ct of pH on Hydrolysis of Chlsroacetic Acid L. F. BERHENKE AND E. C. BRITTON The Doto Chemical Company, Midland, Mich.
A continuous titration method at constant pH has been used to measure the rate of hydrolysis of chloroacetic acid in 3 N aqueous solution at reflux temperature, 104' C. With this method the rate studies have been carried out at various pH values from 1 to 13. At pH 1 the titrating solution was silver nitrate, and at all others it was sodium hydroxide. The data thus obtained indicate: A t constant pH the measured reaction is first order; the rate is nearly constant a t 1.55*0.15% per minute in the pH range 3 to 10; the rate for the free acid (pH 1) is 0.024% per minute, increasing rapidly to 1.75 at pH 3; the rate falls slowly to 1.40% per minute at pH 6 and increases again to 2.00 at pH 11.5; above pH 11.5 the rate increases exponentially to 5.8% per minute at pH 13. Users of chloroacetic acid who need to neutralize the acid should always add the alkali to the acid, not the acid to the alkali. Furthermore, good stirring is essential.
Temperature control was achieved simply by operating under reflux. The concentration of the reacting solutions selected w'as such that their boiling points were about 104" C.; hence t h e temperature of the refluxing solution did not change more than 0.5' C. during a run. The error thus introduced is not large. I n the runs made with the free acid, it proved impractical t o maintain the pH by the addition of alkali, and so the procedure was modified and the chlorides precipitated continuously by the addition of silver nitrate solution. The reactions were carried out in a one-liter, multinecked, round-bottom flask fitted with a stirrer, reflux condenser, thermometer, reagent inlet from the solenoid valve, and two electrodes for the cell system. Heat was supplied by an electrical heater. The actual control of pH during the reactions wtis automatic. The antimony-saturated calomel cell system used for pH meas-. urement was connected to a Leeds & Sorthrup recording and controlling potentiometer. This, in turn, opened or closed a small glass solenoid valve to add sodium hydroxide as requirtd to maintain the set pH value. In the unbuffered range (pH 6 or above) this system gave control xithin * 0.2 pH unit; in the buffered range the control as even closer. For the run a t pH 1 (free acid) the antimony electrode was replaced with a silver electrode and silver nitrate added automatically t o precipitate the chloride formed in the reaction. The behavior of the antimony electrode was checked on a number of runs a t various pH values by withdrawing a sample from the reaction flask at the end of a run, cooling it rapidly t o room temperahre, and immediately measuring the pH with a glass electrode. The p1-I values agreed within 0.2 unit. For the runs a t pH 3 and above, 95 grams of commercial chloroacetic acid (freezing point 61" C.) were dissolved in 200 ml. of water and placed in the reaction flask. This WLIS warmed to 50-55' C., and then 50y0 sodium hydroxide was added to neutralize the acid. . Thetheat of neutralization was sufficient t o heat the reaction mixture to reflux. A slight excess of sodium hydroxide was usually added, since the normal hydrolysis rapidly brought the solution back to the desired pH. Khen this point was reached, the controller was turned on, timing of the run was started, and the volumes of sodium hydroxide required to maintain the p H were recorded. For the runs at pH 3, 4, and 5, only sufficient sodium hydroxide was added to give the deeired pH. For the run on the free acid (pH l),the neutralization with sodium hydroxide was omitted, the reaction heated directly to reflux, and the controller turned on. Timing was started after the chlorides, formed during t'he heating period, had been precipitated.
ANY studies have been made on the hydrolysis of chloroacetic acid and its sodium salt, starting with Buchanan (1, 6 ) in 1871 and continuing through the extensive series of
studies of Dawson ( 9 , 4 , 6 )and co-workers in 1943. Almost from the first there has been a suggestion that the pH of the reaction medium htEs an effect on the rate of halogen elimination. Buchanan (8) showed that the reaction was much faster Kith excess alkali in the reaction system. Kastle and Keiser (8) and Laseen and Eichloff (9) showed the effect of the reactant concentrations as well as the amount and concentration of excess alkali on the rate of halogen removal. Dawson and Pycock also indicate the effect of pH in their work.on hydrolysis of mixtures of chloroacetic acid and its sodium salt. They found a maiiimum velocity of chlorine removal in solutions containing approximately 0.7 M sodium chloroacetate and 0.3 M chloroacetic acid. The later work of Dawson, Pycock, and Smith indicates also that the hydroxyl ion plays an important role in the halogen removal, especially if its concentration is sufficiently high. With these numerous references in the literature to the effect of pH on the rate of halogen removal, it is surprising that no systematic study has been made of the relation of pH t o hydrolysiE rate. Likewise, no case hag been found where an attempt wae made t o neutralize the acid formed during the reaction. Csually, the effect of this acidification has been minimized by determining the rate from only the fist portion of the hydrolysis. Because of the importance of chloroacetic acid, and especially its sodium salt, in chemical synthesis, the present work was undertaken t o obtain more explicit data on the rate of hydrolysis as affected by changes in the acidity or alkalinity of the system. Furthermore, it was desirable to have this data a t approximately the concentrations which would be used commercially.
CALCULATIONS
The method of Guggenheim (Y), which eliminates the use of the uncertain infinity reading, was used t o calculate the specific velocity coefficients for a first-order reaction. Readings wcre therefore made at 5-minute intervals over a period of 2 hours. The constant int,erval between pairs of readings was usually 1 hour. I n a fev cases, noted in Table I, an infinity reading was made, and this value was taken L\S the second reading of the pair.
APPARATUS AND PROCEDURE
Since it was desired to carry out the velocity studies at constant pH, continual neutralization of the hydrochloric acid formed in the reaction was required. If this neutralization is done with sufficient care, the reaction can be followed by noting the volume of sodium hydroxide solution added a t various times.
544
INDUSTRIAL AND ENGINEERING CHEMISTRY
May, 1946
of the reaction mixture with the.standard alkali and by changes in total ionic concentrations. These are valid criticisms of the method, but the errbrs thus introduced are probably not large. EXPERIMENTAL.Those runs made in the buffered range (pH 3 to 6) were subject t o somewhat larger experimental error than the others. This was due t o the fact that the volume of sodium hydroxide which had to be added to cause response in the controller was several times larger than that required outside the buffcred range. Thus, while in the buffered range the actual pH control was better, the experimental error was larger. The value of pH 1 for the hydrolysis of the free acid is an apprordmation. Actual measurement at 25" C., using a glass electrode, gave a value of 1.3 a t the beginning of the run and 0.6 a t the end. CALCULATION. Dawson and co-workers showed that the removal of chlorine from chloroacetic acid is not a monomolecular reaction, and although it is influenced by a number of factors, it frequently behaves like one of the first order. Therefore, for the sake of uniformity, all the specific velocity coefficients reported here were calculated on the assumption that the reaction measured was first order. I n defense of this choice, it might be pointed out that, for most of the runs semilogarithmic plots of concentration against time gave straight lines, as would be expected for first-order reactions.
These differences over a one-hour period are recorded in Table I. These data were then plotted on semilogarithmic paper, and the best straight line was drawn through the points, giving all equal weight. I n a few cases there was,a marked deviation toward the end of the run, and these points were therefore not considered in drawing the best straight line. The first-order specific velocity constants were then calculated from the slope of the line by reading off the half time and using the relation,
k
= ( l / t ) In
l00/50 = 0.693/t
where t = half time Because of t h e slowness of the reaction at pH 1, the data were handled somewhat differently. Infinity titrations were caltulated from solution concentrations, and the logarithms of the concentrations in terms of milliliters of silver nitrate were plotted against time. This was done because the reaction waa followed for such a short time t h a t plotting on semiPH logarithmic paper was not sufficiently Figure 1. Effect of pH on Hydrolysis Rate of Chloroacetic Acid at accurate. The 104' C. and 3 N Concentration velocity coefficient was then calculated from the slope of the line as above, and the*half time was calculated. Table I presents the data for the various runs. Concentration is given in terms of differences in the pairs of readings for the various times. Values in the laat column for specific reaction constant k (the fraction hydrolyeed per minute) are plotted as a function of pH in Figure 1.
DISCUSSION O F RESULT6
The results here presented are generally in conformity with the data already in the literature. The increased speed of hydrolysis at high pH values is again shown. The observation of Dawson afid Pycock of a maximum rate of chlorine removal from a mipture of 70 mole % sodium chloroacetate and 30 mole % chloroacetic acid is confirmed by the maximum a t pH 3 in Figure 1. The amount'of sodium hydroxide required t o neutralize the chloroacetic acid to pH 3 indicates that about two thirds of the acid was neutralized a t this pH. The effect of the pH of the reaction medium on the rate of halogen removal as shown here may also account for some of the rather anomalous results reported in the literature. Drushel and Simpson did not find a constant first-order k for the hydrolysis of sodium chloroacetate. This is no doubt due to the fact that, as hydrolysis proceeds, the solution becomes more acid. Senter (IO) gave the relative rates of hydrolysis of acetic acid, the sodium salt, and the sodium salt with one equivalent sodium hydroxide as 1:3.6:36. These ratios are not confirmed in this work, and the difference may be in the different concentrations used. He found little change in k with concentration for the free acid, but a considerable increase when the concentration of the sodium salt was increased. The relative rates found here are:
'
POSSIBLE ERRORS
THEMETHOD. The method used here may be open t o criticism, since Drushel and Simpson (6) found some discrepancy in the silver and alkali titrations of reaction mixtures such as those used here. Furthermore, some error is introduced by the dilution TABLE I. EFFECTOF pH pH
13 12.5 11.5 10 8 7
0 min. 79.2 190 52.5 65.0 110.5 185.0 81.8 111.0
106.0 98.0 171.5 89.0 5 69.8 4 47.5 3 60.0 0 min. 279 (') 484 a Infinity reading 6
5 min. 60.2 161 48.5 60.0 103.3 175.0
10 min. 45.2 136 44.0 ' 55.0 96.8 164.5 70.0 98.3 93.0 86.0 151.5 76.0 60.5 39.5 49.8 40 min. 276
ON
'
fRelative H rate
1 0.00024 1
HYDROLYSIS RATEOF CHLOROACETIC ACID AT 104'
15 min. 33.2 114 39.8 51.0 89.8 154.0 65.5 91.3 l04:5 86.5 100.0 80.5 91.5 141.5 161.5 73.0 82.0 57.0 64.5 36.5 43.5 45.7 54.8 30 min. 90 min. ... 273 480 ... 472 used a6 second reading of pair.
545
Concentration 20 min. 25 min. 30 min. 23.2 l9,2 11.2 95 80 65 36.3 33.0 29.5 46.5 42.8 39.0 83.8 77.8 71.4 144.0 133.0 125.0 60.5 56.0 51.5 85.8 79.8 74.2 81.0 76.0 70.3 75.0 70.3 65.0 131.5 122.5 114.0 62.7 58.0 69.0 48.5 44.5 53.0 30.5 27.8 33.0 38.5 42.0 35.7 150 min. 180 min. 210 min. .. . 267 466 .. . 459
...
7 0.0146 62
c. AND 3 N
35 min.
40 min. 45 min. 50 min. 4.2 1.9 1.2 31 23 51 40 24.5 22.0 19.8 26.7 33.0 30.5 28.5 36.0 66.1 51.8 61.0 66.0 101.0 94.0 116.5 108.5 41.0 38.2 47.8 44.0 60.5 50.2 68.0 64.0 56.0 52.0 65.3 60.0 52.5 49.2 61.0 56.0 91.5 86.0 98.5 105.5 50.0 46.0 53.0 55.0 35.0 31.5 37.5 41.0 19.5 23.2 25.5 21.0 28.8 26.6 30.7 32.9 240 min. 270 min. 300 min. 330min. 263.5 260 452 445
...
...
. ..
...
...
13 0.058 247
CONCENTRATION
.
55 min. 60 min.
..
.. 14
...
ii8
18 16.2 18.0 25.5 23.5 44.7 48.0 82.0 87.5 31.4 35.0 51.7 48.0 49.2 44.5 46.3 78.5 73:5 41.0 40.5 29.0 26.0 16.5 18.0 23.5 25.0 360min. 390min. 257
Half Time 11.8
19.1 35.0 41.5 45 5 50.0 45.0 49.5 49.0 50.0 48.5 49.5 43.0 39.0 41.0
100 k 5.8 3.65a 1.98 1.67 1.52 1.3ga 1.54 1.40 1.41 1.39 1.43'" 1.40 1.61 1.77 1.69
3010 0.0235 2890 0.024a
INDUSTRIAL AND ENGINEERING CHEMISTRY
546
These data also indicatc why certain practices in the handling of chloroacetic acid have been beneficial. For example, in many preparations in the literature, the use of sodium carbonate has been specified for the neutralization of chloroacetic acid. This has served t o prevent excessively. high alkalinit,ies during- thr neutralizat,ion j n d thus reduced the amount of hydrolysis. The same result' can be achieved by careful neutralization with sodium hydroxide with good stirring and cooling. It further indicates the inadvisability of adding the acid to the jlkali in neutralization. As already mentioned, plots of log conccntrution against time indicated that most of the runs were of the first order. The esception was the run a t pH 12.5; this plot sh0n-s a deviation from a straight line, in the dircction of a more rapid reaction. Thi:
VoI 38, No. 5
deviation started lvhen about 65% had reacted and may have been due to a slight increase in pH. LITERATURE CITED (1)
Buchanan, Ber., 4 , 340 (18i1).
[hid., 4 1 sG3 (1871). (3) D a w o n and Pycock, J . Chtnz. Soc., 1934, i 7 8 , (4) I b i d , , 1936, 153, ( 5 ) D ~ \Pycock, ~ ~ alld~ Smith, ~ , hid., 1943, 517. (0) Drushel and Simpsoii, J . A m . C'hem. Soc., 39, 2456 (L017). ( 7 ) Guggenheim, Phil. .Vag.. [ i ]2 , 538 (1920). ( 8 ) Kastle and Keiser, A m . C'hem. J . , 15, 471 (1393). !9) Lassen and Eirhloff, 342, (1905!. (10) senter, J . them. yoc.,91, 460 (190:). . . (2j
pREsErTzD on the program of the Division of Physical and Inorganic Chemistry of t h e 1945 I\ieeting-in-Print, AMERICAY CHF,XICAI, SOCIETY.
Cyclization of Pseudoionone by Acidic Reagents E. EARL ROYALS
'
Georgia School of Technology, Atlanta, G u .
A study has been made of the cyclization of pseudoionone by various acidic reagents to yield mixtures of a- and j3ionones. Particular attention has been given to the iufluence of the cyclizing agent on the composition of the resulting ionone mixture. The mechanism of the reaction is discussed, and suitable preparative procedures for CY- and 6-ionone are recommended.
SECDOIOKO?;E (I) under the influence of acidic reagents is converted into a mixture of a-iononc (11)and Pionone (111), the artificial violet odor of commerce (3, 10, 11). Because of the great technical irnportancc of the ionones, a number of methods have been developed for the cyclization of pseudoionone. These are dcsrribed mainly in the patent literature (4, 6, 7 ) .
v
p,
CI-I, I
--CEI=< II-t.=O
1
I1
I11
Although the relative proportions of CL- and 8-iononcs obtained on cyclization of pseudoionone are dependent on the cyclizing :tgc:nt used, little attention has been given to this phase of the problem beyond the general statements that concentrated sulfuric acid gives mainly 3-ionone (Q), n-hereas phosphoric, formic (Q), acetic, and benzoic acids (8) produce mainly the alpha isomer. Hibbert and Cannon develdped laboratory procedures for the preparation of an ionone said to be essentially the alpha isomcr (4). Ito studied in some detail the cyclization of pscudoiononc by sulfuric and phosphoric acids, and showed that the proportion of isomeric ionones obtained is dependent on the concentration of the cyclizing acid ( 5 ) . The present invest'igation was undertaken to ex%end existing knowledge as to the effect of the cyclizing agent on the relative proportions of a-and p-ionones obtaincd
on cyclizing pseudoionone, to throw light on the meciianism of the cyclization reaction, and t o dcvelop practical method. for tho production of CY- and 3-iononcs. CYCLIZATION PROCEDURES
Citral was obi ained from Fritzsche Brothers, Inc., and \\-as uwd v-ithout further purification. The acetone and vnrious ncidin reagents were of C.P. gradc. Phosphoric acid of strength grcatvr than 85yGwas preparcd by adding the cnlciilatcd qi:antity o f phosphorus pentoxide t o the ordiiiary sirupy wid, Psoudoionoiic.. n-as prepared according t o the following procerliire ( 1 1 ) : A niisture of 200 grams (1.31 moles) of citral, 200 grams (3.44 moles) of acetone, and 200 cc. of a saturated solution of barium liydrosidn ~ m vigorously s stirred for 3-4 days at room temperature (about 28" C , ) . The organic materid was taken up in ether and waslied successively 17-ithm ~ t e r 5yc , tartaric acid solation, and again \\-it11 n-zter. The ethcr solution \vas dried over sodium sulfate, thc ether n-\.asremoved by distillation, and the residue w a ~distilled under reduced pressure. The fraction boiling arouiid 150-165 C'. (15 nini.) was collected. This material wai fvcc:d of citral a i d condensation products of acetone by steam distilling until about, one liter of distillate 11-a~collected. The residue was c!ricd over nnhj-drous sodium sulfatc and fractionated unticr reduced pressure through an 8-inch column packed n-ith glacis hcli terrially heated. The portion boiling at 144-149' C. (10-11 inin.) and showing ti2: of 1.5230-1.5280 (succesbive fractions) K:IS taken for use as pseudoionone. The literature reports f o r preudoionone prepared by this method (ff), 143-143" C. (12 nirn.) and 1.5275, respectively. The cyclizations of pseudoionone (Table 1) were carried out, in the following general manner: The cyclizing agent as placed in a 200-cc. three-necked flask fitted n-ith a mechanical stirrer, a dropping funnel, and a thermometer reaching into the reaction mixture. The flask wi,s cooled in an ice bath, and the cyclizing agent was vigorously stirred xhile 20 grams of pseudoionone were added dropvise during 15-30 minutes; the rate of addition was adjusted so that the temperature of the reaction mixture did n o t csceed 10-15" C. Khen the pseudoionone had all been added, the cooling bath was removcd, and stirring was continued for a period considered sufficient t o permit complete reaction. T h e reaction mixture was poured into 500 cc. of ice watcr and cs-
