In the Classroom
Got Bio? A Short Course Introducing Students to the Applications of Biochemistry Reid Chamberlain and Amy L. Rogers* Department of Chemistry and Biochemistry, College of Charleston, Charleston, SC 29424; *
[email protected] High school students from across the state of South Carolina convened on the College of Charleston campus, as they have every summer since 1976, for the Governor’s School of South Carolina (1). The students attended four weeks of a subject-concentration course and a global issues course of their choosing, plus a multitude of extracurricular activities designed to enrich their high school careers and prepare them for college. The subject-concentration course, Got Bio?, introduced twenty-three students to the biochemical workings responsible for numerous modern trends including tattoo pigments, tattoo removal, artificial sweeteners, fuel metabolism associated with dieting, and drug metabolism. Hopefully the high school students got the bio picture and gained an appreciation for the world around them; a world that is filled with quick fixes and long-term consequences. Course Format The course was designed to expose students to the applications of biochemistry. The four weeks were filled with intensive biochemistry lectures with fun hands-on activities, laboratories, and two field trips. The course was divided into three areas: chemistry of color; general fuel metabolism, and basic drug design. Lecture topics and complementary laboratory exercises or activities are listed in Table 1. Chemistry of Color The first week of Got Bio? focused on the chemistry of color and its applications in biochemistry. In the classroom, students were exposed to the different phenomena of color. Five basic phenomena were discussed: (i) vibration and excitations as seen in incandescence; (ii) ligand field effect as seen in transition-metal complexes; (iii) molecular orbital theory applied to organic molecules; (iv) energy bands of metals and alloys; and (v) geometric and physical optics (2). The classroom introductions were enhanced by hands-on activities and laboratory experiments to give a better understanding of some of the phenomena of color. Following the first classroom discussion, students assembled outside to observe the link between a soap bubble’s size and its unique coloration. We discussed the geometric and physical optics of constructive and destructive interference of light. Students experienced that when light waves, reflected off both the outside and inside surface of a bubble, are in the same phase (constructive interference), the intensity is enhanced. However, when the light waves are out-of-phase (destructive interference), the waves cancel and the complementary color is seen. The easiest way for students to conceptualize the color of compounds and relate the color to the molecular orbital theory is to study compounds familiar to them. The first labo-
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ratory experience spectroscopically examined three dyes found in food, drugs, and cosmetics (FD&C) using Ocean Optics spectrophotometers (3). Students examined the difference between the absorption and the emission of light, and determined the wavelength of the maximum absorption, λmax, for each dye. Students were able to visualize that when a molecule intensely absorbs a color of light, its complementary color will be seen. Next, we quantitatively determined the concentration of the FD&C dyes found in commercial drinks by plotting a Beer–Lambert law calibration plot of known concentrations of dye and solved for the unknown concentration with the linear regression equation (3).
Table 1. Course Syllabus Topics Introduction to Color
No. of Lectures
Laboratory
1
Geometric and physical optics
Color of Bubbles
Ligand field effect Molecular orbital theory Colored Compounds
2
FD&C dyes
UV–vis of FD&C dyes
Tattoo ink Introduction to Biochemistry
Tattoo pigs’ feet 2
Carbohydrates Lipids Proteins Caloric value of foods
Structures of Artificial Foods
Energy and specific heat: measuring the caloric value of food 1
Artificial sweeteners and artificial fats General Metabolism
Aspartame detection by UV–vis 3
Glycolysis
Glucose oxidase assay
Citric acid cycle Oxidative phosphorylation Lipid metabolism Amino acid metabolism General Drug Design
1
Structures of aspirin and Tylenol
Synthesis of aspirin
General mechanism of activation
Purity of commercial aspirin
Journal of Chemical Education • Vol. 85 No. 5 May 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Classroom
With an understanding of absorption, students were shown structures of the FD&C dyes as well as natural compounds such as β-carotene or lycopene (4) and were asked to describe common characteristics of the molecules. Immediately students could see the alternating double bonds and in some dyes, the azide bond. After further discussion of the molecular orbital theory, students were able to understand that the π to π* transition becomes smaller with greater conjugation giving a more red-shifted absorption spectrum and changing the color of the molecule (5). To study a third phenomenon of color, we explored the chemistry behind tattoos. Students were asked to independently research why tattoo inks have color. Students eagerly reported that metal salts are the pigmentation in tattoo ink, but few understood the chemistry behind the color. The metal salts found in some tattoo inks include compounds such as ferric oxide, ferrous oxide, mercury sulfide, cadmium sulfide, and chromium oxide and are suspended in varying undisclosed viscous media called carriers. Since tattoo inks are not approved by the Food and Drug Administration (6), manufacturers are not obligated to divulge the contents of the proprietary medium. Many of the tattoo dyes are toxic to the body either by the presence of heavy metals or by the degraded products found after irradiation by a laser (7, 8). We discussed the ligand field effect to explain that when a metal binds ligands, the electronic energy levels of the metal become nondegenerate and can therefore absorb visible light. Once the metal salt in a tattoo dye absorbs its characteristic wavelength of light, the complementary color can be seen. Harnessing each student’s interest is a crucial aspect of creating an appropriate environment for learning and appreciating biochemistry. For this reason, we designed a nonclassical, original laboratory experience of tattooing a pig’s foot to illustrate the chemical theories learned in lecture. Groups of four or five students arrived to lab staggered throughout the morning with a blueprint of their chosen tattoo design. With the aid of teaching assistants, students tattooed their design on a pig’s foot using a professional tattoo device as shown in Figures 1 and 2. Given
the permanent nature of a tattoo, this exercise allowed students to realize the importance of safety in a laboratory. Students were excited to discover the role of biochemistry in modern trends by this hands-on application. A suggested protocol can be found in the online supplement. After each group successfully tattooed a pig’s foot, the students and their tattoos were transported to the dermatologic office of Dr. Pearon Lang at the Medical University of South Carolina. Dr. Lang played two key roles in the students’ learning. First, he was able to teach students from a professional prospective by carefully explaining the mechanism of tattoo removal by laser spectroscopy and the importance of selecting the appropriate laser that targets the corresponding wavelength absorbed by the pigments. His second role in the students’ learning was visual as small groups of students watched Dr. Lang perform a laser removal treatment on their respective tattoos. The color from the pigments were not fully removed; however, this did not take away from the students’ appreciation for the procedure. The students left the dermatology office with a keen understanding of the laser tattoo removal process and, of equal importance, the firsthand experience of witnessing such a procedure.
Figure 1. Students using the tattoo device to tattoo their pig’s feet.
Figure 2. Tattooed pig’s foot by group 1.
Fuel Metabolism This section of the course aimed to give the students a general understanding of how the body metabolizes carbohydrates, lipids, and proteins into usable forms. The chemistry behind artificial sweeteners and fats was tied into the lectures to stimulate student discussion. Overall, the students’ understanding of the metabolic processes was expanded through class lectures followed by supporting labs. The students entered the course with varying levels of understanding of fuel metabolism. To address the different levels of student preparation, the lecture series first introduced the chemical structures of carbohydrates, lipids, and proteins (9). The students learned the major structural differences between the compounds and how their structures dictate their chemical
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In the Classroom
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Absorbance at 540 nm
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properties and functions. Next, general metabolic catabolism was introduced with topics including glycolysis, citric acid cycle, oxidative phosphorylation, lipid metabolism, and amino acid metabolism (9, 10). The class emphasized the conservative nature, complexity, and connectivity of the metabolic pathways without the expectation that the students be able to recall every detail. To help students assimilate and appreciate the massive quantity of detailed information given, we provided each student with a card-stock handout of the backbone of each cycle. During lectures, students filled in enzymes, cofactors, reagents, and products on the cycle as the material was covered. With a general understanding of metabolism, the second laboratory experience demonstrated how to calculate the caloric value of carbohydrates, lipids, and proteins using calorimetry. To familiarize students with basic calorimetry, students first calculated the heat lost from heated samples of aluminum shot to the surrounding water in a calorimeter (9). By calculating the heat lost (q) from the metal to the surroundings in calories, the students could quantitatively appreciate the quantity of energy transferred. The second portion of the laboratory experiment illustrated how to use this same technique to measure the quantity of energy a sample of food provided. Students knew that breaking down food produced energy but few understood how this related to the caloric values found on food labels. To demonstrate this, we adapted a calorimetry experiment from Timberlake’s laboratory manual that measured the heat gain of water above a burning sample of food (Cheezy Poofs) (11). First, students calculated the heat gained (q) in kilocalories for the water, which, to a first approximation, equaled the heat lost by the Cheezy Poofs. Students then compared their experimental heat in kilocalories to the number of calories1 calculated from the nutrition label on the bag of Cheezy Poofs. The numbers were not exact owing to heat lost to the aluminum can holding the water and the surrounding air, but it gave students an idea of how nutritionists determine the caloric value of foods. Next, we related the heat to ATP formation. When food is metabolized by the body (similar to burning), energy (measured in kJ or in kilocalories) can be released as heat or it can be harnessed through oxidative processes to synthesize ATP. Following the first law of thermodynamics, the total energy lost as heat during the students’ calorimetry experiment is equivalent to the quantity of energy available for transformation into the body’s usable form of energy, ATP. With an understanding of how food can provide usable energy, the discussion moved to the fate of excess calories due to overeating. Most often, excess energy must be stored by synthesizing either glycogen or fat. Instead of limiting the intake of food, America has created a market for a low-calorie substitute to help with weight control. To delve into the “low-calorie” phenomenon, students were introduced to several structures of artificial sweeteners and asked how they compared to sucrose. Students found it interesting that the structures of some sweetening compounds mimicked sucrose while others were completely different. Students quickly understood that sweet-tasting compounds are not necessarily comparable in structure. Sweetness is perceived when certain molecules bind to a receptor site in the mouth. One way to
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Wavelength / nm Figure 3. Glucose absorbance spectrum of a dilute solution of white grape juice and the glucose standard calibration plot (insert).
measure sweetness is by the human sweetness response. To demonstrate this, students were given four different solutions that contained the same concentration of sweetener. The four solutions were sucrose (table sugar), aspartame (Equal), sucralose (Splenda), and saccharin (Sweet-n-Low). Students were asked to rank the solutions from most to least sweet. Most students ranked the solutions as follows: sucralose, aspartame, saccharin, and sucrose. As a result of this demonstration, students concluded the reason artificial sweeteners are advertised as zero or low calorie was because low concentrations were needed to “sweeten” the food or beverage. Once an appreciation of the potency of artificial sweeteners was obtained, we went to the laboratory to determine the quantity of a sweetening compound that is present in a typical cup of artificially sweetened coffee. Aspartame absorbs in the UV range allowing UV–vis spectroscopy to be used to quantitatively measure the quantity of sweetener added. Having the previous experience of constructing a Beer–Lambert plot, students easily made four solutions of known concentrations of aspartame and graphed an equation of the line. Students received 50 mL of water and were asked to sweeten their “coffee” as they normally would. The “coffee” beverage was water because the caffeine in coffee also absorbs in the UV region and hides the absorbance of aspartame owing to its high concentration relative to aspartame. Based on the absorbance value at λmax for aspartame, a concentration in mol∙L could be calculated for the aspartame that was then converted to milligrams. To conclude the discussion of fuel metabolism, we used a laboratory experiment that emphasized the metabolic specificity of the human body. From lecture students gained an appreciation for the importance of enzymes and how an enzyme can only catalyze specific reactions. But to impress the remarkable specificity of enzymes, we used a glucose oxidase assay purchased from Sigma Chemical. The assay contained the appropriate enzymes to convert d-glucose into a pink-colored product that could be measured via UV–vis spectroscopy. Students ran assays
Journal of Chemical Education • Vol. 85 No. 5 May 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Classroom
on three different unknown solutions that contained glucose (white grape juice, orange juice, and a solution of l-glucose). The absorbance of each unknown was recorded and used to calculate the final concentration of glucose in the unknown solution via the standard curve created as a class (Figure 3). Students found that two of the three unknowns turned pink. Knowing that each unknown contained glucose, some students feared that they had made a mistake. Setting up the lab for unexpected results not only kept the students engaged in the activity, but also forced them to formulate an original explanation for their results. This led to a discussion of the minute structural difference between d- and l-glucose and provided an opportunity to reiterate that the smallest difference in structure can turn off the activity of an enzyme. Students concluded the lab by back-calculating to estimate the quantity of d-glucose found in a carton of white grape juice and orange juice. Drug Design One lecture was spent on basic drug design and how it relates to enzymes. Specifically, we discussed the differences in the structures of aspirin and acetaminophen and how they work in the body. Having discussed specificity of enzymes through metabolism, students could appreciate the effectiveness of a drug when it inhibits an enzyme. We concluded our laboratory experiments by synthesizing aspirin and testing the purity of commercial aspirin (3). Conclusions Relating biochemistry to everyday experiences is an excellent avenue to help students better understand and appreciate this difficult subject. According to student evaluations, students were receptive to the course. Most students found the coursework to be interesting and life-applicable. Others said this course was an amazing and insightful experience that they would recommend to anyone. A number of students noted that the material was complex and difficult to fully understand. Although the material was rigorous, the course provided a detailed overview that emphasized the connectivity of the intricate biochemical processes responsible for everyday observations. Allowing students to discover how biochemistry is part of their lives generates both a greater appreciation for and interest in the field of biochemistry. Acknowledgments We would first like to thank Dr. Pearon Lang for his time and resources in teaching the students about tattoo removal. We would also like to thank Hank Martin at the Medical University of South Carolina for talking to our students about the anatomy
and disorders of the brain. We also thank Linh Doan and Kelly King from the College of Charleston and Katie Horne from Furman University for voluntarily serving as teaching assistants for this course. Note 1. The calories listed on a food package are actually kilocalories. The word is sometimes capitalized to show the difference, but usually not.
Literature Cited 1. Deavor, James P. J. Chem. Educ. 1990, 67, 669–670. 2. Institute for Dynamic Educational Advancement Web Exhibits. http://webexhibits.org/causesofcolor/0.html (accessed Jan 2008). 3. Faber, G.; Martin, E. M.; Potts, G. E.; Tate, M. L.; Rogers, A. L. Introductory Chemistry, 6th ed.; Pearson Custom Publishing: Boston, 2005; pp 59–61. 4. Parkinson, Thomas M.; Brown, Joseph P. Ann. Rev. Nutr. 1981, 1, 175–205. 5. Volland, Walt. Bellevue Community College. http:// www.800mainstreet.com/elsp/Elsp.html (accessed Jan 2008). 6. U.S. Food and Drug Administration Center for Food Safety and Applied Nutrition. http://www.cfsan.fda.gov/~dms/cos-204.html (accessed Jan 2008). 7. Elkins, H. B. The Chemistry of Industrial Toxicology; Wiley: New York, 1950. 8. Vasold, R.; Naarmann, N.; Ulrich, H.; Fischer, D.; König, B.; Landthaler, M.; Bäumler, W. Photochemistry and Photobiology 2004, 80, 185–190. 9. Horton, H. R.; Moran, Laurence, A. M.; Ochs, R. S.; Rawn, J. D.; Scrimgeour, K. G. Principles of Biochemistry, 3rd ed.; Prentice Hall: Upper Saddle River, NJ, 2002. 10. Marks, D. B. Biochemistry, 3rd ed.; Lippincott Williams and Wilkins: Philadelphia, 1999. 11. Timberlake, K. C. Laboratory Manual General, Organic, and Biological Chemistry; Benjamin Cummings: San Francisco, 2002; pp 75–84.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2008/May/abs658.html Abstract and keywords Full text (PDF) Links to cited URLs and JCE articles
Color figures
Supplement
Tattoo procedure
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