The Energetics of Aerobic versus Anaerobic Respiration A Biophysical Chemistry Laboratory Experiment Using an Adiabatic Calorimeter Timothy D. Champion Johnson C.Smith University, Charlotte, NC 28218 Richard W. Schwenz University of Northern Colorado. Greeley. CO 80639 The calorimetry of combustion processes remains a commonly performed experiment in physical chemistry laboratories ( I ) . The theory of bomb calorimetry has been well developed over a period of years (2) to include a variety of terms involved in the reaction as well as examination of the exact shapes of the temperature versus time graphs under a number of different conditions (3). Novel student exercises involving adiabatic calorimetry include the measurement of the energy content of pizza (4). We are in the process of preparinga one-semester physical chemistry course with laboratory for chemistry majors in pre-health and education emphases. Due to the biological interests of the students, we have tried to incorporate laboratory exercises that involve the students' interests in biochemical applications while continuing to illustrate several elementaw thermochemical urinci~lessuch as the measurement and &dculation of heats for processes, Hess's law, and the effect of t e m ~ e r a t u r eon thermodynamic properties. An example exercisiis the difference in energy gained from the aerobic and anaerobic oxidation of glucose. Resplratlon and Fermentallon The method in which glucose is catabolized provides a distinction between different types of life forms. Life forms that use 0 2 to catabolize glucose are referred to as aerobic. Life forms that catabolize glucose without Oz are called anaerobic. Life that may use either form are called facultative (5). In aerobic catabolism (respiration), glucose is oxidized to carbon dioxide and water: C,H,,Os
+602
-
K O 2 + 6H,O
The biochemical pathways involved in respiration are glycolysis, the Kreb's (tricarboxylic acid) cycle and the mitochondrial electron transport system (5). In anaerobic catabolism (fermentation), glucose is oxidized by a biological agent t o ethanol and carbon dioxide:
Fermentation involves only the glycolysis pathway (5). These processes are exergonic. The energy given off is used t o phosphorylate ADP t o ATP, which is the usual form of energy storage in living systems. T o compare the energetics of these two combustion reactions, calorimetry and Hess's law may be applied.' Experlmental In ~ r e ~ a r a t i ofor n student ~erformanceof this e x ~ e r i ment;se;eral trial runs on ethanol and glucose combustion were performed after standardization of the calorimeter us-
'
It has been suggested by a colleague that the example of partial wmbustion of glucose to lactate is another possibility for this sort of exoeriment. 528
Journal of Chemical Education
Table 1. ExDerlmental and Calculated Data in kJ mol-I
Ethanol Glucose
-1319 -2763
*
43
+ 58
-1366.8 -2803.2
'Data from ref 7.
ing benzoic acid. Standard procedures for the measurement of the temperature rise in the water bath surrounding the bomb calorimeter were used ( I ) . Good results and no ignition problems were observed for ethanol, but not for glucose. Examination of pure glucose pellets following several misfires suggested that the pure glucose does not ignite as easily as benzoic acid. The experimental procedure was then modified to use tablets formed from an equimolar mixture of benzoic acid and glucose following the procedure used in determination of the energy content of pizza without the drying steps (4). Calorimetry experiments using these mixed tablets had no ignition problems and gave results that agree both among themselves and with literature values (see Table 1).The heat produced for such a mixed substance tablet is simply the sum of the heats produced by the combustion of benzoic acid and glucose according to AU,,,
=Mass frnctionBAAUBAass f r a c t i ~ & ~ ~ ~ ~ , ~ ~
where BA is benzoic acid and GLU is glucose. Results and Calculations Table 1compares the literature values of AHfor the combustion of glucose and ethanol with those obtained experimentally by a combination of student and instructor use of the bomb calorimeter (10 runs total). The agreement of the experimental and literature values is quite good, with both values being somewhat low (apparently, the measured heat capacity of the calorimeter was low by less than 1%). In addition to the quality of these measurements, it is instructive to have the students use these measurements to calculate other thermodynamic properties. In particular, biochemists are interested in the value of AG a t 37 "C for both types of processes. AG can easily be calculated if AH and A S are both known a t 37 OC. Unfortunately, tabulated values for AH and A S are given a t 25 OC rather than 37 O C (6).There are two ways t o proceed based on known information. First, the adjustment of the tabulated values a t 25 'C to the value a t 37 'C can be done using the again tabulated values of the heat capacities for each substance and applying the standard formulas for the temperature dependence of AH and AS. Or second, the value of the entropy at 25 OC can be calculated using the methods given in the appendix of Klotz (7) that calculate the entropy based on the groups present in the molecules. After the entropy a t 25 " C is obtained, it is a relatively simple matter to calculate AG a t 25 OC and then calculate AG a t 37 OC through the use of the temperature dependence of AG, namely
Table 2.
vides some insight into the energetics of anaerobic versus aerobic processes in different species.
Data lor Rwplratlon Process&
AMss Wlmol ASssaJlml-K A
h kJ/mol
AGuo,,,
A@mw
Literature Clted
Textbook;Scott Forramen:Glenview. IL, 1988.
AG(T,)/T, - AC(T1)lT,= - A H I R ( l I T 2
- 1ITJ
Table 2 compares calculated values for a variety of thermodynamic functions with those obtained in this laboratory. Agreement hetween the calculated and experimental values is again good, as expected from the previous data, and pro-
, coiorimetry:~undomentolr ond~larrice: verlag-Chernie: ~w.;~ o h n e~G. ~ i ~ ~ ~
2. ~
,
Deerfield Beach, FL, 19%. 3. Canagaratna, S. G.; Wiff, J.J. C h m Edvc. 1388.65.126. 4. Stout, R, P.; Nettleton, F. E.; Price, L. M. J Chsm. Educ. 1385.62.5. 5. B O ~ D ~G. I , M. J. them E ~ U C1986.63,566: 1986.63.673: 1988.63.772. 6. Dean. J. A,, Ed. Long& Handbook of Chamisfry, 13th ed.: MeGraru-Hill: New Ymk.
,
1985.
,,
R. M. chemical ~
h
~
~4th ed.; ~ ~ ~
d~cum-~ ~ ~
rninga: ~ e n park, ~ o CA, 1986.
Volume 67
Number 6
June 1990
529
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