Kinetics of the Acid-Catalvzed Cyclization of Citral and Citronellal J
CHARLES C. PRICE' AND MARION LUND DICKMAN Noyes Chemical Laboratory, University of Illinois, Urbana, Ill. Citral and citronellal undergo a series of transformations in dilute aqueous acid solution, the first of which has been shown to be a first-order reaction with respect to the aldehyde concentration and directly dependent on the acid concentration. Analysis has demonstrated that this initial reaction involves the simultaneous disappearanceof the aldehyde group and one double bond. On the basis of these observations, a mechanism has been outlined which may be considered as a cyclization involving an internal Prins reaction.
double bond per citral unit. The dimer was converted to terephthalic acid by permanganate oxidation but because of difficulty in isolating it in a pure condition, it was not further characterized. In early experiments, it was observed that, for solutions a t 25", the initial analysis for unsaturation always indicated 10 to 20% less than the theoretical value for the amount of citral added. In later experiments at 25', it was observed that the addition of phosphoric acid caused an initial increase in citral concentration. Apparently a portion of the citral added must have remained as a film on the surface or absorbed on the glass walls. Addition of acid apparently considerably increased its solubility, leading to dissolution of this material. Because of this phenomenon, it was necessary to extrapolate the data to obtain the actual initial concentration, some 10 to 20% greater than the measured value. Typical data are illustrated in Figure 3. Using the extrapolated value for initial concentration, and based on the disappearance of half the original unsaturation, the rate of the decomposition was found to be first-order with respect to the citral concentration. Typical experiments are summarized in Figure 4. Data similarly obtained for experiments with several different acids, summarized in Table I, indicate that the first-order rate constants observed for different acid concentrations are directly proportional to the hydrogen-ion concentration, &s calculated from the concentration and known value for the dissociation constant for each acid.
I
T H4S long been known that citral and citronellal, like many other terpenes and terpene derivatives, are susceptible to a variety of transformations by treatment under acidic conditions. It was the purpose of the present investigation to establish the nature and the mechanism of the reactions involved in dilute aqueous solution. Commercial citral, presumably principally citral a, was vacuum-distilled before use. Citronellal was purified through the bisulfite addition compound and distillation, boiling point 91-93" C. (16 mm.); n'no 1.4463 [(IS), boiling point 89-91' C. (14 mm.); nD 1.44611. RATE OF DECOMPOSITION OF CITRAL
A. BY UNSATURATION. The concentration of the unsaturation in dilute aqueous citral solutions was determined by reaction with bromine. Tuenty-five milliliters of standardized bromine water were added to 50-ml. aliquots of the citral solution (ca. 0.2 gram per liter) in a glass-stoppered flask. After 2 minutes, an excess of 10% potassium iodide was added and the iodine liberated with 0.025 N thiosulfate. Experiment demonstrated that the addition of bromine to citral under these conditions was complete within one minute and that no further reaction occurred on longer standing. Preliminary measurements of the rate of disappearance of citral in water (ca. 200 mg. per liter) demonstrated that 85 to 90% of the initial unsaturation remained after 17 days. When 5 ml. per liter of 857' phosphoric acid were added to similar solutions, a brown oil precipitated after several days and, at the end of 17 days, the water contained no measurable unsaturation. Preliminary measurements of the effect of temperature on the acid-catalyzed reaction are summarized in Figure 1. Further data a t 45" are summarized in Figure 2. The experiments indicate that the initial reaction measured proceeded by loss of 50%of the initial unsaturation-Le., one of the two original double bonds was destroyed. Further subsequent decrease in unsaturation in the solution was only observed concurrent with precipitation of an insoluble oil. In view of the inhomogeneity of the system during this secondary process, no attempt has been made to interpret this secondary phase of the reaction kinetically. Distillation of this oil led to the isolation of small amounts of (a) cymene, ( b ) a dimeric condensation product and, principally, (c) a brown, thermoplastic resinous residue. Both the dimer ( b ) and the polymeric residue ( c ) were found by titration with bromine in carbon tetrachloride to contain approximately one 1
d [citral) __ __= dt
k' [citral]
k' = k[H+]
The values for k for experiments using lactic acid were in agreement with the other experiments only when the value of 3.1 x 10-4,estimated kinetically (7) for the dissociation constant of lactic acid, was used, rather than the values of 1.38 x 10-4 to 1.55 X 10 - 4 reported by Ostwald (IO) and others (4,9) by the standard methods. B. BY WZ-PHENYLENEDIAMINE. The procedure used was a modification (16) of the Hiltner (8) method. The reagent used was m-phenylenediamine oxalate, prepared as follows: one gram of m-phenylenediamine hydrochloride was dissolved in 45 ml. of 85% alcohol, and 1gram of crystallized oxalic acid in a similar quantity of alcohol of the same strength. The solutions were mixed, and the mixture was diluted to 100 ml. with 85% alcohol, decolorized with 2 to 3 grams of fuller's earth, and filtered t,hrough a double filter. The resulting reagent had a greenishyellow color which darkened on standing, but gave satisfactory results, provided fresh reagent was used. All measurements were made with the Evelyn photoelectric colorimeter, using a 420 mp filter with distilled water defined as giving 100% transmittance. A calibration of the method was obtained by use of a known alcoholic citral solution, aliquots of which were treated with 10 ml. of the m-phenylenediamine reagent and 10 ml. of water. After dilution to 50 ml., the transmittance was measured. The
Present address, University of Notre Dame, Notre Dame, Ind.
257
I N D U S T R I A L A N D E N G I N E E R I N G CHE'MISTRY
258
Vol. 40, No. 2
data from two t:xperiments on the decomposition of citral as measured by this method are summarized in Figure 5 and Table I. The rate is in agrcement, with that mcasurcd by the unsaturation method. In order to determine whether this e conclusively demonstrated the simultaiicous disappearance of a double bond (as measured by unsaturation) and the aldehyde group (as measured by m-phenylenediamine), a few experiments were carried out which demonstrated that the color developed by citral is evidently characteristic of ol,$-unsaturated aldehydes, since citronellal and I I I I I I I I aldol gave essentially no color (97 I 2 3 4 5 6 7 8 9 and 927, transmittance, respectively), TIME IN HOURS whereas crotonaldehyde and citral gave Figure 1. Decomposition of Citral in 0.5% Phosphoric licid approximately equivalent color intenc) A t 25' sity (60 and 55% transmittance, respecA t so tively). It is thus evident that the Z decomposition measured by m-phenylene 0 100 I I diamine might have been either the aldehyde group or the a,pdouble bond. C. BY FUCHSISE~ULFITE. Since this reagent is reported to be fiuitable for the estimation of aldehyde groups in very dilute solution ( I ) , its application to the measurement of the decomposiZ t,ion of citral was investigated. The reagent was prepared bjl i 60 t,reat,ing0.5 gram of fuchsine in 1 liter of water n.ifh 16 grams of 4 sulfur dioxide. A 20-ml. portion of this reagent was added to !aliquot portions of the citral solutions and made up t o 50 nil. L 40 2 4 6 8 IO with aldehyde-free ethanol. After 15 minutes at 15", the color ,\" developed v a s compared in a visual Duboscq-type colorimeter TIME IN HOURS t o a st,andard solut,ion of the same concentration similarly preFigure 2. Decomposition of Citral in pared. In view of the time required to develop the color, this Phosphoric Acid Solution at 45' C. analyt,ical procedure could not be adapted t,o precise rate studies, but it was shovn that, in 0.5% phosphoric acid at 25', the aldehyde group disappeared (9r0 in 1.25 hours, 25% in 2.4 hours,
T.4BLE
I.
R.4TE O F AiCID-CAT.4LYZED DGCOMPOSITIOS O F CITR.41, IS A kSOLUTION ~ ~ ~
Acid
[€I-], 3- X 102 [Citrallo, -11 X 101
k
-1. By unsaturation, 25' C. H&'OI (@.1?") HaP04 ( 0 . 5 % )
@.Y+ 2.40
1.40 1,38
0.54 0.91 1.38 0.45 0.45 0.47 0.65 0.65 0.26 0.41 0.37 0.59 0.64
1.14 1.36 1.39 1.32 1.38 1.16 1.27 1,41 1.43 1.34 1.06 1.47 1.77
9.0 9 ,9
At 450 c. HsP04 HiPOil Hap04 Citric Citric Citric Citric Citric Lactic Lactic Lactic Lactic Lactic
B.
(0.05%) (0.17~) (0.2%;
(0.6%) (0,570) (0.5%) (1.0%) (1. O % ) (0.2%) (0.5%) (0.5%) (1.0%) (1.0%)
HaPOa ( 1 . 0 % )
D.
I
I
I
1
3
5
7
9
II
TIME IN HOURS Figure 3. Rate of Disappearance of Unsaturation of Citral at 25' C. by Bromine Titration 0.1% Hap04 8 0.5 yo H 3PO4
36 35 31 28 29 33 30 36 28
By m-phenylenediamine, 25' C
?&PO& ( 0 . 5 % )
I
32 29 28 31
2.2; 3.39
1.33 1.11
7 1 8.0
1.67 1.54 1.64
7.6 7.3 7.5
By hydroxylamine sulfate, 2.5" C.
HC1 HC1 HC1
8
6 5.08
A t 450 C
HC1
HCl HCI
0.184 0.19 0.26
58
4s
,0:)
46 07 37
~
~
INDUSTRIAL AND ENGINEERING CHEMISTRY
February 1948
259
1.2
1.0
;7 h
_I
Q
K k-
50.8 0 0
-I
0.6
2 I
5
3
7
I1
9
6
4
TIME IN HOURS
TIME IN HOURS
Figure 4. Rate of Disappearance of Citral at 25' C. by Bromine Titration
Figure 6. Rate of Disappearance of Citral at 25' C. by Hydroxylamine Sulfate
First-order kinetics bawd on extrapolated values for initial unsaturatioh
a
1
I
0.1 % HiPo4 0.5% HsPOi
I
I
I
I
I
I
I
I
1
0 0.5% Hap04
I 2 3 TIME IN HOURS Figure 7. Rate of Disappearance of Citronellal by Hydroxylamine Sulfate 0 A t 2 5 O . 0 A t 45'. [H+] 5.5 X 10-6
and 35% in 5.2 hours) at very much the same rate as the decomposition estimated by the other methods used. D. BYHYDROXYLAMINE SULFATE.In view of the ambiguity of the m-phenylenediamine and the impracticability of the fuchsine-sulfite procedures for establishing t,he rate of disappearance of the carbonyl function, an investigation was made of the application of the hydroxylamine sulfate method (6, 18) to the extremely dilute solutions encountered. It was necessary to reduce the barium hydroxide used for titration to 0.01 N , to use not more than a double excess of hydroxylamine sulfate, and to
titrate potentiometrically to a pH of 4.25, where experiment indicated that the end of the titration occurred, rather than colorimetrically with bromophenol blue indicator. To a 50-ml. sample of the citral solution was added 0.5 ml. of 8% hydroxylamine sulfate and, after one minute, the sample was titrated to pH 4.25 with 0.01 N barium hydroxide. Since it was essential to avoid buffering, weak acids could not be used as catalysts in these experiments. Very dilute hydrochloric acid was therefore employed. Samples of the reaction mixture were withdrawn, titrated to pH 4.25, treated with hydrox-
2 4 6 TIME IN HOURS Figure 5. Rate of Disappearance of Citral at 25' C. by rn-Phenylenediamine 1.0% H 8 0 4
9 A t 25'.
-
[H+1 = 7.8 X 10-4
260
INDUSTRIAL AND ENGINEERING CHEMISTRY
ylamine sulfate, and again titrated to pH 4.25. The data at 25", illustrated in Figure 6, and the data at 25" and 45', summarized in Table I, are in agreement xith those obtained by may other the be due methods. to slightThe experimental two high error valuesinfor titrating the ratetheof very 45O small hydrochloric acid concentration used for these experiments. The fact that hydroxylamine titration indicated complete loss of the carbonyl group when only one of the two double bonds had disappeared substantiates the conclusion that the initial reaction involves only one double bond.
Vol. 40, No. 2
CK
CH,
A
CH, 1
CH I
CHz
\CH
5 % HzS04
CHo
6LH
~
+
CH3C02CzH6
/&\
CHs
I/
CHz
()OH I
/\\ Hz
CHs
CHZ
/"\ CHI
CHJ RATE OF DECOMPOSITION OF CITRONELLAL
The procedure employed was the same hydroxylamiue sulfate
Dehydropulegol
Citral
Dehydroisopulegol
hethod as that used for citral. The reaction was so much more rapid that addition of acid led to almost complete decomposition Prins ( 2 1 ) has pointed out the close analogy between the acldby the time the first measurement could be made. The rate catalyzed cyclizations of citral and citronellal outlined above and constant for acid decomposition m s therefore estimated from the acid-catalyzed condensation of aldehydes with olefins which the decomposition in n-ater of pH 5.26. bears his name. The data, summarized in Figure 7 and Table 11, demonstrated that the decomposition of citronellal hydrating R'CHzCHCH2CH20H (A) occurred at a rate about 500 times greater than for conditions' I OH citral, both at 25' and 45 CHzO R'CH2CH=CH2
".
i
+
acid
DISCUSSION
Although the very low solubility of citral and citronellal in aqueous solution made it difficult to attain a high degree of accuracy in the quantitative measurenicnts, the precision was sufficient to establish that unsaturation (by bromine titration) equivalent to one of the two original carbon-carbon double bonds in citral remained when the carbonyl group (by hydroxylamine titration) had entirely disappeared. Further subsequent decrease in unsaturation must be due to a secondary reaction, since it vas observed only concurrent with precipitation of an oil, which retained unsaturation corresponding roughly to one double bond per citral residue. In addition, the rates of disappearance of the double bond and of the carbonyl group could each be satisfactorily analyzed on the basis of kinetics first-order Kith respect to citral and acid-catalyzed. The agreement of the rate constants for the two processes Fvas sufficient to indicate that the two rates were identical. Since it would be a remarkable coincidence for two entirely independent reactions of the carbonyl group and the double bond to occur with the same kinetics and at the same rate, it seems most logical to conclude that the two groups interact. Furthermore, since citronellal with no a,pdouble bond reacts even more readily than citral, it must be the more distant double bond which is involved. These conclusions from the kinetics are aniply substantiated by observations previously reported on the nature of various acid-catalyzed transformation products of citronellal and citral. For example, Barbier and Leser (3) were able to isolate menthoglycol, as well as isopulegol and a dimeric ether, from the action of 5% sulfuric acid on citronellal. CHs
Citronellal
CH,
CH, j
Menthoglycol
Isopulegol
S'erley (14) and Zeitschel and Schmidt (16) were able to isolate dehydropiilegol and dehydroisopulegol from the action of 50% sulfuric acid in ethyl acetate on citral.
\dehydrating+ conditions
R'CH=CHCHzCH,OH (13)
In the usual Prins reaction, the 1,3-glycol (A) may react furthcr with the aldehyde to form a 1,3-dioxane. The course of the reaction leading to an unsaturated alcohol (B) has been discussed by Baker (P), who considered the mechanism to involve attack at the a-methylenic position of the olefin. It seems difficult to account for the catalytic action of acid on the basis of the mechanism proposed by Raker. There is, howcvcr, an alternative course for the reaction which accounts equally well for the products and in addition offers a sound and satisfactory explanation for the dependence of the rate on the acid concentration.
TABLE 11. RATE
O F DECOiVPOSITION O F
CITRONELLAL
IX
AQCEOUSSOLUTION Temp.,
c.
[Ciaronellal]~, .M x 103
45
1.48
25
1.24 1.133
25
\ H A ; ,,V 5.5
x
10-6 5.6 X 10-6 7 . 8 2 X 10 - 4 (HCI)
k , 1. mole-' 2.41 x 3.84 x 4.6 x 1
hour-' 104 103 103
Since the carbonyl osygcn is the most basic portion of either of the two reactive functional groups, it would be the preferred posilion for coordination with a proton froin the acid. The resulting cationic carbonyl carbon atom could t,hen couple TTith a pair of electrons from the olefin to give a new &ionic intermediate (C). The intermediate (C)could either react with water t,o give tht: glycol (A) or lose a proton to give the unsaturated alcohol (B). On the basis of this course for the Prins reaction, one can thcri suggest the following course for the acid-catalyzed cyclization oi citral (or citronellal), which accounts both for the products observed previously and the kinetics reported herein.
CH,
A N D E N G I N E E R I N GCHEMISTR-Y
INDUSTRIAL
February 1948 CHs
261
amount of D1 asLn be represented as a constant (and perhaps a minor) fraction of the total concentration of D1 and Da.
CHI
[Dil
Ki[Dal
(2)
The Concentration of D1 is also going to be regulated by the equilibrium reaction of citral with acid. \CH
I
[DaI
Ka = [H+][citral a]
(3)
[Dll [H+][citral b]
(4)
K,
=
,
[Dl] = Ka[H+][citral b] = KoKl[H+][citrala]
(5)
On the basis of the mechanism suggested, the equilibrium between citrals a and b will be established more rapidly than the . cyclil;cltion and will be given by the following expression (from Equation 5):
CHa
I
\\
C/H,
CH
CHI
CHOH
I
I
\CH
+H+
\CH/ I
Dehydromenthoglycol
[citral b] KO [citral a] =
z
K1
Thus the equilibrium distribution in the cation (between DI and Da) would be the same as that between citrals a and b only if K , were equal to Ka-i.e., only if citral a and citral b had equal base strengths. From Equation 6, the equilibrium concentrations of citrals a and b in terms of the total citral concentration can be expressed as follows:
+ KoKl [citral b] = KoKl[ citral] /Kb + KoK1 [citral a] = K,
[citral]/Kb
(7)
Substitution of (7) in (5) gives an expression for DI in terms of total citral concentration which, when substituted in Equation 1, give kinetics in agreement with those observed.
dx/dt = k[H+] [citral] k = keKsKbKdKb
Citral a
Pa)
The coordination of a proton with the carbonyl group of citral to give the cation (D) accounts satisfactorily for the acid-catalyzed cis-trans isomerization of citrals a and b (6),due to the resonance indicated by the canonical formulas D1, Dz, and Ds. In addition, the intermediate D1 (which may, of course, be present as only a minor contribution to the intermediate D) may cyclize slowly but substantially irreversibly to the cation E, which then can react rapidly with water to form dehydromenthoglycol. Dehydromenthoglycol was not isolated in the present investigation but is postulated as the primary transformation product, (a) from analogy to the menthoglycol isolated from citronellal (9) under similar conditions and (b) from the analytical data reported here, which indicate the presence of one double bond in the primary product. In the substantial absence of water, the cation E loses a proton at either side of the cationic carbon to give dehydropulegol and dehydroisopulegol (1416). If one presumes that the cyclization reaction, DI +E, is the rate-controlling step in the conversion of citral to dehydromenthoglycol, one can then write the rate of reaction as proportional to the intermediate D1. &/dt = ka [Dil
(1)
The contribution of form D1 to the intermediate is going to be regulated by factors unaffected by concentration-Le., the
+ KaKI
The subsequent transformation of dehydromenthoglycol to an insoluble oil apparently involves both dehydration by loss of two molecules of water to cymene and polymerization. The loss of water is not surprising, since one hydroxyl group is tertiary and one is both allylic and secondary. The polymerization undoubtedly involves either the unsaturation already present in dehydromenthoglycol or that in the dehydration products intermediate in the transformation to cymene. LITERATURE CITED
(1) Assoc. Official Agr. Chem., “Official and Tentative Methods of Analysis,” 4th ed., p. 311,1935. (2) Baker, J. W., J. Chem. Soc., 1944,296. (3) Barbier, P., and Leser, G., Compt. rend., 124,1308 (1897). (4) Boeseken, J., Hansen, L. W., and Bertram, S. H., Rec. trav. chim., 35,315 (1916). ( 5 ) Bouveault, L., Bull. 8oc. chim. (a),21,423 (1899). (6) Gettler, J. D.,and Hammett, L. P., J. Am. Chem. Soc., 65, 1824 (1943). (7) Goldschmidt, H., and Burkle, E., Ber., 32,364 (1899). (8) Hiltner, R.S.,J. IND.ENG.CHEM.,1, 708 (1909). (9) Holwerda, B.J., Bwchem. Z., 128,466 (1922). (10) Ostwald, W., 2.physik. Chem., 3 , 191 (1889). (11) Prins, H. J., C h m . Weekblad, 14,627 (1917). (12) Sohultes, Hermann, 2. angew. Chem., 47,258 (1934). (13) Tiemann, F., Ber., 32,818 (1899). (14) Verley, A,, Bull. 8oc. chim. (3),21,408 (1899). (15) Woodman, A.G.,“Food Analysis,” p. 472,New York, McGrawHill Book Co., 1941. (16) Zeitschel, O.,and Schmidt, H., J.prakt. Chem., 133,370 (1932) R ~ C ~ I VJune B D 28, 1946.
