Author Institution: Department of Physical Science, Denison University, Granville, OhioThe decarboxylation of trichloroacetic acid has been found to occur readily at temperatures as low as 25.0°C in the presence of dimethylsulfoxide in water solutions, a fact not previously reported in chemical literature. The reaction rate is dependent upon the concentration of the dimethylsulfoxide. The reaction rate constant increases about sixty percent as the concentration of dimethylsulfoxide (DMSO) is increased from c. 50 percent to c. 86 percent. The reaction is first order and the rate is also influenced by the actual concentration of the trichloroacetate ion. An effect of the DMSO is to increase the concentration of this ion. The paper presents density curves at 20.00°C and 25.00°C (each±0.05 °C) for solutions of DMSO in water

Laque, William E.

Ronneberg, Conrad E.

The Ohio Journal of Science

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A STUDY OF THE DECARBOXYLATION OFTRICHLOROACETIC ACID IN SOLUTIONS OFWATER AND DIMETHYLSULFOXIDE1WILLIAM E. LAQUE AND CONRAD E. RONNEBERGDepartment of Physical Science, Denison University, Granville, OhioABSTRACTThe decarboxylation of trichloroacetic acid has been found to occur readily at tem-peratures as low as 25.0°C in the presence of dimethylsulfoxide in water solutions, a factnot previously reported in chemical literature. The reaction rate is dependent upon theconcentration of the dimethylsulfoxide. The reaction rate constant increases aboutsixty percent as the concentration of dimethylsulfoxide (DMSO) is increased from c. 50percent to c. 86 percent. The reaction is first order and the rate is also influenced by theactual concentration of the trichloroacetate ion. An effect of the DMSO is to increase theconcentration of this ion. The paper presents density curves at 20.00°C and 25.00°C(each±0.05 °C) for solutions of DMSO in water.INTRODUCTIONSince 1955 this laboratory (Denison University) has been concerned with thestudies pertaining to the relative strengths of trichloroacetic acid (HB), tri-fluoroacetic acid (HA), and hydrochloric acid. Inman, Barringer, and Ronneberg(1955) measured the rate constants for the hydrolysis of sucrose in water with vary-ing concentrations of HA, HB, and hydrochloric acid. Considering structure andproperties, HA, HB, and hydrochloric acid might be expected to differ in strengthin solution of comparable concentrations of suitable solvents. In water, however,the leveling effect is great enough to make them essentially equally strong, indi-cating complete ionization of dilute solutions of all three acids in water.In 1963 the authors attempted to differentiate the relative strengths of hydro-chloric acid, HA, and HB. The object was to find a solvent with a dielectricconstant sufficiently low as compared with water to permit the differentiation ofthe relative strengths of these three acids. The solvent chosen to test this hypo-thesis was dimethylsulfoxide [(CH3)2SO] in water solution, hereafter referred toas DMSO. It was chosen because it has a dielectric constant of about 46, ascompared with 78 for water. DMSO is an active proton acceptor.The pH's of various concentrations of HA and HB in 50 percent and 75 percentby weight DMSO and water solutions were measured. From the experimentallymeasured pH's, the apparent ionizations were calculated by comparing the mea-sured hydronium-ion concentration with the hydronium-ion concentration, assum-ing complete ionization. In concentrations greater than 0.025 N, a definite dif-ference was found in the relative strengths of HA, HB, and hydrochloric acid invarious concentrations of DMSO in water solutions.While determining the pH of solutions of the acids, no difficulty was met inobtaining reproducible results for HA, HB, and hydrochloric acid in water.Similarly no difficulty was met in determining the pH of HA in any of the DMSOsolutions. But the measured pH of HB in a DMSO solution was appreciablyhigher than the calculated pH. This was quite unexpected in view of the assump-tion that, in water solutions, HB is as strong an acid as HA. It was then observedthat DMSO solutions of HB were slowly effervescing. This could possibly meanthat carbon dioxide was being given off. If decarboxylation was occurring, theunexpectedly high pH values could be explained. While attempting to measurethe pH of solutions of HB in 75 percent DMSO and water, the rate of evolutionof gas was very noticeably increased. The pH of the solution increased more'Manuscript received October 3, 1968.THE OHIO JOURNAL OF SCIENCE 70(2): 97, March, 1970.98 WILLIAM E. LAQUE AND CONRAD E. RONNEBERG Vol. 70rapidly than when using 50 percent DMSO and water solution. A barium hy-droxide test indicated that the gas was carbon dioxide.All reports in the literature on the decarboxylation of trichloroacetic acid or itssalts have been at elevated temperatures. Verhoek (1934) studied this reactionin various water, alcohol, and aniline solutions at high temperatures. Clark(1955, 1956, 1957) reported on the decarboxylation of trichloroacetic acid at tem-peratures considerably above room temperature in solvents such as glycerol andDMSO. He also studied the decarboxylation of malonic and oxalic acids inDMSO and in glycerol, but this work was also at high temperatures (100° C to166° C). He showed, for example, that malonic acid decarboxylates seven timesas fast in DMSO as in glycerol.According to Verhoek (1934), and as verified by Clark (1955). the decarboxyla-tion of HB at elevated temperatures is monomolecular. A mechanistic interpreta-tion (Fieser and Fieser, 1961) is based on the fact that a chlorine atom linked toa carbon atom is electron-attracting. The three chlorine atoms of the anion causean inductive pull on the electron pair shared by the two carbon atoms to theextent that this pair is captured by the chlorinated carbon, as a lone pair of electronson the ionic oxygen moves in to be shared with the oxygenated carbon. The re-sult is the liberation of carbon dioxide and the formation of a transient carbanion,which combines with protons from various possible sources to form chloroform.The above discussion may be illustrated by the following equations:Cl Cl| ,/O /CH3 | /Or + /CH3Cl-C-C + O = S <F± Cl -C-C + HO = S Equation A| \OH \CH 3 | \ O \CH3Cl ClCl ClI / O I / OCl-C-C -> Cl-Cr + C Equation BI \O" | \OCl ClEquation B is the rate-determining reaction in this decarboxylation (thecarbanion may have other resonance forms). Finally, as shown below, the car-banion picks up a proton to form chloroform and complete the reaction.Cl ClC l - C r + H3O+ -> C l - C - H + H2O Equation CCl ClDMSO in acid solutions is highly protonated. It appeared therefore thatDMSO should likewise react with the intermediate ion, as follows,ClC l - C r + (CH3)2S = OH+ -> CHC13 + (CH3)2S = O Equation DCland thus accelerate the decarboxylation of HB even at room temperatures. Thispaper is concerned with the verification of this hypothesis.EXPERIMENT ABefore the decarboxylation of HB studies could proceed, certain informationpertaining to DMSO and its solutions not available in the literature was necessary.The density of various concentrations of DMSO and water solutions was necessary.Two density curves were made, one at 20.00°C and the other at 25.00°C. ThisNo. 2 TRICHLOROACETIC ACID IN WATER AND DMSO 99was accomplished by finding the mass and volume relationships of solutions rang-ing in concentrations from 0.0 to 100 percent DMSO by weight.The method of preparing the solutions was as follows. Known amounts ofDMSO and water weighed to the closest drop on a balance accurate to ±0.002 gwere mixed. Herewith is an example:To prepare 25 g of 87.50 percent by weight DMSO, use21.88±0.06 g (1 drop) DMSO3.12±0.05 g (1 drop) water, or25.00 ±0.07 g solution (±0.3%).Densities were determined by using a calibrated pycnometer. The followingexamples illustrate the procedure:Composition of solution 87.5 percent±0.3%Volume of pycnometer (20°C) 10.1088 ml±0.004%Mass of solution in pycnometer (20°C) 11.1750 g± 0.004%Density of solution (20°C) 1.1037 g/ml± 0.0007 g/mlThe temperature of the water bath was kept constant to±0.05°C. The precisionwith which the densities were determined was of the order of 0.10 percent. Thedensities of the various solutions of DMSO in water at 20°C and 25°C are shownin Table 1, and the graphs of these data are shown in Figure 1. The inflectionin the curve at about 75 percent can be accounted for by the very polar natures ofboth solvent and solute.TABLE 1Data for densities of DMSO solutionsPercentDMSO SolutionPrecision0.3%0.0012.5025.0037.5050.0062.5068.7575.0081.2587.5093.75100.00Density20.00°C±0.05°CPrecision±0.10%0.9982 g/ml1.01471.03311.05341.07291.08991.09541.10031.10241.10371.10271.1014Density25.00°C±0.05°CPrecision±0.10%0.9970 g/ml1.01341.03081.05021.07041.08591.09311.09661.09951.10031.09841.0962EXPERIMENT BThe apparatus finally devised to follow the rate of evolution of carbon dioxidefrom samples of HB in the presence of DMSO is shown in Figure 2. The reactionvessel, labelled No. 1 in Figure 2, was a 250 ml round-bottom flask. To this a 90°connecting tube with a stopcock was fitted; the connecting tube led to a large testtube, labelled No. 2, filled with distilled water. The purpose of the water was tocollect any DMSO vapor that might escape from the reaction vessel. It shouldbe noted that DMSO is extremely soluble in water. The vapor pressure of pureDMSO at 25°C is only 0.06 mm of mercury. The DMSO solution, as well asthe water, in tube No. 2 was saturated with carbon dioxide from a tank beforeany experimental run. From tube No. 2 the displaced carbon dioxide was di-rected to another large test tube, labelled No. 3 in Figure 2, which was filled withdrierite. Here the moisture was removed. The exit gas was then assumed to bedry carbon dioxide, equivalent in amount to that produced in the reaction vessel.100 WILLIAM E. LAQUE AND CONRAD E. RONNEBERG Vol. 70O)(AcQ1.12OO —1.1OOO —1.O8OO -1.O6OO -1.O4OO —1.O2OO —1.OOOQ99OO25 5O 75Weight Per Cent DMSOFIGURE 1. Density of DMSO solutions.1OOCircled dots—20 °CCrosses — 25 °CThese first three vessels were immersed in a constant temperature bath (25.00°C±0.05°C).From the drying tube, the carbon dioxide was passed to a Schwartz U tube(No. 5 in Figure 2) filled with ascarite, where it was collected and weighed. Be-tween the collecting tube and the drying tube was an open manometer (No. 4)connected to the system by means of a three-way stopcock. This was includedto check occasionally on the pressure within the system. On the other end of thecollecting tube were two smaller tubes. One (No. 6) was filled with ascarite toinsure against the absorption of carbon dioxide from the atmosphere. The other(No. 7) was filled with drierite to prevent the entrance of water vapor from theatmosphere.Trichloroacetic acid is extremely hygroscopic. It was dried before using byplacing in a small glass-stoppered weighing bottle with cap removed in a vacuumNo. 2 TRICHLOROACETIC ACID IN WATER AND DMSO 101desiccator over concentrated sulfuric acid. After being desiccated overnight, thebottle was removed, capped, and weighed. When ready for a run, the acid wasdropped into the reaction vessel No. 1 (fig. 2). The empty vessel, with stopperin place, was then weighed. This gave the weight of the added HB. The effect ofadding as much as 2.0 grams of HB in 100 ml of DMSO plus water solvent did notresult in a noticeable volume change.The sequence of operations for an experimental run on the rate of evolution ofcarbon dioxide was as follows. One hundred milliliters of DMSO solution ofNo. 1No.No. 2FIGURE 2.No. 3Experimental set up.No. 1—reaction vesselNo. 2—DMSO vapor collectorNo. 3—gas drierNo. 4—pressure check, open manometerNo. 5—gas collecting tubeNo. 6—outside carbon dioxide intercepterNo. 7—outside water vapor intercepterknown composition was transferred to the reaction vessel (No. 1 in Figure 2).Carbon dioxide (from a tank) was slowly bubbled through the system until satu-ration of the solution in vessels 1 and 2 was assured, and all air in the system wasdisplaced. The time necessary for this operation was about one and one halfhours. The weighed sample of HB was introduced at a known time. The pres-sure of the system was checked with the manometer and a weighed collecting tube(No. 5) put into place. Two collecting tubes were used in each run. A weighedtube was always in place. The tubes were changed at known time intervals andthe increases in weight ascertained. The change in weight was the weight equiv-alent to the weight of carbon dioxide produced in the reaction vessel.The data from a typical run using 80 percent DMSO solution in water is pre-sented in table 2. The calculations foi k, the specific-rate constant, indicated thatthe reaction was first order. The equation for the first-order reaction was usedin the following form:2.303 / a \k= log —t \a-x/The value of "x" was calculated from the known weight of carbon dioxide collected;"a" was the original molar concentration of HB. From the individual reactionrate constants, the average k was determined. The constant was also determined102 WILLIAM E. LAQUE AND CONRAD E. RONNEBERGTABLE 2Vol. 70Total CO2collectedto time tTimeHoursMols CO2/1collectedto time tMols HB/1decarboxylatedto time tlogReaction rateconstant k0.18320.30530.41040.44850.4710111720220.04170.06940.09300.10180.10650.04170.06940.09300.10180.10650.16730.33080.54780.66370.77020.05520.06930.07420.07640.0806k (average)—0.0711 hr"1k from graph—0.0753 hr"1from the slope of the curve obtained by plotting the log I — I against time t.a-xFrom the slope of the curve for 80 percent DMSO, for example, k was determinedas follows.2.303k = ——• x 1.15 = 0.0755 hr-135 hrs.The half time of the reaction is given by0.693Using the reaction-rate constant for the 86 percent concentration, the half timeis calculated to be 9.2 hrs. It should be noted that the values of k obtained fromaveraging the individual experimental reaction-rate constants and the value fork from the slope of the curve are in satisfactory agreement (±5.00 percent).SolutionPercentDMSO86%86%80%80%80%50%75%75%80%InitialConcentrationHB mols/10.13520.14600.13010.15560.13160.13460.13500.14590.1159TABLE 3Experimental resultsConditions25.0°25.0°25.0°25.0°25.0°25.0°25.0°0.06505 molarB~ ion25.0°0.06505 molarB~ ion35.0°Reaction-RateConstantk-hr-10.577\0.578/0.0728|0.0789 \0.0732J0.008900.2460.3960.738Shown onFig. 3 by:Xt'Xon Fig. 5 by:XReaction-RateConstantfrom Curvek-hr-10.5530.07550.00837ReactionHalf Time1.2 hrs1.2 hrs9.5 hrs8.9 hrs9.5 hrs77.5 hrs2.8 hrs1.7 hrs0.94 hrsFrom Fig. 3 and Fig. 5-2.303Reaction Rate Constant k = log [ -^— Jt & \ a -x /No. 2 TRICHLOROACETIC ACID IN WATER AND DMSO 10386% DMSO1.40 -80% DMSO50% DMSO x0.20I0 10 20 30 40 50 60 70 80 90 100 110 120Time-HoursFIGURE 3. Rates of Decarboxylation.The experimental results obtained in the entire set of runs are shown in Table 3and in Figures 3 and 5. These data indicate that the reaction of decarboxylationis first order and is markedly influenced by the concentration of DMSO. Thereaction-rate constants increase by a factor of 6-7 with a change in concentrationof DMSO from 50 to 86 percent. A comparison of the reaction-rate constants asinfluenced by the mol-fraction of DMSO is shown in Table 4 and in Figure 4.It is reasonable to believe that DMSO becomes highly protonated in the presenceof HB, and in this form reacts with the carbanion (Cl3C-)~ to accelerate thedecarboxylation (refer to Equations A, B, C and D).As hypothesized before, the rate of the reaction is also dependent upon theconcentration of B~ ions. Added trichloroacetate ions therefore should also ac-celerate the decomposition of trichloroacetic acid. This was verified by adding theanion in the form of sodium trichloracetate. These solutions were prepared by104 WILLIAM E. LAQUE AND CONRAD E. RONNEBERG Vol. 70TABLE 4Collected data for runs of all HB at 25.0°C and variousconcentrations of DMSOPercent DMSO Mole FractionDMSOk-hr" Ty250.80.86.0.19710.49700.60530.008900.07490.57777.9 hrs9.3 hrs1.2 hrsfirst neutralizing a known amount of HB with sodium hydroxide. Two runs weremade with 75 percent DMSO solution plus 0.06505 molar B~ ions and 0.1350molar HB, and one with 0.1459 HB. The comparative results are shown inTable 3 and Figure 5. The rates of decarboxylation shown in Curves 5 and 6are 6-10 times faster than when no trichloroacetate ions (Curve 4) were present.These curves (5 and 6) deviate somewhat from each other, but the rate constantsare of the same order of magnitude and considerably greater than for Curve 4where there were no B^ ions present as added trichloroacetate. The*data for0.600.50j< a40c0*>I/Icu 0.30ionwOa>* 0.200.10«•*••///1IJo ? > * - \ 1 10.800.20 0.40 0.60Mole Fraction DMSOFIGURE 4. Reaction constant k plotted against mole fraction of DMSO.No. 2 TRICHLOROACETIC ACID IN WATER AND DMSO 105Curve No. 5 were determined twenty-four hours after determining the data forCurve No. 6. The decrease in rate of decarboxylation was attributed to twofactors: namely, (1) the concentration of HB was somewhat less, and (2) some ofthe trichloroacetate ions had decomposed during the twenty-four-hour intervalbetween these runs.1.000.80* 0.60O)o0.400.204 6Time -Hours8 10FIGURE 5. Effect of Trichloroacetate ions.dots in triangles—75 percent DMSOdots in circles —75 percent DMSO,0.1350 molar HB0.0655 molar B~ ionscrosses —75 percent DMSO0.1450 molar HB0.0655 molar B~ ions.An indication of the magnitude of the energy of activation for the decarboxyla-tion of HB was obtained by determining the rate constant at 35.0°C and thenapplying the Arrhenius equation. One run was made in 80 percent DMSO and0.1159 molar HB (Table 4). The rate constant at 35.0°C was found to be 0.738hr^1. The rate constant for a comparable concentration of HB in 80 percentDMSO at 25.0°C was found to be 0.0777 hr"1. Using these values in the Arrheniusequation gave AE of activation of about 42.3 kcal/mole. More precise valuesfor the AE of activation must await further experimental work. The rate con-stant appears to increase by a factor of about 10 in going from 25.0°C to 35.0°C.106 WILLIAM E LAQUE AND CONRAD E. RONNERBERG Vol. 70SUMMARY1. The decarboxylation reaction of HB in the presence of DMSO in water solu-tions proceeds with significant rates at a temperature of 25.0°C, which isfar below any temperature reported in the literature.2. The reaction rate for decarboxylation of trichloroacetic acid in solutions ofdimethylsulfoxide in water at 25.0°C increases rapidly with increasing con-centration of DMSO.3. The reaction rate is determined by the speed with which the carbanion [CI3C-]-is broken down to form chloroform.4. The addition of trichloroacetate ion to the reaction significantly increases therate of the reaction.5. The rate of decarboxylation increased by a factor of about 10 with a 10-degreerise in temperature.6. It would appear that the energy of activation of the decarboxylation reactionis on the order of 42 kcal/mole.ACKNOWLEDGMENTThe suggestion for this research came from Dr. Conrad E. Ronneberg; sincereappreciation is extended to him for helpful counsel and guidance. Much of thefinancial support for this research came from the Research Corporation, NewYork, N. Y., and from the Denison Research Foundation, which is gratefullyacknowledged. Thanks must also go to the Crown Zellerbach Corporation,Camas, Washington, for supplying the dimethylsulfoxide (DMSO) necessary forthis research.REFERENCESClark, L. W. 1955. The Decomposition of Trichloroacetic Acid Alone and in Glycerol. JournalThe American Chemical Society 77: 3130-3131.. 1956. The Decomposition of Malonic Acid in Glycerol and in Dimethyl Sulfoxide.Journal of Physical Chemistry 60: 825-826.. 1957. The Kinetics of the Decomposition of Oxalic Acid in Non-Aqueous Solvents.Journal of Physical Chemistry 61: 699-700.Fieser, L. F., and M. Fieser. 1961. Advanced Organic Chemistry. (New York and London:Reinhold Publishing Co.).Inman, C. W. Barringer and C. E. Ronneberg. 1955. The Decomposition of TrichloroaceticAcid and Trifluoroacetic Acid in Solutions of Water and Glucose. Journal of ScientificLaboratories (Denison University, Granville, Ohio). 43: Art. 9, 175-182.Verhoek, F. H. 1934. The Kinetics of the Decomposition of the Trichloroacetates in VariousSolvents. Journal of The American Chemical Society 56: 571.

A Study of the Decarboxylation of Trichloroacetic Acid in Solutions of Water and Dimethylsulfoxide

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A Study of the Decarboxylation of Trichloroacetic Acid in Solutions of Water and Dimethylsulfoxide

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