In the Classroom
“As Simple as Possible, but Not Simpler”— The Case of Dehydroascorbic Acid Robert C. Kerber Department of Chemistry, Stony Brook University, Long Island, NY 11794;
[email protected] The injunction to “make everything as simple as possible, but not simpler”, ascribed to Albert Einstein, characterizes a continuing dilemma of the teacher–scholar and of the textbook author. In presenting scientific material, the question becomes which simplifications are justified for the sake of clarity and which represent unacceptable distortions of reality? A related question is how quickly new knowledge should be incorporated into the canon. Is it more urgent or less urgent if it complicates a previously simple picture? Certainly in an age where new editions of textbooks regularly appear every three years or so, it is hard to justify delays of decades in incorporating new knowledge or understanding. But the case of dehydroascorbic acid (the oxidized form of ascorbic acid, vitamin C), invariably described in textbooks as a tricarbonyl compound, 1 (Figure 1), raises questions about our effectiveness in correcting old oversimplifications and incorporating new knowledge. Doubts were expressed about the credibility of structure 1 as long ago as the 1940s, and definitive results disproving that structure were published in the 1970s (vide infra), but still the structure persists in our textbooks. In the 19th century, chemistry “textbooks” were scholarly summaries of the state of contemporaneous knowledge, written by the foremost practitioners. As the body of knowledge grew, these evolved on the one hand into encyclopedias, compendia, and databases for experts and on the other hand, into schoolbooks for students. The latter were necessarily smaller (still not small enough, in many cases; ref 1) and less up-to-date, and they have become increasingly similar to one another as market forces have favored conformity in presentation and coverage. But this author finds it difficult to understand why that conformity should embrace misleading terminology (2) and erroneous
structures. The textbook structure of oxidized ascorbic acid provides a case in point. Dehydroascorbic Acid Ascorbic acid owes its biochemical and commercial importance to its effectiveness as a reducing agent. So the determination of the structure of ascorbic acid in the 1930s soon led to discussions of the structures of the one- and two-electron oxidized forms (3). The latter came to be referred to as “dehydroascorbic acid (DHA)” (4), a rather unsatisfactory name given its lack of acidity. Definitive structure determination was hindered by its facile hydrolysis and failure to form crystals, but the earliest structural suggestion (3) was for a dihydrate, 2. The absence of color in DHA was understood to be inconsistent with the diketolactone structure 1 by early workers (5), but nevertheless representations as 1 came to be common. In 1970, a preliminary communication described the crystal structure of DHA crystals grown from 95% acetic acid containing 0.01 M hydrochloric acid (6). The full paper was published in 1972 (7). The structure revealed was that of a symmetrical dimer, C12H12O12, 3, having a twist-boat dioxane ring assembled from the O2, C2 (a hemiacetal), and C3 (an acetal incorporating the primary alcohol group of the side chain) atoms of each dehydroascorbate unit. An analogous dimeric crystal structure was subsequently obtained for the tetraacetate derivative of 3 (8). Commercial DHA is composed of this dimeric solid, but the dimer is clearly the result of several chemical steps following the initial oxidation of ascorbic acid; a key step is the formation of the second five-membered ring, analogous to formation of furanose rings by carbohydrates. O
O O
2
O
1
3
4
O
O
H
6
HO HO HO
H
1
H
MeO H
OH
5
HO
O
HO HO
O
HO
O
OH
OH H
4
H
O
O
HO O
H
HO HO
O
O
H O
O
OH O
O
OH H
3 Figure 1. Structures associated with dehydroascorbic acid (DHA): textbook DHA, 1; dihydrate, 2; dimer, 3.
H N
O
N
2
H HO
Ar
H O
H
O O
HO
O
OH H
OH H HO
Ar = 4-BrC6H4, 2-HOC6H4 5
6 Figure 2. Structures associated with DHA: C2-benzylated, 4; hydrazones, 5; bicyclic monohydrate, 6.
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1237
In the Classroom
Dissolution of the dimer 3 in DMSO was thought to result in a solution of a monomer, as indicated by cryoscopy (9), but that result may have been perturbed by adventitious water. The infrared spectra in KBr or in DMSO showed no carbonyl group other than that of the lactone (ca. 1785 cm‒1) (10). A crystallographic indication of a monomer structure showing a hemiacetal ring in lieu of a ketone carbonyl group was afforded by study of the product 4 (Figure 2), prepared by C2-benzylation of ascorbate anion, followed by reaction with methanolic HCl (11). On the other hand, crystal structures of two hydrazones 5 show monomeric structures, with neither hydration nor hemiacetal formation (12). This suggests that formation of carbonyl adducts (hydrates, hemiacetals, or acetals) of dehydroascorbic acid is only mildly favored and is reversed by electron donation from the anilino group of the hydrazone. When the dimer 3 is dissolved in aprotic solvents, rapid mutarotation is observed (13). Originally ascribed to equilibration of dimer and monomer, this predominantly reflects formation of
an alternate dimer, epimeric at C2, as shown by NMR spectroscopy. The less symmetrical epimer predominates in solution by 3/1 at room temperature, but only the more symmetrical epimer crystallizes (14). The dimers persist in anhydrous solvents, no monomers being apparent in NMR spectra. But dissolution of 3 in hydroxylic solvents, which requires heating, forms a solution of monomeric, solvated DHA. Fresh solutions can be reduced back to ascorbic acid in moderate to high yield, and solutions made by dissolving 3 give spectra identical to those obtained by direct oxidation of ascorbic acid. The most useful tool for determining the structure of DHA in solution is NMR. Both 13C and 1H NMR have proven useful. With respect to the presence or absence of the hemiacetal ring formed (as in 3 and 4) by addition of the primary alcohol group of the side chain to C3, the 1H NMR spectrum provides a clear indicator, since the diastereotopic C6 protons have distinctly different chemical shifts in the cyclized forms, but generally not in the open forms. Table 1 shows proton NMR data for solutions
Table 1. 1H NMR Data for DHA and Reference Compounds
Compound
Solvent
H4
H5
H6a
H6b
J45
J56a
J56b
Rings
Ref
Ascorbic acid
D2O
4.99
4.10
3.79
3.79
2.0
6.6
6.6
J6a6b –
1
14
5 (p-BrPh hydrazone)
DMSO/CDCl3
4.92
4.00
3.54
3.54
0.9
7.0
7.0
–
1
14
1 (2,3-bis-DNPH)
aq DMSO
5.45
3.95
3.50
3.50
1.2
7.2
7.2
–-
1
15
6
D2O
4.71
4.58
4.17
4.27
0.9
2.6
5.7 –10.4
2
14
6 (2-OMe)
DMSO
4.53
4.41
4.10
4.25
0.7
3.5
5.8 –9.7
2
16
4
DMSO
3.93
4.17
3.72
4.20
0.0
4.5
6.4 –9.6
2
14
6 (2,2,3-triOMe-5-OAc) CDCl3
4.71
5.20
4.06
4.55
0.5
3.7
6.0 –10.5
2
17
Table 2 . 13C Chemical Shifts (δ) for DHA and Reference Compounds
Reported Structure
Probably
C1
C2
C3
2
7
7 (ref 4,18)
172.1
95.1
96.9
171.2
94.5
96.1
1
176.9
202.0
1 (5,6-acetonide)
168.0
202.8
1 (2,3-bis-DNPH)
163.8
--
--
80.4
72.3
Ascorbic acid
174.4
--
--
77.4
70.1
5
165.9
--
82.6
71.7
9
173.2
92.4
96.6
2 (5,6-acetonide)
173.9
91.3
96.0
84.6
2 (6-bromo)
173.6
91.3
96.0
83.5
68.2
2
6
175
95
108
90
73
77
3
6 (ref 18)
173.9
91.8
106.1
88.0
73.3
3 (solid, MAS) 3 (DMF, –50 °C)
175.2
95.4
107.0
93.8
169.1
92.3
106.5
90.8
3 (tetraacetate)
169.5
91.8
104.8
4 (DMSO)
172.8
109.3
10 (DMSO)
172.5
8 (H2O)
171.7
6 (aq DMSO)
173.8
6 (2-OMe)
171.4
a This
1238
-186.5
C5
C6
74.7
68.5
74.3
68.0
187.2
82.9
66.5
188.3
83.5
193.8
98.3a
C4
Rings
Ref
63.1
0
14, 19, 20
62.4
0
4
63.6
1
20
1
20
61.5
1
15
63.4
1
19
62.4
1
14
--
--
--
--
--
1
21
--
--
1
4
--
1
4
2(?)
22
77.1
2
19, 20
73.4
77.8
2
14, 21
73.7
76.8
2
14
87.8
74.0
74.9
2
8
86.9
73.8
74.7
2
14 20
85.9
73.2
75.7
2
108.8
79.6
73.8
75.3
2
23
92.1
106.4
88.6
74.0
76.2
2
4, 14,
95.2
106.8
89.1
74.8
77.0
2
16, 24
--
value reflects the perturbation by the C2 carbonly group.
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In the Classroom O
obtained by oxidizing ascorbic acid or dissolving 3 in hydroxylic solvents, along with reference compounds whose structures are unequivocal. The bicyclic structure 6 (or its alkoxy analogs from alcohol solvents) is indicated not only by the evident magnetic nonequivalence and deshielding of the C6 protons, but also by the small J45 coupling constant (the H−C4−C5−H dihedral angle in 4 is 99°; ref 11). Based on these criteria, the proton NMR data clearly support the retention of the fused furanose ring when the dimer 3 is dissolved in hydroxylic solvents to form the monomeric DHA, 6. The best evidence for the state of the putative carbonyl carbons C1−C3 is afforded by 13C NMR studies. A substantial body of 13C NMR data exists in the literature, but in some cases the same spectra have been assigned to different species, and reinterpretations of earlier assignments have appeared. A selection of available data is presented in Table 2. The carbon spectra provide several criteria for assignment or confirmation of structure. Ketone carbonyls in these compounds (Figures 1– 3) show chemical shifts of 186–203 ppm, while adducts (hydrates, hemiacetals, and acetals) fall in the range (100 ± 10) ppm. The C2 chemical shift is insensitive to the nature of the adduct, all chemical shifts being observed at (94 ± 2) ppm. However, C3 shows non-overlapping ranges for hydrates (97 ± 1), hemiacetals (106 ± 2), and acetals (109 ± 1) ppm. Unfortunately, this does not allow clear distinction between the dimer 3 and the monomeric product of its hydrolysis, 6. The carbon NMR data do allow clear discrimination between products with a furanose ring and those with a free dihydroxyethyl side chain. Formation of the ring is accompanied by deshielding of the resonances of C4 by 8 ppm, C5 by 4 ppm, and C6 by 13 ppm on average, providing clear evidence of cyclization. In 1977, Matusch (20) reported oxidation of ascorbic acid with iodine in dry DMSO that allowed 13C NMR observation (Table 2) of the consecutive formation of 1, which cyclized to 10, and then to the dimer 3. He also reported preparing an analog of 1 in which cyclization was prevented by protection of the 5,6-diol as an acetonide. These are the only reports of keto forms of DHA. As previously mentioned, 3 in dry DMSO undergoes
O
H HO
O
H
O
O O
O
O
OH O
O
H
H
MeO H MeO
H O
O
O O
H
HO HO
O H
HO O
OH H 8
9
OH H 10
Figure 3. Analogs of DHA: acetal, 8; dihydrate, 9; bicyclic monomer, 10.
epimerization at C2, but does not dissociate to 10 or 6 (Scheme I). The epimerization requires a partial dissociation, but further breakup of the dimer evidently requires stabilization of the free α-keto group in the intermediate by formation of an adduct with water or alcohol. In the presence of water or alcohols, di- or tricarbonyl compounds such as 1 and 10 are undetectable by NMR, carbonyl adducts being observed instead. Even the acyclic compound, 2,3-diketogulonic acid, the hydrolysis product of 1 or 2, is seen only as the dihydrate, 7 [CH2(OH)CH(OH)CH(OH) C(OH)2C(OH)2COOH, refs 4, 18]. Since spectra originally attributed to 2 appear to be due to 7, there is no unequivocal evidence of 2, although analogs such as 9 have been characterized. Semi-empirical molecular orbital calculations of the various forms of DHA gave the following relative energies: 1 + 2H2O: 0.0 kcal/mol; the 2-hydrate of 1 + H2O: ‒14.9 kcal/mol; 6 + H2O: ‒22.7 kcal/mol; the dihydrate 2: ‒29.8 (25). Although, according to these results, 2 is the most stable form, its susceptibility to lactone hydrolysis makes it too labile to be observed. A particularly informative set of experiments relating to DHA and analogs in hydroxylic solvents was reported by Kwon, Foote, and Khan (26), who were studying the photosensitized oxidation of ascorbic acid by singlet oxygen in deuterated methanol at temperatures below ‒60 °C. The low temperatures allowed
O OH O
O
OH H
O
O
H HO
HO
HO HO
O
O
H HO O
O
H H
O
HO H O
O
OH
OH
OH
H
O
O
O
O
O
OH H
3ł
3 H2O O
H HO
O OH
H O
OH OH
+
HO HO
O
HO
H O
OH H
O 6
6
Scheme I. Epimerization of dimer 3.
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1239
In the Classroom HO
HO
HO O
1O 2
O
HO
HO
H O
HO
O
O
O
O
HO
O
HO
ź80 °C CD3 OD
OH
HO
O
O
O O
H
O O
O
HO
O
O
O H
HO
OMe
O
HO
OH
O
O
CD3 OD
1O 2
+ DMS
ź80 °C
− DMSO
O
O O
O
O
+ DMS
O
O
− DMSO
O
O
25 °C
OMe OOH
O
HO
H 3 Oá
O OMe OH
O
O HO
OH
OH
6
O
O O
O
HO
O
OH
1O 2
O
O
O
24 h
O
HO
ź78 °C CD3 OD
O
HO HOO
O O
H
O O
O
O OH OCD3
+ DMS
O
O
O
− DMSO
HO HO
OH OCD3
Scheme II. Reactions of singlet oxygen with ascorbic acid and derivatives.
detection of intermediates that would otherwise have been too transient. These intermediates were characterized by 13C NMR data and identified by comparison to earlier spectra, particularly those of Matusch (20). A summary of some of their results is shown in Scheme II. Complete conversion of the mixture of hydroperoxides initially formed from ascorbic acid and singlet oxygen (generated by use of Rose Bengal and filtered xenon lamp radiation at ‒85 °C) to the peroxy analog of 6 required about 35 hours at ‒78 °C. This has been extrapolated to a rate constant of about 10 s‒1 at 25 °C (18). All things considered, it appears that the most appropriate single structure for representing the complex mixture of interconverting structures that constitutes aqueous “dehydroascorbic acid” is the bicyclic hemiacetal hydrate 6 (4, 25, 27).
O
źH
OH HO
O
?
H
H źO
O
H
O
HO
H
Scheme III. Reaction to generate the radical-anion.
1240
H
HO OH
źO
O O
HO
ź
O
O
O
OH H
Partially Oxidized Ascorbate Intermediates Conversion of ascorbic acid to DHA requires loss of two electrons and two protons and, as discussed above, formation of a second ring. This is intrinsically a multistep process whose details must be relevant to the function of ascorbic acid as a biological or chemical reducing agent. The principal intermediate in the ascorbate–DHA conversion process is the free-radical species C6H6O6⋅−, sometimes referred to as “semidehydroascorbate”. A review summarizes the state of knowledge of this species up to 1980 (28). It had been found that the neutral radical C6H7O6⋅, which would result from simple one-electron transfer from ascorbate ion, the principal form of ascorbic acid (pKa 4.17; ref 23) at physiological pH, was a strong acid. UV and ESR spectra of the deprotonated radical-anion were therefore pH-independent over a broad range, and the initial oxidation process may be described as a hydrogen atom transfer from ascorbate (29). The ESR spectrum showed a moderate coupling (1.76 G) to H4 and small couplings (0.19 and 0.07 G) to H6 and H5, respectively (30). The larger coupling to H6 and the IR spectrum of the radical-anion led some researchers to propose a cyclized structure for the radical species (31) (Scheme III). In support of this direct involvement of the side chain is the finding that ascorbic acid-5,6-acetonide, in contrast to ascorbic acid, undergoes only partial oxidation by iodine in DMSO (20).
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In the Classroom
Notwithstanding this interesting suggestion, the bicyclic radical-anion structure would presumably allow less electron delocalization than the monocyclic form. Indeed, recent density functional calculations have indicated that cyclization of the open C6H6O6⋅− radical ion would have an enthalpy of +20.0 kcal/mol (isolated molecule), compared to the neutral 1, which would cyclize with an enthalpy of ‒10.6 kcal/mol (32). Several other studies have applied theoretical methods to the structure of the C6H6O6⋅− intermediate. The experimental hyperfine coupling constants for C1−C4 and H4 have been satisfactorily reproduced by density functional calculations that incorporated hydrogen bonding to O2 and O3 (33). The spin density is greatest on the diketone moiety, O2−C2−C3−O3 (34). A moderately consistent set of bond lengths was obtained using these density functional calculations (33) and an earlier set of calculations using STO-3G MO calculations with a 4-31G basis set (35). The principal fate of the radical-anion is disproportionation to ascorbate and DHA. Kinetic study has shown that the disproportionation occurs through a dimer of unknown structure that reacts with proton donors to form the products (28, 36). Recent theoretical studies have led to the proposal that disproportionation of C6H6O6⋅− would be endothermic but for the stabilizing formation of the hemiacetal linkage in the structure 6 of DHA (32). In addition to these results that bear on the structural questions, the extensive biochemistry of ascorbate species remains an active area of research, but these studies lie outside the scope of this summary. What Does It Matter? It might be argued that the textbook structure 1 is intended as a symbolic structure rather than as a literal representation of the arrangement of atoms in the DHA molecule. This raises the question of how a student (or even a more experienced reader) is supposed to distinguish a symbolic structure from a structural structure. If chemical structures are the basic language of chemistry, an inaccurate or oversimplified symbolic structure is a hindrance to communication and possibly to scientific progress. An historic example of the consequences of misdescribed structures was related in The Double Helix (37). In his autobiographical account of the discovery of the structure of DNA in the 1950s, J. D. Watson described weeks of futile effort trying to make sense of the pairing of the purine and pyrimidine bases using textbook representations of their structures. The breakthrough came when crystallographer Jerry Donohue took exception to the conventional enol representations of guanine and thymine on which Watson had relied. In Watson’s words, My immediate retort that several other texts also pictured guanine and thymine in the enol form cut no ice with Jerry. Happily he let out that for years organic chemists had been arbitrarily favoring particular tautomeric forms over their alternatives on only the flimsiest of grounds. In fact, organicchemistry textbooks were littered with pictures of highly improbable tautomeric forms.
The use of more realistic keto structures for these nucleic acids quickly thereafter led to the recognition of the canonical base pairing scheme that forms the basis of molecular genetics.
Thankfully, modern-day textbooks have generally cleaned up their representations of the tautomeric structures of heterocycles. But, in addition to DHA, a number of oversimplified or inaccurate structural representations persist. These include the use of 19th century “dot structures” for the conjugate acids of organic amines (for example, C5H5N⋅HCl or “pyridine hydrochloride” for pyridinium chloride) and representation of amino acids at neutral pH as un-ionized molecules (H2NCHRCOOH instead of zwitterions, H3N+CHRCO2−) (38). The case for simplicity at the expense of clarity and accuracy in these examples is weak. Pedagogic Use of DHA Structure Even as we strive to make our textbooks more accurate and timely, it remains useful to remind students that incorporation of material into a textbook does not ensure its unimpeachability and that they should approach their reading with reasonable skepticism. We have constructed a workshop exercise for second-semester organic chemistry from the DHA example. It is assigned for the week when carbonyl addition products are first encountered in the organic course. It provides an interesting application of the material and a foreshadowing of carbohydrate chemistry, which comes later in the course. A copy of this group workshop (with answers) is provided in the online material. Literature Cited 1. Kerber, R. C. J. Chem. Educ. 1988, 65, 719. 2. Kerber, R. C. J. Chem. Educ. 2006, 83, 223–227. 3. Herbert, R. W.; Hirst, E. L.; Percival, E. G. V.; Reynolds, R. J. W.; Smith, F. J. Chem. Soc. 1933, 1270–1290. 4. Tolbert, B. M.; Ward, J. B. Dehydroascorbic Acid. In Ascorbic Acid: Chemistry, Metabolism, and Uses; Seib, P. A., Tolbert, B. M., Eds.; Advances in Chemistry Series 200; American Chemical Society: Washington, DC, 1982; pp 101–123. 5. (a) Kenyon, J.; Munro, N. J. Chem. Soc. 1948, 158–161. (b) Albers, H.; Müller, E.; Dietz, H. Z. Physiol. Chem. 1963, 334, 243–258. 6. Hvoslef, J. Acta Chem. Scand. 1970, 24, 2238–2239. 7. Hvoslef, J. Acta Cryst. 1972, B28, 916–923. 8. Hvoslef, J.; Hope, H.; Murray, B. D. Carbohydrate Res. 1986, 147, 11–19. 9. Teichmann, B.; Ziebarth, D. J. Prakt. Chem. 1966, 33, 124–129; for a contrary finding see Albers, H.; Müller, E.; Dietz, H. HoppeSeylers Z. Physiol. Chem. 1963, 334, 243–258. 10. Teichmann, B.; Ziebarth, D. Acta Chim. Acad. Sci. Hungar. 1966, 49, 311–318. 11. Hvoslef, J.; Nordenson, S. Acta Cryst. 1976, B32, 1665–1669. 12. (a) Hvoslef, J.; Nordenson, S. Acta Cryst. 1976, B32, 448–452. (b) Imai, T.; Yokoyama, Y.; Sekine, A.; Uekusa, H.; Ohashi, Y. Acta Cryst Sect. 1, Structure Reports Online 2003, 59 (9), 1241–1243. 13. Müller-Mulot, W. Hoppe-Seylers Z. Physiol. Chem. 1970, 351, 56–60. 14. Hvoslef, J.; Pedersen, B. Acta Chem. Scand. 1979, B33, 503–511. 15. Kishida, E.; Nishimoto, Y.; Kojo, S.; Anal. Chem. 1992, 64, 1505–1507. 16. Hvoslef, J.; Pedersen, B. Acta Chem. Scand. 1980, B34, 285–288.
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29.
Egge, H. Tetrahedron Lett. 1969, 10, 801–803. Sisley, M. J.; Jordan, R. B. Inorg. Chem. 1992, 31, 2137–2143. Berger, S. Tetrahedron 1977, 33, 1587–1589. Matusch, R. Z. Naturforsch. 1977, B32, 562–568. Schmidt, M.; Albert, K.; Brindle, R.; Maichle-Mössmer, C.; Eger, K. Z. Naturforsch. 1993, B48, 1372–1380. Kang, S.-O.; Sapper, H.; Lohmann, W. Z. Naturforsch. 1982, C37, 1064–1069. Lu, P.-W.; Lillard, D. W., Jr.; Seib, P. A.; Kramer, K. J.; Liang, Y.-T. J. Agric. Food Chem. 1984, 32, 21–28. Sawai, Y.; Moon, J.-H. J. Agric. Food Chem. 2000, 48, 6247–6253. Kurata, T.; Nishikawa, Y. Biosci. Biotechnol. Biochem. 2000, 64, 1651–1655. Kwon, B.-M.; Foote, C. S.; Khan, S. I. J. Am. Chem. Soc. 1989, 111, 1854–1860. Kwon, B.-M.; Foote, C. S. J. Am. Chem. Soc. 1988, 110, 6582–6583. Pfeilsticker, K.; Marx, F.; Bockisch, M. Carbohydrate Res. 1975, 45, 269–274. Bielski, B. H. J. Chemistry of Ascorbic Acid Radicals. In Ascorbic Acid: Chemistry, Metabolism, and Uses; Seib, P. A., Tolbert, B. M., Eds.; Advances in Chemistry Series 200; American Chemical Society: Washington, DC, 1982; pp 81–100. (a) Sawyer, D. T.; Chiericato, G., Jr.; Tsuchiya, T. J. Am. Chem. Soc. 1982, 104, 6273–6278. (b) Njus, D.; Jalukar, V.; Zu, J.; Kelley, P. M. Am. J. Clin. Nutr. 1991, 54, 1179S–1183S. (c) Njus, D.; Kelley, P. M. Biochem. Biophys. Acta 1993, 1144, 235–240.
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30. Laroff, G. P.; Fessenden, R. W.; Schuler, R. H. J. Am. Chem. Soc. 1972, 94, 9062–9073. 31. Sapper, H.; Pleyer-Weber, A.; Lohmann, W. Z. Naturforsch. 1981, C37, 129–132. Lohmann, W.; Pagel, D.; Penka, V. Eur. J. Biochem. 1984, 138, 479–480. 32. DiLabio, G. A.; Wright, J. S. Free Radical Biol. Med. 2000, 29, 480–485. 33. O’Malley, P. J. J. Phys. Chem. B 2001, 105, 11290–11293. 34. Costanzo, F.; Sulpizi, M.; Vandevondele, J.; Della Valle, R. G.; Sprik, M. Molecular Physics 2007, 105, 17–23. 35. Abe, Y.; Okada, S.; Horii, H.; Taniguchi, S.; Yamabe, S. J. Chem. Soc., Perkin Trans II 1987, 715–720. 36. Bielski, B. H. J.; Allen, A. O.; Schwarz, H. A. J. Am. Chem. Soc. 1981, 103, 3516–3518. Nishikawa, Y.; Kurata, T. Biosci. Biotechnol. Biochem. 2000, 64, 476–483. 37. Watson, J. D. The Double Helix; Atheneum: New York, 1968; Chapters 25, 26. 38. Coleman, W. F. J. Chem. Educ. 2006, 83, 1103
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2008/Sep/abs1237.html Abstract and keywords Full text (PDF) Links to cited JCE articles Supplement Carbonyl addition reactions workshop exercise
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