Rose M. Lanes
and Donald G. Lee1 Pacific Lutheran University Tacoma, Washington 98447
I
I
Chromic Acid Oxidation of Alcohols A simple experiment on reaction rates
Despite the emphasis that a number of modern organic chemistry textbooks place on the effect of substituents on reaction rates there is a general scarcity of simple laboratory experiments that can be used to illustrate this phenomenon. I n an attempt to improvise such an experimentwe have found the chromic acid oxidation of substituted ethanols, X-CH2CH20H, to be well suited for this purpose. This reaction has the following advantages: (1) during the course of the reaction the solution undergoes a definite color change (bright yellow to pale green) that permits the progress of the reaction to be followed visually, (2) the rates are subject to large changes when the electronic nature of the substituents is varied, (3) the rates can be adjusted to give a convenient reaction time simply by changing the acidity of the reaction medium, (4) the entire experiment can be carried out using equipment that is readily available in the average laboratory, (5) it is a well known and important reaction with a readily available literature (1-7), (6) an understanding of the mechanism of the reaction does not require an advanced knowledge of organic chemistry on the part of the student, and (7) the results present the students with a set of facts that demand an explanation and that encourage many to pursue some further reading and independent thought. I n the experiment, as it was used here, pairs of students determined the relative rates of oxidation by adding exact quantities of each alcohol to a solution of chromium(V1) in dilute sulfuric acid and visually monitoring the course of the reaction. Since it was not possible to add all of the alcohols simultaneously, the rates of only two substituted ethanols could be compared during any one run. However, since the number of alcohols used was not large, the number of runs required to organize the total group in order of increasing rate of reaction was not burdensome. I n fact, most students found it a challenge to amange the alcohols in the correct order using the least possible number of runs. All of the students (the experiment has been used to date with over seventy students) were easily able to determine the correct order of reactivity in one laboratory period (2.5 hr); however the interpretation of the results required more effort. I n a pre-laboratory conference certain features of the reaction were described. I n particular, the work of Westheimer and his coworkers in establishing the reaction sequence by use of induced oxidations and deuterium labeling was discussed (1). Their work showed that the reaction 'Author to whom inquiries should be sent at the Department of Chemistry, University of Saskatchewan Regins Campus, Regina, Saskatchewan, Canada.
involved initial formation of a chromate ester which then underwent decomposition by a two-electron transfer with cleavage of the a-carbon-hydrogen bond being the rate determining step; i.e.,
-
n K
+ H,CrO, 3 RCH2OdrOH + H,O fast I1
RCHZOH
(1)
-
II + 2RCH=OH-+2Cr1" + RCH fast
2CrY
(4)
(Aldehydes are also susceptible to oxidation by chromic acid but under the conditions of the present experiment where the primary alcohols are present in excess little second stage oxidation would occur.) The students were also told, in order to prevent them from falling into a trap, that the effect of substituents on the equilibrium step (1) was not great. This is well illustrated in the table by comparing the exact oxidation rates for the alcohols (which were determined spectrophotometrically by one of us, R. M. L., as part of an undergraduate research project) with the equilibrium constants, which are available in the literature. It was also stressed that since step (2) is rate determining the effect of suhstituents should be explained in terns of carbon-hydrogen bond cleavage and electron transfer. Since the exact explanation of the effect of substituents on this reaction is still a matter of some doubt, the laboratory reports, as well as a subsequent examination question based on this experiment, were graded on the coherence of the discussion rather than on the basis of any "correct" explanation. The decrease in reaction rates observed for those alcohols bearing electron withdrawing suhstituents is difficult to reconcile with a proton transfer step. On the other hand, although more reasonable on the basis of these results, it is apparent that a simple hydride transfer step is not acceptable either: Oxidants which effect hydride transfers attack ethers readily (lo), but chromic acid does not oxidize most ethers a t an appreciable rate (11). A more satisfactory interpretation of the observed results is that electrons are transferred from carbon to chromium via a cyclic transition state, IV, that can be realized by any one of three possible electron flows; I (representing hydride transfer), I1 (representing proton transfer), or I11 (representing homolytic cleavage of the carbon-hydrogen bond) (2,12). Volume 45, Number 4, April 1968
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On the basis of this model the type of electron transfer would be determined by the degree of overlap of the carbon-hydrogen-oxygen and the carbon-oxygen-chromium orbitals in the transition state and could be dependent on the nature of R. I n any case, introduction of electron withdrawing substituents into the structure of the alcohol would make it more reluctant to release its electrons to chromium(V1) and thus reduce the rate of reaction. As can be seen from the table the rate of oxidation of methanol is abnormallv low. This has also been noted by RoEek (IS) and & probably a result of the low reactivity of a methyl group as compared to methylene groups. Hence it seems best, to us, not to include methanol in the series of alcohols employed for this experiment. Oxidation of Substituted Ethonols bv Chromic Acid [alcohol] = 9.23 X 10-P M; [H$Od = 3.56 M; [KCrzOd = 3.56 X lo-"; Temperature = 25.0 =t 0.1"C
Alcohol
k, (sec-I).
C,. b i n )
was ready, the two partners simultaneously added the chromic acid solutions to the alcohols, stoppered the test tubes, shook them briefly and proceeded to determine the relative reaction rates by a comparison of the intensity of color in each tube a t intervals. They then repeated the procedure with as many different pairs as required to arrange all of the alcohols in order of increasing ease of oxidation. The alcohols which gave us the best results were 1-propanol, ethanol, 2-methoxyethanol, 2-chloroethanol, and 2,2,2-trifluoroethanol. We also experimented with the use of 3-chloro-1-propanol but abandoned it because it was too difficult to purify and because good pseudo first order kinetics were not obtained when the exact rate was determined spectrophotometrically. Ethylene glycol can also be used if the concentration is halved to compensate for the presence of two hydroxyl groups in this molecule, hut 1,3-propanediol reacts abnormally fast because it can form a cyclic diester with chormic acid (14). Although it was not absolutely necessary we found it helpful to provide each group of 20 or 30 students with a colorimeter so that they could occasionally verify electronically what they were observing visually. (A wavelength of either 350 mp or 444 mp can be used.) Actually the only instance in which it is necessary to consider using instrumentation is with the ethanolpropanol pair. Since the rates for these two vary by only about 25%, it is sometimes difficult to be certain which solution is changing more rapidly.
KL Exod Rote Meosurementr
a
a
Pseudo first order rate constants. Equilibrium constant for ohromate ester formation. Glycol concentration reduced to 4.62 X lo-' M. Reference (8). Reference (7).
Although we have used this reaction solely as a demonstration of substituent effects on rates it is probable that it could be used to illustrate other rate phenomena as well. For example, a comparison of the rates a t different acidities or different temperatures could be used to show acid catalysis and temperature effects on reaction rates. Experimental Determinotion of Relative Rater
Each pair of students first prepared an oxidizing solution by addition of about 0.5 g of sodium dichromate dihydrate and 50 ml of concentrated sulfuric acid to 200ml of water. After allowing it to cool to room temperature, each student carefully filled a 10 ml graduated cylinder to the mark with the bright yellow chromic acid solution. He then secured exactly 3.0 ml of a 0.40 M solution of one of the alcohols in a 15 X 125 mm test tube, and his partner similarly obtained a diierent alcohol. (The alcohol solutions were prepared prior to the laboratory period and dispensed from 50-ml burets. Mouth-sucked pipets should not be used because some of the alcohols are detrimental to ones health if ingested.) When all 270
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The exact rates reported in the table were determined using a Beckman DU spectrophotometer fitted with a thermostated cell compartment. The oxidation medium was prepared by addition of 1.00 ml of a M solution of potassium dichromate to 3.91 X 10.0 ml of 3.91 M H2SOP. Aliquots (3.00 ml) of this solution were then pipetted into 1.00-cm cells which were thennostated at 25.0 1 O.l°C for several minutes. An excess of alcohol (2.77 X lo-* moles) was added from a microliter syringe fitted with a Chaney adaptor, and the reaction rates were followed by noting the absorbance (350 mp) a t intervals of time. The pseudo first order rate constants were then obtained from a plot of log absorbance against time. All of the alcohols were purified by fractional distillation before use. Acknowledgment
The authors are pleased to acknowledge several helpful discussions with Drs. Laurence Huestis and Charles Anderson. Literature Ciled
(1) WESTHEIMER, F. H., Chem. Rev., 45, 419 (1949). (2) STEWART, R., "Oxidation Mechanisms," W. A. Benjamin, Inc., New York, 1964, p. 3 7 4 8 . (3) WIBER~,K. B., in "Oxidation in Organic Chemistry" (Edilw: WIBER~,K. B.), Academic Press, Inc., New York, 1965, p. 142 ff. (4) WATERS,W. A,, "Mechanism of Oxidation of Organic Compounds," John Wiley & Sons, Inc., New York, 1964, p. 588.
(5) HOUSE, H. O., "Modern Synthetic Reactions," W. A. Benjamin, Inc., New York, 1965, p. 7E90.
(6) ROBERTS,J. D.,
AND CASERIO, M. C., "Basic Principles of Organic Chemistry," W. A. Benjamin, Inc., New York,
1965.. . 0. 401-405. (7) LEE, D. G.,in "Oxidation: ~ ~ ~and Application h ~ in i Organic Synthesis," (Editw: Augustine, R. L.) Marcel Dekker, Inc., New York, in pms.
(8) KLXNINQ,U., (1961).
AND
SYMONS, M. C . R., J . Chem. Soc., 3204
(9) STEWART, R., AND LEE, D. G., Unpubli~hedresults. (10) BARTER,R. M., AND LITTLER,J. S., J . Chem. Soe., (B), 205 (1967). AND WESTHEIMER, ~ (11)~ BROWNELL, ~ , R., LEO,A., CAANQ, Y. F. H., J . Am. Chem. Soe., 8 2 , 406 (1960). (12) STEWLILT, R., AND LEE,D. G., Can. J . Chem., 42,439 (1964). (13) R O ~ E KJ.,, Coll. Czech. Chem. Comm., 25, 1052 (1960). J.,, AND WESTHEIMER, F. H., J . Am. Chem. Soe., 84, (14) R O ~ E K 2241 (1962).
w.,
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