234

The Inhibitory Effect of Water on the C o 2 + and C u 2 + Catalyzed Decomposition of Methyl Linoleate Hydroperoxides 1 H. Chen 2, D.J. Lee* and E.G. Schanus 3 Department of Food Science and Human Nutrition, Washington State University, Pullman, Washington 99164-6376

The inhibitory effect of water on the decomposition of m e t h y l linoleate hydroperoxides CMLHP) catalyzed by Co 2+ and Cu 2+ w a s studied in a model s y s t e m using p r o ton nuclear magnetic resonance (NMR) spectroscopy. M L H P were prepared by photoxidation and purified by chromatographic methods. Proton N M R spectroscopy was used to measure reaction rates by monitoring changes in the intensity of the O O H signal. The rate c o n s t a n t of the reaction was obtained by plotting the natural logarithm of M L H P concentration vs time. In the first part of the study, no transition metals were added to the model system, so that the effect of water could be attributed to the interaction between water and M L H P only. The rate constant of the reaction IK) was found inversely proportional to the concentration of water. There w a s a downfield chemical shift of both hydroperoxide and water peaks in the N M R spectra when water w a s added. A s temperature increased to 40~ the difference in K between the s y s t e m s with 0% and 2% water disappeared. It is proposed that the hydroperoxides were solvated with water which retarded their decomposition. When Co 2+ w a s added to the model system, K decreased as the concentration of water increased from 0% to 1.5%. A s temperature increased from 18~ to 40~ differencesbetween the K for 0 % and 2 % water disappeared. A similar phenomenon was observed in reactions catalyzed with Cu 2+. These findings would support a mechanism in which the protective effect of water involves both the solvation of O O H and hydration of the metal catalyst. Lipids 27, 234-239 (1992).

Lipid autoxidation, a free radical chain reaction, is a recognized problem in preserving food stuffs and involves unsaturated lipids reacting with oxygen. The mechanism of lipid autoxidation has been studied by several research groups and reviewed (1-4). The overall reaction consists of four major stages: initiation, propagation, chain branching and termination. Decomposition of lipid peroxides plays an important role in autoxidation of unsaturated lipids. First, during the propagation stage, hydroperoxide decomposition occurs with aikoxy and/or peroxyl radicals as reaction 1Based on a paper presented at the Symposiumon Metals and Lipid Oxidation,held at the AOCS Annual Meetingin Baltimore, MD, April 1990. 2Current address: Department of FoodScienceand Technology,109 FS&T Building, University of California-Davis, Davis, CA 95616-0488. *To whom correspondence should be addressed. 3Current address: Bioproduct, Inc., Warrenton, OR 97146. Abbreviations: acetone-d6, fully deuterated acetone; HPLC, highperformance liquid chromatography; K, apparent rate constant; MLHP, methyl linoleate hydroperoxides;NMR, nuclear magnetic r e s o n a n c e ; r 2, linear regression correlation coefficient;UV, ultraviolet. LIP!DS, Vol. 27, no. 3 (1992)

products being formed. These radicals in turn abstract hydrogens from fatty acids to form more hydroperoxides and perpetuate the chain reaction. Chain propagation accelerates autoxidation because the reactions of alkoxy and peroxyl radicals with lipid molecules are faster than ab initio formation of free radicals by light, heat, or other energy sources. Second, the volatile breakdown products from the reaction are directly responsible for undesirable flavors in rancid foods. The autoxidation of lipids and the generation, synthesis and decomposition of lipid peroxides have been studied extensively during the past two decades (5-17}. Transition metals, such as iron, copper and cobalt, are found in many foods and are important elements in the decomposition of lipid peroxides. Because of their unpaired electronCs) in 3d or 4s orbital, transition metals readily lose or gain an electron so they serve as excellent catalysts of the reaction. The role of transition metals in the decomposition of lipid peroxides has been investigated and reviewed {18-22}. Water retards lipid autoxidation in many dehydrated and low moisture foods (23,24}. The shelf life of oat flakes containing 2-6% water was less than two weeks compared with eight months for flakes with 10% moisture {25). Several hypotheses have been offered to explain this inhibitory effect {26-28). The presence of water could slow autoxidation by the following mechanisms: (a) lowering diffusion rate of oxygen; (b} lowering effectiveness of metal catalysts; (c) promoting non-enzymatic browning which produces antioxidants; and {d) excluding air from the surface of food. Because of the complexity of food systems, however, the effect of water observed in the studies cited above does not support or rule out any single one of these possibilities. The effect of water on lipid autoxidation has been investigated in model systems using methyl linoleate as substrate {29-32}. Based on systematic studies, Karel C33) suggested the following events occurred during autoxidation which retarded degradation of lipid peroxides: Ca) hydrogen-bonding of amphipolar hydroperoxides at the lipidwater interface lowered the total effective hydroperoxide concentration; and (b) hydration of metal reduced the effectiveness of the catalysts. This study was undertaken to provide data to validate earlier hypotheses and to increase our understanding of how water interacts with hydroperoxides and transition metals to retard lipid autoxidation. A new model system was developed that contained methyl linoleate hydroperoxides {MLHP), transition metals and various amounts of water in fully deuterated acetone (acetone-d6). Proton nuclear magnetic resonance (NMR} was used to monitor the reaction by measuring the area of the OOH peak. The area of this peak is proportional to the concentration of MLHP in the system, and thus changes as decomposition occurring can be measured. The inhibitory effect of water was investigated by: Ca) studying the interaction between water and lipid peroxides; and (b) investigating the interaction between water and transition metals.

235

LIPID AUTOXIDATION WITH WATER MATERIALS AND METHODS

Materials. The following chemicals were used to prepare and purify M L H P : methyl linoleate (NuChek Prep, Elysian, MN); rose bengal (cert.) (Eastman Kodak, Rochester, NY); silicic acid (n-hydrate, J.T. Baker Chemical Co., Phillipsburg, N J), hexane and methanol. Fully deuterated acetone (acetone-de) (Wilmand Glass Co., Buena, NJ) was used as solvent in the model system. Cobalt(II)chloride, COC12"6H20 (crystal) and Cu(II)chloride, CuC12-2H20 (crystal) (J.T. Baker) were used as catalysts. Preparation of MLHP. The photoxidation method reported by Schenk and Schults-Elte (34) was modified slightly to synthesize MLHP. Three grams of methyl linoleate, 0.05 g rose bengal and 70 m L of methanol were mixed in a 150-mL Erlenmeyer flask. The mixture was transferred to a 100-mL fritted glass funnel connected to an oxygen tank. Two 300-W spotlights were used as an energy source for the photochemical reaction. Temperature was maintained at 2~ using a refrigerated water b a t h covered with a glass baking dish containing water to act as an infrared light filter. Water was added to compensate for evaporation during the 18-hr reaction. When the reaction was stopped, the methanol was removed under vacuum in a rotary evaporator. Purification of M L H P was done by column chromatography (35). Silicic acid (70 g) was activated in a l l 0 ~ oven for 12 hr, cooled in a desiccator, transferred to a 500-mL beaker and mixed with enough hexane to form a slurry. The slurry was packed into a 400 m m • 25 mm glass column and washed with 500 m L of hexane. Products from the photochemical reaction were loaded onto the column in a minimal volume of hexane/diethyl ether (15:5, v/v). Five hundred milliliters of hexane/diethyl ether (95:5, v/v) were used to elute non-polar compounds and unreacted methyl linoleat~ M L H P were eluted with 500 m L of hexane/diethyl ether (80:20, v/v). A drop-controlled fractionator was used to collect M L H P in disposable culture tubes. Thin-layer chromatography served to identify tubes containing M L H P (36). 1H NMR spectrometry was employed to confirm the presence of M L H P in combined fractions. High-performance liquid chromatography (HPLC) was used to further purify MLHP. Ten microliters of the M L H P mixture were loaded onto a silica gel semipreparative column (10 mm • 300 mm). The elution solvent was 1.2% isopropanol and 98.8% hexane (v/v). The flow rate of the solvent was 4 mL/min. The wavelength of the ultraviolet (UV) detector was set at 245 nm. Solvent in the collected M L H P fraction was evaporated with a r o t a r y evaporator. The purified M L H P was transferred to small glass vessels, flushed with nitrogen, and stored at -22~ Preparation of model system. The model system was prepared in an NMR tube using 50 ~L of acetone-d6 as solvent. To investigate the interactions between M L H P and water, the solvent was replaced with various amounts of distilled water (v/v) to arrive at a final concentration of water of 0%, 0.5%, 1.0%, 1.5%, or 2.0%. Interactions between transition metals and water were studied by replacing a like volume of acetone-d6 in the system with a solution of cobalt(II)chloride or copper(II)chloride dissolved in acetone-d 6. Experimentalprocedure. The proton NMR spectrum of M L H P shown in Figure 1 was recorded on a J E O L FX90Q FT-NMR spectrometer (Tokyo, Japan) operating at 89.90

TMS

i 11

.

L 10

.

, 9

.

i . . fl

i 7

.

i 6

.

i

.

5

f 4

2

!

Chemical Shift (ppm)

FIG. 1. Proton NMR spectrum of MLHP (89.90 MHz). Peak of OOH appeared at 10.40 ppm. Peak at 3.65 ppm was assigned to OCH 3 and was used as reference for the measurement of the peak area of OOH.

MHz. Complete identification and interpretation of the spectra of methyl linoleate and M L H P have been previously reported by Gunstone and Norris (37), Chan (3), Chan and Levette {15}, and Frankel et al. (8). The OOH peak appeared downfield at 10.40 ppm. The peak at 3.65 ppm can be assigned to OCH 3, and was chosen as a reference peak for measuring OOH peak areas. The area for OOH protons was integrated using the operating disk FAFT72 in the J E O L computer. Since the initial concentrations of OOH were known, the areas of OOH at any given time could be converted into molar concentrations. For each data point, three readings were taken. The average standard deviation was 0.011. All measurements were conducted at 18~ unless specified otherwise. To determine the apparent rate constant for the reaction, K, the natural logarithm of the concentration of M L H P was plotted against t i m e The slope of the straight line is K. All statistical analyses, such as linear regression correlation coefficient (r2) and confidence level, were carried out using the statistical program, MSUSTAT, on an IBM computer. RESULTS AND DISCUSSION

For convenience of discussion, our results are presented in two parts: (i) solvation of M L H P ; and (ii) hydration of metals. In the first part, transition metals were not involved so t h a t the effect of water could be attributed to the solvation of M L H P only. In the second part, water, M L H P and metals were present in the same system; so the effect of water could be attributed to solvation of M L H P and to hydration of metals. Solvation of MLHP. The apparent rate constant, K, for the decomposition of MLHP, as affected by various concentrations of water, is shown in Figure 2. K decreased linearly as the concentration of water increased, indicating t h a t water acts as an inhibitor of the reaction. A downfield shift of the OOH and water signals in the NMR spectra was observed when water was added to MLHP. The extent of the shift was dependent on the concentration of water in the system {Fig. 3 and Table 1), The LIPIDS, Vol. 27, no. 3 (1992)

236

H. CHEN E T AL.

O.O3.

10.8 10.7 -

0.02 9

__R

0,01'

_u E

10.5-

o

109 -

== 10.3

0.00

0',5

0.0

110

Conc.

of

115 Water

(%,

210

215

10

9

8

7

6

Cb~nical

5

Shift

4

3

2

I

0

(p~m)

FIG. 3. Proton N M R spectrum of M L H P (89.90 MHz). The model system contained 2% H20. Peak for OOH was shifted from 10.40 ppm to 10.65 ppm (Fig. 1). Peak for water appeared at 3.25 ppm.

TABLE 1 Change in Chemical Shift of NMR Peak of Water After Mixing with MLHP Conc. of H20 (%, v/v) 0.25 0.50 0.75 1.00 1.50 2.00

Chemical shift (ppm)a Before mixing After mixing 2.85 2.90 2.93 2.95 3.00 3.10

2.90 2.95 3.00 3.10 3.15 3.25

,

9

,

0,5 1.0 Cone. of

9

,

1.5 Water

.

,

9

,

.

2.0 2.5 (%, v/v)

I

3.0

FIG. 4. Change in chemical shift of N M R peak of OOH v s increased concentration of water. Standard deviations of the measurements are shown by the error bar.

d6 was the cause. When the chemical shift was plotted against water concentration (Fig. 4), the curve obtained was s o m e w h a t bimodal. Water concentration above 1.0% caused less change per unit increase t h a n at concentrations below 1.0%. This finding s u p p o r t s the m e c h a n i s m proposed by Karel (4) and by Barclay and Ingold (5) t h a t polar hydroperoxides form water-peroxide complexes which alter the p a t h of the reaction. This would be expected if hydrogen bonding occurred between the peroxide group and two molecules of water. The molar concentration corresponding to 1% water in the s y s t e m is approximately twice t h a t of M L H P in the system. The addition of more water would not change the equilibrium of the reaction toward the bound s t a t e as rapidly as when water concentrations were lower and more binding sites on the hydroperoxide molecules were free. The preferential bonding between water and one of the isomers in the mixtttre could also account for the observed non-linear response to increased water concentrations. However, the fact t h a t the oxidation rate and response to water of the 13-hydroperoxy-cis-9, trans-ll-octadecadienoate isomer, separated by using the m e t h o d by Chan and Levett {15), was nearly identical to t h a t of the M L H P mixture does not support this hypothesis. The chemical shift of the N M R p e a k for water was observed to revert b a c k toward initial values as the reaction progressed and the concentration of hydroperoxide decreased {Table 2). The formation and subsequent decay of water-OOH hydrogen

Change 0.05 0.05 0.07 0.15 0.15 0.15

aStandard deviation range of the measurements was +_0.007 to 0.014.

observed change in chemical shift can be explained by interaction between the O O H protons and water molecules. H y d r o g e n bonding would be expected to alter shielding of these protons which would result in the observed downfield shifts. This explanation is s u p p o r t e d by the nonlinear change in chemical shift with increasing a m o u n t s of water in the s y s t e m (Fig. 4), in c o n t r a s t to the linear change t h a t would be expected if replacement of acetoneLIPIDS, Vol. 27, no. 3 (1992)

.-.

0.0

v/v)

FIG. 2. Plot of apparent rate constant (K) v s concentration of water in M L H P decomposition in the absence of metals. Duplicate experiments were carried out (r2 = 0.94, confidence level >99%}.

11

10.6 -

TABLE 2 N M R Chemical Shifts of Water and Hydroperoxide Peaks in Various S y s t e m s

Chemical shift (ppm)a System Acetone-d 6 + Acetone-d 6 + (100 ppm) Acetone~ds + Acetone~d6 + Acetone-d s + Acetone-d 6 + (100 ppm)

OOH

2% H20 2% H20 + Co2+

H20 3.15 3.85

Co2+ (100 ppm) + MLHP MLHP 2% H20 + MLHP 2% H20 -F MLHP + Co2+

10.4 10.4 10.6

3.20

10.6

3.40

aStandard deviation range of the measurements was + 0.007 to 0.014.

237

LIPID AUTOXIDATION WITH WATER

-2.5 -3.0 -3.5 -4.0 -4.5 -5.0 -5.5 3.1

3.2

3.3 lIT,

X

8:5

3:4

10E-3

(K)

FIG. 5. Plot of K v s temperature for reactions in the absence of metals: O, 0% H20; O, 2% H20. Standard deviation of the measurements are shown by the error bar.

bonds could explain this observation. The results support the proposals of Karel (4) and Barclay and Ingold (5) that polar hydroperoxides form water-peroxide complexes which slow their decomposition. To further support our findings, the effect of temperature on the reaction with 0% and 2.0% water was investigated. At 18~ the apparent rate constant (K) for ML H P decomposition in the absence of water was twice that of the one with 2.0% water {Fig. 5). The reaction rate of the system with water was more temperature~dependent than the one without water, and the difference in K disappeared at 30~ and the plotted lines met at 40~ A possible explanation of this is that at 18~ most of the peroxide moieties were complexed with water, as indicated by the chemical shifts of the water and peroxide peaks. As temperature increased, the hydrogen bonds were weakened and the equilibrium shifted toward more free water and more free peroxide moieties. This would explain the near-linear temperature response of the 2% water system. Unfortunately, neither the NMR signals for water at different temperatures, nor the effect of temperature at lower water concentrations have been determined so far; so this explanation remains tentative. The influence of temperature on hydrogen bonding has been studied previously. Luck (38) reported that as temperature increased, hydrogen bonding between water and alcohol decreased; but this observation was made over a much wider temperature range than used in the present study. Hydration of metals. The effect of water on the reaction catalyzed by Co 2+ (Fig. 6) was attributed to both solvation and hydration. Adding 0.5% water to the model

system reduced K by 55%. Water concentrations above 0.5% had decreasing effects on reaction rates, which were reduced by 67% and 80% by 1.0% and 1.5% water, respectively. No additional inhibitory effect was found when 2.0% water was added to the system. The supposition that hydration of the metal ion reduces the reaction rate is supported by these results. When 100 ppm Co 2+ was added to a system containing acetoned6, 2% water and no MLHP, a chemical shift in the NMR spectrum of the water peak to 3.85 was observed, This indicated the water was complexing with the metal ion. Decomposition of hydroperoxides catalyzed by metal ions as proposed by Chan (3) involves the exchange of electrons; a) ROOH + M In§ ~ ROO* + H + + M n+, and; b) ROOH + M n+ RO* + O H - + M In+l). Hydration of metals can reduce their effectiveness as catalysts {19}. Pokorny (18) reported that transition metal ions, such as Co 2+, are normally surrounded by water molecules and can form polymeric micelles. This can inhibit electron transfer from the outer layer of the metal to the oxidizing form of molecule. Increasing amounts of water in the model system could increase the degree of hydration with subsequent reduction in reaction rates. At 1.5% water, the molar concentration of added water is one-half that of the Co2+ ion and is the same when maximum inhibition occurs. This supports the formation of a complex containing two cobalt ions per water molecule. Experiments to determine the effects of temperature on reactions catalyzed by Co2+ and Cu 2+ with 0% and 2.0% water in the system were conducted. Results {Fig. 7) were similar to those obtained with a non-catalyzed system {Fig. 5). At 18~ there were differences in the rate constants between the systems with and without water. The

030 9 A

025 020.c

0.15 9 0.10" 0.05' 0.00 10

2'0

3'0

Temperature

0.10 -

5O

4'0 (oC)

B

0.08 0.30"

0.06 -

0.25'

0.04 -

0.20" 0.02 5.,

0.150.00

0.100.05

10

2~0

"

3'0

Temperature

'

4'0

9

5'0

(oC)

-

0.00 0.0

0.5

Conc.

1.0

1.5

of Water (%,

2.0

2.5

v/v)

FIG. 6. Change in K v s concentrations of water in systems containing 100 ppm Co 2+. Standard deviations are shown by error bars.

FIG. 7. Plot of K v s temperature for reactions catalyzed by Co 2+ a n d C u 2 + .(A) O , 0 % H 2 O a n d C o 2 + ; 9 , 2% H2O and Co 2 + .(B) O, 0% H20 and Cu 2 + ; o , 2% H20 and Cu 2 + . Concentration of Co 2 + and Cu 2 + : 10 ppm. Standard deviation of the measurements are shown by the error bar.

LIPIDS, Vol. 27, no. 3 (1992)

238

H. CHEN E T AL. s y s t e m s with 2.0% water were more temperature-dependent and with b o t h catalysts the K of the reactions increased faster t h a n in s y s t e m s with no water. Near 30 ~C, or slightly above, the rates in the Co 2+ catalyzed reactions were nearly the s a m e for b o t h systems. This effect of temperature on the reactions would be expected if there is an interaction of water and m e t a l which inhibits the reaction at 18 ~ b u t which is disrupted as t e m p e r a t u r e s increase. The observed differences between the reactions catalyzed with Co 2+ and Cu 2+ could be due to the fact t h a t they catalyze redox reactions by different mechanisms. Copper normally catalyzes via direct electron transfer while cobalt more easily forms complexes with polar compounds, such as oxygen and hydroperoxides. The differences in heat of hydration could also account for p a r t of the observed differences. Hydration heat of Cu 2+ is approximately 11 kcaYmole greater t h a n t h a t for Co 2+ (39). The d a t a obtained in this s t u d y do not enable us to determine which, if either, explanation is correct, b u t the addition of water inhibited the reactions catalyzed with copper as well as those catalyzed with cobalt. Transition metals do not exist in the free s t a t e b u t are always surrounded by solvents, such as water, and oxidation products, such as lipid peroxides, to form polymeric micelles (18). The complex formed can stop or slow the electron transfer between the metals and the organic s u b s t r a t a Hill and McAuley (21) have described how such a complex would interfere with the oxidative decarboxylation of carboxylic acid. Waters (20) indicated t h a t the ease of electron transfer depends on the equilibrium cons t a n t for the ligand-displacement process. A MLHP-metalwater complex m a y actually have been formed in the model s y s t e m during M L H P degradation. Indirect evidence for the formation of such a complex is found in the N M R d a t a presented in Tables 2 and 3. The d a t a presented in Table 2 indicate a) t h a t water was bound to Co 2+ through coordination bonding, which caused a downfield shift of the proton in the presence of water; b) t h a t there was no direct interaction between the peroxide O O H groups and Co 2+ since the chemical shift of the O O H proton did not change when the two were mixed in the absence of water; and, c) t h a t water inter-

TABLE 3 Change in Chemical Shift of N M R Peak of Water vs Time in Decomposition of M L H P Catalyzed by Co 2+

Time (hr) 0.00 0.25 1.00 2.00 2.50 5.00 6.00 8.00 9.00 10.0 24.0

0.5% H20 2.95 3.50 3.35 3.20

Chemical shift (ppm)a 1.0% H20 1.5% H20 3.00 3.25 3.10

3.10 3.30 3.30

3.25 3.40 3.35

3.10 3.20 3.00

2.0% H20

3.00

3.30 3.30

3.15 3.15

3.25

astandard deviations range of the measurements was +0.007 to 0.014. LIPIDS, Vol. 27, no. 3 (1992)

acted with M L H P and Co 2+ causing a downfield shift of b o t h the protons of water and the OOH groups. Two possible explanations for these findings are; i) Co 2+ was hydrated to form a water-metal complex and, in an independent reaction, water was bound to O O H groups via hydrogen bonding to form a water-hydroperoxide complex. The physical barrier these water complexes create would slow electron transfer from c a t a l y s t to s u b s t r a t e and inhibit the reaction; ii) previous studies and our experimental d a t a indicate the possibility t h a t Co 2+ was bound to water with water molecules surrounding the Co 2+ ion. H y d r o g e n bonding of the water molecules with the oxygens of the O O H group of M L H P occurred to form a m e t a l - w a t e r - M L H P complex. This is s u p p o r t e d by the chemical shift d a t a presented in Table 3. In the reaction with 0.5% water, the chemical shift of the water p e a k was 2.95 p p m before Co 2+ was added (0 hr). After mixing (0.25 hr), the p e a k was downshifted to 3.50 ppm. This is obviously due to a change in the chemical environment of the water protons, which could be caused by watercobalt complex forming hydrogen bonds with the peroxide O O H group to form a water-MLHP-Co 2+ complex. As M L H P decomposed, the a m o u n t of hydrogen bonding between water and M L H P lessened and the protons in water became more shielded causing the N M R water p e a k to shift back toward its original position. In our experiments, at the end of the reaction the chemical shift of the water p e a k was 3.00 ppm. Similar observations were m a d e at other water concentrations. The association between the change in the chemical shift of the water p e a k and the decomposition of M L H P can be explained b y a mechanism involving a m e t a l - w a t e r - M L H P complex. I t is difficult to explain the chemical shift returning to the original value as the reaction progressed if the binding of water with Co 2+ and hydroperoxides are separate events. The possible role of breakdown products remains to be determined. This s t u d y d e m o n s t r a t e d t h a t in a model s y s t e m the decomposition of M L H P was inhibited by water concentrations of 0.5% to 2.0%, and N M R spectral data provided evidence for water molecules complexing with hydroperoxides and binding with the m e t a l catalysts, Co 2+ and Cu 2+. We propose t h a t formation of a metal-water-hydroperoxide complex is involved in the inhibition by water. F u r t h e r studies are needed to verify these findings and their interpretation. ACKNOWLEDG M ENTS We thank the Collegeof Pharmacy and the Department of Chemistry at Washington State University for the use of their proton NMR spectrometer. We also thank Dr. J. Hunt and Dr. H. Dogden, Professors of Chemistry at Washington State University, for their helpful suggestions and criticisms during this study. This study was supported in part by the National Science Foundation, Grant Na CBT-88504518, and the Washington State University Agricultural Research Center, Project 0666. REFERENCES 1. Porter, N.A., and Dussault, P.H. (1988) in Free Radicals in Synthesis and Biology (Minisci, F., ed.) pp. 407-421, Kluwer Academic Publishers, London. 2. Ingold, K.U. (1963) inLipids and Their Oxidation (Schultz, H.W., and Day, E.A., eds.) pp. 93-121, AVI Publishing Company, Westport.

239 LIPID AUTOXIDATION WITH WATER 3. Chan, H.W.S. (1987)Autoxidation of Unsaturated Lipids (Chan, H.W.S., ed.) pp. 1-16, Academic Press, London. 4. Karel, M. (1980) in Autoxidation in Food and Biological Systems (Simic, E.M., and Karel, M., eds.) pp. 191-206, Plenum Press, New York. 5. Barclay, L.R.C., and Ingold, K.U. (1981) J. Am. Chem. Soc. 103, 6478-6485. 6. Ewing, J.C., Cosgrove, J.P., Giawalva, D.H., Church, D.F., and Pryor, W.A. (1989) Lipids 24, 609-615. 7. Neff, W.E., Frankel, E.N., and Miyashita, K. (1990) Lipids 25, 33-39. 8. Frankel, E.N., Neff, W.E., and Miyashita, K. (1990} Lipids 25, 40-47. 9. Miyashita, K., Frankel, E.N., Neff, W.E., and Awl, R.A. (1990} Lipids 25, 48-53. 10. Pryor, W.A., Prier, D.G., Lightsey, J.W., and Church, D.F. (1970) in Autoxidation in Food and Biological Systems (Simic, E.M., and Karel, M., eds.) pp. 1-16, Plenum Press, New York. 11. Hiatt, R. (1971) in Organic Peroxides (Swern, D., ed.) pp. 4-40, Wiley-Interscience, New York. 12. Porter, N.A., Funk, M.O., Gilmore, D., Isacc, R., and Nixon, J. (1976) J. Am. Chem. Soa 98, 6000-6005. 13. Porter, N.A., Lehman, L.S., Weber, B.A., and Smith, K.J. (1981) J. Am. Chem. Soa 103, 6447-6455. 14. FrankeL E.N., Neff, W.E., Rehwedder, W.IL, Khambay, B.P.S., Garwood, R.F., and Weedon, B.C.L.. (1977) Lipids 12, 901-907. 15. Chan, H.W.S., and Levett, G. (1977) Lipids 12, 99-104. 16. Frankel, E.N., Neff, W.E., Rohwedder, W.K., Khambay, B.ES., Garwood, R.F., and Weedon, B.C.L. (1977) Lipids 12, 1055-1061. 17. Kochi, J.K. (1973} in Free Radicals (Kochi, J.K., ed.) Vol. I, pp. 591-683, Wiley, New York. 18. Pokorny, J. (1987) in Autoxidation of Unsaturated Lipids (Chan, H.W.S., ed.) pp. 141-206, Academic Press, London. 19. Uri, N. (1961}inAutoxidation andAntioxidants (Lundberg, L.O., ed.) Vol. I, pp. 55-106, Interscience, New York. 20. Waters, W.A. (1971) J. Am. Oil Chem. Soc. 48, 427-433. 21. Hill, J., and McAuley, A. (1968) J. Chem. Soa (London) A 1, 169-1173.

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LIPIDS, Vol. 27, no. 3 (1992)