Ashley M. Bryan and Patrick G. ~ ~ a f s s o n SUNY a t Albany Albany, New York 12203
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Activation Energy Determination
I
A n organic chemistry experiment
A u t h o r s of introductory texts in organic chemistry, in an effort to correlate reactions in terms of structural and mechanistic theories, are forced to introduce the concept of activation energy at an early stage in the student's career. Too often the ideas in thermodynamics are presented as mathematical formulations and arc accepted with limited comprehension. Consequent,ly the student tends to regard the concept of activation energy as a nebulous term unless he is required to work with this concept and possibly even to determine its value for one reaction. The decomposition of N-bromobenzamide in base provides a good experiment for the determination of kinetic data since it can he completed in one laboratory period. The rate of decomposition can readily be followed by a potassium-iodide-thiosulfate titration following quenching of the reaction in acetic acid. The Experiment
The N-bromobenzamide lyas prepared in advance1 of the experiment. A mixture of 75 ml of 2 N NaOH and 75 ml distilled a a t e r in an erlenmeyer was placed in a thermostated bath at the desired temperature. N-bromobenzamide (1.7 g) was added with vigorous stirring. Aliquot,~(10 ml) of the homogeneous solution were withdrawn by pipet at specified times and released into a solution of 4 ml of acetic acid and 20 ml of 5% potassium iodide. The liberated iodine was titrated with 0.02 N sodium thiosulfate to purple end-point with
starch indicator. The stoichiometry of the reaction is presented as follows
+
10
15
20
25
+
++ B21-r + I.
Thus the volume of thiosulfate used is proportional to the amount of undecomposed N-bromohenzamide remaining in the aliquot. I n these laboratories we have found it convenient to carry out the study with students working in pairs, each determining the rate constant k for a particular temperature. Four or five temperatures in the range between 0°C and 40°C have provided satisfactory rates
5
-
HOBr HC + Z I HsO 2s*oa2- I*4 S,06s-
'HAUSER, CHARLES R., AND REFROW, W. B., JR.,J . Am. Chem. Soc., 59, 121 (1937).
*HAUSER, CHARLES R., AND KANTOR, SIMON W., J. Am. Chem. Soe., 72, 4284 (1950).
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TIME ( M I N S I
Figure 1.
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Plot of log of thiorulfate volume versus time.
Journal o f Chemicol Education
Figure 2.
Plot of log k versus 1 /T used for tho determination of Ea.
for decompositions carried out in 1 N sodium hydroxide. The rate constant was calculated,
and a first-order plot was obtained by substituting the volume (in ml) of thiosulfate required to titrate the first aliquot for "a" and volumes required to titrate successive aliquots taken after elapsed time t (min) for (a - 2). Since the decompositions were proceeding at different rates, the students were instructed to obtain appropriate times for the third and subsequent aliquots in the following manner. The k value calculated from the times of the first two aliquots was divided into 0.25 to give the time for taking the third sample. Appropriate times for removing other samples were obtained by dividing the following numbers by k , 0.4, 0.6, 0.8, 1.1, 1.5, and 2.2. The value of k based on the second aliquot is genefally found to he less accurate than subsequent values. The students were asked to obtain graphical and
calculated values of k from their data. Plots of log of volume of thiosulfate versus time were obtained as in Figure 1. The k values obtained for each temperature were tabulated along with values obtained by other sections and classes from previous years. The cumulated data provided the students with an excellent opportunity to learn how raw data may be refined by statistical methods and also to evaluate their performance. A computer program was written to analyze each student's data. The computer output included calculated values for k, standard deviations of the mean, the observed probable error of the weighted mean, and the theoretical error of the weighted mean. The data from the experiment were also fitted to a least squares computer program to give the best straight line. The students were then required to utilize the refined data to obtain a plot of log k versus 1 / T (Fig. 2) from which the value of EA could be evaluated. Determination of the activation of the reaction was also obtained by solving the integrated form of the Arrhenius equation and the values compared with those obtained from the graphical data.
Volume 46, Number 4, April 1969
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