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Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
Tetrahedron, 1963, Vol. 19 Suppl. I, PP. 197 to 212. Pergamon Press Ltd.
KINETICS AND EQUILIBRIUM OF THE NITROFORM-ALDEHYDEALCOHOL REACTIONl C.
S. RONDESTVEDT, JR.,
M.
STILES
and A. L.
KRIEGER2
Abstract- The reaction of nitroform.van aldehyde, and an alcohol yields a trinitroether in a reversiblereaction. The probable mechanism involves rapid formation of the hemiacetal, conversion to an alkoxycarbonium ion, and reaction with nitroformate anion. RCHO
+ HNF + R'OH
=
= HNF
RCHOH(OR')
R~HOR' ,
NF-
1\
RCH(OR')C(N0 2h The product is attacked by excess alcohol and is converted reversibly to acetal. Reaction of the acetal with nitroform is the most convenient method of preparing trinitroethers in high yield, since the water can be disposed of during the acetal preparation. The kinetics of the acetal-nitroform and the trinitroether-alcoholreactionhave been studied in dioxane solution. The rate expressions obtained are interpreted in terms of a mechanism involving several equilibria centering around the alkoxycarbonium ion RCHOR'.
+ THE reaction of nitroform with an aldehyde and an alcohol (or a mercaptan or amides) was first discovered by Hartman and Tawney of the United States Rubber Company. Because of its potential value in preparation of aliphatic polynitro compounds, the synthetic scope of the reaction was extensively investigated by Tawney's group. The over-all reaction (Eq. 1), though three components participate, is unlikely to be a termolecular process. More reasonably it involves some or all of the well-known equilibria shown below. Equations (2) and (3), the hemiacetalacetal equilibria, require no comment. The well-known example ofEq. (4) is the formation oftrinitroethanolfrom nitrofonn and formaldehyde. Most other nitroThis work was carried out during 1953-1955 with support from the U. S. Navy Bureau of Ordnance, through a sub-contract from the United States Rubber Co. The authors are indebted to Dr. P.O. Tawney for many helpful discussions. 2 Present addresses: C. S.Rondestvedt, Jr., Jackson Laboratory, E.I. du Pont de Nemours and Co., Inc., Wilmington 99, Delaware; M.StiIes, University of Michigan, Ann Arbor; A. L. Krieger, Ansco, Binghamton, New York. Inquiries and requests for reprints should be sent to C. S. Rondestvedt.
1
197
C.S.RONDESTVEDT, JR.,
M. STILES and
AL.KRIEGER
form-aldehyde adducts are not stable; on attempted isolation, they revert to the addends. OR' RCHO + R'OH + HNF
RCHO + R'OH
~
~
I
RCHNF + H20
OH I ReHOR'
RCH(OH)OR' + R'OH RCHO + HNF
~
~
(2)
RCH(OR')2 + H 2 0
RCHOHNF
(3)
(4)
R'OH + HNF ~ R/OH~ + NF '3
(5)
The system n-butyraldehyde-ethanol-nitroform was selected for study, since all components and the product 1,1,I-trinitro-2-ethoxypentane (I) were convenient to handle and purify. Preliminary quantitative study of reaction (6) PrCHO + HNF + EtOH
~
PrCH(OEt)NF + H20
(6)
I
did not afford reproducible results, a difficulty quickly traced to small variable amounts of water; Eq. (6) is reversible. It soon appeared that use of excess ethanol as solvent impaired the yield of I, and butyraldehyde diethyl acetal (II) was found in the equilibrium mixture. Separate experiments showed that I reacts reversibly with ethanol fairly rapidly, yielding II and nitroform. Since I + EtOH ~ PrCH(OEt)2 + HNF
(7 a)
II
II + HNF ..!!l....... I + EtOH
(7 b)
reaction (6) was manifestly complicated, a detailed study of the rates arid equilibria of the component reactions was launched. It was expectedthat a thorough understanding of the model system would permit designing experimental conditions for the synthesis of new types of compounds useful as propellants and high explosives. The system acetal
+ nitroform ~
trinitro ether
+ alcohol
The acetal-nitroform reaction. Anhydrous dioxane was selected as the "inert" solvent for most of the quantitative work, although carbon tetrachloride re3
For simplicity, nitroforrn (trinitromethane, (02N)SCH) is abbreviated HNF; the trinitromethyl group is then NF. 198
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
ceived some attention.' The reactions were followed by titration for nitroform. It was essential to work under rigorously anhydrous conditions. The reaction (7b) is first-order in HNF and in II. For various initial concentrations, plots of initial rate Vo vs, [HNF]o for various [Il], (Fig. 1) and of vo![II]o vs. [HNFl o(Fig. 2) (where subscript zero represents initial concentration) 20,-----------------,
15
"2 ...E
~ :s"
,0
8
6
" o
~
_ ___'_ _ 0'05
_.J._ __ ' __ ___'__ ____L__ _- '
0"0
0"5
0·20
0·30
(HNFlo m FIG. 1. Reaction of nitroform (HNF) with diethyl butyral (Ae) in dioxane, 50°. Initial rate vs. initial nitroform concentration at three different acetal initial concentrations
were linear. The values of the bimolecular rate constant k~ obtained from Fig. 2 are 3·6 x 10-2 and 1·4 x 1O-21/mole!min at 50·0 and 40'0°, respectively. Numerical calculation from the integrated form of the bimolecular equation, in which the reaction is regarded as irreversible but only the first portion of the data before curvature sets in is used, yielded values of 102k~ varying from 3·78 to 4,50, average 4·17, at 50°, and 1·51 and 1·65 at 40°. The Arrhenius activation energy was estimated to be about 19 kcal/mole, 4
Alcohols and acids are obviously unsuitable, and hydrocarbons are insufficiently polar. Probably other ethers would work just as well. More basic solvents like tetrahydrofuran, dimethylformamide, and dimethyl sulfoxide would be of considerable interest. 199
C.S.RoNDESTVEDT, JR., M.STILES
and
A.L.KRIEGER
A brief study of the reaction (7 b) in carbon tetrachloride at 50" showed that the order was less than 1 with respect to II and greater than 1 with respect to nitroform. The initial rates for comparable initial concentrations were slower in carbon tetrachloride than in dioxane. An attempt to treat (7b) as a second-order reversible reaction by means of the integrated form of Eq. (8) yielded poor fits. The plots showed curvature for
8,-------------.,..-----, 050° C ., 40° C 7
6
~
5
o
Q; u
e
"'-.4 ~
'"Q
3
2
Slope: k =I-4. 10-~L/mole/min
o ! : - - _ - - I_ _- l . ._ _---'-_ _-L._ _--l.-_----l
0·05
0'10
0'20
0-15
0'25
0·30
(HNF)o
FIG.
2. Reaction of nitroform (HNF) with diethyl butyral (acetal) in dioxane, 50° and 40°. Initial rate/acetal initial vs. initial nitroform concentrations
the first 15-20 % of reaction, but thereafter became straight lines whose slopes gave an average value of k~ = 3·53 x lO-21jmole/min at 50°. - d [HNF]/d t = k~ [HNF] [II] - k~l [I] [EtOH].
(8)
The ethanolysis of I. - Reaction (7 a) is cleanly first-order in 1. When the integrated form of Eq. (8) was plotted for different runs, excellent straight lines were obtained. Surprisingly, the values of from the slopes varied considerably for different values of [EtOH]o. The order with respect to ethanol is unity when the ratio [EtOH]/[I] < 0·7. Under these conditions, k~l (50°, dioxane) = 5·5 x 10-3 l/rnole/min. When the ratio is increased, the order with respect to ethanol decreases and approaches zero with ratios greater than about 1·8. The reaction (7a) in carbon tetrachloride was too slow to measure at 50°, and at 75° decomposition occurred.
r,
200
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
The equilibrium constant. - Eleven determinations in both directions at 50° in dioxane yielded K = [I] [EtOH]/[II] [HNF] = 7·35 ± 0·27 "'" k~/k~l' The magnitude of this constant is such that I in excess ethanol would be converted almost completely to II and nitroform. Excess alcohol is certainly not the solvent of choice for reaction (1) in preparative use. At 40° in dioxane, K = 5·2 and 7·7 for different initial concentrations; the poor agreement suggests that a side reaction, perhaps that discussed in the next section is more important at 40° than at 50°. In carbon tetrachloride at 50°, the value of K is 12'0, more favorable for the production of 1. However, the equilibrium is attained much more slowly than in the more polar dioxane. Decomposition of I in dioxane. -Pure I in dioxane at 50° dissociated to the extent of 2-5 mole % in two days, as shown by titration of the liberated nitroform. The decomposition then stopped, indicating that an equilibrium had been reached. It is supposed that the vinyl ether l-ethoxy-l-butene (III) is formed by analogy to the well-known acid-catalyzed reversible dealcoholation of acetals.?'
I
~
HNF
+ CH sCH 2CH=CHOEt
(9)
III
III was not identified positively since it was present only in small quantity. Vapor-phase chromatography was not available at the time of this work. This decomposition, whatever its exact nature, does not seriously affect the rate or equilibrium constants determined for reaction (7) at 50°, except at low ratios of [EtOH]/[I], and these were the runs which gave values for K deviating the most from the average. The salt effect. - The effect of added salts on the ethanolysis of I is shown in Table 1. It is clear that the initial rate of ethanolysis is considerably enhanced by low concentrations of salt. At the same ionic strength 0·002 M, nitroformate ion depresses the rate about 15%. The study could not be carried to higher ionic TABLE
1. EFFECT OF SALTS ON INITIAL RATE OF ETHANOLYSIS OF I (Reaction (7a)] IN DIOXANE AT 50'0°; [EtOH]o.::=: 0'12 M; [I]o = 0'18 M Benzyltrimethyl-
ammonium
Perchlorate Perchlorate Perchlorate Nitroformate a
[Salt], M
0·0000 0'0007 0'0020 0·0032" 0,0020
IDS Vo, moles/ Ijmin 0,91 1,18 1'4 ± 0·1 1'65
1'17
The solubility limit in dioxane at 50°.
4a H.Shechter and H.L.Cates, J. Org. Chem. 26,51 (1961), report the preparative addition of HNF to vinyl ethers, but do not mention that the reaction is reversible.
201
C.S.RONDESTVEDT,
JR., M.STILES and A.L.KRIEGER
strengths because of the very low salt solubilities in anhydrous dioxane. Benzyltrimethylammonium bromide and nitrate are even less soluble than the perchlorate, and the perfluorobutyrate could not be obtained pure. The salt effect was not studied in carbon tetrachloride. 17·0
15·0
"
Q
10·0
o Added ccid 5'0
o
0'010
0'020
0·030
0·040
c-oso
(Nitroforrnl, x (acetol)o
FIG. 3. Reaction of nitroform with diethyl butyral (acetal) in dioxane, 50°, with and without added p-toluenesulfonic acid
Acid catalysis. - Intuitively one expects a reaction of an acetal or an ether to be acid-catalyzed. However, neither the forward nor reverse reaction (Eq.7) was influenced by anhydrous p-toluenesulfonic acid or perfluorobutyric acid in dioxane. The rates in either direction were identical to those without added acid, within experimental error, Figure 3 shows the results with the forward reaction. Some rather erratic rate enhancement was noted in carbon tetrachloride solvent. Isopropyl alcohol. - The system with isopropyl alcohol (Eq. 10) was studied briefly. Isopropyl alcohol reacted with I in dioxane slower than did ethanol. Qualitatively the same type of mixed order dependence on [i-PrOH]o was noted as with ethanol. The initial rate of the reaction of diisopropyl butyral (Iv) in 202
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
carbon tetrachloride was about ten times faster than that of the ethyl acetal II with nitroform in the same solvent. The kinetics were more erratic than with II. CH gCH2CH2CH(Oi-Pr)2 + HNF IV ;.=
(10)
CH 3CH 2CR 2CH(Oi-Pr)NF + i-PrOH V
No valid equilibrium constant could be obtained in carbon tetrachloride. Values for the expression [V] [i-PrOH]![HNF] [IV] varied from 27·4 to 64·3 for various initial concentrations. However, the equilibrium concentration of' V in reaction (10) is higher than that of I in reaction (7) (50 CCIJ. Mechanism of the reaction (7). - The following scheme is consistent with the facts. 0
,
PrCH(OEt)2
k,
+ HNF -;---'"k PrCROEt + NFI
-1
(11)
HOEt
II k,
PrCHOEt -;---'"k PrCROEt + EtOH
I
-I
(12)
®
VI
HOEt ®
PrCHOEt ffi
k,
+ NF- +-k~ PrCR(OEt)NF
(13)
-8
I
The first step is a transfer of a proton from undissociated nitroform! to the acetal. The conjugate acid splits into ethanol and a solvated carbonium ion, which in the final step reacts with nitroforrnate anion to form trinitro ether. It is probable that all ionic species exist as ion-pairs in the slightly polar solvent dioxane, rather than free ions. The reverse reaction is an ionization to an ion-pair, doubtless solvated by ethanol. Ethanol displaces nitroformate, and the solvated carbonium ion collapses to acetal. Assuming no accumulation of either unstable positive ion, one may apply the method of Christiansen" to Eqs. (11)-(13) and obtain the espressions (14) 5
Nitroform is a rather strong acid in water, but it is apparently not significantly dissociated in dioxane. The characteristic yellow color of nitroformate anion, which is readily perceived at 10- 4 M in water, cannot be seen in dry dioxane. Some degree of proton transfer to dioxane is probable, perhaps (5+ {y-
r",
O
O:H:NF
.......... _ /
6
Hammett, Physical Organic Chemistry pp. 107-108. McGraw Hill, N. Y. (1940).
14 Times N. R.
203
C.S.RONDESTVEDT, JR., M.STILES
and
A.L.KIUEGER
and (15), where V = V + - V_. V+ = _ d[HNF] = d[I] = k Ik2ka[II] [HNF] dt dt k zka + k_Ik_ z [EtOH] + k_Ik a[NP-]
(14a)
(14b) In the reverse direction, the initial rate is given by Eq. (15). (V-)0
= =
k_Ik_zk_ a [EtOH] [I] k 2ka + k_Ik_z[EtOH] + k _Ik a[NP-]
(15)
».,«,«, [EtOH] [1] k a(k 2
+ k_1[NFD + k_Ik_z[EtOH]
Various limiting cases may now be considered. Case 1. Assume k z ~ k_1[NF-]. A. Assume further that k _1k_z [EtOH] ~ kzk a• Then (15) reduces to (16). (16)
That is, the rate of ethanolysis of I is independent of ethanol at high [EtOH]o. B. The reverse of condition A converts (15) to (17). (V-)0 = k_1:--;'C a [EtOH] [I] = k 4 [EtOH] [I] z a
(17)
That is, at low (EtOH]o> ethanolysis of I is first order in ethanol. C. For the forward reaction (7 a), the rate at any time is given by (18) and the initial rate by (19). V+ = klkzk a[II] [HNF] (18) kzk a + k_1k_ z [EtOH] (V +)0 = k l (II] [HNF] .
(19)
According to (19), the initial forward rate is bimolecular and Eq. (11) is ratedetermining. This will be true at any time if condition B applies, as it does at low [EtOH]o' Case II. Assume k 2 ~ k ., [NF -]. Then (15) reduces to (20). (V_)o
A. If k_"2 [EtOH]
~
=
k _zk_ a[EtOH] [I] k a[NP-] + k_z[EtOH]
k« [NP -], (20) reduces to (17) with k:« = k 4 • 204
(20)
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
B. If the opposite is true, (20) becomes (21), which is equivalent to (17) if [NF-] quickly attains a constant value, or if a nitroform salt is added.
(V ) -
= 0
7c k_ 2 S k s [NP-]
[EtOH] [I]
(21)
C. For the forward reaction, the rate at any time is given by (22), and initially by (23) which is of the form (19). V+ =
(V1·)0 =
k lk 2k s [II] [NHF] klks[NP-] + L 2 [EtOH]
k 1k2 cdNP-]
J,
[II] [HNP]
(22)
(23)
Either Case I or II fits the data on the reverse reaction, that the observed rate is first order in [EtOH]o when it is low, and zero order when it is high. Either case will also fit the bimolecular kinetics observed with the initial forward reaction. The limiting equation (21) was tested by four runs in dioxane at 50° with the following initial concentrations: [I] = 0·18 M; [C6H s CH2NMes NF -]0 = = 0·0020 M; [EtOH]o = 0·017-0'060 M. A plot of' V, vs. [EtOH]o was linear, as required by Eq. (21). Eq. (20) was rearranged to (24). (24) By assuming that the stationary concentration of [NP-] was the same for all values of [1]0 and [EtOH]o which had been investigated, Fig. 4 was obtained. The scatter at low [EtOH]ois a result of the fact that a small error in Vo causes a large error in its reciprocal; the lengths of the bars are estimates of the error in l/Vo' From the intercept, k.; = 1-4 X lO-s min:", and k s[NP-]lk_ 2 = 0·2. This value of L s agrees well with the limiting value of Vol[I] at low [EtOH]o, 1·2 x lO-s min -1. Mixed kinetics are to be expected when the two terms in the denominator of Eq. (20) are about equal; i.e., when [EtOH] ~ 0·2; this is the range where we observed mixed kinetics. The absence of catalysis by strong acids is surprising, since in polar solvents only the proton (specific lyonium ion catalysis) is effective in converting an acetal or hemiacetal to VI. The strong acid would reduce [NP-] by converting the ion to molecular nitroform. This would help to insure that Case I was valid. However, a drastic reduction in [NP-] would retard reaction (13), probably enough to make it rate-determining in the forward direction. Perhaps an increase in rate of production of VI is balanced by a retardation of its conversion to 1. 7 Since dioxane is a much poorer ionizing solvent than alcohols or 7
A reaction of VI with undissociated nitroform could replace Eq, (13) if acid is added, and this could well become rate-determining. VI + HNF ~ I + H-.
205
C.S.RONDESTVEDT, JR., M.STILES and A.L.KRIEGER
water, it is not unlikely that molecular nitroform may function to form VI as an ion pair. The unexpected rate-determining proton transfer in Eq. (11) which is required by the observed kinetics points out the difference in the solvents.8 Since carbon tetrachloride is an even poorer ionizing solvent than dioxane, reactions which involve ionizations should be retarded by a change to carbon tetrachloride. This effect was observed. 5·0
r--------------------.,
3·0 c:
E ~
';::,.
'"Q
H
2'0
a
5
10
15
20
25
30
FIG. 4. Reaction of 1,1,1-trinitro-2-ethoxypentane (I) with ethanol (EtOH). Plot of Eq. (24)
The mechanism written is quite similar to that accepted for the hydrolysis and other reactions of acetals. The isopropyl acetal IV should and does react faster than the ethyl acetal II with nitroform, just as. diisopropyl formal is hydrolyzed six times faster than diethyl formal." The intermediate i-PrOCHR + is stabilized somewhat more than EtOCHR+ by the greater inductive effect Time did not permit an experiment with deuteronitroform to provide further evidence for rate-determining proton transfer. 9 A.Skrabal and H.H.Eger, Z. physik. Chern. 122,349 (1926).
8
206
Kinetics and equilibrium of the nitroforrn-aldehyde-alcohol reaction
of the isopropyl group." On the other hand, isopropyl alcohol is known to form acetals from aldehydes more slowly than ethanol and with a less favorable equilibrium; this parallels the lesser reactivity of isopropyl alcohol toward 1. The positive salt effectobserved for reaction (7a) is consistent with the operation of a rate-determining ionization such as (13). The smaller rate with nitroformate ion compared to perchlorate is probably a result of a common-ion depression of the ionization. Alternative explanations, that the two salts are ionized to different extents in dioxane, or that they have different activity coefficients, cannot be excluded by our data. It is unfortunate that the salt effect could not be studied with different salts and at higher concentrations. Sulfur analogs. Hartman and Tawney showed that a mercaptan would react with an aldehyde and nitroform (Eq. 25). We observed that butyraldehyde diethyl (25) RCHO + HNF + R'SH ~ RCHNF + H 20
I SR' VII thioacetal did not react with nitroform in dioxane at 50° during three days. The mixed oxygen-sulfur acetal l-ethoxybutyl ethyl sulfide reacted with nitroform to yield a mixture of materials probably containing VII (R = Pr, R' = Et), but it could not be obtained in analytical purity. On the other hand, the cyclicmixed acetal 2-propyl-l,3-oxathiolane did not react detectably with nitroform. The equilibrium for cyclic acetals is farther to the right than for acyclic acetals. A carbon-sulfur bond is not cleaved by nitroform, at least under these mild conditions. Excess butyl mercaptan at 50° converted I to a mixture apparently containing thioacetal, I, and the trinitro thioether VII, but the reaction was far from clean. It had been previously observed by Tawney that no reaction occurred at room temperature. 2-Mercaptotetrahydropyran reacted with nitroform to liberate hydrogen sulfide, but VIII could not be isolated. The parallel experiment with 2-hydroxytetrahydrothiapyran could not be performed; water failed to add to the double bond of 2,3-dihydrothiapyran, and an alternative synthesis also failed. Semiquantitative experiments with the system butyraldehyde-butyl mercaptan-nitroform showed that the equilibrium lay further on the side of product than that for the ethanol system studied in detail. The hemiacetal as a reactant. At this juncture it appeared that the formation of I via Eq. (6) proceeded by a moderately rapid nitroform-catalyzed acetalization of the aldehyde, followed by a somewhat slower conversion of acetal to trinitraether by the path already discussed. However, since acetal formation in alcoholic solution proceeds through the intermediate resonance-stabilized 10
C.K.Ingold, Structure and Mechanism in Organic Chemistry p. 334. Cornell University Press (1953).
207
c. S. RONDESTVEDT, JR.,
M. STILES and A. L.KRIEGER
(doubtless solvated) carbonium ion VI, already proposed in Eq. (12) for the nitroformolysis of II, it seemed equally reasonable that VI formed from hemiacetal would react directly with nitroformate ion or molecular nitroform (Eqs. 13-14) as well as follow the longer route to acetal and back again. This is equivalent to asking whether VI reacts faster with ethanol or with nitroform (or its ion). The question could be answered by a knowledge of the rate of acetal formation.There is no information in the literature on the rate of acetal formation or hydrolysis in dilute solution in non-polar solvents, despite the fact that these reactions in aqueous or alcoholic solutions are beloved topics for kinetic studies. In alcoholic solution, reaction (2) (hemiacetal formation) is general acid-catalyzed, i.e, molecular or undissociated acids are active as well as the lyonium ionY The subsequent conversion of the hemiacetal to intermediate VI is specific lyonium ion-catalyzed in polar solvents." Acetal hydrolysis in water solution" and in 60 %dioxane's is specific oxonium ion-catalyzed. In hydrolysis, the rate is faster in D 20 than in H 2 0 , indicating that protonation of the acetal occurs prior to the rate-determining step." It is risky to extend these conclusions obtained in ionizing solvents to nonpolar media like dioxane. Time did not permit a thorough study of the rate of acetal formation in anhydrous dioxane, and therefore a few semi-quantitative experiments were performed. We estimated the equilibrium constant K 2 , Eq. (26), for the hydrolysis of II in dilute dioxane solution containing limited amounts of water. The aldehyde formed K 2 = [II] [H 20]/[PrCHO] [EtOH]2
(26)
was measured by oximation ;16 this method, of course, includes hemiacetal with the aldehyde. The apparent values of K~Oo ranged from 0·17 to O'30 l/mole. To correct for the amount of hemiacetal at equilibrium, rough data were obtained for the formation of hemiacetal. About 15 % of the aldehyde was converted to hemiacetal when a dioxane solution 0·2 Min butyraldehyde and 0·3 M in ethanol was allowed to stand at 22°.I' The constant K2 was then measured at the same temperature; the aldehyde concentration determined by oximation was reduced by the amount of hemiacetal present. The values of K~2° ranged from 11 For
example, see G. W. Meadows and B.B.Darwent, Canad. J. Chern. 30, 501 (1952); Trans. Farad. Soc. 48, 1015 (1952). 12 Deyrup, J. Amer. Chem. Soc. 56, 60 (1934). 13 J.N.Bronsted and Grove, J. Amer. Chem. Soc. 52,1394 (1930); J.N.Bronsted and WynneJones, Trans. Farad. Soc. 25, 59 (1929); Olson and Tong, J. Amer. Chern. Soc. 66, 1555 (1944); Grove, Ibid. 52, 1404 (1930); Mi Kilpatrick and Chase, Ibid, 53, 1732 (1931). 14 E. N. Prilezhaeva, E. S.Shapiro, and M. F. Shostakovskii, Zh. Obshch. Khim. 18,1663 (1948). 15 J.e.Hornel and J. A.V. Butler, J. Chem. Soc. 1361 (1936). 16 R. H. Buchanan, Austr. J. Appl. Sci. 2, 276 (1951). 17
The attenuation of the carbonyl peak in the infrared was measured. The temperature of the spectrometer room was maintained at 22°. At the time of this work, no equipment was available for thermostating the samples at 50°. 208
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
0·64 to 0·85 l/rnole, average 0·73l/mole. In runs at 50°, there should be less than 15 % hemiacetal, since the increased temperature would promote its dissociation into aldehyde and ethanol. It did not seem possible to study the conversion of acyclic hemiacetals to trinitroethers, since the former are unstable with respect to aldehyde and alcohol. For this reason, a stable hemiacetal was selected. 2-Hydroxytetrahydropyran (IX) exists almost entirely in the hemiacetal form, and very little of the hydroxy aldehyde is present at equilibrium." The bimolecular rate constant (Figs. 4 and 5) for the reaction of IX with nitroform (Eq. 27) at 50° in dioxane was 3·48 x 1O-2l/mole/min. The equilibrium constant measured in the forward direction was 83 ± 5, and about 48 hours was required to attain a constant nitroform titer. The reverse reaction, hydrolysis of VIII, was so slow under the same conditions that constant nitroform titers were not obtained in 10 days. It should be noted that pure Vfll in dry dioxane steadily liberated nitroform at 50°, perhaps with formation of dihydropyran by an elimination analogous to Eq. (9)for 1. For comparison, the reaction of the acetal XI with nitroform was followed under the same conditions. The initial rate was only about 2 % of that with the hemiacetal IX: accurate results were not obtained because the decomposition of VIII apparently interfered. r-.
.r-.
(O)-OH + HNF;=, IX /
A
I,
(o)ttJ;=' (O)-NF + H X
-,
I_OMe +
HNF
r-. 20
(27)
VIII
;=' X ;=' VIII + MeOH
"0/ XI
Mechanism of the aldehyde-nitroform-alcohol reaction. A reasonable path for the direct production ofthe trinitroether in a non-polar solvent involves the following steps. First, rapid formation of hemiacetal catalyzed by molecular nitroform. Second, reaction of hemiacetal with molecular nitroform to form alkoxycarbonium ion VI as an ion pair with nitroformate ion, plus water. Finally, internal reaction of the ion pair to yield product with displacement of water from the solvation sphere of the ion pair. According to this model, formation of the acetal requires the diffusion of ethanol from the dilute solution into L. E. Schniepp and H. H. Geller, J. Amer. Chem. Soc. 68, 1646 (1946), observed 95 % of the hemiacetal form in aqueous solution at room temperature. C. D. Hurd and W. H. Saunders, Ibid. 74, 5324 (1952), observed 94 % hemiacetal in 95 % ethanol or in 3: 1 dioxane-water at 25° and 93 % at 35°. They noted that equilibrium was established very slowly in dry dioxane and explained this by the necessity of having both an acid and a base, as has been shown necessary for the mutarotation of carbohydrate derivatives.P 19 E.R.Alexander, Principles of Ionic Organic Reactions pp. 159-160. Wiley, N. Y, (1950).
18
209
C.S.RONDESTVEDT, JR., M.STILES
and
A.LKRIEGER
the solvent cage surrounding the ion pair, and this appears to be relatively slow, judging from the comparative rates of reaction of IX and XI. This model is consistent with the results with mercaptans. The direct reaction of aldehyde, nitroform, and a mercaptan yields trinitrothio ester in high yields," but any thioacetal formed constitutes a reaction dead end. Trinitro alcohol. An alternative explanation must be disposed of. A priori, the equilibrium (4) might participate in reaction (1). The proposal is unattractive for the following reasons. The only stable, isolable ~-trinitroa1cohol is the formaldehyde-nitroform adduct trinitroethanol. Not only is this compound not etherified by ethanol, but also formaldehyde is unique among aldehydes in failing to give trinitroethers according to reaction (1).20 Since it was possible that a more labile trinitro alcohol might indeed participate, the reaction of butyraldehyde with nitroform according to Eq. (4) was studied briefly at 22° by infrared spectroscopy. In dioxane, the equilibrium lay well to the right, and it was attained in about half an hour; the hydroxyl band of the trinitro alcohol appeared at 3·05 ft. In carbon tetrachloride, the reaction was much slower, and the hydroxyl band was at 2·78/t. These equilibria should be established much faster at higher temperatures, and should lie further to the left. It thus appeared that a substantial portion of the aldehyde and nitroform would be present during the reaction as trinitroalcohol, the concentration of which would be much higher than that of hemiacetal. The acetal-nitroform reaction (7b) was then studied in the presence of added butyraldehyde. A marked reduction in rate would be expected if the effective concentration of nitroform were reduced by conversion to trinitroaIcohol. Within experimental error, the rate was identical to that obtained in the absence of added aldehyde. In any case, formation of trinitroether from trinitroalcohol is theoretically unattractive. First, if Eq. (28) represented the etherification, the formation of the intermediate XII would require a second mole of nitroform and should thus be second order in nitroform. Second, XII should be destabilized by the powerfully electron-attracting trinitromethyl group, while the alkoxycarbonium ion VI involved in the route through hemiacetal is stabilized. The situation is similar to that involved in the conversion of a cyanohydrin to an amino nitrile." The cyanohydrin dissociates to aldehyde and HCN, an aminohydrin is formed and converted to amino nitrile; in this way a cyanocarbonium ion is avoided. ®
RCHOHNF + HNF ::=::= RCHNF
+ H 20 + NF-
R'OH--+
XII
RCH(OR')NF + H+ (28)
Cyclic mechanism. An alternative cyclic mechanism obviates the necessity of forming ions in the non-polar solvent. If the aci tautomer of nitroform were the active species, the entire reaction from hemiacetal to product could proceed
20 P.Hartman and P. O.Tawney, unpublished results. 21
T.O.Stewart and C.H.Li, J. Amer. Chem. Soc. 60,2782 (1938).
210
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
by cyclic electron shifts within a transition state like XIII, in which the two nonparticipating nitro groups are not shown. With respect to the acetal-nitroform reaction (7b), formation and dissociation of XIII b would obey the observed second-order kinetics and should be insensitive to acid catalysis. A negligible salt effect would be expected, and this could be tested especially with solvents capable of dissolving larger quantities of salts than does dioxane. For the ethanol-trinitroether reaction (7a), formation of XIII b requires second-order kinetics, but this is observed only at low concentrations of ethanol. At higher ethanol concentrations, the ionization process [reverse of (13)] must be dominant. In the presence of salts, the observed acceleration could reflect the transformation of the cyclic mechanism into the ionization mechanism. R'O' OEI
"H,
"-"0
RCH' ("iiI-o /~7 .....c"
XIII a, R' XIII b, R'
= =
H Et
Swain has observed "polyfunctional" catalysis in the mutarotation of tetramethylglucose, and has pointed out that one advantage to a cyclic mechanism is that the substrate need never acquire an electric charge.P If this is true in water, an excellentionizing solvent, it should befar more true in poorly ionizing dioxane. The cyclic mechanism can be tested experimentally. It would first be desirable to find more versatile solvents for kinetic studies, which will dissolve salts, and in which a variety of different alcohol-aldehyde combinations could be examined without erratic kinetics.<1 The rate of enolization of nitroform in these solvents must be determined, for it is possible that enolization might be rate-determining in at least some solvents. The rates of hemiacetal and acetal formation and their catalysis by acids in non-polar solvents, and the rate of formation of trinitroalcohol, should be determined for comparison. With these data in hand, a kinetic study of the alcohol-aldehyde-nitroform system should yield meaningful results. It is probable then that the reaction can be extended to other carbon acids such as tricyanomethane, tris-alkanesulfonylmethanes, and the like. EXPERIMENTAL Materials. Dioxane was purified by refluxing with hydrochloric acid and then distilled and refluxed with sodium until no further hydrogen was evolved. It was then transferred to an apparatus consisting of a boiling flask surmounted by a reflux condenser and a takeoff built so that al1 the inner surfaces were continually washed by the hot liquid. The dioxane was stored in this apparatus over calcium hydride, then boiled for 1hr immediately before with22
C.G.Swain and J.F.Brown, J. Amer. Chern. Soc. 74, 2534 (1952). 211
C.S.RONDESTVEDT, JR., M.STILES and A.L.KRIEGER drawal for use. The apparatus was protected from moisture by a drying tower filled with phosphorus pentoxide. Ethanol was obtained by distilling commercial absolute ethanol from magnesium ethoxide immediately before use. n-Butyraldehyde was Eastman White Label Grade distilled in a nitrogen atmosphere and protected from the air. For the preparation of 1,1,I-trinitro-2-etlzoxypentane (I), the route via the acetal II was most efficient. Nitroform was mixed with a slight excess of diethyl butyral and kept at 70-75° for 1hr. The reaction was initially exothermic. The pressure was reduced somewhat to permit distillation of the ethanol, and finally the product was collected at 69° (0'8 mm), n~5 1'4424,87% yield. (Found: C, 33'51; H, 5'17; N, 16'58; Calc. for C7HlS07Ns: C, 33'47; H, 5'22; N, 16'73 %). 1,1,I-Trinitro-2-isopropoxypentane (V). An equimolar mixture of diisopropyl butyral and nitroforrn was maintained at 50° (100mm) for 6hr in a distilling flask, and then distilled. A light yellow liquid was obtained, b.p. 70° (0'3 mm) 66 % yield. It was washed twice with ice water and redistilled for analysis. (Found: C, 35'77;H, 5'71; N,16'46; Calc. forCBH!i,07Ns: C, 36'23; H, 5'70; N, 15'85 %). 'l-Trtnitromethyltetrahydropyran (VIII). Equimolar quantities of nitroform and 2-methoxytetrahydropyran (XI) were heated at 60-80° for 3hr. Distillation yielded 85 % of material, b.p, 78-80° (0·3mm). Redistilled material boiled at 70'0° (0'2mm), 1l~5 1'4705. Material of identical physical constants and infrared spectrum was prepared similarly from 2-hydroxytetrahydropyran (IX).2s ~ Although no explosions were encountered with any of the compounds described in this paper, prudence dictates that all experiments be conducted with appropriate precautions against explosions. Rate measurements. The reaction vessel was a large test tube fitted with a neck to hold a rubber serum stopple, a ground joint for attaching to the dioxane still, and a side arm protected by a Drierite tube. The vessel was flamed out in a stream of dry nitrogen. After the delivery tip of the dioxane still had been flushed with several portions of hot dioxane, the vessel was attached and the required quantity of dioxane was distilled in. A weighed quantity of one reactant was added, the side arm was sealed off in a flame, and the vessel was equilibrated in the thermostat. The second component was added from a weighed syringe. Aliquots for titration were removed by forcing 5ml dry nitrogen into the flask from a dry syringe and then withdrawing 5ml solution. The aliquots were run into 25ml carbon dioxide-free water to precipitate insolubles, and the nitroform was titrated rapidly with standard alkali to a phenolphthalein endpoint. The endpoint was stable for at least Zrnin. Pure I in aqueous dioxane corresponding to the concentration at the endpoint did not liberate detectable quantities of nitroform during the time required for the titrations. 23
Shechter and Catesv" prepared VIII from HNF and dihydropyran in 81% yield, b.p. 103° (1.0 mm), ll~o 1-4708.
212
KINETICS AND EQUILIBRIUM OF THE NITROFORM-ALDEHYDEALCOHOL REACTIONl C.
S. RONDESTVEDT, JR.,
M.
STILES
and A. L.
KRIEGER2
Abstract- The reaction of nitroform.van aldehyde, and an alcohol yields a trinitroether in a reversiblereaction. The probable mechanism involves rapid formation of the hemiacetal, conversion to an alkoxycarbonium ion, and reaction with nitroformate anion. RCHO
+ HNF + R'OH
=
= HNF
RCHOH(OR')
R~HOR' ,
NF-
1\
RCH(OR')C(N0 2h The product is attacked by excess alcohol and is converted reversibly to acetal. Reaction of the acetal with nitroform is the most convenient method of preparing trinitroethers in high yield, since the water can be disposed of during the acetal preparation. The kinetics of the acetal-nitroform and the trinitroether-alcoholreactionhave been studied in dioxane solution. The rate expressions obtained are interpreted in terms of a mechanism involving several equilibria centering around the alkoxycarbonium ion RCHOR'.
+ THE reaction of nitroform with an aldehyde and an alcohol (or a mercaptan or amides) was first discovered by Hartman and Tawney of the United States Rubber Company. Because of its potential value in preparation of aliphatic polynitro compounds, the synthetic scope of the reaction was extensively investigated by Tawney's group. The over-all reaction (Eq. 1), though three components participate, is unlikely to be a termolecular process. More reasonably it involves some or all of the well-known equilibria shown below. Equations (2) and (3), the hemiacetalacetal equilibria, require no comment. The well-known example ofEq. (4) is the formation oftrinitroethanolfrom nitrofonn and formaldehyde. Most other nitroThis work was carried out during 1953-1955 with support from the U. S. Navy Bureau of Ordnance, through a sub-contract from the United States Rubber Co. The authors are indebted to Dr. P.O. Tawney for many helpful discussions. 2 Present addresses: C. S.Rondestvedt, Jr., Jackson Laboratory, E.I. du Pont de Nemours and Co., Inc., Wilmington 99, Delaware; M.StiIes, University of Michigan, Ann Arbor; A. L. Krieger, Ansco, Binghamton, New York. Inquiries and requests for reprints should be sent to C. S. Rondestvedt.
1
197
C.S.RONDESTVEDT, JR.,
M. STILES and
AL.KRIEGER
form-aldehyde adducts are not stable; on attempted isolation, they revert to the addends. OR' RCHO + R'OH + HNF
RCHO + R'OH
~
~
I
RCHNF + H20
OH I ReHOR'
RCH(OH)OR' + R'OH RCHO + HNF
~
~
(2)
RCH(OR')2 + H 2 0
RCHOHNF
(3)
(4)
R'OH + HNF ~ R/OH~ + NF '3
(5)
The system n-butyraldehyde-ethanol-nitroform was selected for study, since all components and the product 1,1,I-trinitro-2-ethoxypentane (I) were convenient to handle and purify. Preliminary quantitative study of reaction (6) PrCHO + HNF + EtOH
~
PrCH(OEt)NF + H20
(6)
I
did not afford reproducible results, a difficulty quickly traced to small variable amounts of water; Eq. (6) is reversible. It soon appeared that use of excess ethanol as solvent impaired the yield of I, and butyraldehyde diethyl acetal (II) was found in the equilibrium mixture. Separate experiments showed that I reacts reversibly with ethanol fairly rapidly, yielding II and nitroform. Since I + EtOH ~ PrCH(OEt)2 + HNF
(7 a)
II
II + HNF ..!!l....... I + EtOH
(7 b)
reaction (6) was manifestly complicated, a detailed study of the rates arid equilibria of the component reactions was launched. It was expectedthat a thorough understanding of the model system would permit designing experimental conditions for the synthesis of new types of compounds useful as propellants and high explosives. The system acetal
+ nitroform ~
trinitro ether
+ alcohol
The acetal-nitroform reaction. Anhydrous dioxane was selected as the "inert" solvent for most of the quantitative work, although carbon tetrachloride re3
For simplicity, nitroforrn (trinitromethane, (02N)SCH) is abbreviated HNF; the trinitromethyl group is then NF. 198
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
ceived some attention.' The reactions were followed by titration for nitroform. It was essential to work under rigorously anhydrous conditions. The reaction (7b) is first-order in HNF and in II. For various initial concentrations, plots of initial rate Vo vs, [HNF]o for various [Il], (Fig. 1) and of vo![II]o vs. [HNFl o(Fig. 2) (where subscript zero represents initial concentration) 20,-----------------,
15
"2 ...E
~ :s"
,0
8
6
" o
~
_ ___'_ _ 0'05
_.J._ __ ' __ ___'__ ____L__ _- '
0"0
0"5
0·20
0·30
(HNFlo m FIG. 1. Reaction of nitroform (HNF) with diethyl butyral (Ae) in dioxane, 50°. Initial rate vs. initial nitroform concentration at three different acetal initial concentrations
were linear. The values of the bimolecular rate constant k~ obtained from Fig. 2 are 3·6 x 10-2 and 1·4 x 1O-21/mole!min at 50·0 and 40'0°, respectively. Numerical calculation from the integrated form of the bimolecular equation, in which the reaction is regarded as irreversible but only the first portion of the data before curvature sets in is used, yielded values of 102k~ varying from 3·78 to 4,50, average 4·17, at 50°, and 1·51 and 1·65 at 40°. The Arrhenius activation energy was estimated to be about 19 kcal/mole, 4
Alcohols and acids are obviously unsuitable, and hydrocarbons are insufficiently polar. Probably other ethers would work just as well. More basic solvents like tetrahydrofuran, dimethylformamide, and dimethyl sulfoxide would be of considerable interest. 199
C.S.RoNDESTVEDT, JR., M.STILES
and
A.L.KRIEGER
A brief study of the reaction (7 b) in carbon tetrachloride at 50" showed that the order was less than 1 with respect to II and greater than 1 with respect to nitroform. The initial rates for comparable initial concentrations were slower in carbon tetrachloride than in dioxane. An attempt to treat (7b) as a second-order reversible reaction by means of the integrated form of Eq. (8) yielded poor fits. The plots showed curvature for
8,-------------.,..-----, 050° C ., 40° C 7
6
~
5
o
Q; u
e
"'-.4 ~
'"Q
3
2
Slope: k =I-4. 10-~L/mole/min
o ! : - - _ - - I_ _- l . ._ _---'-_ _-L._ _--l.-_----l
0·05
0'10
0'20
0-15
0'25
0·30
(HNF)o
FIG.
2. Reaction of nitroform (HNF) with diethyl butyral (acetal) in dioxane, 50° and 40°. Initial rate/acetal initial vs. initial nitroform concentrations
the first 15-20 % of reaction, but thereafter became straight lines whose slopes gave an average value of k~ = 3·53 x lO-21jmole/min at 50°. - d [HNF]/d t = k~ [HNF] [II] - k~l [I] [EtOH].
(8)
The ethanolysis of I. - Reaction (7 a) is cleanly first-order in 1. When the integrated form of Eq. (8) was plotted for different runs, excellent straight lines were obtained. Surprisingly, the values of from the slopes varied considerably for different values of [EtOH]o. The order with respect to ethanol is unity when the ratio [EtOH]/[I] < 0·7. Under these conditions, k~l (50°, dioxane) = 5·5 x 10-3 l/rnole/min. When the ratio is increased, the order with respect to ethanol decreases and approaches zero with ratios greater than about 1·8. The reaction (7a) in carbon tetrachloride was too slow to measure at 50°, and at 75° decomposition occurred.
r,
200
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
The equilibrium constant. - Eleven determinations in both directions at 50° in dioxane yielded K = [I] [EtOH]/[II] [HNF] = 7·35 ± 0·27 "'" k~/k~l' The magnitude of this constant is such that I in excess ethanol would be converted almost completely to II and nitroform. Excess alcohol is certainly not the solvent of choice for reaction (1) in preparative use. At 40° in dioxane, K = 5·2 and 7·7 for different initial concentrations; the poor agreement suggests that a side reaction, perhaps that discussed in the next section is more important at 40° than at 50°. In carbon tetrachloride at 50°, the value of K is 12'0, more favorable for the production of 1. However, the equilibrium is attained much more slowly than in the more polar dioxane. Decomposition of I in dioxane. -Pure I in dioxane at 50° dissociated to the extent of 2-5 mole % in two days, as shown by titration of the liberated nitroform. The decomposition then stopped, indicating that an equilibrium had been reached. It is supposed that the vinyl ether l-ethoxy-l-butene (III) is formed by analogy to the well-known acid-catalyzed reversible dealcoholation of acetals.?'
I
~
HNF
+ CH sCH 2CH=CHOEt
(9)
III
III was not identified positively since it was present only in small quantity. Vapor-phase chromatography was not available at the time of this work. This decomposition, whatever its exact nature, does not seriously affect the rate or equilibrium constants determined for reaction (7) at 50°, except at low ratios of [EtOH]/[I], and these were the runs which gave values for K deviating the most from the average. The salt effect. - The effect of added salts on the ethanolysis of I is shown in Table 1. It is clear that the initial rate of ethanolysis is considerably enhanced by low concentrations of salt. At the same ionic strength 0·002 M, nitroformate ion depresses the rate about 15%. The study could not be carried to higher ionic TABLE
1. EFFECT OF SALTS ON INITIAL RATE OF ETHANOLYSIS OF I (Reaction (7a)] IN DIOXANE AT 50'0°; [EtOH]o.::=: 0'12 M; [I]o = 0'18 M Benzyltrimethyl-
ammonium
Perchlorate Perchlorate Perchlorate Nitroformate a
[Salt], M
0·0000 0'0007 0'0020 0·0032" 0,0020
IDS Vo, moles/ Ijmin 0,91 1,18 1'4 ± 0·1 1'65
1'17
The solubility limit in dioxane at 50°.
4a H.Shechter and H.L.Cates, J. Org. Chem. 26,51 (1961), report the preparative addition of HNF to vinyl ethers, but do not mention that the reaction is reversible.
201
C.S.RONDESTVEDT,
JR., M.STILES and A.L.KRIEGER
strengths because of the very low salt solubilities in anhydrous dioxane. Benzyltrimethylammonium bromide and nitrate are even less soluble than the perchlorate, and the perfluorobutyrate could not be obtained pure. The salt effect was not studied in carbon tetrachloride. 17·0
15·0
"
Q
10·0
o Added ccid 5'0
o
0'010
0'020
0·030
0·040
c-oso
(Nitroforrnl, x (acetol)o
FIG. 3. Reaction of nitroform with diethyl butyral (acetal) in dioxane, 50°, with and without added p-toluenesulfonic acid
Acid catalysis. - Intuitively one expects a reaction of an acetal or an ether to be acid-catalyzed. However, neither the forward nor reverse reaction (Eq.7) was influenced by anhydrous p-toluenesulfonic acid or perfluorobutyric acid in dioxane. The rates in either direction were identical to those without added acid, within experimental error, Figure 3 shows the results with the forward reaction. Some rather erratic rate enhancement was noted in carbon tetrachloride solvent. Isopropyl alcohol. - The system with isopropyl alcohol (Eq. 10) was studied briefly. Isopropyl alcohol reacted with I in dioxane slower than did ethanol. Qualitatively the same type of mixed order dependence on [i-PrOH]o was noted as with ethanol. The initial rate of the reaction of diisopropyl butyral (Iv) in 202
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
carbon tetrachloride was about ten times faster than that of the ethyl acetal II with nitroform in the same solvent. The kinetics were more erratic than with II. CH gCH2CH2CH(Oi-Pr)2 + HNF IV ;.=
(10)
CH 3CH 2CR 2CH(Oi-Pr)NF + i-PrOH V
No valid equilibrium constant could be obtained in carbon tetrachloride. Values for the expression [V] [i-PrOH]![HNF] [IV] varied from 27·4 to 64·3 for various initial concentrations. However, the equilibrium concentration of' V in reaction (10) is higher than that of I in reaction (7) (50 CCIJ. Mechanism of the reaction (7). - The following scheme is consistent with the facts. 0
,
PrCH(OEt)2
k,
+ HNF -;---'"k PrCROEt + NFI
-1
(11)
HOEt
II k,
PrCHOEt -;---'"k PrCROEt + EtOH
I
-I
(12)
®
VI
HOEt ®
PrCHOEt ffi
k,
+ NF- +-k~ PrCR(OEt)NF
(13)
-8
I
The first step is a transfer of a proton from undissociated nitroform! to the acetal. The conjugate acid splits into ethanol and a solvated carbonium ion, which in the final step reacts with nitroforrnate anion to form trinitro ether. It is probable that all ionic species exist as ion-pairs in the slightly polar solvent dioxane, rather than free ions. The reverse reaction is an ionization to an ion-pair, doubtless solvated by ethanol. Ethanol displaces nitroformate, and the solvated carbonium ion collapses to acetal. Assuming no accumulation of either unstable positive ion, one may apply the method of Christiansen" to Eqs. (11)-(13) and obtain the espressions (14) 5
Nitroform is a rather strong acid in water, but it is apparently not significantly dissociated in dioxane. The characteristic yellow color of nitroformate anion, which is readily perceived at 10- 4 M in water, cannot be seen in dry dioxane. Some degree of proton transfer to dioxane is probable, perhaps (5+ {y-
r",
O
O:H:NF
.......... _ /
6
Hammett, Physical Organic Chemistry pp. 107-108. McGraw Hill, N. Y. (1940).
14 Times N. R.
203
C.S.RONDESTVEDT, JR., M.STILES
and
A.L.KIUEGER
and (15), where V = V + - V_. V+ = _ d[HNF] = d[I] = k Ik2ka[II] [HNF] dt dt k zka + k_Ik_ z [EtOH] + k_Ik a[NP-]
(14a)
(14b) In the reverse direction, the initial rate is given by Eq. (15). (V-)0
= =
k_Ik_zk_ a [EtOH] [I] k 2ka + k_Ik_z[EtOH] + k _Ik a[NP-]
(15)
».,«,«, [EtOH] [1] k a(k 2
+ k_1[NFD + k_Ik_z[EtOH]
Various limiting cases may now be considered. Case 1. Assume k z ~ k_1[NF-]. A. Assume further that k _1k_z [EtOH] ~ kzk a• Then (15) reduces to (16). (16)
That is, the rate of ethanolysis of I is independent of ethanol at high [EtOH]o. B. The reverse of condition A converts (15) to (17). (V-)0 = k_1:--;'C a [EtOH] [I] = k 4 [EtOH] [I] z a
(17)
That is, at low (EtOH]o> ethanolysis of I is first order in ethanol. C. For the forward reaction (7 a), the rate at any time is given by (18) and the initial rate by (19). V+ = klkzk a[II] [HNF] (18) kzk a + k_1k_ z [EtOH] (V +)0 = k l (II] [HNF] .
(19)
According to (19), the initial forward rate is bimolecular and Eq. (11) is ratedetermining. This will be true at any time if condition B applies, as it does at low [EtOH]o' Case II. Assume k 2 ~ k ., [NF -]. Then (15) reduces to (20). (V_)o
A. If k_"2 [EtOH]
~
=
k _zk_ a[EtOH] [I] k a[NP-] + k_z[EtOH]
k« [NP -], (20) reduces to (17) with k:« = k 4 • 204
(20)
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
B. If the opposite is true, (20) becomes (21), which is equivalent to (17) if [NF-] quickly attains a constant value, or if a nitroform salt is added.
(V ) -
= 0
7c k_ 2 S k s [NP-]
[EtOH] [I]
(21)
C. For the forward reaction, the rate at any time is given by (22), and initially by (23) which is of the form (19). V+ =
(V1·)0 =
k lk 2k s [II] [NHF] klks[NP-] + L 2 [EtOH]
k 1k2 cdNP-]
J,
[II] [HNP]
(22)
(23)
Either Case I or II fits the data on the reverse reaction, that the observed rate is first order in [EtOH]o when it is low, and zero order when it is high. Either case will also fit the bimolecular kinetics observed with the initial forward reaction. The limiting equation (21) was tested by four runs in dioxane at 50° with the following initial concentrations: [I] = 0·18 M; [C6H s CH2NMes NF -]0 = = 0·0020 M; [EtOH]o = 0·017-0'060 M. A plot of' V, vs. [EtOH]o was linear, as required by Eq. (21). Eq. (20) was rearranged to (24). (24) By assuming that the stationary concentration of [NP-] was the same for all values of [1]0 and [EtOH]o which had been investigated, Fig. 4 was obtained. The scatter at low [EtOH]ois a result of the fact that a small error in Vo causes a large error in its reciprocal; the lengths of the bars are estimates of the error in l/Vo' From the intercept, k.; = 1-4 X lO-s min:", and k s[NP-]lk_ 2 = 0·2. This value of L s agrees well with the limiting value of Vol[I] at low [EtOH]o, 1·2 x lO-s min -1. Mixed kinetics are to be expected when the two terms in the denominator of Eq. (20) are about equal; i.e., when [EtOH] ~ 0·2; this is the range where we observed mixed kinetics. The absence of catalysis by strong acids is surprising, since in polar solvents only the proton (specific lyonium ion catalysis) is effective in converting an acetal or hemiacetal to VI. The strong acid would reduce [NP-] by converting the ion to molecular nitroform. This would help to insure that Case I was valid. However, a drastic reduction in [NP-] would retard reaction (13), probably enough to make it rate-determining in the forward direction. Perhaps an increase in rate of production of VI is balanced by a retardation of its conversion to 1. 7 Since dioxane is a much poorer ionizing solvent than alcohols or 7
A reaction of VI with undissociated nitroform could replace Eq, (13) if acid is added, and this could well become rate-determining. VI + HNF ~ I + H-.
205
C.S.RONDESTVEDT, JR., M.STILES and A.L.KRIEGER
water, it is not unlikely that molecular nitroform may function to form VI as an ion pair. The unexpected rate-determining proton transfer in Eq. (11) which is required by the observed kinetics points out the difference in the solvents.8 Since carbon tetrachloride is an even poorer ionizing solvent than dioxane, reactions which involve ionizations should be retarded by a change to carbon tetrachloride. This effect was observed. 5·0
r--------------------.,
3·0 c:
E ~
';::,.
'"Q
H
2'0
a
5
10
15
20
25
30
FIG. 4. Reaction of 1,1,1-trinitro-2-ethoxypentane (I) with ethanol (EtOH). Plot of Eq. (24)
The mechanism written is quite similar to that accepted for the hydrolysis and other reactions of acetals. The isopropyl acetal IV should and does react faster than the ethyl acetal II with nitroform, just as. diisopropyl formal is hydrolyzed six times faster than diethyl formal." The intermediate i-PrOCHR + is stabilized somewhat more than EtOCHR+ by the greater inductive effect Time did not permit an experiment with deuteronitroform to provide further evidence for rate-determining proton transfer. 9 A.Skrabal and H.H.Eger, Z. physik. Chern. 122,349 (1926).
8
206
Kinetics and equilibrium of the nitroforrn-aldehyde-alcohol reaction
of the isopropyl group." On the other hand, isopropyl alcohol is known to form acetals from aldehydes more slowly than ethanol and with a less favorable equilibrium; this parallels the lesser reactivity of isopropyl alcohol toward 1. The positive salt effectobserved for reaction (7a) is consistent with the operation of a rate-determining ionization such as (13). The smaller rate with nitroformate ion compared to perchlorate is probably a result of a common-ion depression of the ionization. Alternative explanations, that the two salts are ionized to different extents in dioxane, or that they have different activity coefficients, cannot be excluded by our data. It is unfortunate that the salt effect could not be studied with different salts and at higher concentrations. Sulfur analogs. Hartman and Tawney showed that a mercaptan would react with an aldehyde and nitroform (Eq. 25). We observed that butyraldehyde diethyl (25) RCHO + HNF + R'SH ~ RCHNF + H 20
I SR' VII thioacetal did not react with nitroform in dioxane at 50° during three days. The mixed oxygen-sulfur acetal l-ethoxybutyl ethyl sulfide reacted with nitroform to yield a mixture of materials probably containing VII (R = Pr, R' = Et), but it could not be obtained in analytical purity. On the other hand, the cyclicmixed acetal 2-propyl-l,3-oxathiolane did not react detectably with nitroform. The equilibrium for cyclic acetals is farther to the right than for acyclic acetals. A carbon-sulfur bond is not cleaved by nitroform, at least under these mild conditions. Excess butyl mercaptan at 50° converted I to a mixture apparently containing thioacetal, I, and the trinitro thioether VII, but the reaction was far from clean. It had been previously observed by Tawney that no reaction occurred at room temperature. 2-Mercaptotetrahydropyran reacted with nitroform to liberate hydrogen sulfide, but VIII could not be isolated. The parallel experiment with 2-hydroxytetrahydrothiapyran could not be performed; water failed to add to the double bond of 2,3-dihydrothiapyran, and an alternative synthesis also failed. Semiquantitative experiments with the system butyraldehyde-butyl mercaptan-nitroform showed that the equilibrium lay further on the side of product than that for the ethanol system studied in detail. The hemiacetal as a reactant. At this juncture it appeared that the formation of I via Eq. (6) proceeded by a moderately rapid nitroform-catalyzed acetalization of the aldehyde, followed by a somewhat slower conversion of acetal to trinitraether by the path already discussed. However, since acetal formation in alcoholic solution proceeds through the intermediate resonance-stabilized 10
C.K.Ingold, Structure and Mechanism in Organic Chemistry p. 334. Cornell University Press (1953).
207
c. S. RONDESTVEDT, JR.,
M. STILES and A. L.KRIEGER
(doubtless solvated) carbonium ion VI, already proposed in Eq. (12) for the nitroformolysis of II, it seemed equally reasonable that VI formed from hemiacetal would react directly with nitroformate ion or molecular nitroform (Eqs. 13-14) as well as follow the longer route to acetal and back again. This is equivalent to asking whether VI reacts faster with ethanol or with nitroform (or its ion). The question could be answered by a knowledge of the rate of acetal formation.There is no information in the literature on the rate of acetal formation or hydrolysis in dilute solution in non-polar solvents, despite the fact that these reactions in aqueous or alcoholic solutions are beloved topics for kinetic studies. In alcoholic solution, reaction (2) (hemiacetal formation) is general acid-catalyzed, i.e, molecular or undissociated acids are active as well as the lyonium ionY The subsequent conversion of the hemiacetal to intermediate VI is specific lyonium ion-catalyzed in polar solvents." Acetal hydrolysis in water solution" and in 60 %dioxane's is specific oxonium ion-catalyzed. In hydrolysis, the rate is faster in D 20 than in H 2 0 , indicating that protonation of the acetal occurs prior to the rate-determining step." It is risky to extend these conclusions obtained in ionizing solvents to nonpolar media like dioxane. Time did not permit a thorough study of the rate of acetal formation in anhydrous dioxane, and therefore a few semi-quantitative experiments were performed. We estimated the equilibrium constant K 2 , Eq. (26), for the hydrolysis of II in dilute dioxane solution containing limited amounts of water. The aldehyde formed K 2 = [II] [H 20]/[PrCHO] [EtOH]2
(26)
was measured by oximation ;16 this method, of course, includes hemiacetal with the aldehyde. The apparent values of K~Oo ranged from 0·17 to O'30 l/mole. To correct for the amount of hemiacetal at equilibrium, rough data were obtained for the formation of hemiacetal. About 15 % of the aldehyde was converted to hemiacetal when a dioxane solution 0·2 Min butyraldehyde and 0·3 M in ethanol was allowed to stand at 22°.I' The constant K2 was then measured at the same temperature; the aldehyde concentration determined by oximation was reduced by the amount of hemiacetal present. The values of K~2° ranged from 11 For
example, see G. W. Meadows and B.B.Darwent, Canad. J. Chern. 30, 501 (1952); Trans. Farad. Soc. 48, 1015 (1952). 12 Deyrup, J. Amer. Chem. Soc. 56, 60 (1934). 13 J.N.Bronsted and Grove, J. Amer. Chem. Soc. 52,1394 (1930); J.N.Bronsted and WynneJones, Trans. Farad. Soc. 25, 59 (1929); Olson and Tong, J. Amer. Chern. Soc. 66, 1555 (1944); Grove, Ibid. 52, 1404 (1930); Mi Kilpatrick and Chase, Ibid, 53, 1732 (1931). 14 E. N. Prilezhaeva, E. S.Shapiro, and M. F. Shostakovskii, Zh. Obshch. Khim. 18,1663 (1948). 15 J.e.Hornel and J. A.V. Butler, J. Chem. Soc. 1361 (1936). 16 R. H. Buchanan, Austr. J. Appl. Sci. 2, 276 (1951). 17
The attenuation of the carbonyl peak in the infrared was measured. The temperature of the spectrometer room was maintained at 22°. At the time of this work, no equipment was available for thermostating the samples at 50°. 208
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
0·64 to 0·85 l/rnole, average 0·73l/mole. In runs at 50°, there should be less than 15 % hemiacetal, since the increased temperature would promote its dissociation into aldehyde and ethanol. It did not seem possible to study the conversion of acyclic hemiacetals to trinitroethers, since the former are unstable with respect to aldehyde and alcohol. For this reason, a stable hemiacetal was selected. 2-Hydroxytetrahydropyran (IX) exists almost entirely in the hemiacetal form, and very little of the hydroxy aldehyde is present at equilibrium." The bimolecular rate constant (Figs. 4 and 5) for the reaction of IX with nitroform (Eq. 27) at 50° in dioxane was 3·48 x 1O-2l/mole/min. The equilibrium constant measured in the forward direction was 83 ± 5, and about 48 hours was required to attain a constant nitroform titer. The reverse reaction, hydrolysis of VIII, was so slow under the same conditions that constant nitroform titers were not obtained in 10 days. It should be noted that pure Vfll in dry dioxane steadily liberated nitroform at 50°, perhaps with formation of dihydropyran by an elimination analogous to Eq. (9)for 1. For comparison, the reaction of the acetal XI with nitroform was followed under the same conditions. The initial rate was only about 2 % of that with the hemiacetal IX: accurate results were not obtained because the decomposition of VIII apparently interfered. r-.
.r-.
(O)-OH + HNF;=, IX /
A
I,
(o)ttJ;=' (O)-NF + H X
-,
I_OMe +
HNF
r-. 20
(27)
VIII
;=' X ;=' VIII + MeOH
"0/ XI
Mechanism of the aldehyde-nitroform-alcohol reaction. A reasonable path for the direct production ofthe trinitroether in a non-polar solvent involves the following steps. First, rapid formation of hemiacetal catalyzed by molecular nitroform. Second, reaction of hemiacetal with molecular nitroform to form alkoxycarbonium ion VI as an ion pair with nitroformate ion, plus water. Finally, internal reaction of the ion pair to yield product with displacement of water from the solvation sphere of the ion pair. According to this model, formation of the acetal requires the diffusion of ethanol from the dilute solution into L. E. Schniepp and H. H. Geller, J. Amer. Chem. Soc. 68, 1646 (1946), observed 95 % of the hemiacetal form in aqueous solution at room temperature. C. D. Hurd and W. H. Saunders, Ibid. 74, 5324 (1952), observed 94 % hemiacetal in 95 % ethanol or in 3: 1 dioxane-water at 25° and 93 % at 35°. They noted that equilibrium was established very slowly in dry dioxane and explained this by the necessity of having both an acid and a base, as has been shown necessary for the mutarotation of carbohydrate derivatives.P 19 E.R.Alexander, Principles of Ionic Organic Reactions pp. 159-160. Wiley, N. Y, (1950).
18
209
C.S.RONDESTVEDT, JR., M.STILES
and
A.LKRIEGER
the solvent cage surrounding the ion pair, and this appears to be relatively slow, judging from the comparative rates of reaction of IX and XI. This model is consistent with the results with mercaptans. The direct reaction of aldehyde, nitroform, and a mercaptan yields trinitrothio ester in high yields," but any thioacetal formed constitutes a reaction dead end. Trinitro alcohol. An alternative explanation must be disposed of. A priori, the equilibrium (4) might participate in reaction (1). The proposal is unattractive for the following reasons. The only stable, isolable ~-trinitroa1cohol is the formaldehyde-nitroform adduct trinitroethanol. Not only is this compound not etherified by ethanol, but also formaldehyde is unique among aldehydes in failing to give trinitroethers according to reaction (1).20 Since it was possible that a more labile trinitro alcohol might indeed participate, the reaction of butyraldehyde with nitroform according to Eq. (4) was studied briefly at 22° by infrared spectroscopy. In dioxane, the equilibrium lay well to the right, and it was attained in about half an hour; the hydroxyl band of the trinitro alcohol appeared at 3·05 ft. In carbon tetrachloride, the reaction was much slower, and the hydroxyl band was at 2·78/t. These equilibria should be established much faster at higher temperatures, and should lie further to the left. It thus appeared that a substantial portion of the aldehyde and nitroform would be present during the reaction as trinitroalcohol, the concentration of which would be much higher than that of hemiacetal. The acetal-nitroform reaction (7b) was then studied in the presence of added butyraldehyde. A marked reduction in rate would be expected if the effective concentration of nitroform were reduced by conversion to trinitroaIcohol. Within experimental error, the rate was identical to that obtained in the absence of added aldehyde. In any case, formation of trinitroether from trinitroalcohol is theoretically unattractive. First, if Eq. (28) represented the etherification, the formation of the intermediate XII would require a second mole of nitroform and should thus be second order in nitroform. Second, XII should be destabilized by the powerfully electron-attracting trinitromethyl group, while the alkoxycarbonium ion VI involved in the route through hemiacetal is stabilized. The situation is similar to that involved in the conversion of a cyanohydrin to an amino nitrile." The cyanohydrin dissociates to aldehyde and HCN, an aminohydrin is formed and converted to amino nitrile; in this way a cyanocarbonium ion is avoided. ®
RCHOHNF + HNF ::=::= RCHNF
+ H 20 + NF-
R'OH--+
XII
RCH(OR')NF + H+ (28)
Cyclic mechanism. An alternative cyclic mechanism obviates the necessity of forming ions in the non-polar solvent. If the aci tautomer of nitroform were the active species, the entire reaction from hemiacetal to product could proceed
20 P.Hartman and P. O.Tawney, unpublished results. 21
T.O.Stewart and C.H.Li, J. Amer. Chem. Soc. 60,2782 (1938).
210
Kinetics and equilibrium of the nitroform-aldehyde-alcohol reaction
by cyclic electron shifts within a transition state like XIII, in which the two nonparticipating nitro groups are not shown. With respect to the acetal-nitroform reaction (7b), formation and dissociation of XIII b would obey the observed second-order kinetics and should be insensitive to acid catalysis. A negligible salt effect would be expected, and this could be tested especially with solvents capable of dissolving larger quantities of salts than does dioxane. For the ethanol-trinitroether reaction (7a), formation of XIII b requires second-order kinetics, but this is observed only at low concentrations of ethanol. At higher ethanol concentrations, the ionization process [reverse of (13)] must be dominant. In the presence of salts, the observed acceleration could reflect the transformation of the cyclic mechanism into the ionization mechanism. R'O' OEI
"H,
"-"0
RCH' ("iiI-o /~7 .....c"
XIII a, R' XIII b, R'
= =
H Et
Swain has observed "polyfunctional" catalysis in the mutarotation of tetramethylglucose, and has pointed out that one advantage to a cyclic mechanism is that the substrate need never acquire an electric charge.P If this is true in water, an excellentionizing solvent, it should befar more true in poorly ionizing dioxane. The cyclic mechanism can be tested experimentally. It would first be desirable to find more versatile solvents for kinetic studies, which will dissolve salts, and in which a variety of different alcohol-aldehyde combinations could be examined without erratic kinetics.<1 The rate of enolization of nitroform in these solvents must be determined, for it is possible that enolization might be rate-determining in at least some solvents. The rates of hemiacetal and acetal formation and their catalysis by acids in non-polar solvents, and the rate of formation of trinitroalcohol, should be determined for comparison. With these data in hand, a kinetic study of the alcohol-aldehyde-nitroform system should yield meaningful results. It is probable then that the reaction can be extended to other carbon acids such as tricyanomethane, tris-alkanesulfonylmethanes, and the like. EXPERIMENTAL Materials. Dioxane was purified by refluxing with hydrochloric acid and then distilled and refluxed with sodium until no further hydrogen was evolved. It was then transferred to an apparatus consisting of a boiling flask surmounted by a reflux condenser and a takeoff built so that al1 the inner surfaces were continually washed by the hot liquid. The dioxane was stored in this apparatus over calcium hydride, then boiled for 1hr immediately before with22
C.G.Swain and J.F.Brown, J. Amer. Chern. Soc. 74, 2534 (1952). 211
C.S.RONDESTVEDT, JR., M.STILES and A.L.KRIEGER drawal for use. The apparatus was protected from moisture by a drying tower filled with phosphorus pentoxide. Ethanol was obtained by distilling commercial absolute ethanol from magnesium ethoxide immediately before use. n-Butyraldehyde was Eastman White Label Grade distilled in a nitrogen atmosphere and protected from the air. For the preparation of 1,1,I-trinitro-2-etlzoxypentane (I), the route via the acetal II was most efficient. Nitroform was mixed with a slight excess of diethyl butyral and kept at 70-75° for 1hr. The reaction was initially exothermic. The pressure was reduced somewhat to permit distillation of the ethanol, and finally the product was collected at 69° (0'8 mm), n~5 1'4424,87% yield. (Found: C, 33'51; H, 5'17; N, 16'58; Calc. for C7HlS07Ns: C, 33'47; H, 5'22; N, 16'73 %). 1,1,I-Trinitro-2-isopropoxypentane (V). An equimolar mixture of diisopropyl butyral and nitroforrn was maintained at 50° (100mm) for 6hr in a distilling flask, and then distilled. A light yellow liquid was obtained, b.p. 70° (0'3 mm) 66 % yield. It was washed twice with ice water and redistilled for analysis. (Found: C, 35'77;H, 5'71; N,16'46; Calc. forCBH!i,07Ns: C, 36'23; H, 5'70; N, 15'85 %). 'l-Trtnitromethyltetrahydropyran (VIII). Equimolar quantities of nitroform and 2-methoxytetrahydropyran (XI) were heated at 60-80° for 3hr. Distillation yielded 85 % of material, b.p, 78-80° (0·3mm). Redistilled material boiled at 70'0° (0'2mm), 1l~5 1'4705. Material of identical physical constants and infrared spectrum was prepared similarly from 2-hydroxytetrahydropyran (IX).2s ~ Although no explosions were encountered with any of the compounds described in this paper, prudence dictates that all experiments be conducted with appropriate precautions against explosions. Rate measurements. The reaction vessel was a large test tube fitted with a neck to hold a rubber serum stopple, a ground joint for attaching to the dioxane still, and a side arm protected by a Drierite tube. The vessel was flamed out in a stream of dry nitrogen. After the delivery tip of the dioxane still had been flushed with several portions of hot dioxane, the vessel was attached and the required quantity of dioxane was distilled in. A weighed quantity of one reactant was added, the side arm was sealed off in a flame, and the vessel was equilibrated in the thermostat. The second component was added from a weighed syringe. Aliquots for titration were removed by forcing 5ml dry nitrogen into the flask from a dry syringe and then withdrawing 5ml solution. The aliquots were run into 25ml carbon dioxide-free water to precipitate insolubles, and the nitroform was titrated rapidly with standard alkali to a phenolphthalein endpoint. The endpoint was stable for at least Zrnin. Pure I in aqueous dioxane corresponding to the concentration at the endpoint did not liberate detectable quantities of nitroform during the time required for the titrations. 23
Shechter and Catesv" prepared VIII from HNF and dihydropyran in 81% yield, b.p. 103° (1.0 mm), ll~o 1-4708.
212