Issuu on Google+


Nicole Loeven and Molly Woods Wells College 170 Main Street Aurora, NY 13026 Warfarin, enantiomer, TLC, column chromatography, recrystallization. ABSTRACT: Warfarin is a commonly used anticoagulant prescription drug. Anticoagulants are a class of drugs that are used to prevent the clotting of blood. The Warfarin molecule is a racemic mixture of both (R,R) and (S,S) conformations. The purpose of this synthesis was to create Warfarin in a way that is both efficient and environmentally friendly. This was done by two procedures, one with column chromatography on a 0.1 mmol scale and another with recrystallization on a 2.0 mmol scale. The results of a 1HNMR spectrum showed that there was a higher keto form present in the product than the ketal form of Warfarin.

INTRODUCTION Warfarin is a commonly used anticoagulant, or blood thinner. The name Warfarin comes from Wisconsin Alumni Research Foundation and the suffix -arin was added onto the end1. There are both (R,R) and (S,S) enantiomers of Warfarin that affect potency and have different affects when taken with other herbal medicines and drugs, such as aspirin, antibiotics, antidepressants, and antiplatelet drugs2.. The following experiment is an environmentally friendly way of making Warfarin. This is considered better for the environment because the reaction was done on a very small scale. Both enantiomers were made in two different procedures. See Figure 1. One procedure involved column purification on a 0.1 mmol scale and another involved recrystallization on a 2.0 mmol scale. The results of the each experiment were analyzed by TLC and 1HNMR spectroscopy.

Figure 1. The reaction of 4-hydroxycoumarin and trans-4-phenyl-buten-2-one to make both (R,R) and (S,S)-warfarin3..

EXPERIMENTAL METHODS Experimental Procedure 14.: (0.1 mmol scale with column purification) 16mg (0.10mmol) of 4-hydroxycoumarin, 15mg (0.11mmol) of trans-4-phenyl-3-buten-2-one, 2.1mg (0.010mmol) of (R,R)-1,2-diphenylethylenediamine (dpen), 0.2mL of anhydrous tetrahydrofuran, and 57ÎźL (1.0mmol) of acetic acid was added to a vial. The vial was mixed to dissolve the solids. Another vial was mixed with the same amount of chemicals but with 2.1mg (0.010mmol) (S,S)-1,2-

diphenylethylenediamine (dpen). This was left to sit for a week at room temperature. After one week TLC was performed on both reaction mixtures using silica gel plates (CH2Cl2 mobile phase and anisaldehyde stain). A packed column was used for purification. The crude product from each reaction was dissolved in a minimal amount of CH2Cl2. Column chromatography (silica gel, 1% Methanol/CH2Cl2) was used to purify the product. See Image 1. Experimental Procedure 2: (2.0 mmol scale with recrystallization) In a 4-dram vial, 0.325 grams of 4-hydroxycoumarin (2.0 mmol), 0.307 grams of trans-4-phenyl-3-buten-2one (2.1 mmol), 0.0402 grams of (R,R)-1,2-diphenylethylenediamine (dpen) catalyst, 4.0 mL anhydrous THF, and 1.14 mL of acetic acid were added. The vial was and left at room temperature for one week. The same amount of each chemical was added to a second 4-dram vial only with the (S,S)-1,2-dpen catalyst. See Image 2.


Image 1. These are the vials with the complete reaction mixtures.

A TLC was performed on both reaction mixtures (CH2Cl2 and anisaldehyde stain) after one week. The liquid in the reaction vials were removed of solvent and acetic acid by rotary evaporation. The evaporation gave an oil, which was recrystallized and collected by vacuum filtration. The crystals were air-dried giving 0.54 grams of purified (R,R) product and 0.51 grams of purified (S,S) product. The percent yield of the (R,R) product was 87.5% and the (S,S) product was 82.7%.

RESULTS The TLC results showed that Warfarin was present in the reaction solutions of both procedures from known data of Rf values and color. It also showed that some unreacted starting material was in the product. The 1HNMR spectrum showed that there was a higher keto form present in the Warfarin product than the ketal form. See Figure 2.

Figure 2. Keto-Ketal Isomerism of Warfarin3..

This synthesis of Warfarin was environmentally friendly because it was successful on a small scale. It was completed without heating or stirring for long periods of time. There was also no need for additional pressure for the reaction to proceed. However experimental procedure 1 was unsuccessful because there was too little product made to analyze. Experimental procedure 2 was the better method with high percent yields.

1. Holmes, R. W.; Love, J. J. Am. Med. Assoc. 1952, 148, 935-937. 2. NYU Medical Center. Patients and Family Resource Center. 2011. Managing Your Warfarin (Coumadin速) Therapy. 4,10-12 3. Wong, C. T.; Sultana, C. M.; Vosburg, D. A. J. Chem. Educ. 2010, 87, 194-195. 4. Kim, H.; Yen, C.; Preston, P.; Chin, Org. Lett. 2006, 8, 5239.

Chan, E. et al. Br. J. Clin. Pharmac. 1994, 37, 563. Halland, N.; Hansen, T.; Jorgensen, K. A. Angew. Chem. Int. Ed. 2003, 42, 4955. West, B. D.; Preis, S.; Schroeder, C. H.; Link, K. P. J. Am. Chem. Soc. 1961, 83, 2676.


Zakarra Butts*, Antonaia Merritt Wells College, Aurora, NY 13026 KEYWORDS Diels-Alder, 2,4-Hexadien-1-ol, Maleic Anhydride, cis-1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofurancarboxylic Acid, COSY NMR, 13C NMR, 1H NMR, thin-layer chromatography (TLC) ABSTRACT: Diels-Alder reactions are used to produce a cyclic product through cycloaddition of a diene and a dienophile. The goal of this experiment was to produce an endo product from the diene and the dienophile. The reactants used were 2,4hexadien-1-ol and maleic anhydride; this reaction produced cis-1,3,3,4,5,7-hexahydro-5-methyl-3-oxo-4-isobenzofuran-carboxylic acid, the endo product. The product was analyzed by thin-layer chromatography (TLC) along with various NMR spectrums (1H NMR, 13C NMR, and COSY NMR).

Scheme 1. Reaction of E,E-2,4-Hexadien-1-ol with Maleic Anhydride Introduction In 1950, Hermann Diels and his student, Kurt Alder, received the Nobel Prize in Chemistry for their discovery of cycloaddition reaction, known today as the Diels-Alder reaction.2 A cycloaddition is a reaction in which two reactants produce a cyclic product; the reactants are called a dienophile and a diene.2 The dienophile is an alkene that has an electron withdrawing group attached to it.2 The electron withdrawing group is next to the double or triple bond in the dienophile; this makes the C-C double/triple bonded carbons less electron rich, allowing it to be more easily attacked by the diene.2 The diene has to be in an s-cis, or all-cis, conformation in order for the Diels-Alder reaction to occur.2 If the diene was not in an s-cis conformation the reaction would not occur because there would be no overlap of p orbitals with the dienophile.2 Diels-Alder reactions are efficient because they occur in one step rather than in multiple steps 2. The reaction of a dieneophile and a diene results in a stereospecific product which is either endo or exo.2 Endo products are the typical products of a Diels-Alder reaction because they result in the most overlap between the orbitals of the diene and the dienophile.2 The Diels-Alder reaction researched was between E,E-2,4-hexadien-1-ol and maleic anhydride in toluene. E,E2,4-hexadien-1-ol is the diene and maleic anhydride is the dienophile in this reaction. In order to determine that the

product of this reaction was an endo product, three different NMR’s were obtained. The first NMR was a 13C NMR which shows how many unique carbon environments are in the product. The second NMR obtained was a 1H NMR which shows the unique hydrogen environments of the product. Finally, the third NMR obtained was a COSY NMR; COSY NMR is a 2D NMR of the hydrogen coupling in the product, cis-1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofuran-carboxylic Acid . Methods Maleic Anhydride

E,E-2,4-hexandien-1-ol

Toluene

Melting Point ( ̊ C)

52.6

28-33

-95

Boiling Point ( ̊ C)

202

80

111

Mass (g/mol)

98.06

98.14

92.14

Density (g/mL)

1.48

0.871

0.87

Table 1. Physical Properties 0.40g (0.0040858 moles) of maleic anhydride was added to 2.2mL (0.0040758 moles) of E,E-2,4-hexadien-1ol. The solution was dissolved in 5mL of toluene in a round bottom flask. The solution was refluxed under heat for 5 minutes; a TLC was performed in ethyl acetate solvent to ensure all beginning products were not present in the solution. The solution was cooled to room temperature, placed in an ice bath for 10 minutes, sealed and reacted for a week. The crystals that formed were collected through vacuum filtration and rinsed with cold toluene. The product was air dried, weighed and the melting point was taken along with a 1H NMR, 13C NMR and a COSY NMR, with a solvent of acetone D6.


Results Figure 1 shows the 1H NMR of the product, showing that there are eight unique hydrogen interactions, even though the product has a total of ten different hydrogens off of carbons. The eight peaks compared to the ten hydrogens found in the product results from two sets of hydrogens being in the same environment.

Figure 3. 13C NMR of cis-1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofuran-carboxylic Acid

Figure 1. 1H NMR of cis-1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofuran-carboxylic Acid

Figure 2 shows the product with its hydrogens. H7 and H6 correspond with the peak at 5.6ppm; H4 corresponds with the peak at 4.5ppm; H8 corresponds with the peak at 4.2ppm; H3 corresponds with the peak at 3.4ppm; H1 and H2 correspond with the peak at 3.1ppm; H9 corresponds with the peak at 2.6ppm; H5 corresponds with the peak at 2.0ppm; the methyl group corresponds with the peak at 1.2ppm.

Figure 4 shows the COSY NMR of the product shows that there are six different hydrogen coupling interactions in the final product. This allows the hydrogens that have the same interactions to be identified from one another in the product. The COSY of the product revealed interactions of hydrogens at 2.1ppm and 3.0ppm, 5.7ppm and 3.2ppm, 3.0ppm and 1.15ppm, 4.5ppm and 3.25ppm, 2.0ppm and 6.0 ppm, and at 3.5ppm and 5.5ppm. These interactions show that the product obtained is the endo product of the reaction.

Figure 2. Cis-1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofurancarboxylic Acid Figure 4. COSY of cis-1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofuran-carboxylic Acid

Figure 3 shows the 13C NMR of the product, this reveals that there are eight different unique carbon interactions in the product, even though the product has ten carbons. Having only 8 peaks and 10 carbons in the product means that two sets of carbons are in the same environment, resulting in the eight peaks.

The reaction resulted in 0.54 g of product, cis1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofurancarboxylic Acid, with a melting point of 157 ĚŠC-160 ĚŠC. The theoretical yield of product is 0.7997 g; the actual yield was 0.5400 g, resulting in a 67.53% yield.


Conclusion/Discussion The interactions between E,E-2,4-hexadien-1-ol and maleic anhydride resulted in the product, cis1,3,3,4,5,7-Hexahydro-5-methyl-3-oxo-4-isobenzofurancarboxylic acid. For this reaction, 0.5400 g was yielded (67.53%) with a melting point of 157째C-160째C. Ideally this reaction should have resulted in an 81% yield of product and a melting point of 156째C -160째C.1 The difference in yield percentages could be a result of product being lost or not

formed during the reflux. The TLC of the product revealed that there were still reactants present in solution. The combinations of the different NMR spectrums verified that the product was endo rather than exo.

1 Mcdaniel, Keith F., and R. Matthew Weekly. "The Diels-Alder Reaction of 2,4-Hexadien-1-ol with Maleic Anhydride: A Novel Preparation for the Undergraduate Organic Chemistry Laboratory Course."Journal of Chemical Education 74.12 (1997): 1465. Print. 2 McMurry, John. "Conjugated Compounds and Ultraviolet Spectroscophy." Organic Chemistry: Hybrid Edition. Belmont: Brooks/Cole Cengage Learning, 2012. 400-425. Print.


Alexandria F. Roberson, Bri’Anna S. Horne Wells College, 170 Main Street, Aurora, NY 13026 KEYWORDS: Bisphenol A, Bisphenol Z, Hormone, Fat Solubility, Endocrine Disruptor, Electrophilic Aromatic Substitution ABSTRACT: There is controversy over health and environmental concerns from bisphenol A and bisphenol Z11. They are well known endocrine disruptors that are used in polycarbonate plastics and epoxy resins. Both bisphenol A and Z were synthesized and the fat solubility of the two were compared. The fat solubility was found by determining the partition coefficient. Bisphenol Z was the more fat-soluble chemical.

The difference in the solubility of a solute in two solvents is measured by the partition coefficient, KPp. One of the most useful examples is the octanol/water partition coefficient. KPp measures the hydrophilicity/hydrophobicity of a solute. The molecule is more hydrophobic when it has a larger Kp value. Kp can be transformed to log P (= log Kp). Determination of the log P of a compound is important in the field of medicinal chemistry, because a drug must be able to pass through the hydrophobic cell membrane to reach its target4.

Figure 1. Synthesis of BPA

(2)

Figure 2. Synthesis of BPZ Introduction Bisphenol A (BPA), Figure 1, 4-4’-(propane-2,2-diyl)diphenol, is used to synthesize various plastics and epoxy resins that are used in everyday life. Bisphenol Z (BPZ), Figure 2, 4-4’-(cyclohexane-1,1-diyl)diphenol, is produced in massive quantities as a precursor to polycarbonate plastics and epoxy resins, which has raised recent concerns that BPZ could be as hazardous as BPA1. Bisphenol Z and bisphenol A undergo an electrophilic aromatic substitution reaction2. Bisphenol A and Z are endocrine disrupters which bare structural similarities to other hormones1. Endocrine disruptors are chemicals that can mimic and interfere with the natural endocrine systems in mammals3. BPA and BPZ imitates our body’s own hormone system, which creates concern because it can pose harmful effects such as cancer, tumors, birth defects, and developmental disorders1. As a result, many states in the U.S. have taken legislative initiative in passing laws that ban the use of bisphenol A in products, which releases the harmful chemical when exposed to heat. However, recent studies argue that these endocrine disruptors are not nearly as dangerous as previously thought, stating that the average amount of BPA exposure is millions of times lower than the amount needed to have significant or noticeable biological changes3.

This experiment intends to provide evidence, through the synthesis and fat solubility comparison of BPZ and BPA, whether BPA or BPZ are more harmful chemicals when released into the environment. This will support or weaken the claims on the danger of BPA in mammals, and provide insight into the danger of BPZ. Experimental details: To a round bottom flask was added 12 mmol of phenol (1.12g), 3 mmol of cyclohexanone (0.31 mL) and 0.5 mL of concentrated HCl to synthesize BPZ. In the synthesis of BPA, 3 mmol of acetone (0.27 mL) was used as a substitute for cyclohexanone. The mixture was refluxed for one hour. The solid product was collected by vacuum filtration and washed with 30 mL of water and 10 mL of toluene. 30 mL of boiling water was added to the adduct. The phenol was removed under vacuum filtration and washed with hot water. The fat solubilites of bisphenol A and bisphenol Z were compared. The bisphenol A standard was created with 0.002 g (40 mg/L) of solid in 50 mL of H2O. 3 mL of this standard was added to a test tube containing 3 mL of 2-octanol. The concentration of BPA over the range from 0 to 80 mg/L obeys the equation A = 0.01325 x C where C is the concentration4. The partition coefficient (Kp) was calculated in bisphenol A and bisphenol Z as follows: Kp = [BPA (octanol)]/ [BPA (aq)] and logP = logKp3. Kp= 30mg/L ÷ 10mg/L. 3 mL of the bisphenol Z standard (0.002 g (40 mg/L) in 50 mL of H2O) was added to 3 mL of 2-octanol in a test tube. The concentration of BPA over the range from 0 to 80 mg obeys the


equation A = 1.3332 x C where C is the concentration4. The absorbance of the bisphenol Z standard was A276=1.208 in a concentration of 40 mg. Kp = 39.09 mg/L á 0.91 mg/L. Results: The absorbance of the bisphenol A standard was A276= 0.154 in a concentration of 10 mg/L. The Kp of bisphenol A is 2.9997 moles and logP=0.477 moles). The Kp of bisphenol Z is 42.96 moles and logP=1.630 moles. The Kp of both products indicates that bisphenol Z is more fat-soluble than bisphenol A. The logP value being a larger, positive value suggests that the chemical has more hydrophobic parts because of a larger Kp value. Octanol has more fat than the aqueous layer, making the chemicals fat-soluble. The logP value being larger and a negative value makes the chemical have more hydrophilic parts because of a smaller Kp value. This aqueous layer is more water-soluble than the octanol layer. The logP value for bisphenol A was 0.477 and the logP value of bisphenol Z was 1.630. Bisphenol Z has the larger logP value making it more fatsoluble. Conclusions: Bisphenol Z is more fat-soluble because the structure contains more carbons, making the molecule have hydrophobic parts. This is taken into account with making medicines. Chemists do not want drugs to be too fat-soluble because then it can be stored in fat throughout the body, which can be released in high doses if rapid weight loss occurs. The normal range for solubility of chemicals is -5 to 5. -5 to 0 is water-soluble and 0 to 5 is fat-soluble4. Both bisphenol A and bisphenol Z fall relatively low in the fat-soluble area. This suggests that the chemical poses a minor threat. This conclusion supports the argument that BPA's effect in the environment poses a low threat at the average level of exposure.

We thank Dr. Lauren O’Neil and fellow peers for the helpful comments on revision and the help of instruction to perform the experiment. We also thank Wells College for the laboratory space, equipment, and chemicals.

1. 2. 3. 4.

Gregor, R. W. J. Chem. Educ. 2012, 89, 669-671. Bull. Korean Chem. Soc. 2011, 32, No. 8 Delclos, K. B. Toxicol. Sci. 2014, 139, 174-197. Lipinski, C.A. et al. Adv. Drug. Dev. Rev. 1997, 46, 3-26.


Lindsay Achzet, Marie Valliere, Ashley Roser. Wells College, 170 Main Street Aurora, NY 13026 KEYWORDS Malachite Green, Crystal Violet, Grignard Reagent, Paper Chromatography ABSTRACT: Use of dyes within the fish and medical industries has become increasingly beneficial towards research. Specific dyes, produced by coupling with a Grignard reagent, were tested for their polarity, which would impact their use in such industries. The Rf values of Malachite Green and Crystal Violet were high which indicated that these substances were relatively nonpolar and thereby quite similar in polarity. This results in the possibility of interchangeable use between these two substances in the fish and medical industries. INTRODUCTION Dyes have numerous uses within the medical and scientific communities. Malachite Green is used as an anti-microbial, anti-parasitic, and anti-fungal by the fish-farming industry1. Crystal Violet is used to dye mammalian cells2. These dyes are very similar in structure. Malachite Green shows carcinogenic properties and is now a banned substance in the fish industry despite its beneficial properties. Due to the similar nature of Malachite Green and Crystal Violet, it’s a possibility that Crystal Violet, a non-carcinogenic substance, can be used instead of Malachite Green in the fish industry. The objective is to determine the polarity of Malachite Green and Crystal Violet through the use of paper chromatography. By determining the relative polarity of each dye, it may be possible to distinguish whether these two substances can be used interchangeably within the fish industry, as well the medical industry. MATERIALS AND METHODS To synthesize the Grignard reagent, 2.5 g (0.0125 mol) of 4-bromo-N,N-dimethylaniline, 0.40 g (0.0165 mol) magnesium (Mg) turnings, 30 mL of dry tetrahydrofuran (THF) and 2-3 crystals of iodine were added to a 50 mL round- bottom flask, heated and refluxed for thirty minutes. The reaction flask was cooled to room temperature. To synthesize Malachite Green, 0.21 g (0.00154 mol) of methyl benzoate was added to 1.0 mL of THF. This ester solution was added to the Grignard reagent and heated to reflux. The flask was cooled and 5 mL of 5% HCl solution was added. To synthesize Crystal Violet, 0.30 g (0.00254 mol) of diethyl carbonate was added to 1.0 mL of THF. This ester solution was added to the Grignard reagent and heated to reflux. The flask was cooled and 5 mL of 5% HCl solution was added. The percent yield of Malachite Green was 32.896%. The percent yield of Crystal Violet was 51.29%. A paper chromatography chamber was organized using 200 proof ethanol as a solvent and 1:1:1: 95% ethanol, 1-butanol, 2.0 M ammonia (aq) as a solvent. The polarities of the Malachite Green and Crystal Violet were compared to the polarities of Indigo Carmine and Allura Red. The R f was calculated for each dye. RESULTS

Table 1. Rf Values of Compounds in 200 Proof ethanol and 1:1:1 95% ethanol, 1-butanol, 2.0 M ammonia Compound

Solvent

Rf Values

Allura Red

Ethanol

0.696

1:1:1

0.471

Ethanol

0.120

1:1:1

0.267

Ethanol

0.927

1:1:1

0.956

Ethanol

0.945

1:1:1

0.956

Indigo Carmine

Malachite Green

Crystal Violet

DISCUSSION The objective was to observe whether Crystal Violet dye could be substituted for Malachite Green within the fish industry. Based on their relatively similar polarities, it is very possible that Malachite Green could be replaced with Crystal Violet. The average Rf value for Malachite Green was 0.9415, which represents a substance with low polarity. The average Rf of crystal violet was 0.9505, which also represents a low polarity substance. Compared to the dyes Indigo Carmine (average Rf of 0.387) and Allura Red (average Rf of 0.5835), Malachite Green and Crystal Violet are very similar in polarity. Therefore based upon polarity, these two substances may be used interchangeably in the fish industry, and possibly the medical field. There are other factors that could affect the use of Crystal Violet instead of Malachite Green as well, which need to be further examined. If Crystal Violet were to be used instead of Malachite Green, it would be extremely beneficial to the


fish industry as Crystal Violet is a non-carcinogen and Malachite Green is a carcinogen.

Acknowledgment should be made to Dr. Lauren O’Neil who helped plan, prepare, and execute this experiment.

.

1Chan, D.; Tarbin, J.A.; Stubbings, G.; Kay, J.; and Sharman, M. Food Additives and Contaminants 2012: 29, 66-72.

2Kirilova, E.; Ivanova, I. Scientific Journal of Riga Technical University 2011: 23, 29-33.


Kyle Admire* and Hayden Shuster Wells College, 170 Main Street, Aurora, NY, 13026 KEYWORDS Room Temperature Ionic Liquid, RTIL, Mannich Reaction, Green Chemistry ABSTRACT: Room Temperature Ionic Liquids (RTILs) are one of the many facets of green chemistry. These liquids have the benefit of use and reuse through filtration and washes. This is unlike many traditional solvents, such as ethanol, which can only be used in one reaction. One of the first uses of RTIL was as a solvent in the Mannich Reaction. Two Mannich Reactions, one with ethanol and one with RTIL as the solvent, were run side-by-side to compare the percent yield in a traditional solvent and “green” solvent. The RTIL solution was effective for the reaction, as well as having a high percent return of the mother liquor (94.7%). One drawback the RTIL is that it became very difficult to collect the precipitate after the first wash, requiring a diethyl ether wash to extract the solid.

Green chemistry is the quickly growing field involving chemistry that reduces or eliminates waste by using reusable solvents. One such solvent, room temperature ionic liquids (RTILs) may be a step forward within the green chemistry field. The liquid ionic state of RTILs is caused by the large organic cations that weaken the electrostatic attractions by the ions, therefore reducing the melting point. RTILs may grow green chemistry by replacing traditional organic solvents such as ethanol. The major advantage that RTILs have over traditional solvents is that they are capable of being recycled. RTILs are capable of running a reaction repeatedly, without loss of the liquid. The Mannich reaction, used to create beta-aminocarbonyls (see scheme 1), produces steroids, vita

Figure 1: The [BMIM] [BF4] where [BMIM] is 1-butyl-3-methylimidazolium cation.

mins, and antibiotics. The Mannich reaction can be performed in the traditional solvent ethanol or in RTIL. By comparing the yield of the Mannich reaction in ethanol to the yield in RTIL, the effectiveness and quatity of recovery of RTIL’s may be assessed1.

10.0 g of 1-methylimidazole (122 mmol), 14.69 g of 1chlorobutane (159 mmol) and 6.5 mL of acetonitrile was added to a 100 mL round bottom flask with a water-condenser under N₂. The reaction was heated to reflux at 80°C with stirring for 48 hours. After cooling to room temperature, 20 mL ethyl acetate was added to the mixture. The mixture was stirred vigorously for 1-2 min and the layers were allowed to separate. This was repeated 3 times. The upper layer (ethyl acetate) was removed with a pipette. The liquid was further concentrated by rotary evaporation at 60 °C for 10 minutes.

13.4 g (122 mmol) of sodium tetrafluoroborate was added to 48 mL deionized water, and slowly added to the [BMIM] [Cl] solution. The mixture was stirred at room temperature for 1.5 hours. While this was stirring, 2.12 mol of freshly purified benzaldehyde, 0.861 mol of 3-pentanone, 0.771 mol of ammonium acetate, and 2.5mL of ethanol were mixed in a 25 mL round bottom flask. The flask was parafilmed and stirred at room temperature for 24 hours. After 1.7 hours, the water was removed from the RTIL mixture by rotary evaporation at 80 °C. The mixture was diluted with 25 mL of dichloromethane and 8.0 g of anhydrous magnesium sulfate was added. The mixture was


stirred for ten minutes and filtered by gravity. The filtrate was concentrated by rotary evaporation at 40°C until constant weight. The RTIL was a light yellow, oily liquid. 2.12 mol of freshly purified benzaldehyde, 0.861 mol of 3-pentanone, 0.771 mol of ammonium acetate, and 2.5g of RTIL was added to a 25 mL round bottom flask. The flask was parafilmed and stirred in a water bath at 30°C for 24 hours. The RTIL solution was diluted with 20 mL of deionized water and stirred for 5 minutes. A white precipitate was collected by vacuum filtration, and the filtrate was saved. The precipitate was purified by recrystallization with 15-20 mL of hot ethanol and the product was collected by vacuum filtration. 20 mL of deionized water was added to the ethanol solution, which was stirred for 5 minutes. The off-white precipitate was collected by vacuum filtration and purified by recrystallization with 27 mL hot ethanol. This product and the filtrate from the vacuum filtration were retained. The RTIL was recovered by transferring the mother liquor collected from the reaction mixture into a separatory funnel and extracting it three times with diethyl ether. The aqueous solution was collected and concentrated by rotary evaporation.

The Mannich Reaction was repeated in only the RTIL and the yield of the reaction in recycled RTIL was determined. The precipitate did not separate from the RTIL by vacuum filtration, and could only be removed through extraction with diethyl ether. The diethyl ether/precipitate mixture was concentrated at room temperature and the precipitate was collected.

2.5182g of RTIL was synthesized. The Mannich reaction was performed with both the RTIL and with ethanol, a traditional organic solvent. The Mannich reaction with ethanol produced 0.0070g of product and the mannich reaction with RTIL produced 2.91g of product, which indicates much higher yields when using RTIL over traditional solvent. Re-running the Mannich reaction with the recovered RTIL produced no product and the amount of recovered RTIL weighed 2.3842g (94.7%) of the RTIL was recovered indicating it is recyclable, however further tests are needed to generate larger yields of product upon using recycled RTIL.

Scheme 1: The Mannich Reaction

Recovery of the mother liquor was very effective, giving 94.7% recovery. However, the extraction of the precipitate became very difficult after the first run of the reaction. After the first use of the liquid, vacuum filtration soon became insufficient to remove the would-be precipitate. One potential reason why it would be so difficult to remove the precipitate is that during the recovery step, some diethyl ether remained with the RTIL mother liquor. Even if the layers had a chance to separate, there would be no change in the efficacy of the second vacuum filtration. This makes the diethyl ether wash such an important step because it is when the purification of the mother liquor actually occurs. In order to combat this, we could attempt the Mannich Reaction with a different RTIL. RTIL’s are “designer solvents” (Mak et al) meaning that we could tailor an RTIL to our needs which would include a lower complexity to produce (primarily for use in a classroom setting), higher percent yield, and simpler means of recovering the precipitate after the first wash.

The authors would like to acknowledge Dr. Lauren O’Neil, the professor who made all of this possible. Additionally Angelo Papagelos and Melena Hagstrom, the teaching assistants who answered our questions, quandaries and concerns. REFERENCES Mak, K. K. W., et al. Journal of Chemical Education. 2006. 6. 943946.


Shelby L. Bourn, Charlee J. Weidman. Wells College 170 Main Street Aurora, NY 13026

ABSTRACT: Malachite green and crystal violet dyes were synthesized by a Grignard reaction and the malachite green dye was characterized using UV-Vis spectroscopy and solvatochromism with differing solvents. 2.5384 g of malachite green and 3.0625 g of crystal violet were synthesized in the experiment. The solvent that caused the biggest solvatochromic shift was acetone because the electron transition did not occur. INTRODUCTION:

Organic dyes such as malachite green and crystal violet can be synthesized by a Grignard reaction. Malachite green, which is a triphenylmethane dye, has many uses worldwide such as dyeing silk, wool, leather and cotton. The dye can also be used as an anti-fungal agent in aquaculture. This dye is extremely toxic to mammalian cells and can cause DNA damage as well as tumors. Due to its toxicity malachite green is banned by the United States Food and Drug Administration2. Crystal violet may be used to dye many diverse products such as detergent, paper, leather, antifreezes, and can even be a component in inks of pens. Crystal violet can also be utilized in veterinary medicine for dermatological drugs and expelling fungi and parasites from the body. However, in high doses the dye can be highly toxic to living organisms1. The visible light spectrum is from 380 nm to 780 nm. The shorter wavelengths, towards 380 nm, is where colors such as violet and blue are seen. The longer wavelengths, towards 780 nm, is where colors such as red can be seen. The ultraviolet region has shorter wavelengths than the visible region. Therefore, the UV-Vis spectrophotometer measures from roughly 200 nm to 700 nm covering both the ultraviolet and visible regions3. The solvents that are used in the sample for the spectrophotometer will have an effect called a solvatochromic shift. The shift is affected by whether the solvent is polar or nonpolar and whether it is protic or aprotic.

This study focuses on the synthesis of malachite green and crystal violet dyes and characterizes these dyes using UV-Vis spectroscopy and analyzing any solvatochromic shifts in the data. EXPERIMENTAL DETAILS: To make the Grignard reagent, a drying tube was prepared. Then, the condenser and round-bottom flask (RBF) were washed with tetrahydrofuran (THF) to dry them. Once the condenser and RBF were both washed, 2.5 g of 4-bromo-N,N-dimethylaniline (0.0015 m), 0.40 g Mg turnings, 30mL of dry THF, and 2-3 small crystals of iodine were added to the RBF. The RBF was clamped to the condenser and the set up was then heated to 75˚C in a water bath. A gentle reflux was maintained for 30 minutes. The RBF was cooled to room temperature. To synthesize the malachite green dye 0.21g of methyl benzoate and 1.0 mL of THF were added to a vial. Using a Pasteur pipette, the solution was added drop by drop, into the Grignard reagent. After the solution was added the condenser was replaced and the RBF was heated to reflux for 5 minutes. The flask was cooled to room temperature. The mixture was poured into a beaker and 5mL of 5% HCl solution was added. To synthesize the crystal violet dye 0.30g of diethyl carbonate (0.0125 m) and 1.0 mL of THF was added into a vial. The procedure for the synthesis listed for the synthesis of malachite green dye was repeated. Next, a ultra-violet spectrum was obtained for the malachite green dye in four different solvents. 0.5 mg of the malachite green dye was added to water, acetone, dimethylformamide (DMF), and ethanol. The solutions were run through the UV spectrophotometer to give the UV spectrum for each solution. RESULTS 2.5384 g of malachite green and 3.0625 g of crystal violet were synthesized. The theoretical yield of malachite green was 0.547 g, the actual yield acquired was 2.558 g. The theoretical yield for crystal violet was 1.019 g, the actual yield acquired was 3.063 g. The λmax for the first electron transition in the malachite green dye and water was 218 nm. The λmax for the same transition in malachite


green dye and acetone was nonexistent, meaning there was no peak in the area. With the malachite green dye and ethanol the λmax of the electron transition was 218 nm. The λmax for the malachite green dye and DMF electron transition was 280 nm. (Fig. 1) Figure 1. Solvent

λmax

Water

218

Acetone

No peak

Ethanol

218

DMF

280

CONCLUSION Due to the wetness of the product, the actual yield that was found for the dyes was much higher than the theoretical yield. The malachite green dye with water as a solvent was used as a control. The λmax for the control was 218 nm. The solvent that shifted the λmax the most, 62 nm, was the dye with DMF. Ethanol effected the λmax very little, there was no change in the positioning of the peak. The dye with acetone had no peak at the transition point. These effects have to do with whether or not the solvent was polar or non-polar and protic or aprotic. The control group, the water, was a polar protic solution. Acetone, which showed no peak, was a polar aprotic solution. DMF, was a very polar aprotic solution and ethanol is a solution that is less polar than water and protic. ACKNOWLEDGMENT We would like to thank Wells College for providing a laboratory with the proper supplies need to complete the experiment. Thank you to the teaching-assistants who were there to help, Angelo Papagelos and Jamyra Young. Thank you to Dr. O’Neil for all her assistance and fellow classmates for peer-review aid. REFERENCES 1Depci,

T.; Kul, A.; Onal, Y.; Disli, E.; Alkan, S.; Turkmenoglu, Z. Physiochemical Problems of Mineral Processing. 2012, 48 (1), 253-270. 2Kumar,

J.; Sarkar, S.; Hossain, U.; Chakraborty, P.; Kumar, R.; Bhattacharya, S. Indian J Med Res. 2013, 1163-1173. 3McMurry,

J. Organic Chemistry. 2012, 352


Kirsti Bruce and Marisa Smith Wells College, 170 Main Street, Aurora, NY Warfarin, Blood Thinner ABSTRACT: Experiments are being conducted to get an understanding how Warfarin can be made. One way is by column purification, while the other one is recrystallization. Scientists are trying to determine which procedure will give the most pure substance and which procedure would produce a substance that will result in a better profit in the end. Both procedures were completed within a three week period to get a better understanding of which process produces pure warfarin. Tests and measurements were completed to compare the two products. INTRODUCTION: With age comes health related issues and some may need to take warfarin. Warfarin is used to prevent blood clots from growing larger in the blood stream and the blood vessels1. This drug is usually prescribed to patients who have irregular heartbeat, prosthetic heart valves, and those that have suffered heart attacks and strokes1. It prevents the formation of clots by decreasing the production of factors II, VII, IX, and X that the liver promotes. It also helps prevent clots by reducing the anticoagulant proteins C and S2. Warfarin is also given to patients with deep vein thrombosis (DVT). These patients are given warfarin because it prevents the extension of the clots and it reduces the risks of emboli2. An Emboli is a blood clot, air bubble, or other object that is carried into the bloodstream that lodges in a vessel. Warfarin was approved by the FDA (The Food and Drug Administration) in June of 19542. It has two brand names, Coumadin and Jantoven1. The enantiomers of warfarin are different in potency, metabolism, and the effects when taken with other medicines3. Now experiments are being performed to see if there is a way to produce warfarin that will not have complications with other medications3. These experiments are also being done to assess which procedure gives a more pure substance and which one will produce the bigger profit. The two experiments being tested now are column purification and recrystallization with (R,R)-diphenylethylenediamine to see which will produce a more profitable yield. EXPERIMENTAL PLAN: To a ½ dram vial was added 0.0167g (0.103 mmols) of 4-hydroxycoumarin, 0.0152g (0.104 mmols) of trans-4-phenyl-3-buten-2-one, 0.0023g (0.0108 mmols) (R,R)-diphenylethylenediamine, 0.2mL of anhydrous tetrahydrofuran and 0.057mL acetic acid. The vial was swirled and allowed to react for 14 days. To a 4/6 dram vial 0.3248g (2.0 mmols) of 4-hydroxycoumarin, 0.3086g (2.11 mmols) trans-4-phenyl-3-buten-2-one, 0.0429g (0.20 mmols) of (R,R)-diphenylethylenediamine, 4.0mL of anhydrous tetrahydrofuran, and 1.14mL of acetic acid was added. The vial was swirled and left to react for a week. Refer to figure 1 for

the reaction and figure 2 for the whole mechanism that occurs in both vials.

Figure 1: The reaction that occurs in both experiments

After a week of reacting, the 4/6 vial, a CH2Cl2 TLC chamber was made and Anisaldehyde stain was used to see the spots after it was ran (figure 3). The solvent and acetic acid were removed from the vial by rotary evaporation. A small amount of residue was dissolved in boiling acetone. Then boiling water was added drop wise until mixture turned cloudy. Crystals were heated until they dissolved. The liquid was cooled to room temperature and put into an ice bath. Recrystallization occurred and the crystals were collected and carefully rinsed with ice cold 4:1 acetone/water. The crystals were left to dry. After two weeks of reacting, the ½ dram vial, a TLC was run using a gel plate (figure 4), CH2Cl2, and Anisaldehyde stain. Solvent and acetic acid were removed by a stream of air. A small amount of crude produce was set aside as a TLC sample. A microscale column was prepared with glass wool, sand, and slurry of silica in CH2Cl2. The product was dissolved in CH2Cl2 and added to the column. Eight, 1 mL fractions were then collected using 1% methanol/ CH2Cl2. Fractions 1,3,5,7, and 8 were spotted on the TLC plate with the crude product spot. Fractions 2,4 and 6 were inconclusive. The active spots were checked with UV light. The product was concentrated by rotary evaporation. 1H NMR spectrums were obtained for both products and the yield of each product was calculated. RESULTS: From the recrystallization experiment 0.39g of product was obtained along with 0.01g of product from the column experiment. The NMR spectrum from recrystallization indicated warfarin (figure 5) whereas; the column NMR had one peak that indicated just the solvent and not the product. The recrystallization TLC plate showed one pure product and the column TLC plate showed impure products.


CONCLUSIONS: From the masses it is evident that the recrystallization leads to more of a product being produced. The 1H NMR spectrum of the recrystallization supports this by giving a spectrum that has a product with 3 peaks. The 1H NMR for the column showed a spectrum with a peak indicative of chloroform and no peaks of warfarin which agrees with the 0.01g of impure product that was collected. The TLC plate from the recrystallization showed one spot of product. The TLC plate from the column showed three different spots of impure products. With these observed results it can be concluded that the recrystallization purification procedure is more profitable and produces the purest products. ABBREVIATIONS FDA, Food and Drug Administration. TLC, thin layer chromatography. NMR, Nuclear Magnetic Resonance.

April 29, 2014. What is Warfarin. NewsMedical.net. Oqbru, O. 2014. Warfarin, Coumadin, Jantoven. MedicineNet.com 2014. A Green, Enantioselective Synthesis of Warfarin.

Figure 3: TLC plate from Figure 4: TLC plate from column recrystallization.

Figure 2: The mechanism for the synthesis of Warfarin


Sarah Paddock; Norma Valdez Wells College 170 Main Street Aurora, New York 13026 KEYWORDS (Grignard reagent, malachite green, crystal violet) ABSTRACT: Grignard reagents were used to synthesize the dyes: malachite green and crystal violet. The dyes were used to determine the dying properties of polyester, nylon, silk, and cotton. The dye adhered to polyester and cotton due to charge attraction that silk and nylon lack. INTRODUCTION In 1912, Victor Grignard won the Nobel Prize in Chemistry for his method of hydrogenating organic compounds by means of finely divided metals.1 Grignard reagents are prepared by a reaction of organohalides with magnesium metal and reacted with carbonyl compounds to yield alcohols.2 Just as carbonyl reduction involves addition of a hydride ion nucleophile to the C=O bond, Grignard reactions involve an addition of a carbanion nucleophile.2 Crystal violet and malachite green are well known dyes that undergo Grignard reactions.2 The synthesis of Grignard reagents: malachite green and crystal violet, are shown in Figure 1 and Figure 2. Alternative dyeing methods include arenediazonium salts coupled to activate aromatic rings, such as phenols and arylamines, to yield brightly colored azo compounds.2 Diazonium coupling reactions are typical electrophilic aromatic substitutions in which the positively charged diazonium ion is the electrophile that reacts with the phenol or arylamine.2 Azo compounds make various red, orange, and yellow colors. The purpose of dyeing polyester, nylon, silk, and cotton using malachite green and crystal violet was to determine the dying properties of the fabric.

Figure 1. The Grignard reagent synthesis of malachite green.

Figure 2. The Grignard reagent synthesis of crystal violet. METHODS The Grignard reagents used to yield malachite green and crystal violet were prepared by conducting the same basic procedures twice. A condenser and RBF were rinsed with 5 mL of THF and a drying tube was placed on the top of the condenser. 2.5 grams (0.0125 moles) of 4-bromoN,N-dimethylaniline, 0.4 grams (0.0165 moles) of Mg turnings, 25 mL (0.31 moles) of THF, and 3 crystals of iodine were added to the RBF. The mixture was heated in a water bath for 30 minutes maintaining a gentle reflux. To synthesize malachite green, 0.21 grams (0.00154 moles) of methyl benzoate and 1.0 mL (0.012 moles) of THF were added to the Grignard reagent. The mixture was refluxed for an additional 5 minutes and 5 mL (0.204 moles) of 5% HCl was added. To synthesize crystal violet, 0.30 grams (0.00254 moles) of diethyl carbonate and 1.0 mL (0.012 moles) of THF were added to the Grignard reagent. The mixture refluxed for an additional 5 minutes and 5 mL (0.204 moles) of 5% HCl was added. Four fabrics; polyester, nylon, silk, and cotton were dyed. Small quantities of the dyes were added to separate beakers with 20 mL of boiling water. The four fabrics were dyed using two different methods (1) fabrics were soaked in Na2CO3 and dried prior to dyeing and (2) fabrics were not soaked prior to dyeing. The fabrics were dyed in the


boiling bath for 5 minutes and washed with soap and water. This process was repeated for the 16 pieces of cloth. RESULTS The qualitative analysis of the dyeing processes of malachite green and crystal violet resulted in deep vibrant colors on the polyester and cotton fabrics but light coloring or no coloring on the nylon and silk fabrics. When soaked in Na2CO3 and dyed with malachite green the polyester and cotton fabrics resembled a deep turquoise color, the nylon and silk fabrics showed no coloring (Table 1). When not soaked in Na2CO3 and dyed with malachite green the polyester and cotton fabrics resembled a light blue color, the nylon and silk fabrics showed no coloring (Table 1). When soaked and not soaked in Na2CO3 and dyed with crystal violet the polyester and cotton fabrics resembled a deep dark blue color, the nylon and silk fabric resembled a light blue color (Table 2). The resulting pH for malachite green was 8 and the resulting pH for crystal violet was 10. Fabric

Soaked in Na2Co3

Not Soaked in Na2Co3

Polyester

Deep Turquoise

Light Blue

Nylon

Washed Out

Washed Out

Silk

Washed Out

Washed Out

Cotton

Deep Turquoise

Light Sky Blue

Table 1. Fabrics Dyed in Malachite Green Fabric

Soaked in Na2Co3

Not Soaked in Na2Co3

Polyester

Deep Dark Blue

Deep Dark Blue

Nylon

Light Bold Blue

Light Blue

Silk

Light Bold Blue

Light Blue

Cotton

Deep Dark Blue

Deep Dark Blue

treated with acid to create a positive dipole charge on the double bonded nitrogen. Because malachite green and crystal violet contain a partial positive charge and are basic solutions, the dying properties of the four different fabrics can be determined. Polyester and cotton have a partial negative charge allowing the dyes to easily attach to the fabric. It has been reported that the adsorption of positively charged cationic surfactants increases with the negative charge of polyester and cotton fibers in basic solutions.4 Concluding that polyester and cotton did easily adhere to the dye. Nylon and silk have an overall neutral charge making it difficult for the dyes to attach to these fabrics. The possible solution for the effect of pH on adsorption of the monascorubrin was due to the ionic interactions of the pigment anions with the protonated amino groups on the nylon and silk fibers.5 Concluding that nylon and silk did not easily adhere to the dye.

Dr. Lauren O’Neil, Angelo Papagelos, and Jamyra Young for outstanding planning techniques, laboratory assistance, organization and patience.

RBF, round bottom flask. THF, tetrahydrofuran.

1. 2. 3. 4.

Table 2. Fabrics Dyed with Crystal Violet CONCLUSION The synthesis that malachite green and crystal violet experienced was the reaction of organohalides with magnesium and reacted with a carbonyl compound.2 The carbonyl compound was then attacked by a carbanion and

5.

Nobel Lectures, Chemistry 1901-1921, Elsevier Publishing Company, Amsterdam, 1966. McMurry, J. Organic Chemistry; 8th ed; Nelson Education, Ltd.: Canada. 2012. 635, 971-972 Meagly, R. and Taber B. Journal of Chemical Education. 1996, vol 73, issue 3, 259 Grancaric, A.M., T. Pusic, I. Soljacic, and V. Ribitsch, Influence of Electrokinetic Potential on Adsorption of Cationic Surfactants, Textile Chemist Colorist 29(12): 33 (1997). Yen, P. H., and Chen, K. M., Preparation and Properties of Novel Low-foaming Dyeing Properties of Anionic Derivatives of Polyoxyethylenated Stearylamine, J. Soc. Dyers Colour 115, 88-91 (1999).


Julie Cavanaugh and Dino Constantine Wells College, 170 Main St, Aurora, NY 13026 ABSTRACT: Four grams of used printer paper was turned into ethanol by saccharification, then the glucose product was isolated. TLC of the glucose product and fermentation were done. Distillation was completed a few days later in order to get our final ethanol product of 0.9654g or 0.209M. Recycling the used printer paper produced more ethanol than recycling the same amount of newspaper through the same process.

Ethanol is most commonly used as an additive to automotive fuel, but is also used for many other purposes such as beauty products and cleaning products. In 2012, the United States used over 12.9 billion gallons of ethanol. Currently, corn and grain are the two main sources of cellulosic material used to produce ethanol. Many efforts are being made in order to make ethanol production more efficient and environmentally friendly. For our experiment we tested a method of making ethanol out of cellulosic material that involved saccharification, fermentation and distillation to produce a solution containing mostly water and ethanol. Last semester we made ethanol from newspaper and this semester we decided to make ethanol from used computer paper. Repeating the experiment with a variant source of cellulosic material will show us which cellulosic material produces the highest percent yield of ethanol. We believed that the computer paper would produce more ethanol than the newspaper because there is more cellulose by weight in used printer paper than in newspaper since newspaper is more saturated with ink than used printer paper.

Four grams of computer paper was cut into small squares and put into a 250mL beaker with 12mL of 75% H2SO4(aq). The mix was stirred until a paste was made. 20mL of 80°C water was added. The paste was moved into a 125mL Erlenmeyer flask. The solution was heated for 90 minutes, then cooled and stoppered until next week. The hydrolysate was vacuum filtrated with celite over the filter. The filtrate was transferred into a 250mL Erlenmeyer flask and was treated with 11.5g of Ca(OH)2. The CaSO4 precipitate was filtered by suction filtration. The paste was agitated with 25mL water and filtered again using new filter paper. The pH was tested and 5M HCl or 5M NaOH was added until the pH was 7. The mix was heated to around 100 degrees Celsius. The sugar solution was allowed to evaporate for several days. Thin layer chromatography was done to test the purity of the sugar produced compared with a glucose standard solution. The sugar was dissolved in 15mL fermentation medium containing 0.2g KH2PO4, 0.2g NH4Cl, and 0.1g MgSO4 per 100 mL. 0.2g of

dried Fleishmann’s Yeast was added to the solution and the solution was left in a growth chamber to ferment for 4 days. The resulting mixture was decanted into a 50mL round bottom flask. A fractional distillation apparatus was assembled with copper sponge used as column packing material. The solution was fractionally distilled and liquid was collected until 100 degrees Celsius was reached.

The mass of the distillate was 11.76g and the volume was 11.89mL. The mass of ethanol produced was 0.9654g or 0.209M. The amount of cellulose from the paper used was assumed to be 2g based on the calculations used of the experiment done with newspaper, which means the percent yield of ethanol was 48.3% from the 4g of used printer paper used.

Compared to the first trial's results, which had a yield of 0.24g of ethanol, roughly four times the mass of ethanol was produced this trial. These results agree with our hypothesis that stated using computer paper would produce a higher percent yield of ethanol. In this trial, saccharification was completed in much less time than in the first trial. The mixture of H2SO4 and paper was homogenized in 30 minutes. After the 30 minutes, no change in the mixture was observed for the remaining hour of heating and agitating with a stir bar. The liquid produced had similar viscosity to water with grey settling matter. However, in the trial using newspaper, a viscous black sludge was produced after saccharification and heating for 90 minutes. Vacuum filtration was easier during the second trial because the liquid was less viscous and could be filtered in less time. By using computer paper instead of newspaper, the energy costs due to the extra 60 minutes of heating during saccharification could be cut. The dried glucose had a mild smell of caramel and was a clear yellow color, compared to the glucose from the first trial that presented a slightly darker yellow color with a slightly more burned smell. This burned odor is probably due to the prolonged heating that is necessary when doing this experiment with newspaper. This extra heating is also wasteful because it


decomposes sugars that would have been used in the fermentation process. Using used printer paper instead of newspaper to produce ethanol would reduce electricity

We would like to thank Wells College for the use of their lab and equipment and Professor Lauren O'neil for her assistance and encouragement

and heat usage, the length of the vacuum filtration process, and time needed to complete saccharification by roughly 60 minutes.


Rachel Nichols and Kylie Nishioka Wells College 170 Main St. Aurora, NY 13026 ABSTRACT: The purpose of this experiment was to compare the yields and stereochemistry of products obtained from the reaction of different dienes that react with maleic anhydride. The Diels-Alder reactions of Furan and E,E2,4-hexadien-1-ol with maleic anhydride were tested and compared. The reaction involving Furan was found to have 40.3% yield, 100% of which was the exo adduct. This implies the reaction was under thermodynamic control. The reaction involving E,E-2,4-hexadien-1-ol was found to produce the endo adduct, with a 36.3% yield. For these cases the use of a different diene did not cause a significant difference in the percent yield. Introduction: Diels-Alder reactions are often taught in undergraduate

organic chemistry courses. The most common Diels-Alder reaction done in an undergraduate laboratory involves the use of a cyclopentadiene and maleic anhydride. This reaction works quite well and generates an easily isolated product. The concepts of the mechanism, stereochemical results, and overall scope and limitations of the reaction are observed during this experiment (1). An aromatic cyclic compound such as furan can behave as a conjugated diene in the case of a Diels-Alder reaction with a dienophile such as maleic anhydride. The result is a bicyclic compound with either exo or endo stereochemistry depending on whether or not the reaction is under thermodynamic or kinetic control, respectively.

Mechanism 1: Reaction of Furan with Maleic Anhydride

The reaction of trans, trans- 2,4- hexadien-1-ol and maleic anhydride in refluxing toluene for 5 minutes generates a white crystalline material upon cooling. This produc can be isolated by vacuum filtration, with yields ranging from 50% to 90%. The reaction can also be carried out in microscale amounts (50mg of each reactant) or in gram scale amounts without significantly diminishing the yields. Anaylsis of the 1H NMR and COSY allows for the identification of the product:

Equation 1. Mechanism of E,E-2,4-hexadien-1-ol reacting with Maleic Anhydride, with intermediate product The product forms an intermediate anhydride, which is followed by the intramolecular cleavage of the anhydride by the alcohol. The stereochemical outcome can be explained by the endo rule to give all cis, perfectly positioning the hydroxymethyl group to attack the reactive anhydride (1). Procedure: Preparation of Furan: 1.2g (0.0122 mol) of maleic anhydride was dissolved in 10 mL of anhydrous ethyl ether in a beaker and warmed gently. Any ethyl ether that evaporated was replaced and the solution was allowed to cool before the addition of 1 mL (0.01375 mol) of furan. The vial was capped and covered in parafilm to allow the reaction to take place at room temperature. The crystallized adduct was collected by vacuum filtration and then recrystallized by heating to just boiling in 5 mL of hexanes. 1-3 mL ethyl acetate was added to dissolve with stirring. The crude product was collected by vacuum filtration and washed on filter with cold hexanes. The product was air-dried and the mass and melting point were measured. E,E-2,4-hexadien-1-ol Preparation: A mixture of 0.040g (0.004075 mol) of maleic anhydride and 0.040g (0.004079 mol) of E,E-2,4-hexadien-1ol in 5 mL of toluene was heated under reflux for 5 minutes. Thin layer chromatography, using ethyl acetate as developing agent, was performed to determine that no starting material remained as the reaction progressed.


The solution was then cooled to room temperature, during which time a white crystalline product deposited on sides of the beaker. The reaction mixture was cooled in an ice/water bath for 10 minutes to ensure complete crystallization of the product. The product was isolated by vacuum filtration and washed with cold toluene. When dry, the mass and melting point of the product were measured. Results: The reaction of furan and maleic anhydride gave a 40.3% yield, 100% of which was the exo adduct, as the melting range was 114-116째C. The reaction of E,E-2,4hexadien-1-ol and maleic anhydride gave a 36.3% yield, all of which was assumed to be the endo adduct because the melting point was 158-160째 Celsius. 1HNMR

spectra and 2D NMR spectra of the two products were taken and shown as follows:

Figure 3. E,E-2,4-hexadien-1-ol H NMR Spectrum

Figure 4. Furan H NMR Spectrum Figure 2. E,E-2,4-hexadien-1-ol COSY Spectrum

Discussion:

Figure 3. COSY Spectrum of Furan

This experiment was designed to compare the yields and adducts obtained in Diels-Alder reactions utilizing the dienophile maleic anhydride and two different dienes: furan and E,E-2,4-hexadien-1-ol. When maleic anhydride and furan were allowed to react, the resulting product was found to have a melting point of 114-116째C, which is characteristic of the exo adduct. Further proof for this was found in the H NMR scan of the product, with three of the hydrogen environments on the scan correlating with the three hydrogen environments of the exo adduct. 0.004953 moles of product was obtained, giving a 40.3% yield. When maleic anhydride and E,E-2,4-hexadien-1-ol were allowed to react, the resulting product was found to have a melting point of 158-160째C, which is characteristic of the endo adduct. H NMR and 2D NMR


COSY scans further revealed that the adduct obtained appeared to be endo. Furthermore, the COSY scan revealed hydrogen interactions at 1 and 2.75 ppm, 2.5 and 5.5 ppm, and at 3.5 and 5.5 ppm, all indicative of this product having endo stereochemistry. 0.001478 moles of product was obtained, giving a 36.3% yield. When the use of the two different dienes is compared, it can be seen that the percent yield was similar, with no diene giving a significantly greater yield than the other (40.3% versus 36.4%). Also noteworthy is that while the furan reaction yielded the exo adduct, reaction with E,E-2,4-hexadien1-ol gave the endo adduct.

We would like to thank Wells College for the use of their laboratories and equipment. We would also like to thank Angelo Pa-

pagelos for helping us to obtain our spectra and Melena Hagstrom for assisting us during our lab period. Thank you to Dr. O’Neil for showing us this experiment and for your time and patience, without which we could not have done this experiment.

mL; milliliter, g; grams, C; Celsius

The Diels-Alder Reaction of 2,4-hexadien-1-ol with Maleic Anhydride. McDaniel, K; Weekly,R. Journal of Chemical Education. December 1997. Vol 74. No. 12. pp. 1465-1467


Jessica M. Gulvin and Emily R. Guzman 170 Main Street Wells College Aurora NY KEYWORDS Biodiesel, energy, diesel, John Deere, diesel conversion, corn oil, Tomion Farms ABSTRACT: Biodiesel is an alternative energy that has been developed to combat the growing carbon dioxide levels throughout the world. The conversion from the use of petroleum based diesel fuel to the use of biodiesel fuel produced from animal fats and waste oil is an attempt to decrease a cause of global warming. However, biodiesel use involves a mechanical complication on diesel engines and takes an astronomical amount of time to produce. The study of the production of diesel from both animal fat and waste vegetable oil concludes that the production of biodiesel from animal fat and waste vegetable oil is not practical for tractor use on Tomion Farms, a common family farm, by this production method. In addition, it is not an efficient use of natural resources. INTRODUCTION The US biodiesel industry set a new record in 2013 producing 1.8 billion gallons of biodiesel1. Biodiesel can be used to run diesel engines, clean oil spills, degrease tools, heat homes, to generate electricity, to add lubricity to diesel fuel, and help remove paint and adhesives2. Biodiesel can also be used as a crop adjuvant and for cooking and illumination instead of kerosene2. Currently 90% of the biodiesel produced in the US is made by the transesterfication of soybean oil but animal fat proves to be a more economical alternative. Soybean oil is 33 cents pound and chicken fat is only 19 cents pound3. However, animal fat, lard and soybean oils are not waste products. These materials are also used in the production of wax paper, crayons, soap, livestock feed, pet food, shampoo, food, and among many other things4. The use of seed and vegetable oils for biodiesel also limits resources that are needed for other processes. The production of biofuel, therefore, directly results in a rise in feedstock prices for farmers. The use of animal fats and seed or vegetable oils to produce diesel prevent other industries from the benefits that these resources could have provided them. The production of biodiesel from yellow grease, a type of waste cooking or vegetable oil, is much less expensive than soybean oil but its supply is limited and it has other uses such as animal feed additives and in the production of soaps and detergents5. Yellow grease is half the price of soybean oil5. Biodiesel production from yellow grease comes closer to meeting the of cost petroleum diesel than is biodiesel from soybean oil. The use of waste vegetable oil (WVO) is a likely source but, over time, the triglyceride breaks down significantly and is not profitable in biodiesel production. In addition, the available supply of yellow grease limits its use for biodiesel production to 100 million gallons per year or 6,523 barrels or less per day5. Biodiesel is produced by a transesterification reaction in which the triglyceride is reacted with an alcohol. In this reaction, an organic acid ester is converted to a different

ester of the same acid, biodiesel6. Transesterification removes the glycerin stem from the molecule, resulting in a smaller desirable molecule that is operational as engine fuel (figure 1). Regular diesel fuel is manufactured through fractional distillation wherein diesel is produced when boiled between 250-350⁰C7. An average of 10 gallons of diesel fuel is produced from 1 barrel, 42 gallons, of crude oil8.

Figure 1: Transesterification of biodiesel The flash point of diesel is between 60-80⁰C and the boiling point is between 200-350⁰C, whereas biodiesel has a flash point of 100-138⁰C and a boiling point of 315400⁰C7. Therefore, biodiesel is safer to store, but harder for an engine to ignite and is not as adaptable in diesel engines7. It also results in a shorter shelf life as the fuel is unable to resist chemical changes during long-term storage9. Biodiesel does not burn as hot as diesel10 and results in a significant power loss when used in a diesel engine. Tomion Farms, a popular family farm located in Penn Yan


New York, uses diesel in all tractors to successfully plant and harvest produce. For the small average farmer, biodiesel proves to be both uneconomical and unreliable. METHODS11,12 Biodiesel was prepared from waste vegetable oil from the Wells College dining hall. The WVO was filtered through double-layered cheese cloth to remove unwanted particulates. 10mL of 2-propanol, 1.00mL WVO, and 3 drops of phenolphthalein were titrated with 1.85mL of 10M NaOH until a light pink colored solution was produced. 10mL WVO, 2mL methanol, and 0.114mL of 10.0M KOH were mixed and heated in a Erlenmeyer flask at 50⁰C for 30 minutes. The solution was centrifuged for 5 minutes. The glycerol (bottom layer) was discarded and the biodiesel (top layer) was combined with 2mL of 0.1M acetic acid and centrifuged for 5 minutes. The glycerol was discarded and the biodiesel was mixed with 2mL of 0.1M acetic acid. The mixture was centrifuged for 5 minutes. The final product was heated for 15 minutes at 80-90⁰C. The resulting mass was 6.6461g and color was a gold/dark colored. Biodiesel was also prepared from Smithfield Recipe Ready™ cooking lard. The lard was melted and filtered through double-layered cheese cloth. 100g lard, 30mL methanol, and 4g KOH were combined and heated at 60⁰C for 2 hours with stirring. The solution was transferred to a separatory funnel and the glycerol (bottom layer) was separated from the biodiesel (top layer). The remaining biodiesel was washed with 100mL of 50% v/v weak acid water solution using phosphoric acid. The glycerol layer was drained and the biodiesel was washed with 75 mL H2O. An additional wash with 50mL H2O was completed to neutralize the solution. The mixture was dehydrated using magnesium sulfate anhydrous. The solution was filtered by gravity to remove the absorbent and the final product was weighed to be 54.18g having a clear/white color. Characterization tests were performed on biodiesel from lard, biodiesel from WVO, and traditional diesel. The density of each diesel sample was determined. Rubber degradation was determined by cutting fuel line into 3 small weighed pieces. Each piece was placed in a preweighed Erlenmeyer flask containing either WVO biodiesel, lard biodiesel, or road diesel. The rubber was left in the diesel product for 3 days. The rubber was removed and its weight was recorded. Each Erlenmeyer flask containing a diesel type was weighed and recorded. The acidity/basicity of each product was determined by adding 5 drops of a single product into 1mL of deionized water and testing it with pH paper. Cloud point was determined by putting 3mL of each diesel into a graduated cylinder. The starting temperature was recorded and the diesel samples were stored in a -13oC freezer for 20 minutes. The temperature was recorded immediately after removing the product from the freezer. When the gel melted the temperature was also recorded. A combustion test was performed using glass rods and cotton balls to make three

mini torches. The flammability of each diesel fuel was tested by dipping a torch into a diesel sample and setting it on fire with a match in the hood. The time the torch burned, the color, and the presence of smoke was recorded for each sample. RESULTS Biodiesel production from waste vegetable oil resulted in 6.6461g (9.49mL, 0.0025gal) of gold/dark colored product and had a preparation time of 1 hour 37 minutes. Production of biodiesel from lard had a preparation time of 22 hours 28 minutes for 54.18g (58.26mL, 0.0154gal) of clear/white product. The degradation of rubber revealed a 0.026g increase in rubber weight and 0.061g decrease in WVO biodiesel, 0.0446g increase in rubber weight and 0.0876g decrease in lard biodiesel, and 0.0187g increase in rubber weight and 0.029 decrease in diesel. The density of each product was 0.7g/mL for WVO, 0.93g/mL for lard, and 0.803 for diesel. pH was 7.5 for WVO, 7 for lard, and 7 for diesel. The cloud point, the moment when the solution un-gelled, of WVO was -1⁰C to 5⁰C and -2⁰C to 5⁰C for lard. Diesel did not gel in the freezer and reached a temperature of -2.5⁰C. The combustion test showed that WVO biodiesel was flammable, burned blue, the flames were small, and very little smoke was given off. The lard biodiesel was also flammable and gave off a yellow/orange flame. The small amount of smoke that was given off was black. When diesel was burned the flame was yellow/orange and a lot of black smoke was given off. The duration of the burn was not recorded due to excessive flames and smoke. DISCUSSION Tomion Farms specializes in growing home-grown fruits and vegetables. The farm uses an annual average 500 gallons of diesel fuel per month, or 16.1 gallons per day. To produce the needed biodiesel to run the tractors for this farm for one day of work, 644,000g (1,419.78lbs) of lard would be required and it would take 434.7 days in this laboratory setting to produce. Production of biodiesel from WVO would require 10,454.5mL (2.76gal) of WVO from the Wells College dining hall and take 978.8 days to produce. In a large-scale production factory 20 gallons of biodiesel are made every 6 hours, which would take 4.83hours to produce enough biodiesel for the farm13. The production of the needed diesel from petroleum would require 61.18gal of crude oil, or 1.46 barrels8, and takes an average of 0.38s to produce14. Mechanically, biodiesel is destructive to the injection system of the diesel engine, specifically the injection pumps and fuel injectors15. Biodiesel has different solvent properties than regular diesel. Through repeated exposure, it can degrade or seep through seals, gaskets, hoses, glues, and plastics16. This was not visible in the rubber degradation test due to a minimal exposure time, but it was visible that fuel hose absorbs diesel materials. Fuel lines composed of brass, bronze, copper, lead, tin, or


zinc can accelerate the oxidation of biodiesel and create deposits throughout the engine and fuel systems16. It has also been known to break down deposits of residue within the fuel lines resulting in particulates, which clog the system17. In addition, in temperatures below 17⁰C, biodiesel gels and hinders the moving components. The use of animal fats and WVO is unrealistic in the production of biodiesel due to the wide use of these components in other necessary products. In comparison, diesel is produced from petroleum that is used for other crude products, but these products distill out at different temperatures and thus comprise a different fraction of crude distillate. Biodiesel is an unhealthy and impractical alternative to diesel fuel for the tractors at Tomion Farms as it hinders the engines components and is not nearly as time effective to make when compared to regular diesel fuel. The only benefit for the use of biodiesel is lower emissions and less environmental pollution. However, biodiesel’s foul effect on the motors of tractors and trucks, in

addition to its long production time, makes petroleumbased diesel the preferred choice for a small farm system.

A special thanks to Dr. Lauren O’Neil for her guidance and the use of her lab, Jeremiah Tomion of Tomion Farms for a diesel mechanic perspective and farm diesel use information, the Wells College Dining Hall for their donation of WVO, Richard Gulvin for his donation of road diesel and fuel hose, and Angelo Papagelos and Jamyra Young for their support during labs.

WVO; waste vegetable oil, mL; milliliter, M; molar, g; gram, H2O; water, KOH; potassium hydroxide, Ca(OH)2; calcium hydroxide, NaOH; sodium hydroxide, gal; gallon.

9. 1. 2. 3. 4. 5. 6. 7. 8.

World Grain Staff; Biodiesel industry sets a new record, World Grain. 2014 Ali, A; Biodiesel Production From Residual Animal Fat Using Various Catalysts, Pakistan Journal of Science. 2012 Kage, B; Startup hopes to convert chicken fat into biodiesel, Natural News Network. 2013 Mohammed, R; Animal Rendering Products In More Places Than You Think, Rense. 2003 Rodich, A; Biodiesel performance, costs, and use, Energy Information Agency. 2013 Frisby, J; Biodiesel Fuel, University of Missouri Extension, 1993 Srinivas; Difference Between Diesel and Biodiesel, Know it’s Difference. 2011 US Energy Information Administration. 2013

10. 11. 12. 13. 14. 15. 16.

Gerpen, J. et al.; Basics of the Transesterification Reaction. Biodiesel Production. 2005 Wedel, R; Marine Biodiesel. Cyto Culture Environmental Biotechnology. 1999 Mata, T; Sustainable Production of Biodiesel from Tallow, Lard and Poultry Fat and It’s Quality Evaluation. LEPAE. 2009 Lab #7: Preparation and Evaluation of Biodiesel from Waste Vegetable Oil, Buffalo State University. 2013 Biodiesel Equipment, Processors, Acc. & Information. 2002. Processing & Refining Crude Oil, Chevron Pascagoula Refinery. 2014 Corkwell, K; Diesel Fuel Injector Performance: The difference is in The Additives. Biodiesel Magazine. 2008 Deere Company; What Every Biodiesel User Needs to Know, John Deere. 2014


Chris A Castro and Thomas G Rions-Maehren Wells College, 170 Main Street, Aurora, NY 13026 ABSTRACT: Dyes are important in society for their ability to give clothing distinct colors. Dyes are also important to chemists for their unique UV-Vis absorptions. The indigo dye molecule, despite being known for its color showed a major peak in the UV region, while showing a minor peak in the visible region which accounts for its color. The amount of energy of these absorptions can be altered by the solvent in which the dye is dissolved. This is known as solvatochromism. Solvatochromatic effects can determine information about the excited state of a molecule. A bathochromic shift was observed in the first excited state of the molecule, showing increased polarity, and a hypsochromic shift was observed in the second excited state of the molecule, showing decreased polarity.

Dyes have played an important role in society for millennia due to their unique and often vivid colors. These very same dyes are important to chemists as well. Their color results from the absorption of light in the visible spectrum which allows an electron to enter an excited state. Solvents may also have an effect on the UV-Vis absorption of dyes. This is known as solvatochromism. Solvents can cause a bathochromic or red shift, in which the absorbed wavelength of light is increased. They may also cause a hypsochromic or blue shift in which the wavelength of light absorbed is decreased1. Increasing the wavelength of light decreases its energy and â„Žđ?‘? vice versa according to the equation: đ??¸ = . Therefore đ?œ† a bathochromic shift corresponds to a decrease in the energy required to excite an electron. This occurs if Îźgr (the dipole moment of the ground state) is smaller than Îźex (the dipole moment of the excited state) in a polar solvent, and a hypsochromic shift results when Îźgr is greater than Îźex in a polar solvent1. This is caused by the Franck-Condon principle, which states that electrons move very rapidly compared to nuclei2. This means that when an electron is excited in a solute molecule, the solvent will not immediately be rearranged to suite the polarity of the excited solute2. This effect was first used to quantitatively describe the polarity of a solvent1. Several methods for this have been devised, but he most commonly used scale today for quantifying solvent polarity is the ET(30) scale1. Dyes have been the focus for much research in the field of solvatochromism due to their unique UV-Vis absorptions. One such study on a dye known as Nile Red showed a bathochromic shift when dissolved in water3. This allowed the authors to determine important information regarding the structure, polarity, and molecular orbitals of the excited molecule3. This study has similar goals but with Indigo dye. Wyman and Zarnegar propose a scheme for the excited states of indigo dye 4 (scheme 1). This shows that the excited states of the dye

contain hydrogen bonds to oxygen, and it could be hypothesized that this would result in a bathochromic shift for both excited states. SCHEME 1.

DYE SYNTHESIS A solution containing 12.0 mL acetone and 1.50g (9.9x 10-3 mol) of O-Nitrobenzaldehyde was refluxed for 20 minutes. A solution containing 0.98g of NaOH in 10mL of H2O was slowly added to the refluxing solution. The solution was cooled to room temperature and 10mL of H2O was added to the solution. The solution was recrystallized in an ice bath, and the crystals were filtered by vacuum and washed with H2O and acetone.

0.50mM solutions were prepared in various solvents using the recovered crystals. Solutions were diluted as needed to make recovering a spectrum feasible. UV-Vis spectra of all of the solutions were recovered with a wavelength range of 200-700nm. The Îťmax values were determined in both the UV and visible ranges of each solution.


TABLE 1. Max absorptions and polarities of the solvents tested 0.92g (3.5x10-3mol) of crystals were recovered Solvent

UV λmax (nm)

Gas Phase

Vis λmax (nm)

ET(30) (kcal/mol)

362

Acetone

322

600

42.2

DMF

292

609

63.1

Ethanol

239

690

43.8

H2O

205

695

51.9

(70.8% recovery). The results from the spectra are displayed in table 1. ET (30) data was taken from literature1, and the gas phase value was obtained through TD-DFT. CHART 1.UV Absorbance vs. Solvent Polarity

UV λmax (nm)

380

Gas Phase

330

Acetone

280

DMF Ethanol

230

Water

180 0

20

40

60

80

ET(30) CHART 2.Visible Absorbance vs. Solvent Polarity

700

Water

680

Vis λmax (nm)

meaning various shades of blue and indigo were emitted. This explains the unique color of indigo dye. Chart 1 shows that as the polarity of the solvent increased, the λmax decreased (a hypsochromic shift). This would disprove the hypothesis stated in the introduction that μex is greater than μgr in the second excited state of the molecule. The decreased wavelengths show that the transition energy from the ground state to the second excited state is increased with the polarity of the solvent. This means that the excited molecule is less polar than its ground state. The ground state molecule (structure 1 in scheme 1) has an electron withdrawing carbonyl group and an electron donating nitrogen that resembles the nitrogen in pyrrole in the heterocycle that cause a dipole moment in the molecule. The excited molecule (structure 3 in scheme 1) has an electron donating hydroxyl group and an electron withdrawing nitrogen that resembles the nitrogen in pyridine. Since both molecules have an electron donating group and an electron withdrawing group on the heterocycle, the effect of the hydrogen bonding must be balanced out by the “pyridine” nitrogen, resulting in a net reduction in the dipole moment of the molecule. Chart 2 shows a bathochromic shift, meaning that the first excited state of the molecule (structure 2 in scheme 1) does have a larger dipole moment than the ground state of the molecule. This may be due to both electron withdrawing groups being on one of the heterocycles, while both electron donating groups are on the other heterocycle. Another interesting result was that the absorbance in the UV range was always greater than that in the visible range. This would suggest that the higher energy transition, where both of the carbonyl oxygens have been converted to alcohols (structure 3 in scheme 1), is favored over the first energy transition (structure 2 in scheme 1).

Ethanol

660 640

1.

620

DMF

600

2. 3.

Acetone

580 0

20

40

60

80

ET(30)

4.

5.

The wavelengths that were absorbed in the visible region for all of the solvents correspond to orange light5,

Connors, K.A. Chemical Kinetics: the study of reaction rates in solution; VCH: New York, 1991; pp 435-438 Reichardt, C. Chem. Rev. 1994, 94, 2319-2358 Wiley Online Library: Modeling Solvatochromism of Nile Red in Water. Wileyonlinelibrary.com DOI 10.1002/qua.22655 (accessed April 13, 2014) Wyman, G.M.; Zarnegar, B.M. J. Phys. Chem. 1973, 77, 1204-1207 Visible and Ultraviolet Spectroscopy. http://www2.chemistry.msu.edu/faculty/reusch/virttxtjml/spectrpy/uv-vis/spectrum.htm (accessed April 14, 2014


Mikayla R. Kravetz and Marisa J. Paradise Wells College, 170 Main St, Aurora, NY 13026 Green Chemistry, Aldol Condensation, Dry grind, Cyanoamides ABSTRACT: Solvent free methods were used in reactions of benzaldehyde, p-hydroxybenzaldehyde, o-nitrobenzaldehyde and p-chlorobenzaldehyde with cyanoamides by a Knoevenagel condensation in order to analyze substituent effects on benzaldehyde carbonyl carbon electrophilic reactivity. No change occurred in the reaction involving o-nitrobenzaldehyde, indicating reduced electrophilic activity of the benzaldehyde. 1H NMR data for the products of the remaining three reactions were inconclusive, indicating that a more conventional approach involving solvents may be necessary to give a product for comparison with products from solvent-free reactions.

The use of green chemistry methods, including solvent-free reactions, is useful in reducing energy and materials costs and lowering the amount of wastes produced in organic synthesis.1 Knoevenagel condensation is one method of producing carbon-carbon bonds that has been achieved using solvent-free methods.1,2 Different substituted benzaldehydes were used in order to investigate substituent effects on reactivity and yields Benzaldehydes with electron-withdrawing substituents ortho and para to the aldehyde should have a greater partial positive charge on the carbonyl carbon, resulting in greater electrophilic behavior of the carbonyl carbon; higher yields were therefore predicted for o-nitrobenzaldehyde relative to benzaldehyde. Electron-donating substituents decrease the partial positive charge on the carbonyl carbon and thus decrease the electrophilic character of the carbonyl carbon; lower yields were therefore predicted for p-hydroxybenzaldehyde and p-chlorobenzaldehyde compared to benzaldehyde.

To form N-tert-butyl-2-cyanoacetemide in-situ to react with benzaldehyde to produce the Knoevenagel product, 9.93 mmol of ethyl cyanoacetate and 9.93 mmol of tert-butylamine were mixed in a mortar with a pestle with one drop of triethylamine catalyst for ten minutes. 9.95 mmol of benzaldehyde was added to the mortar and this mixture was ground with a pestle for five minutes until it solidified. The product was weighed and analyzed by 1H NMR. The same procedure was used for three different substituted benzaldehydes. Aniline was used in place of tert-butlyamine to form the cyanoamide reactant in-situ. 9.19 mmol ethyl cyanoacetate and 10.30mmol aniline were mixed with 10.00 mmol of p-hydroxybenzaldehyde with a drop of triethylamine. 9.00 mmol ethyl cyanoacetate and 10.00

mmol aniline were mixed with 10.10 mmol o-nitrobenzaldehyde and a drop of triethylamine. 8.25 mmol ethyl cyanoacetate and 10.00 mmol aniline were mixed with p-chlorobenzaldehyde and a drop of triethylamine. For each mixture, grinding continued until the products solidified. Masses of each product were measured, and each of the products was analyzed by 1H NMR. Results After 3.5 minutes of grinding, a rust-orange, semi-solid substance was formed; this weighed 0.8560 g and was designated Product A (Figure 1a). The reaction between aniline and 4-chlorobenzaldehyde formed a maroon, semi-solid substance after approximately 15 minutes of grinding; this weighed 2.6607 g and was designated Product B (Figure 1b). The reaction between aniline and p-hydroxybenzaldehyde solidified in two seconds of grinding into a dark yellow, semi-solid substance; this weighed 2.6905 g and was designated Product C (Figure 1c). Moles of the products could not be calculated from these masses because their exact composition was unknown. No evident change occurred in the attempted reaction with o-nitrobenzaldehyde. 1H NMR spectra indicate significant amounts of starting materials remaining in reaction mixtures for Products A, B and C. Discussion The lack of spectral evidence that any Knoevenagel products were formed indicates that it is not possible to conclude from these results any difference in reactivity between the three benzaldehydes which did undergo an obvious reaction. However, the o-nitrobenzaldeyde did not appear to react, indicating that this electon-withdrawing substituent may have reduced the electrophilic behavior of the carbonyl carbon by an unknown mechanism. This phenomenon is consistent with the results of McCluskey at et al.


(2002), who achieved a yield of only 12% with p-nitrobenzaldehyde while all other benzaldehydes used gave yields greater than 50%.2

Figure 2. Mechanism in curved-arrow formalism for synthesis of Knoevenagel products.

Thank you to Dr. O’Neil for her guidance and knowledge, to Dr. Schwab for her suggestions prior to the beginning of research, and to Jamyra Young and Angelo Papagelos for their assistance throughout the experiment.

Figure 1. Structures of theoretical Knoevenagel products. Attempting a Knoevenagel Condensation by more conventional, solvent-based methods prior to experimentation with solvent-free methods may produce a higher yield of a product with greater purity. If this is done before the green chemistry solvent-free method is attempted, physical properties including color and appearance as well as NMR spectra of the products of the two different methods may be compared. Figure 2 depicts a proposed mechanism for the attempted Knoevenagel condensations.

1. Green Chemistry Task Force Committee, DST. Oxidation Reactions. Monograph on Green Chemistry Laboratory Experiments. N.d. 38-46 2. McCluskey, A.; Robinson, P. J.; Hill, T.; Scott, J. L. and J. K. Edwards. Green Chemistry approaches to the Knovenagel condensation: comparison of ethanol, water and solventfree (dry grind) methods. Tetrahedron Lett. 2002, 43, 31173120.


Journal of the Organic Laboratory, Volume 4