Muriel E. Finlayson and Donald G. Lee University of Saskatchewan Regina, Canada
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Oxidation of Ethanol by ~ h r o m i u m ( ~ l ) A kinetics experiment f o r freshmen
T h e oxidation of alcohols by chromium(V1) is a reaction of considerable importance and has been widely studied (1, 2). The progress of the reaction can be follo~ved visually or spectrophotometrically because the color of the solution changes from orange (HCrOa-) to green (Cr3+) (3). AS a consequence of this dramatic color change it has found application in the Breathalyzers presently being used throughout Canada. A measured volume of a person's breath is passed into a standardized chromium(V1) solution and any ethanol present reacts according to eqn. (1).
The decrease in absorbance is then a direct measure of the amount of alcohol present in his lungs. Consequently, this is a reaction with which some freshman students have had experience and in which most are often more than passively interested. It is therefore, in terms of the current vernacular, a reaction with relevance. Furthermore, we have observed that it can successfully be used to teach some of the basic principles of kinetics. Specificallywe asked our students to follow the rate of reaction titrametrically and to determine the order of the reaction with respect to HCrOn-. I n a prelaboratory conference it was pointed out that the complete rate law for the reaction was of the form V = k,lHCrO,-l*[CHsCH~OHl'[H+l' However, under conditions where the concentrations of ethanol and acid are large (and therefore approximately constant during the reaction) the rate law reduces to V = k[HCrOa-1". The problem is then t o determine a value for x (0, 1, or 2) by following the rate of reaction titrametrically (see experimental section for details) and by making the appropriate rate plots
(4, 6).
We have found that the experiment is one in which students have a good deal of interest and which gives clear cut results without the necessity of using equipment more elaborate than a pair of burets. Results obtained by an experienced demonstator are plotted in Figure 1; a typical student's results are plotted in Figure 2. After using this experiment with a large number of students we have observed that only 2'3, fail to identify the reaction as being first-order in HCr04-. I n addition to sewing as an introduction to the subject of kinetics, this experiment provides students with practice in the use of a large number of skills including weighing, titrating, preparation of standard solutions, standardizing solutions, and han-
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Figure 1. Kinetic d ~ t aobtained by an experienced worker. The ordinote has "nib of l/mole for the second-order plot, moler/l for the reroorder plot and dirnenrionlerr logarithmic "nib far the flrrt-order plot.
Volume 48, Number 7, July 1971
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473
an analytical balance and make up the solution in a 500-ml volumetric flask. Calculate the concentration of HCrOdpresent in the solution. 3. Rinse and fill an acid buret with the HCr04- solution. Rinse and fill a base buret with NsnSIOa sqlution (supplied in the laboratory). 4. Fmm the acid buret measure exactly 10.00 ml of HCrOdsolution into a clean 250-ml erlenmeyer flask. Add 4 ml of 3% asueous XI which will reduce HCrOl- and turn the saintidn red-brown. 2HCrO.(oran@)
+ 61- + 14H+
4
312 (brown)
+ 2Cr8+ + 8HzO (green)
Titrate with S n A 0 3 6duli011 to a peel) color and add 2 ml of 10C; aqueous T h y d e ~ ~(which e' qhould turn thp aolution dnrk blue,. Conrinuc titratinp, to a pie blurgreeu end point. The reaction here is
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Figure 2. Typisol student rerults. The ordinate has units of I/mols for the recond-erder plot, mobs/\ fw the zer-rder plot and dimemionless logarilhmic unitr for the Rrrt-order plot.
dling of stoichiometric equations. We have found that
it fits in well a t the end of a one-semester course. Experimental Section
Essentially the students were asked to prepare a solution of HCrOp-in 3.6 IMHCl, add some ethanol to it, and follow the rate of reaction by titrating an aliquot every ten minutes for 80 min. They then prepared plots of [HCr04-1, log [HCr04-] and l/[HCr04-] against time to determine if the order of the reaotion with respect to HCr0,- was 0, 1, or 2. Specifically they were given the following experimental procedure. I. Prrp~lre500 ml of 3.6b t HCL in a graduated 600-ml heakrr. 2. Prepare 500 ml of -0.0037 A1 K2ChOlrn 3.6 .If HCl (1% lute solutions of K2Crr0, dissociate in aqueous acid as folH~10 2HCrOl-.) Weigh the K2Cr10Ton lows: Crr07"
+
-
1 Available from the Fisher Scientific Co., Montreal. starch can he used a s an alternative.
474
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Journal of Chemical Education
Soluble
+ 2S%Os'-
+ 21-
+ SrOs+
~ i itep 4, the valrulnted 5. 1Jving the rwo balanced c q u a t i o ~ in molarity of H ( W i and thc volww of Ka,S90r solution r c quircd to rnlure 10.011ml oi H(:rO),i wlutim, rnlculnre the concentration of the Na2S10zsolution. 6. Using a graduated cylinder measure 200 ml of the HCrOlsolution into a 250-1111 erlcnmever flask. Then usine a. volumetric pipet add 2.0 ml of ah"s01ute ethanol and Gix well. Note the time of addition; it will be the sera time for this experiment. The concentration of HCrOa- a t zero time will, of course, be that calculnted in step 2. 7. Rinse and fill the acid buret with this CPH~OH:KCI~O,: HC1 mixture. 8. Refill Na,SlOs buret and repeat titrstions (step 4) every 10 min for 80 min. 9. Summarize your ~xperimentaldata as H rahlc giviug t l ~ e lime and a,rrwpordiw, volumes of NaAOa and conrcutrations, log&of wnwatration, and reviprocsl* of rlncrrtlrations of HCrOl-. 10. Plot each function ([HCrO,i], log[HCrOl-I and I / [HCrO,-1) against time on the same piece of graph psper and identifv the best straight line thereby establishing the order of the reaction with L ~ e c to t HC& 11. Calculate the rate constant for the reaction and the initial rate of reaction, assuming it to be fimt-order in ethanol and in [H+].
Literature Cited (1) W m m m ~ x s s F. , H., Chsm. RBD.,45. 419 (1949). (2) STEWART. R.. "Oxidation Mechanisms." W. A. Benjamin, Inc.. New York. 1964, pn. 37-48. (3) L m e s , R. M., ~ x LBE, o D. G., J. EDUO., 15, 269 (1968). (4) G-Y. H. B.. A N D H A I ~ X T G,. P.. JR.. "Basic Prinoiplea of Chemistry,'' W. A. Benjamin. Ino., New York, 1967, pp. 360-1. (5) L ~ e a u o x ,W. H., AND P A R ~ O N T.~ D.. , "General Chemistm." John Wilay & Sona,Inn., New York, 1966, D. 119.
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