General Acid Catalysis A Flexible Experiment, Adaptable to Student Ability and Various Teaching Approaches R. S. Buimer, E. Senogles, and R. A. Thomas James Cook University of North Queensland, Townsville, Australia The acid-catalyzed hydrolysis of N-vinyl pyrrolidone, first reported by Breitenbach and co-workers (1)and more recently studied in our lahoratory (2),provides a simple and interesting spectrophotometric kinetic experiment which can he used to introduce general acid catalysis, solvent isotope effects, and other aspects of ionic reactions in solution to final year undergraduates. By having different student groups (usually student pairs) investigate different features of the reaction and pooling the results for evaluation and class discussion, the Brlinsted equation and the concept of linear free-energy changes can he covered also. N-vinyl pyrrolidone (VP) is hydrated in aqueous solutions of strong acids according to the following reaction sequence (1):
and the experimentally-determined second-order velocity constant ( k )is identified with h1. (h) Alternatively, the rate-determining step may he the addition of water to the carhonium ion, which is formed in aprior fast equilibrium reaction:
HlO.fest
VPH~H NVPHOH ~
+H~O+
If this mechanism applies, the rate of reaction is given by:
CHCH
where K is the equilibrium constant for reaction (6). In this case the reaction rate should again he proportional to [VPI[H30+], although the experimentally determined velocity constant is now identified with k7K. If mechanism (a) applies and proton transfer to VP occurs from ot,her acidic s ~ e c i e Dresent s in the solution. as well as from the hydronium ion, the reaction rate will depend upon the concentration of each of these acids and will be given hy an expression of the form:
CH,CHOH,
I
I
At O0C, 1-hydroxyl-1-(2-pyrrolidonyl) ethane (VPHOH) is stable; hut a t higher temperatures (-3O0C), it decomposes slowly into acetaldehyPe and pyrrolidone. The latter suhsequently reacts with VPH (1,2),in competition with reaction (2). However, reaction (2) is predominant a t the temperature of the experiment, and VPHOH is the main product. General Acid and Specific Hydrogen-ion Catalysis The product VPHOH can arise through two distinct kinetic mechanisms:
VP+H~O+~VPH+H~O
(4)
k,
followed by the rapid addition of water to give VPHOH. This mechanism leads to the rate-equation:
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Journal ot Chemical taucallon
-
where HA, k m . , etc. are the velocity constants of reactions of the type VP + H A VPH A-, and the reaction is said to exhihit general acid catalysis (3). On the o t . e r hand, if mechanism (h) applies, the concentration of VPH depends only on the hydrogen-ion concentration in the solution, regardless of the source of the wroton. Hence the reaction rate
+
In this case the reaction is said tdexhibit specific hydrogenion catalysis (3). I t is possible to distinguish between general acid and specific hydrogen-ion catalysis by carrying out the reaction in a range of buffer solutions a t constant pH (i.e., same [Acid]/ [Salt] value) but containing different concentrations of buffer (4). The ionic strength of each solution is adjusted to a con-
stant value hv t h e addition of a n inert electrolvte. ex. KC1. If the reaction rate increases with the concentration of acid HA, this is evidence for general acid catalysis a n d t h e value of HA-the acid catalysis constant-can he determined from a plot of the observed pseudo-first-order velocitv constant (h') versus carhoxylic acid concentration as given by: k'
+
= ~ I [ H J O + ]~ H A [ H A ]
(10)
On the other hand, if no effect on the reaction rate is observed then this is consistent with specific hydrogen-ion catalysis. Solvent Isotope Effect A distinction between the two mechanisms discussed above can he made also by comparing reaction rates measured i n ~ o l i e ,sa. smaller velocitv water and in deuterium oxide. If (a) . . a .. constant is expected in D 2 0 due t o t h e normal isotope effect on t h e rate-determinine sten. If (h) aonlies. t h e velocitv con-
Figure 1. Plot of pseudo-first-order velocity constant versus total buffer ooncentration. [ H A ] [A-1, for cyanaacetate buffer of [ H A ] l [ A - ] = 4(1), 1.5(2) and 0.67(3),Ionic strength (adjusted by the addition of KCI) = 1 mol dm+ Temperature = 25% Insert shows plots of slopes of the straight lines drawn in Figure 1 against fraction of cyanoacetic acid in buffer.
+
The Bronsted Catalysis Law For those reactions which are subject t o general acid catalysis, a n expected relationship exists ( 6 ) between t h e ionization constant (KA) of t h e acid and its efficiency as a catalyst. This relationship is summarized by the Bronsted Catalysis Law: HA = G A K A ~ (11) where G A a n d u are constants characteristic of the reaction, the solvent and temperature. Hence
Experimental' Students are supplied with a small amount of N~vinylpyrrolidone (Haven Chemical Ca.) which has been carefully distilled previously, and a middle fraction has been retained. This is stared in a refrigerator to prevent polymerization. Students make up the buffer solutions and measure the pH of each. The extent of reaction is followed by measuring the disappearance of VP speetrophotometrically. A Varian 534 spectrophotometer is used in our undergraduate experiments, with the eell compartment thermostatted. Acid or buffer solutions are introduced into two matched silica cells within the speetrophotometer. When thermal equilibrium is established (-10 min), a small quantity (sufficient to give a convenient absorbance value after temperature re-equilibra~ tion) of an aqueous solution of VP (-5 X lo-' mol dm-") is introduced into the sample eell from a micropipette. This is quickly and efficiently mixed by gently hrthhling a stream of air through the solution using a Pasteur pipette. After the re-establishment of temperature equilibrium (-5 m i d absorbance measurements are recorded on a eslibrated chart recorder. Reactions conducted in hydrochloric acid are studied at 233 nm, where VP has a maximum absorption coefficient (c = 15,800 dm" m ~ l -cm-'1. ~ Reactions in buffer solulions are followed at hieher wavelengths (245 nm cyanoacetate; 250 nm formate, acetate, 255 nm chloroacetate),where absorption by the buffer is insignificant. With hydrochloric acid and cyanoacetateand chloroaeetate buffer solutions the rate is sufficiently fast at 25'C for the reaction to be followed to high conversion, and the results are analyzed according to the firstorder integrated rate equation. With other buffers the rates are too slow for this procedureand a limited change in absorbance is measured, from which the reaction rate and velocity constant are evaluated by the initial rate technique. 0
~
Results T h e following results were obtained by our 1979 class of third-year students. Each group spent 4-6 hours, i n turn, on t h e experiment. Group I: These students studied t h e reaction with four hydrochloric acid solutions in the range 10-2 - 10-1 rnol dm-3 and then with one hydrochloric acid solution (0.06 - .10 mol
' Further details are available from the authors upon request.
Figure 2. Plot of logarithm of acid catalysis constant versus pK* forcyanoacetic acid (pKa = 2.47), chloroaceticacid (pK, = 2.871, formicacid (pK, = 375)and acetic acid (pKa = 4.76). Temperature = 25% k,, data determined at ionic strength of 1.0 mol d N 3 . pKa values from reference (9). dm-3) made u p in D 2 0 . T h e reaction was confirmed t o be second order with h = 2.6 X 1 0 P d r n b o l - I s-' a t 25°C. T h e solvent isotope effect was evaluated a s K H ~ O + / ~ D = 4.3, ~O which establishes t h a t the proton transfer reaction (4) is rate determinine. Group 2: This group studied the hydrochloric acid catalyzed reaction (IHCII = 0.08 mol dm-" a t different temoeratures in the range i 5 - 4 5 ' ~ , and a t 25'C with different ionic strenaths (UDt o 1 mol dm-"v t h e addition of KC1). T h e decompositi;n of V P H O H t o aietaldehyde and pyrrolidone and the subsequent reactions ( I , 2 ) , which occur a t higher temperatures, do not interfere with absorbance measurements a t t h e wavelength of t h e investiaations. Furthermore. since thew renvriun. tnke place aiter thv ri~te-~lt,tt rlninin: s t r p [as drtermined t ~ y ( ; r ~ , uI ps i ~ ~ d r n t[he? s ~ , have ~ ~ o r i i vgrn ( t the oi VP. l'h? hrrhenius parameters a e r c r . i t ~ , $disapprwancr f
+
Groups 3 and 4 studied the reaction in cyanoacetate buffer solutions of various concentrations a n d buffer ratios. T h e results are summarized in Fieure 1.which shows olots of t h e pseudo-first-order velocity o n s t a n t versus t h e k t a l buffer concentration a t various huffer ratios. T h e data support direct Volume 58
Number 9
September I981
739
catalysis of the reaction by the huffer and a plot of the slopes of the straight lines in Figure 1against the proportion of cyanoacetic acid in the huffer (see Fie. 1 insert). shows that it is the acid, rather than its conjugate base, which is the catalyst. Hence a general acid catalvsis mechanism is confirmed and hHAis evaluated as 1.1(4) x lo-" dm3 molkl s-'. G r o u ~ 5.6. s and 7 studied the reaction in the oresence of
It was assimed that, as with the cyanoacetate buffers, only the acid catalyzed the reaction. The value of HA was obtained from the slope of the graph of k' versus [HA] according to (10). The ordinate interceut correlated well with the h'value calculated, for the same hydrogen ion concentration and ionic strength as the huffer solution, from the results of Group 1and 2 students. In Figure 2 the combined results of the students are plotted 112). as a Brijnsted relation.. eiven hv* ean. . . . From this cu was evaluated as 0.5. This value can he compared with the literature value (2) and values determined for the carhoxvlic acid catalyzed hydrolysis of vinyl ethers, which proceed by ;similar mechanism (8). ~~~
-
Flexibility of the Experiment By varying the acids used, the experiment may he made suitable for either second- or third-year students. For instance, limiting the acid to HC1 only, i.e., the experiment performed hvour erouo 1. and bvomittine the use of D70, it is converted
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Journal of Chemical Education
into a second-year-level kinetics experiment determining orders with respect to VP and HsO+, and the rate constant h ~ . Usine., soectroohotometric analvsis in kinetics is often a novel . technique for students a t this level. Using one huffer a t constant huffer ratio hut varying concentrations would lead to results which a second-year student could discuss qualitatively and could form a nucleus on which later teaching could he based. As a third-year experiment, it could he performed as we have suggested, using groups and "pooling" results for discussion, or alternatively as a longer project for one pair of students. In any of the forms described, we suaaest that it could he used as a primary source of information'& the topic of acid-hase catalysis, to he followed by tutorial-type discussion. This contrasts with the conventional "lecture first" followed hy the "experiment-as-illustration" method of teaching. Literature Cited (11 Breifenhach. 1. W..Galinnusky, F.. Nesvadhs. H.. and Wdi. 8.. Mnnotsh., 87, 580 ,,o*n, ,.""",.
Sanoe1ei.E.. andThomas. R. A . J. Chem. Soc.. P r i k i n 11.825 11980). h.G.."Kinetiaand Mechanirrn".5th Ed.. John Wiiey pp. 204-2111.
~ hG.,"KinetiaandMechanlsm."5Lh , Ed.,JohnWiley
.""
Bell, R. P.. "Acid-& ,"9 A = , . , .
18-219.
Cafsly~i*"%d Ed., O x h d Oniuer~ityPress. London, 1949,pp.
(61 Froii,Anhw A , and Pearson. Ralph G.,"Kketinand Mechanirrn: gthEd., John Wiley and Sonr. Inc., New York, LOLO, pp. 209-214. (71 Laidlor. Keith J.. "Reaction Kinetirr. Volume Two.Reactions in Solution," Pergamnn Prcrr,Lundun. 1 9 6 3 , ~ 11-15,pp. ~. 2626. (8) K r e s ~ ~ , A . . l . . C h eH. n , L., Chisng. Y.Murril1.E.. P8yne.M. A , and Sapatv8.D. S.,J. Amsr Chrm. S o c , 83,413 11971). 19) K r r ~ r . A . . I . . a n dChian& Y., J. Amer Chem Snc.95,803 11973).
