitude, and this is
`the type of rate enhancement observed experimentally for
`lidocaine degradation.
`The rate constants in Table [II for lidocaine degradation at
`80~100°C can be used to estimate lidocaine stability at tem-
`peratures used for autoclaving. Stability determinations car-
`
`Table V. Efi‘ect of Additives on the Rate of Lidocaine Degradation
`at 100°C”
`
`
`Additive
`Amount
`kw, (sec'l)
`
`9.85 x 10"
`0.15 M"
`Acetate buffer (p. = 0.15)
`1.06 x 10‘”
`0.009 M”
`Acetate buffer (p. = 0.15)
`8.67 x 10“
`20% (v/v)
`Propylene glycol/water
`3.34 x 10‘7
`80% (v/v)
`Prepylene glycol/water
`8.00 x 10”7
`10 ppm each
`FeClleuCl,
`1.69 x 10"
`50 ppm each
`FeC12/CuClz
`
`
`—None 5.65 x 10“BW
`
`‘ The pH of all solutions (except for acetate buffer, pH 4.2) was
`determined at room temperature and found to be pH 7.6.
`b Tbtal buffer concentration.
`
`Page 4
`
`Page 4
`
`