William P. Schaefer
California Institute of Technology Pasadena, California
A Kinetics Experiment for First Year Chemistry
During the past several years, the California Institutc of Technology has attempted to provide laboratory assignments in the freshman chemistry course that illustrate a variety of quantitative measurements (1). This has been done because of the conviction that every scientist or engineer must he able to design and carry out various kinds of quantitative experiments; he must also be able to analyze and interpret the results of experiments reported by others. For these reasons, quantitative determinations are used to demonstrate precise measurements of mass and volume; a colorimetric and a coulometric determination are assigned and a gasometric experiment is being developed. The kinetics experiment described here is used to introduce our students, all of whom must take first year chemistry, to still another type of measurement, that is, the use of radioactive tracers. Two considerations governed the choice of the type of experiment: appropriateness and cost. Since this work is done in the chemistry course, a strictly physical exercise, such as the determination of the half life of an isotope or the maximum energy of beta particles, was not considered suitable. Many of the possible chemical measurements could be performed with or without using radioisotopes; for example, a solubility product can usually be determined using either radiochemical or classical techniques. Such experiments were also less desirable; it seemed best to concentrate on an experiment which could be done only by using labeled atoms. The exchange of iodine atoms between organic and inorganic iodides is such an experiment. Five alkyl iodides (n-propyl, isopropyl, wbutyl, isobutyl, and sec-butyl iodides) have been used. Students determined the rate constants for the reaction, RI+I*-eRI* +I-
This experiment is also suitable from the standpoint of low cost. Iodine-131 is a readily available isotope and it requires no elaborate handling or extra-sensitive counting equipment. A simple survey meter can he used to record the activity of samples, but a Geiger tube detector and very simple scaler give much more reproducible counting rates. Instruments are available for less than $500 (2). A scintillation detector has been used with good success but is considerably more expensive. We have emphasized certain research aspects of this experiment to our students and find that these features add greatly to their interest. These can be appreciated Contribution No. 2963 from the Gates and Crellin Laboratories of Chemistry, California. Institute of Technology, Pasadena, California.
558 / Journal o f Chemical Education
after reviewing briefly the information available concerning iodine exchange reactions. The isotope 1281 was used as early as 1935 (9) to establish beyond doubt the connection between exchange and optical inversion. This study and others (4-7) were made more difficult because of the short half life (25 min) of the lZ8I.The eight-day iodine isotope, 13'1, has been used in some more recent studies. Most of the data concerning exchange reactions are recorded in Streitwieser's comprehensive discussion of displacement reactions a t saturated carbon atoms (8). There are three reasons for wanting to repeat some of this work. Most of the data available pertain to reactions in alcohol or alcohol-water solvents; a few studies have been carried out using acetone (9, 10) or acetonitrile (11, 192) as solvents. It is now recognized that solvolysis as well as exchange reactions are possible in alcohol solutions; for this reason hydroxylic solvent systems should be avoided. Second, inorganic salts are increasingly associated in solvents of low dielectric constant; the degree of association depends upon the ionic strength as well as the nature of the solvent. As Winstein has shown (lo), it may be possible to correct measured apparent rate constants for this association if the ionic strength and exact solvent composition are known; these data are often not mentioned in the early papers. Third, some of the early workers did not prove that the exchange reaction they studied obeyed the second-order rate law they assumed. It is clear that solvolysis or exchange reactiom of many alkyl iodides proceed by both S Nand ~ S Nmechanisms. ~ The overall rate of exchange must he apportioned properly between these two pathways in order to obtain proper rate constants for comparisons. We hope eventually to he able to use student data to determine corrected second-order rate constants for iodide exchange in acetonitrile. Students participating in the initial phases of the investigation seem enthusiastic about this aspect of their freshman chemistry work. The Experiment
Radioactive sodium iodide is obtained as an aqueous solution containing a small amount of sodium sulfite. This is added to inactive sodium iodide dissolved in technical grade acetonitrile to give the desired level of activity, and the total iodide concentration is then adjusted to the desired value. Because the acetonitrile is not specially dried nor purified, the small additional amount of water introduced with the radioactive iodide is not thought to be important. (Technical acetonitrile which is yellow is not a suitable solvent.) All manipulations involving radioactive solutions are done over an enameled tray placed in a hood.
Because of limited space in the hood, the temperature of the reaction mixtures is not controlled, but it is recorded. A 125-ml conical flask is used as a reaction vessel. Forty milliiliters of 0.1 M NaI in acetonitrile are run in from a buret and the alkyl iodide added by pipet. One-half milliiliters of n-butyl or n-propyl iodide, or 2.0 nd of isobutyl, sec-butyl, or isopropyl iodide give suitable reaction rates. Then 1.0 ml of the active NaI solution is added and the mixture is swirled; immediately a 5.0-ml sample is removed and run into a mixture of 5.0 ml CCL and 7.0 ml H,O in a small separatory funnel. (These amounts of CCh and H 2 0 should be closely adhered to; otherwise troublesome emulsions may form.) The funnel is shaken and the CCla layer drained off; then 0.2-ml portions of the aqueous layer are pipetted onto 4.5-cm filter-paper circles in aluminum planchets and counted using an end-window Geiger tube. The activity of the active NaI used is adjusted to give about 1000-2000 counts per minute from 0.2-ml portions a t this first counting. Further samples are taken throughout the period as the exchange proceeds. The alkyl iodides obtained commercially are sometimes discolored. They can he purified by shaking thcm with a slightly basic aqueous solution of sodium sulfite, drying the resulting water-white compound wii;h CaSO,, and distilling once through a short column. Their purity can easily be checked by vapor phase chromatography. Results
Counting liquid samples in the manner described is quite inefficient but our measurements show that counting rates are reproducible to 5-lo%, which is sufficiently accurate for student purposes. Table 1 presents some preliminary results, as well as representative rate constants taken from the literature for comparison. Student values differ from the literature values; the best agreement is found for n-butyl iodide, the only iodide for which a rate constant determined in acetonitrile was available for comparison. Differences in the other cases may be due to different solvents, different temperatures, or different ionic strengths. The variations in rate constants from student to student are generally smaller than the difference between the student average and the reported value. Table 1.
Compound n-propyl iodide
Iodide Exchange Rate Constants
Student Values (1 Mo1.-' sec-') 1.25 f 0.47 X lo-' (41 results)
dt
=
k 1 [I*-] [RI] - IRI'I 11-1
'
2.303 = I IRK 11-llt log
+
[RIl
"2 [[RI] + [I-]] [I la
Toble 2.
Typical Student Data and Results n-Butyl Iodide
(22OCI
Time sample removed
im, I",
..Y
- [I-]
The appropriate concentrations are calculated, the ratio [I*-],/[I*-lo evaluated from the ratio of the counting rates a t time t = t and t = 0, and k is calculated. I n actual practice, students are instructed to prepare two reaction mixtures and remove samples from them for counting a t 20-30 min intervals. By alternating from one to the other they are able to get three samples from each mixture in a 3-hr period; the data from these samples are used to calculate six values of the rate constant. Typical results for one student are shown in Table 2. Some comments on the validity of calculating the rate constants this way should be made. First, the assumption that [I-] is equal to the formal concentration of NaI may be wrong, although there is evidence (11) that NaI is a strong electrolyte in acetonitrile. If the NaI is associated, then the association constant must be known in order to calculate the actual I- concentration to use in the equation for the rate constant. Second, the value obtained for the rate constant will in general depend on the ionic strength of the reaction mixture, since the activity of the reacting I- will be affected by ionic strength (13). Thus, if calculated rate constants for different alkyl iodides are to be intercompared, they should all be corrected to some standard ionic strength. Third, calculation of the second-order rate constant by means of the given equation ignores the possibility that some exchange may be proceeding by a first-order process. This can be checked by performing the experiment with several
.
1 C
I
Brackets indicate molar concentrations, which are assumed equal to formal concentrations. The asterisk indicates a radioactive iodine atom. [RI*] and [I*-] are assumed very small with respect to [RI] and [I-]. The equation is rearranged and integrated to give
Corrected counting rate (o./min)
Calculated rate constant (I m-' sec-' X 109
195 136 9'3 Between tl and 1%
2.3 2 5 2.6
Literature Value 1
1.20 =t0.31 X (70 results)
--
3 . 7 X 10-"/m seo at 35°C in alcohol, p =
4 . 3 X 10-a I/m see at 2QPCin acetonitrile, " = 0.1 19) (2 results) / mat 5 . 5 X i ~ - ~ ~ isec isopropyl iodide 6 X 25'C in acetone, p = n,,."F. 1,.9,) 5.9 X 10"l/m sec at isobutyl iodide 1.1 X lo-' (5 results) 40°C in alcohol, p =
n-butyl iodide
The student values in Table 1 were obtained from the experimental data as follows. For these first experiments, the reaction was assumed to be second order with no complications so that
Fwst d a y tn t1 tz
1 :53 P.M. 2: 18 P.M. 2:40 P.M.
,", rm,
sec-butyl iodide 1.4 X lo-' (6 results) 7.1 X l/m sec at 45'C in ~lclcohol,p = 1..5~ (.9i ~. - ,
Average of students' determinations, all done in acetonitrile, 0.1 M NaI. T = 22-25T.
Average reported: 2 . 3 X 10-a l/msec. Volume 41, Number 10, October 1964
/ 559
different concentrations of the alkyl iodide in question, temperature fluctuations can be ignored or not. For a t constant ionic strength. The rate constants obboth normal and isoalkyl iodides, the second-order tained will not be the same unless the contribution to rate constant roughly doubles in going from 5 to 35OC; the rate of exchange from a firsborder process is negligithus temperature is an important variable. We are ble. It appears from first observations that there is a studying the effect of ionic strength on the value of the significant firsborder contribution to the exchange in rate constant, and we are checking to see that the reacthe case of sec-butyl iodide, and tertiary iodides will tions all obey the second-order rate law assumed; almost surely have large first-order contributions; the these latter studies will not be complete for some time. experiment is thus most suitable for primary iodides. All these aspects of the experiment excite student In fact, tertiary iodides not only exchange by an S N ~ interest; because of this interest and because the experimechanism, they also eliminate H I and form the ment introduces both chemical kinetics and the use of alkene; this complicates the interpretation of experiradioisotopes, we are incorporating it enthusiastically mental data even more. in our freshman program. Recently, a well-type scintillation detector has been The author thanks Walter Deal and David Lichinsky used for these measurements with appropriate modificafor prelin~inaryexperimental work, and Professor Jiirg tions in the experimental technique. Because larger Waser for criticism and encouragement. samples (2-5 ml) can he counted, because both countLiterature Cited ing efficiency and geometry are improved, and because (1) ei. SWIIT, E. H., J. CHEM.EDUC.35,248 (1958). faster counting rates are feasible (because of shorter (2) See LEWIN, S. Z., J. CHEM.EDUC.,38, A135 (1961) and "dead time"), the precision of the experimental deterfollowing articles. minations can be improved while a t the same time the (3) HUGAES,E., JULIOSBERGER, F., MASTERMAN, S., TOPLEY, total amount of radioactivity in the experiment is B., AND WEISS, J. Chem. SOL,1935,1525. (4) MCKAY,H., Nature, 139,283 (1937). decreased one thousandfold. This results in a signifi(5) MCKAY,H., J. Am. Chem. Soe., 65,702 (1943). cant increase in the safety of the experiment. (A C., AND LIND,S., J. Am. Chem. Soe., (6) HULL,D., S~BFLETT, selected group of students who performed this experi58, 535, 1822 (1936). ment with solutions sufficiently active for Geiger countJ. Am. Chem. Soc., 64,940(1942). (7) SEEWG,H., ANDHULL~D., (8) STREITWIESER, A,, Jr., Chem. Rev., 56, 571 (1956); also ing wore film badges during all their work (6-8 hr) "Solvolytio Displacement Reactions," MeGraw-Hill, New and no exposure to radiation was detected. HowYork, 1962. ever, the main danger of 1311 occurs when it is ingested, (9) DE LA MERE,P. B. D., J. Chem. Soc., 1955, 3196. so any decrease in the amount used is welcome.) L., SMITH,S., STEVENS, I., AND (10) WINSTEIN,S., SAVEDOFF, With this better detector we are studying three GALL,J., Tetrahedra Lettem, 1960, No. 9, 24. (11) H~DGSON, G., EVILNS, H., AND WINKLEU,C., Can. J. Chem., aspects of exchange reactions as part of a freshman 29, 60 (1951). honors research program. We are determining the B., Compt. Rad., 235,953 (1952). (12) MAY,S., AND GIRANDEL, temperature coefficient of the rate constant from ex(13) See, for example, HINE, J., "Physical Organic Chemistry," periments done in thermostatted vessels to find out if 2nd ed., McGraw-Hill, New Yark, 1962, p. 74.
Qualitative Organic (?) Chemistry (Department of Redfaced Apologies by the Editor's Office)
We wish we could restore the hair that must have been torn out by the roots from frustrated heads in attempts to solve the puzzle published on page 476 of the September 1964 issue. Abject apologies and complete exoneration are due Pmfessor James G. Traynham (LSU, Baton Rouge, La.) for the irksome errors which slipped into the copy in the Editor's office. We like to think that we do not make errors often-but when we do, they are real whizzem! The problem is repeated here, with the corrections in italics. Now we can expect readers to discover that Friedel has the acetone m d the gas chromatograph is in the green lab. Please forget about the lower half of page 476. We c a n n o t i t haunts us! In a certain chemistry building there are five laba, all in a row, each painted a different color and each containing a special item of apparatus. A different chemist works in each lab and keeps the main stock of one solvent and one reagent. The following information is sufficient to tell who has the acetone and where the gas chromatogrzph is located. Fischer works in the red lab. Grignard has the infrared spectrophotometcr. Ethyl ether is in the green lab. Kekuld has the benzene. The green lab is immediately to the right (your tight) of the ivory lab. The chemist with the phenylhydrazine has the magnetic stirrer. The iron(II1) chloride solution is in the yellow lab. The alcohol is in the middle lab. Friedel works in the first lab on the left. The chemist who has the phenyl isocyanate works in the lab nmt to the lab with the melting point apparatw. Iron(II1) chloride solution is kept in the lab next to the lab with the sublimation furnace. Tollens' reagent is in the same lab as the chloroform. Lucas keeps the supply of Lucas' reagent. Friedel works next to the blue lab.
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Journal of Chemical Education
