ACID CATALYSIS I N LACTONE FORMATION* BY HUGH STOTT TAYLOR AND HAROLD WILBERFORCE CLOSE**
The catalytic activity of acids in a variety of reactions, the velocities of which are sufficiently slow to be measurable, has been the object of continued study since the formulation of the dissociation theory. The results of such study have been used to check conclusions regarding the nature of solutions, obtained by other methods of investigation. Arrhenius showed the parallelism between conductivity and catalytic activity in weak acids. He illustrated the principle of isohydry by reaction velocity studies with acetic acid solutions containing sodium acetate. He showed the abnormality of strong electrolytes by showing that, with strong acids, the proportionality between hydrogen ion concentration and reaction velocity effect was not exact, and by the neutral salt effect with such strong acids. These abnormal catalytic effects with strong acids have at various times been utilised to support (a) the Arrhenius theory of ‘active’ and ‘inactive’ molecules and its modern successor the ‘radiation’ theory of catalytic action (b) the solvate theory of electrolytes, in which the abnormalities were attributed either to the solvation of the ions or of the undissociated salt or acid (c) the dual theory of catalytic activity in which, to both hydrogen ion and undissociated acid molecule, a catalytic activity was ascribed (d) the theory that only the non-hydrated hydrogen ion had any catalytic activity (c) the theory that only alcoholated hydrogen ions had catalytic activity in esterification and (f) the activity theory, which attempts to correlate reaction velocity with the thermodynamic concentrations or activities of the reacting species without specifying the actual molecular nature of those species’. There were two main objects in mind in undertaking the present experimental investigations. We desired to ascertain whether the catalytic conversion of hydroxy-acids to lactones exhibits the same characteristics which are shown by various hydrolytic reactions catalysed by acids. We also desired, by study of this reaction in detail, especially in various solvents, t o learn, if posS:ble, what is the active agent in acid catalysis. Henry2 investigated this reaction under the influence of acids, and came to the conclusion that the reaction velocity is proportional to the hydrogen ion concentration. An examination of his results indicates that the proportionality is not exact, but that, as the concentration of catalysing acid is raised, the reaction velocity increases faster than the hydrogen ion concentration. I t was decided to undertake a systematic study of the reaction by vary*Contribution from the Laboratory of Physical Chemistry, Princeton University. **Abstracted from the thesis of Harold Wilberforce Close presented in partial fulfillment of the requirements for the Ph. D. degree, Princeton University, April 1922. Presented at the Pittsburgh meeting of the A. C. s., September 1922. A useful summary of these various points of view is to be found in “A Treatise of Physical Chemistry,” Vol. 11, Chapter XII, p. 779; Chapter XIV, p. 914 (1924). Z. physik. Chem., 10, 96 (1892).
1086
HUGH STOTT TAYLOR AND HAROLD WILBERFORCE CLOSE
ing in turn each of the following factors:-(a) the nature and concentration of the catalysing acid (b) the nature and concentration of the neutral salt (c) the temperature (d) the solvent.
Experimental Preparation of Materials, Valerolactone. CH, - CH - CH, - CH2 - CO. Hydr2xyvaleric acid is L--O-__-_I less stable than its lactone. Consequently, a quantity of the lactone was prepared and was subsequently converted to the hydroxyacid in small quantities as required for each experiment. The lactone was prepared as follows: One hundred grams of cane sugar were heated on a water bath for twenty hours in contact with half a liter of I :4hydrochloric acid. This charry mass was then filtered and after the addition of a solution of 35 g. of caustic soda to the filtrate it was evaporated to I O O cc. and the sodium chloride which had separated out was filtered. The crude levulinic acid so obtained was next extracted with ether and fractionally distilled, the quantity passing over between 140' and 170' under a pressure of 15-30 mm. being collected, Thirty cc. portions of this product were then reduced in alkaline solution by use of a I O percent sodium amalgam. The reduction was carried out in a vessel surrounded by ice and the mixture was constantly stirred. TWOhundred grams of amalgam were added in lots of 2 0 g. every few hours, small quantities of z :I hydrochloric acid being also added from time to time as required to prevent the mixture from becoming thick and viscous. The mercury was removed when reduction was completed and 7 5 cc. of concentrated hydrochloric acid were added. After boiling under a reflux for ten minutes the mixture was cooled and extracted with ether. The ether extract was dried over freshly ignited potassium carbonate, the ether distilled off and the residue repeatedly distilled until a constant boiling point testified to the purity of the product. The valerolactone thus obtained was a colorless sweet-smelling liquid, boiling at 207'. I.
a.
b. Hydroxy-valeric A c i d . CH, - CH - CH2 - CH2- CO. The hydroxy-
l
I
OH OH acid was prepared by converting the lactone into a soluble salt of the acid, from which the metal was subsequently removed by precipitation, and the free acid liberated. The first method used was to prepare the barium salt and the precipitate the barium as the sulphate. Since this involved the use of sulphuric acid in removing the barium, an alternative method was tried in which the lead salt was prepared by treating the lactone with lead hydroxide and then precipitating the lead with hydrogen sulphide. This was found to give more reproducible results because the hydrogen sulphide could be run in until precipitation was complete, and then the excess entirely removed by shaking under reduced pressure. The method was later simplified by substituting litharge for the lead hydroxide. Litharge can be obtained of more
ACID CATALYSIS I N LACTONE FORMATION
1087
uniform purity and is more stable than lead hydroxide. According to the latter method which was used in all cases except where especially noted to the contrary, thIee cc. of lactone and seven grams of litharge were warmed with zoo cc. of water and then shaken until the color had piactically disappeared, Hydrogen sulphide was then bubbled through until precipitation was complete. After removal of hydrogen sulphidp as already described, the solution of hydroxy-valeric acid was used without delay for velocity determinations. ,'\ - C H ~ O H e. o-Hydroxy-methyl-benzoic Acid. COOH. This acid was pre-
I 1
/\ - CH2
\/-
1 1
h 0 by dissolving the latter in a solution con\/-CO taining a small excess of caustic soda. The solution of the sodium salt of the hydroxy-acid was treated with hydrochloric acid until precipitation was complete. The white precipitate was filtered and washed thoroughly until silver nitrate showed that the chlorides had all been removed. The residue was then sucked dry, transferred to a porus plate, and kept in a vacuum desiccator until used. The dry acid is stable. Titration with alkali showed the acid to be pure. d. Ether. For our purposes it was necessary to use absolute ether and to take all precautions against absorption of moisture. The question as to how well we succeeded in excluding moisture will be referred to later, A good grade of ether was used; was washed three times with dilute caustic soda, and finally eight to twelve times with water. It was then allowed to stand over calcium chloride and subsequently over sodium wire until shortly before it was used. It was then distilled. This ether was kept in a filtei-flask with a tightly-fitting, paraffined stopper, and the side tube was closed with a tube of phosphorus pentoxide. e. Hydrogen Chloride in Ether. Into the ether prepared as described a stream of hydrogen chloride gas was led, the gas having previously been dried by passing through two calcium chloride towers followed by a long tube of phosphorus pentoxide. The acid solution was subsequently diluted with absolute ether until it was normal. Precautions against absorption of moisture were taken as in the case of the ether.
pared from phthalid
Measurement of Reaction Velocity. In the case of the hydroxy-valeric acid, IOO cc. of the solution, after removal of the hydrogen sulphide as already described, was placed in a 2 0 0 cc. flask,the catalyst added, and the flask filled to the mark. The mixture was immediately transferred to an Erlenmeyer flask and hung in a thermostat, the accuracy of which was &to. oz°C. At intervals, 20 cc. portions were pipetted off and titrated in the usual way with baryta of convenient strength. Usually seven readings were taken 2.
I088
HUGH STOTT TAYLOR AND HAROLD WILBERFORCE CLOSE
in addition to the initial and final determinations. The time intervals varied, of course, with the speed of the particular reaction. The determinations were always made at least in duplicate and often four determinations were made in cases where the experimental error was unusually large. When the formula for a simple unimolecular reaction was used in determining the velocity constant of the conversion of hydroxyvaleric acid, the value of the constant was found to fall rather consistently. This was due to the fact, as Henry pointed out, that the conversion to valerolactone is not complete. Consequently the formula for opposed reactions was used and the results obtained were more satisfactory. The following example will serve to illustrate the results obtained by use of the different formulas. The value I ,0 7 5 in the second formula was obtained from the observation that equilibrium is established in the reaction when 92. j per cent of the oxyvaleric acid is changed to lactone.
No. 17 B. Jan. 14.
TABLE I Catalyst, . O I N. HC1+N.CaC12 Baryta ,069 N. Temp. 2 j°C.
Minutes
Time in Minutes
Titer in cc. Baryta
102.4 136.0 182.4 216.2 259.7 289.7 334.2
17.66 15.23 14.55 13.68 13.10 12.42 11.97 11.38 4.61
49 32
5
Seconds
40 5 40
52
5
25
50
9 39 23
20
20
50 End-point Average
X
XI = ‘log? t a-x
2.43 3.11 3.98 4.56 5.24 5.69 6.28
a
= 13.05~~.
a Kz= 1-logt ~e1.075~
8 . 7 3 X IO-^ 8.68 8.66 8.63 8.58 8.58 8.53
9 . 4 5 X IO-^ 9.43 9.45 9.45 9.44 9.49 9.46
8.62
9.45
It is obvious from the above figures that the values for K1, show a tendency to fall, which is not the case with Kz,in which the effect of the reverse reaction is taken into account. Experimental Results Since an example has already been cited which illustrates the method of measurement of the reaction velocity, only the final results need be reported in the tables which follow. a. Variation in Catalysing Acid. Two different catalysing acids of widely differing strengths were used, and in different concentrations. The temperature was also varied. The figures for hydrogen ion concentration are taken either from Kohlrausch and Holborn, from Bray and Hunt1, or from Dawson and Reiman2, as indicated in the table. J. Am. Chem. Soc., 33, 781 (1911). J. Chem. SOC.,107, 1426 (1915).
1089
ACID CATALYSIS IN LACTONE FORMATION
Series and Method
Ser. 1
1
ca+ymg Acid
)-I
TABLE I1
10.01HCll 25’
0.025 Lead Hydro- 0 . 0 5 0 . IO xide Method
K
Temp.
’
”
’’ )’
”
)’
I
I
x
Cond. 104
4
3
2
Mean
from
6.821 6.831 6 . 8 : 6 . 8 4
6 . 8 3 0 . 0 0 9 7 I B f H 703 o.o0989K+H691 17.18 1 7 . 1 1 17. I L :I6.90> : 7 . 1 4 0 . 0 2 4 K + H 7 1 4 3 4 . 5 8 3 4 . 0 4 3 4 . 3 f 3 4 . 2 1 i 4 . 3 0 0 . 0 4 7 B-kH7.30 6 8 . 1 6 7 . 4 6 7 . 1 6 8 . 1 i7.7 0.092 B + H 7 3 6
-
~
-
0,00971 B +H 5 2 . 5 o.oo989K+H51.6 2.580.047 B + H 5 4 . 9 5 . 2 4_ 0 . 0 9 2~ B + _ - H 5-7 . o 5.830.0082 D + R 7 1 o
-
0.5I
” 2 . 5 6 2.60 Lithargeo.05 ” ’) 5 . 2 8 5 . 2 1 Method 0 . 1 0 ” ---Ser. S 0 . o j CH2 ClCO OH j .84
25’
Lithargeo.10 Method 0 . 2 0
”
”
)’
’’
5.83 8.76 8.72 13.43 13.45
8.740.0122 D + R 7 1 6 r3.440.0183 D + R 7 3 4
TABLE I11 Series
I
Method
Conc. of Temp. Neutral Salt
KX104 I
I
3 1 1 2 ---
1 3 . 2 733.82 32.81 . 2jN , 5 , 8 83 6 . 2 8 36 .o; . SON 17.7337.7337.9s 1.N 12.48 43.18 43. I L
0
.os N.HC1 Baryta Method Ser. 2 .I
N
__-25OC.
3
.o o j N . OIN . o2N
I
CHzClCOOH
Litharge Method Ser. S
8.67 7.13 6.14 4.57
--25°C.
3
. oo j N .I N CHzClCOOH Litharge Method Ser. 4 3oOC.
.I N CHzClCOOH Litharge Method
8.54 7.04 5.96 4.63
.OIN .02N
8.76 7.34 6.15 4.90
8.72 7.40 6.33 4.80
--3
4.12 . ooj N 1.67 . OIN 9.87 . o2N 7.40
14.21 11.75 9.97 7.46
4
- C, Aver.
-
Cond. Data from
~
33.230.047 K+H 3 6 . 0 8 0 . 0 4 4 K+H 37.870.04265K+H 42.930.0392 K+H
--8.610.0122 D+R 7.090.01035D+R 6.050.00884D+R 4.600.00667D+R
8 . 7 4 0 . 0 1 2 2 D+R 7.37o.o1035D+R 6.240.00884DSR 4.850.00667DfR -____-
c4.17 [I.
71
0.0122 0.0102 j
D+R D+R
9.920.00884D+R 7.430.00667DfR
-
1090
HUGH STOTT TAYLOR AND HAROLD WILBERFORCE CLOSE
It will be observed that the value of K/CH rises steadily with increasing concentrations of acid. In other words, the hydrogen ion concentration as determined by conductance methods does not increase with increased acid concentration as fast as the reaction velocity. The conversion of hydroxyvaleric acid to lactone is thus seen to be comparable in this respect to the other examples of acid catalysis, in which, as is well known, the same disproportionality is found. .Variation in Neutral Salt. ( I ) Change in Concentration of Salt. Table I11 shows the results obtained when different concentrations of neutral salts were added to reaction mixtures in which the concentrations of the catalysing acid remained constant. Results obtained by different methods are not strictly comparable as it was found that those obtained by the baryta method were consistently somewhat higher than when the lead methods were used. Considering first the influence of varying amounts of potassium chloride upon the reaction velocity in the presence of .05N. hydrochloric acid, it is obvious that the reaction velocity increases markedly, whereas the hydrogen ion concentration, as determined by applying the principle of isohydric solutions and the law of mass action decreases. The conversion of hydroxyvaleric acid to lactone resembles the other hydrolytic processes in this respect also. Turning now to the case of monochloracetic acid, it is obvious that addition of sodium acetate, even in very small concentration, produces a marked fall in the reaction velocity. The ratio K/CHis not constant, there being a tendency for the value of the ratio to fall slightly. In Series 2 , a drop in the value of K/CH as the salt concentration increases is followed by a rise. The series was repeat
