calculation Equation 9 is persuasive evidence, in the views of the authors, that the oxidation of some constituent in S.jbetida seed oil (or mixed methyl esters derived from S. foetida seed oil) is bimolecular, being of the first order with respect to silver nitrate and also with respect to some seed oil constituent. The confirmation of the data to Equation 1 is also considered as persuasive evidence that there are no interfering side reactions. From the fact that there is a progressive decrease in the cyclopropenoid ring (from the infrared data), [also shown by Johnson et al. (9)] it seems very evident that it is the cyclopropene ring that is reacting with silver nitrate in these experiments. It seems apparent that reliable estimations of the equivalents of cyclopropenoid fatty acid esters present in vegetable oils may be obtained through observations of the quantities of silver nitrate consumed in kinetic experiments such as outlined above. It is concluded, also, that the component fatty acid content of a seed oil is a function of the history of the seed and oil, as well as the genetics of the seed. It is well known that the seed oil of S . foetida has a marked tendency to solidify. Moreover, the data presented here indicate that the assay for the cyclopropenoid ring in the mixed methyl esters derived from S . foetida seed oil increased substantially when the solution containing the esters was passed through an alumina column (presumably polymeric material was removed). We also observed that the assay of a sample of the mixed methyl esters decreased during two years of storage at - 15 "C from 36 cyclopropenoid fatty acid ester (calculated as methyl sterculate) to about 1 % . In addition, the loss of cyclopropenoid fatty acids which we observed on the esterification of the acids by the method of Metcalfe and Schmitz ( 4 ) leads to the suggestion that the component cyclopropenoid fatty acids in seed oils, such as S . foetida or G. hirsutum would be reduced should the liberation of free fatty acids because of lypolysis occur.
No measurable quantity of silver nitrate was consumed by peanut, soybean, or corn oils. While these oils undoubtedly contain plasmalogenes, the concentrations in the oils of fatty aldehydes arising from their hydrolysis was probably too small to be measured by the technique used. In the case of cottonseed oil, however, the data indicated the presence of component cyclopropenoid fatty acids to the extent of 0.2--0.6z, calculated as malvalic acid (C18H3202).The assumption is made that the silver nitrate was consumed by the cyclopropane ring (and not by free fatty aldehydes), the argument being that it is not expected that concentrations of free fatty aldehydes in cottonseed oil will be substantially greater than those present in the other seed oils. In this connection, it is noted that aliphatic aldehydes (acetal and proponal) did not reduce silver nitrate under the conditions of these several experiments. It is not expected, therefore, that fatty aldehydes that might accidently be liberated from plasrnologens present in these oils will interfere with the analyses. It has been our observation that silver in the colloidal form, as it appears in the reaction medium, is susceptible to oxidation by the nitric acid liberated during the reaction. The reaction occurs slowly, however, and can be neglected for the shorter periods of time used in the determinations reported. Estimations of silver nitrate consumed will be low if the time of reaction is prolonged.
z
ACKNOWLEDGMENT The technical assistance of E. J. McCourtney and W. B. Carney is acknowledged.
RECEIVED for review April 3, 1972. Accepted October 5 , 1972,
Substituent Effects on Acid Dissociation Constants of N,N-Substituted Dithiocarbamic Acids K . I. Aspila, C. L. Chakrabarti,' and V. S. Sastri Department of Chemistry, Carleton Uniacrsity, Ottawa, Ontario KIS 5B6, Canada The acid dissociation constants of several N-substituted dithiocarbamic acids have been determined by three different methods, and the values show selfconsistency. Comparison of the experimentally obtained values for Kc with the values reported in the literature show that the literature values are anomalous. pK, values for the dialkyl series show the order to be ;-propyl > n-butyl > n-propyl > ethyl > methyl, and in the cyclic derivatives the order is pyrrolidine = hexamethylene > piperidine. Explanation based on factors such as inductive effect, steric strain, and solvation of dithiocarbamate has been advanced to rationalize the changes in pK values of N-substituted dithiocarbamic acids.
RECENTSTUDIES in this laboratory have provided important information on many aspects such as the decomposition mechanism (1-3) and monobasic or dibasic character ( 4 ) of To whom all correspondence should be addressed. ___ (1) K . I. Aspila, V. S. Sastri, and C. L. Chakrabarti, Tulunra, 16, 1099 (1969).
dithiocarbamic acids. However, there exist some disparities in the literature values (5) of acid dissociation constants of substituted dithiocarbamic acids. Further work (6) on the acid dissociation constant of diethyldithiocarbamic acid yielded consistent values for K , for this particular reagent. The present work has been aimed at determination of the acid dissociation constants of a variety of N-substituted dithiocarbamic acids with the hope of determining the effect of the substituent at the nitrogen center on the value of dissociation constant, a comparison between these values and the literature values, the effect of ionic strength (p) on K, values, and (2) S. J. Joris, K. I. Aspila, and C. L. Chakrabarti, ANAL.CHEM., 42, 647 (1970). (3) K. I. Aspila, S. J. Joris, and C . L. Chakrabarti, ibid., 43, 1529 (1971). (4) S. J. Joris, K. I. Aspila, and C. L. Chakrabarti, ibid., 41, 1441 (1969). (5) A. Hulanicki, Talanta, 14, 1371 (1967).
( 6 ) K. I. Aspila, S. J. Joris, and C. L. Chakrabarti, J. Phys. Chem., 74, 3625 (1970).
ANALYTICAL CHEMISTRY, VOL. 45, NO. 2, FEBRUARY 1973
363
“74 =
.06-
.04
-0-
“
I
= =
i-PrZ
I I
n - Pr2
.,
-
D
1.0
1.2
c
-
,021
I
0.2
0
0.4
IONIC
0.6
0.8
I
STRENGTH ( p )
Figure 1. Primary salt effect on limiting rate constant, klim (sec-l), for the decomposition of several DTC acids Temp. = 25.0 f 0.1 “C, pH E 1.0,p is adjusted with KCI A MezDTCH EtzDTCH o n-PrZDTCH 0 i-PrZDTCH O n-ButzDTCH also obtaining K , values by extrapolation to zero ionic strength. Dithjocarbamic acids decompose according to the mechanism (7)
‘R
R
Sl-
S
Table I. Primary Salt Effect on Decomposition of Several DTC Acids, Measured between an Ionic Strength 0 and 1.0 Using KCI and LiCl as Electrolytes pH = 1.0, Temp = 25.0 =k 0.1 “C Increase in stability Using KCl as Using LiCl as DTCH electrolyte electrolyte MerDTCH 7.6 18.6 EtrDTCH 9.2 19.7 n-PriDTCH 11.5 ... i-PriDTCH 0.0 ... n-ButrDTCH 64 150 PyrrDTCHb 11.2 ... 5 increase in DTC acid stability has been determined between p = 0 (extrapolated) and p = 1.0. PyrrDTCH stands for cyclic N,N-tetramethylenedithiocarbarnic acid.
R
give the magnitude of primary and secondary salt effects, respectively, on the rate constants. K , can be evaluated from the rate constant, and is related to K, through Equation 3, K, =
[H+][DTC-] [DTCH]
=
YDTCH
K,
YH +YDTC-
(3)
where K, = acid dissociation constant based on activity, y are molar activity coefficients, and [ ] stands for molar concentrations. Equation 3 can be written as
(4) In order to use Equation 4 to evaluate K,, pH meter readings of H+ ion activities have to be converted into H+ ion concentrations by Equation 5 (8) ~
R
\ NH1+
R
= *
c
rrn*.do(mlc>
(5)
where -ye* = mean molar activity coefficient; yrn+= mean molal activity coefficient; m = molality of HC1; do = density of pure water; c = molar concentration of HC1. The ratio mjc can be evaluated by an approximation (9), such as Equation 6.
/
and the pseudo-first order rate equation is
R =
kdH+l [DTCIt = k,,,[DTC], (1) [H+I Kc where R = rate of decomposition; k z = limiting rate constant; K , = acid dissociation constant based on concentration; [DTCI1 = total dithiocarbamate concentration; [H+] = hydrogen ion concentration; k,,, = apparent rate constant. At pH lower than 3, Equation 1 becomes k z = kapp, since k , is negligible in comparison to H+ ion concentration. At pH 5 or above, Equation 1 becomes
+
where d = density of solution at ionic strength p (obtained from reference IO); w b = molecular weight of HC1; c = molar concentration of HC1. Combination of Equations 5 and 6 leads to ~
c
= *
Yrn*.(do/d)
(7)
y c + can
be used to calculate actual hydrogen ion concentrations at different ionic strengths.
where aH + = activity of hydrogen ion. since [H+]
