In the Classroom
Motivating Students in Sophomore Organic Chemistry by Examining Nature’s Way—Why Are Vitamins E and C Such Good Antioxidants? Bruce Beaver Department of Chemistry and Biochemistry, Duquesne University, Pittsburgh, PA 15282;
[email protected] Many articles in this Journal have pointed out that, from the students’ perspective, it is challenging to be an effective teacher in chemistry classes below the collegiate junior level (1). This same theme was cited in a recent review of undergraduate education in science, mathematics, engineering, and technology by an advisory committee to the National Science Foundation Directorate for Education and Human Resources (2). The NSF report prominently cited a study by Seymour and Hewitt (3) of 335 science, mathematics, and engineering (SME) students at seven campuses that were leaders in producing new SME baccalaureates. Half of these students majored in biology, physical sciences, and mathematics. All students were considered well prepared for their SME courses, having math SAT scores above 649. After 3 years of tracking, Seymour and Hewitt were able to identify two separate cohorts of students: a group that changed to non-SME majors and a group that continued as SME majors with some graduating during the study period. Among the former group, 90% reported that they felt the SME courses were poorly taught. Most surprisingly, 80% of the sample that graduated with SME degrees felt that faculty teaching performance in SME courses was poor. The NSF report states that “In their explanations for the poor teaching they had experienced, students constantly referenced faculty preoccupation with research as the overt reason for their failure to pay serious attention to the teaching of undergraduates and for specific inadequacies in attitude or technique.” Clearly, if the views expressed by the students surveyed for the above reports are representative of the national student body, undergraduate SME courses are in dire need of reform. The obvious question is how to reform these classes. It is probably best to start by reviewing what the students cite as indicators of good teaching. The NSF report indicates that students value openness, respect for students, the encouragement of discussion and the sense of discovering things together. Most interestingly, Seymour and Hewitt’s study found “Student condemnation of the faculty obsession with research changed dramatically, however, when students were allowed to observe or participate in that research.” I suggest that when exposed to faculty research interests, students sense at least two important things: faculty are more enthusiastic when discussing things that interest them and from this enthusiasm the students get some sense of the thrill of “discovery”. While it is unrealistic to expect the majority of the SME undergraduates to become engaged in research, it should be possible for faculty to share with their classes some of their excitement about research. For the past eight years I have used such a teaching technique for part of my sophomore organic class. This portion of the course, which I consider the capstone, is composed of 5 lectures in which we examine the antioxidant functions of vitamins E and C. In these lectures I am reinforcing such important concepts as structure– 1108
reactivity relationships, resonance and inductive effects, the Hammond postulate, solubility, and basic kinetics. In addition, these lectures introduce new concepts such as the importance of the relative rates of competing reactions in determining the course of reactivity, stereoelectronic effects, and the captodative effect. Students have repeatedly commented that this portion of the course is particularly interesting and have requested similar treatment of additional topics. Over the years discussions with students have led me to believe that a primary reason why many students find this topic interesting is that it illustrates how the basic principles of organic chemistry can actually enhance our appreciation for the wonders of the natural world. In fact, some educational theorists argue that we all possess an innate desire to understand the workings of our world, a desire that is most evident when we are toddlers (4). These theorists go on to assert that our educational system unwittingly stifles this curiosity through various “motivational” practices. In this article I summarize my basic treatment of the antioxidant properties of vitamins E and C and have included rather extensive references. It is my hope that interested faculty can use this article, and references where necessary, to share this inherently exciting topic with their classes. In addition, my experience has been that some of the more motivated students will actually consult some of these references to gain additional insights! Background The reaction of molecular oxygen (O2) with organic molecules under mild conditions1 is usually referred to as autoxidation. Substances that can limit the extent of autoxidation are usually referred to as antioxidants. Because of the unique reactivity of oxygen and the fact that the earth’s surface is continually bathed in an atmosphere composed of 21% oxygen, autoxidation is ubiquitous. All organic matter will eventually be degraded by autoxidation or light-assisted autoxidation. Some of the things for which autoxidation has been found to be particularly problematic are plastics, coatings, plants, food, fuels, and all living entities. That autoxidation affects humans (and other mammals) was apparently postulated at some point after 1956 when Denham Harman first proposed the so-called free-radical theory of aging (5, 6 ). It was hypothesized that aging is the result of free radicals produced during the normal course of metabolism, continuously and randomly damaging biopolymers. Eventually, the ability of these biopolymers to maintain homeostasis becomes so impaired that death ensues. Thus, the free-radical theory of aging predicts that if a mammal were to have an optimal antioxidant system, damage to biopolymers would be minimized, thus extending maximum life span.
Journal of Chemical Education • Vol. 76 No. 8 August 1999 • JChemEd.chem.wisc.edu
In the Classroom
However, after many years of research, it has been shown that mammalian maximum life span cannot be significantly increased with antioxidants. The good news from this research is that the mean life span for mammals can be increased. This means that a larger percentage of a particular population can achieve their maximum life span potential. In 1987, in the light of these results, William Pryor proposed a disease-specific freeradical theory of aging (7). This theory states that free radicals and free-radical reactions are involved in the etiology and development of many of the chronic diseases that contribute to shorten the maximum life span potential for a species. For humans these chronic diseases include emphysema, atherosclerosis, and cancer. In the 1980s the autoxidation of low-density lipoprotein (LDL) was strongly implicated as an initiator of atherosclerosis in humans. LDL is the major cholesterol-bearing protein in blood plasma, and autoxidation of the lipid portion of this protein has been proposed to be the initiating event that alters LDL structure into a form that ultimately deposits on arterial walls (8). During the last decade, great increases have been achieved in our understanding of the important role played by antioxidants in the maintenance of optimal health. Clearly, two important biological antioxidants are vitamins E and C. To develop an appreciation of the remarkable aptitude of these molecules as antioxidants, the details of the mechanism of autoxidation will be discussed. A Mechanism for Autoxidation To provide a basic understanding of the mechanism of autoxidation we will examine a biologically important example of this process, the conversion of linoleic acid into its isomeric hydroperoxides (9). COOH
OOH (CH2)6COOH H11C 5
+
HOO
(CH2)6COOH C5H11
In order to understand the mechanistic details of this process, we must first appreciate the uniqueness of molecular oxygen. The ground state of oxygen is a triplet, which means that the electronic configuration of molecular oxygen has two unpaired electrons. The triplet state of molecular oxygen affects its chemical reactivity in at least two ways: First, like most molecules with an unpaired electron, oxygen reacts rapidly (rate constant ~108 M {1 s{1) with most carbon-centered free radicals. This is because an exothermic process, formation of a carbon–oxygen bond, can occur with very little activation energy. Second, under mild conditions oxygen generally does not react with organic molecules in the singlet state (i.e., with
paired electrons). Consequently, most organic molecules must be transformed into radical species (initiation) before they can react with molecular oxygen. These two facets of the reactivity of molecular oxygen are illustrated in steps 1–4 in Scheme I, which depicts a generic mechanism for autoxidation. k1 k2 k3 R? + InOOH In → In? → InO2? → – O2
R H
(Steps 1–3)
k
4 R? + O2 → ROO?
(Step 4)
k5
ROO? + R– H → ROOH + R? k6
2ROO? → non-radical products k
7 ROO? + ArOH → ROOH + ArO?
k8
ROO? + ArO? → non-radical products
(Step 5) (Step 6) (Step 7) (Step 8)
Scheme I
In step 1, some event transforms a generic initiator molecule (In) into a radical species. For instance, initiation of the autoxidation of linoleic acid in benzene only requires gentle heating of a small amount of 2,2′-azobis(2,4-dimethylvaleronitrile). Thermolysis of the weak C–N bonds produces two alkyl radicals and molecular nitrogen. heat
R–N=N–R → 2 R? + N2 ? R? = (CH 3)2CHCH2C(CH3)CN
In step 2, a portion of the initiator radicals rapidly react with O2 to form a peroxyl radical. In step 3 these peroxyl radicals abstract a hydrogen atom from linoleic acid to yield an initiator hydroperoxide and a linoleic acid radical. A rapid reaction between O2 and the linoleic acid radical yields the linoleic peroxyl radical in step 4. In step 5 the linoleic peroxyl radical abstracts a hydrogen atom from another linoleic acid molecule to yield linoleic hydroperoxide and another carboncentered free radical. This process sets up a chain reaction; that is, steps 4 and 5 repeat themselves many times. Autoxidation occurs even in the absence of added initiators. The origin of the initiation for such “non-initiated” autoxidations remains mysterious. This is one of several fundamental aspects of autoxidation that still need to be resolved (10, 11). Nevertheless, in vivo, it is agreed that many of the radical species produced during the normal course of metabolism can initiate biological autoxidation (12). It is very important to appreciate why the mechanism for autoxidation is generally a chain mechanism. First, because of the facility with which O2 reacts with most carbon-centered radicals (i.e., large k4), they are rapidly converted into peroxyl radicals. Second, the rate of hydrogen atom abstraction (step 5) by these peroxyl radicals propagates the chain competitively with reactions that terminate the chain, such as the coupling of two peroxyl radical in step 6. In addition, it is important to appreciate that the peroxyl radical that forms is relatively stable and is therefore quite a selective oxidant. This relationship between radical reactivity and selectivity mirrors the situation for radical initiated halogenation of alkanes. Just as a bromine atom is much more selective than a chlorine atom in hydrogen atom abstractions,2 so the peroxyl radical is more selective than other oxygen radical species (e.g., hydroxyl and alkoxyl radicals). By far, the
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weakest C–H bond in linoleic acid is one of the bis allylic bonds in the molecule. In step 5 one of these bonds is broken in the rate-determining step for the autoxidation. Experimentally k5 has been found to be 62 M{1 s{1 (13, 14). The significance of the chain nature of the autoxidation process has two important consequences. First, because of the chain nature of steps 4 and 5, a rare initiation event can produce a large amount of oxidation. The chain continues until a termination event occurs, such as step 6, which stops (breaks) the chain. Once termination has occurred another initiation event is needed to restart the chain mechanism. Second, a corollary of the chain nature of the autoxidation mechanism is that not much of an antioxidant is needed to inhibit the process. An effective antioxidant requires preferential reaction with peroxyl radicals and also production of reaction products that do not restart the chain. Many antioxidants are phenols. They rapidly donate the phenolic hydrogen atom to the peroxyl radical to yield a hydroperoxide and a resonancestabilized phenoxyl radical as in step 7 (Ar = aryl group), Scheme I. As shown in step 8, the phenoxyl radical can react with a second peroxyl radical to form non-radical products that do not restart the chain. Thus, in the best-case scenario, one molecule of antioxidant can inhibit much autoxidation by inhibiting the formation of two chains.
H2O
H2O
organic phase
polar head
non polar tail
Vitamin E
Representative phospholipid molecule
O O
O
O O-P(O)-O(CH2)2N(CH3)3
Figure 1. Model phospholipid bilayer showing location of vitamin E.
Where Are Vitamins E and C Located in the Human Body? One of the primary components of fat tissue in humans is phosphorus esters of linoleic and arachidonic acid.3 For instance, these molecules are major components of low-density lipoprotein (LDL). These fatty acid moieties both contain weak bis allylic C–H bonds that are prone to autoxidation. As previously mentioned, the autoxidation of LDL is implicated as an initiator of atherosclerosis in humans (8). Vitamin E is the primary vitamin antioxidant for fat tissue in humans and is thus important in preventing the development of atherosclerosis. Vitamin E refers to a family of fat-soluble phenols of which the most active form is α -tocopherol. The fat solubility of vitamin E is due to its long hydrocarbon (phytyl) “tail”, shown below.
Table 1. Structure and k7 Value for Some Phenolic Antioxidants No. 1 2 3
OH R1
R1
R2
4
CH3
320 × 104 CH3 O
H3C
C16H33
CH3
570 × 104
CH3
5 HO
O
CH3
CH3
O
H3C
polar head
1.4 × 104 94 × 104 39 × 104
R2
HO
HO
H3C
R1 = C(CH3)3; R2 = H; R3 = CH3 R1 = CH3; R2 = H; R3 = OCH3 R1 = R2 = CH3; R3 = OCH3
R3
CH3
CH3
CH3
non-polar tail
Within the fatty membrane vitamin E is believed to be oriented so that its polar “head” group (phenol OH) is directed toward the interface of the polar membrane exterior with the aqueous environment. By bobbing up and down, vitamin E molecules maintain contact with both the membrane’s nonpolar interior and the aqueous environment on the exterior. In Figure 1 a schematic representation of this is shown with a phospholipid bilayer representing a simplified biomembrane. In the aqueous environment that bathes all membranes, vitamin C (ascorbic acid) has been shown to be the major 1110
k7/(M{1 s {1)
Structure
6
CH3
2870 × 104
HO
CH3 O
CH3
O
7 HO
HO
110 × 104 O CHOHCH2OCO(CH2)14CH3
Journal of Chemical Education • Vol. 76 No. 8 August 1999 • JChemEd.chem.wisc.edu
In the Classroom
vitamin antioxidant (15). That vitamin C is soluble only in an aqueous environment is readily understood by examination of its molecular formula, C6H8O 6. Three of the oxygens in vitamin C are in alcohol functional groups that promote solubility through hydrogen bonding with water. How Effective Are Vitamins E and C as Antioxidants? Since a major function for vitamins E and C in humans is as antioxidants, it would be interesting to compare just how effective these compounds are at preventing autoxidation compared to a typical commercial antioxidant. The prototypical commercial antioxidant is butylated hydroxytoluene (BHT), whose structure is shown in Table 1, compound 1. This compound and simple derivatives of it are commonly used as antioxidants for jet fuels, paper products, rubber, and polymers. Table 1 is a compilation of experimentally measured rate constants for OH hydrogen atom abstraction by a typical peroxyl radical (k7 in Scheme I) in an organic solvent at 30 ˚C (16 ). Examination of Table 1 reveals that the k 7 value for α-tocopherol (4), is more than two orders of magnitude larger than the corresponding value for BHT (1). To examine watersoluble vitamin C in this organic system, it was necessary to employ an organic-soluble analog. Compound 7 in Table 1, ascorbyl palmitate, is soluble in organic solvents because of the long hydrocarbon tail added by esterification of the terminal OH in vitamin C. Although not as reactive as α-tocopherol, ascorbyl palmitate is almost two orders of magnitude more reactive than BHT. Thus, vitamins E and C are much more reactive with peroxyl radicals than typical commercial antioxidants.
2 is about one third as reactive as α-tocopherol suggests that having a para methoxy (with respect to the phenolic OH) is an important structural criterion. Introduction of two meta methyl groups into structure 2 yields compound 3, which structurally approximates α -tocopherol. Surprisingly, however, structure 3 is less reactive than structure 2. A clue that can help us rationalize this difference in reactivity is provided by the X-ray crystal structure (16a) for compound 3. In the solid state, steric interactions between the para oxygen and the meta methyl groups force the para methoxy group into a conformation in which the nonbonding electrons on the para oxygen are orthogonal to the p orbitals on the aromatic ring. The angle between the nonbonding electron pair on the p-methoxy group and the aromatic p orbitals is referred to as the theta (θ) angle (shown below in an abbreviated manner for compound 5). Presumably in solution a similar effect is operational where the theta angle for compound 3 is such that there cannot be optimal electron donation by resonance. This analysis is further supported by the fact that when an oxygen nonbonding pair is forced into more favorable conformations, for resonance overlap with the p orbitals on the aromatic ring, dramatic increases in reactivity are seen. The X-ray crystal structure of analogs of α -tocopherol reveal that the presence of the sixmember ring orients an oxygen nonbonding pair with the aromatic p orbitals so that they are only 17° from complete overlap and presumably better electron donation by resonance. Analogously, in compound 5, oxygen in the five-member ring reduces the θ angle (shown below) to just 6°, with a concomitant increase in k7.
Why Are Vitamins E and C Such Effective Antioxidants? A fundamental premise underlying the way modern organic chemistry is taught is that there are relationships between molecular structure and chemical reactivity. Such an approach will be used here to provide additional examples of these structure–reactivity relationships. A major factor controlling the effectiveness of a particular antioxidant is the magnitude of its rate constant for reaction with a peroxyl radical, k7 in Scheme I. The larger the ratio k7/k5 is, the better an antioxidant is at protecting a particular substrate from autoxidation. By far the most important structural feature common to all antioxidants is that hydrogen atom donation to a peroxyl radical yields a resonance-delocalized free radical. The importance of this structural facet can be illustrated by comparing the gas-phase homolytic OH bond dissociation energy for methyl alcohol (110 kcal/mol) with phenol (87 kcal/mol). In addition, substituents on the resonance-delocalized phenol radical can provide additional stability. Reaction 7 for all of the compounds listed in Table 1 is most certainly exothermic (17). The Hammond postulate2 will be invoked to relate the stability of the antioxidant radical with the stability of the (late) transition state for hydrogen atom transfer in step 7. Thus, the more stable the antioxidant radical, the lower the activation energy for hydrogen atom transfer and the larger k7. Examination of Table 1 allows us to dissect the structural factors that affect k7. Comparing the k7 value for compound 1 with compound 4 (α-tocopherol) reveals that the latter is more than 200 times more reactive with peroxyl radicals. The fact that compound
theta angle
O
O
Adding another aromatic ring onto compound 5 yields compound 6, which is the most reactive phenolic antioxidant known to date. Presumably, the additional resonance delocalization provided by the second aromatic ring further lowers the activation energy for step 7 and thus increases k7. Vitamin C analog 7 is a much more effective antioxidant than BHT (1). We can explain this fact by suggesting that the radical formed by hydrogen atom abstraction from 7 contains more resonance delocalization than the equivalent radical derived from BHT. As shown below it is believed that the β-hydroxy, with respect to the carbonyl, is the primary site of hydrogen atom donation (16c). O
HO O
O
O
O
R
O
HO
R
O
(8)
O
HO O
R
(9)
Radical 8 is a particularly good resonance structure because it is a captodative radical (18). Such radicals are stabilized by both an electron-donating group (OH) and an electronwithdrawing group (C=O). The net result of the captodative effect is to provide additional resonance stability for 8 and is
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pictorially represented by resonance structure 9. This additional stability is presumably reflected in the transition state for hydrogen atom abstraction for compound 7. In nature vitamin C is designed to function as an antioxidant in an aqueous environment. Structurally, vitamin C should be viewed as a homologous carboxylic acid. Thus, at physiological pH (~7) ascorbic acid is totally ionized (10). Electron spin resonance studies (19) in aqueous solution suggest that loss of a hydrogen atom from ascorbate results in the formation of radical anion 11, another example of a captodative radical as shown in 12. O
-H
O
HO
O
O
O
CHOHCH2OH
O
O
O
(10)
CHOHCH2OH
O
CHOHCH2OH
O
O
(11)
(12)
As antioxidants in nature vitamin E and C are believed to act cooperatively (20) and possibly synergistically (16a). In various micellar systems vitamin C located in the aqueous phase is known to reduce vitamin E that is contained in the organic phase. In fact, model studies suggest that in the absence of vitamin C, vitamin E bound to LDL can actually enhance LDL peroxidation (i.e., act as a prooxidant) (21). These observations suggest that an aqueous ascorbate free radical (11) is more stable than the tocopherol radical in the organic phase. Finally, the importance of free radical structure with respect to antioxidant function can be further illustrated by the fact that captodative resonance structures can be written for many of the radicals formed by the oxidation of compounds that have been proposed to have a antioxidant function in vivo. In fact, the tocopherol radical can be written as a captodative radical as shown in 13 below. Other examples include the urate radical 14 (22) and the ovothiol radical cation 15 (23). CH3
O H N
O HN O
O CH3 (13)
R'
N S
O
R H3C
CH3
H3C
N H
N
N+ H
(14)
(15)
Notes 1. Usually mild conditions refers to air-saturated solutions heated to temperatures less than 100 °C. 2. A detailed explanation of this topic can be found in any organic chemistry text book. 3. The structure of arachidonic acid is shown below: COOH
1112
Literature Cited 1. Hollis, W. G. Jr. J. Chem. Educ. 1996, 73, 61, and references within. 2. Shaping the Future: New Expectations for Undergraduate Education in Science, Mathematics, Engineering, and Technology. A review of undergraduate education by the advisory committee to National Science Foundation Directorate for Education and Human Resources, NSF 96-139; National Science Foundation: Washington, DC, 1996. 3. Seymour, E; Hewitt, N. Talking About Leaving: Factors Contributing to High Attrition Rates Amoung Science, Mathematics, and Engineering Undergraduate Majors; Bureau of Sociological Research, University of Colorado: Boulder, CO, 1994. 4. Ward, R. J.; Bodner, G. M. J. Chem. Educ. 1993, 70, 198, and references within. 5. Harman, D. J. Gerontol. 1956, 11, 298. 6. Harman, D. Proc. Natl. Acad. Sci. USA 1981, 11, 7124. 7. Pryor, W. A. In Modern Biological Theories of Aging; Warner, H. R.; Butler, R. N.; Sprott, R. L., Eds.; Raven: New York, 1987; pp 89–112. 8. Steinberg, D.; Parthasarathy, S.; Carew, T. E.; Khoo, J. C.; Witztum, J. L. N. Engl. J. Med. 1989, 320, 915. 9. For simplicity only the isomeric trans,trans hydroperoxides are shown. In addition, isomeric trans,cis hydroperoxides are also formed. See Porter, N. A.; Lehman, L. S.; Weber, B. A.; Smith, K. J. J. Am. Chem. Soc. 1981, 103, 6447. 10. Benson, S. W.; Nangia, P. S. Acc. Chem. Res. 1979, 12, 223. 11. For a hypothesis addressing “non-initiated” autoxidation, see Beaver, B. D.; Treaster, E.; Kehlbeck, J. D.; Martin, G. S.; Black, B. H. Energy Fuels 1994, 8, 455. 12. Slater, T. F. J. Biochem. 1984, 222, 1. 13. Howard, J. A.; Ingold, K. U. Can. J. Chem. 1967, 45, 793. 14. Handbook of Antioxidants; Denisov, E., Ed.; CRC: New York, 1995. 15. Frei, B.; England, L.; Ames, B. Proc. Natl. Acad. Sci. USA 1989, 86, 6377. 16. These literature rate constants were measured in styrene at 30 °C with azobis(isobutyronitrile) as initiator. For other details see (a) Burton, G. W.; Ingold, K. U. Acc. Chem. Res. 1986, 19, 194. (b) Burton, G. W., Doba, T.; Gabe, E. J.; Hughes, L.: Lee, F. L.; Prasad, L.; Ingold, K. U. J. Am. Chem. Soc. 1985, 107, 7053. (c) Barclay, L. R. C.; Dakin, K. A.; Zahalka, H. A. Can. J. Chem. 1992, 70, 2148. (d) Barclay, L. R. C.; Vinqvist, M. R.; Mukai, K.; Itoh, S.; Morimoto, H. J. Org. Chem. 1993, 58, 7416. 17. Most hydroperoxides have OH bond strengths ~88 kcal/mol, whereas most phenols have OH bond strengths between 78 and 85 kcal/mol. See Foti, M.; Ingold, K. U.; Lusztyk, J. J. Am. Chem. Soc. 1994, 116, 9440. See also Wayner, D. D. M.; Lusztyk, E.; Ingold, K. U.; Mulder, P. J. Org. Chem. 1996, 61, 6430. 18. Viehe, H. G.; Janousek, Z.; Merenyi, R.; Stella, L. Acc. Chem. Res. 1985, 18, 148. Also see Bordwell, F. G.; Zhang X.-M.; Alnajjar, M. S. J. Am. Chem. Soc. 1992, 114, 7623. 19. Laroff, G. P.; Fessenden, R. W.; Schuler, R. H. J. Am. Chem. Soc. 1972, 94, 9062. 20. Pryor, W. A.; Cornicelli, J. A.; Devall, L. J.; Tait, B.; Trivedi, B. K.; Witiak, D. T.; Wu, M. J. Org. Chem. 1993, 58, 3521. 21. Bowry, V. W.; Stocker, R. J. Am. Chem. Soc. 1993, 115, 6029. 22. Simic, M. G.; Jovanovic, S. V. J. Am. Chem. Soc. 1989, 111, 5778. See also Ames, B. N.; Cathcart, R.; Schwiers, E.; Hochstein, P. Proc. Natl. Acad. Sci. USA 1981, 78, 6858. 23. Holer, T. P; Hopkins, P. B. J. Am. Chem. Soc. 1988, 110, 4838. See also Turner, E.; Hager, L. J.; Shapiro, B. M. Science 1988, 242, 939.
Journal of Chemical Education • Vol. 76 No. 8 August 1999 • JChemEd.chem.wisc.edu
