process is well illustrated by the variety of systems that have been successfully analyzed for monomer content. The factors which govern the choice of solvent to be used are the solubility of the latex, the retention times of the monomers expected, and the desirability of analyzing a nonaqueous system. Usually, with a new type of latex, these considerations 1% ould be evaluated in the above order, although the second and third factors are necessarily intertwined. The flexibility of the chromatographic system itself has been sufficient so that, in the work thus far required, only the three solvents mentioned have been necessary. If other solvents are needed in the future, the likelihood of small amounts of impurities in the solvent which can seriously impair the effectiveness of the chromatogram must always be considered. The use of substrates other than polyglycol would certainly broaden the range of separations possible, such as separation of ethyl acrylate-methyl methacrylate or acrylonitrile-ethyl acrylate. We have noticed, though,
that acetic anhydride produces a number of spurious peaks when eluted from some of the more popular packing materials, notably plasticizers, rendering them unsuitable for use. The complete monomer determination in those latex systems containing organic acids-usually acrylic or methacrylic acid-as starting ingredients is not supplied by this procedure. The retention of such compounds is severe on the polyglycol packing and on most partitioning liquids. In addition, peak shape is almost always poor, raising the limit which can be detected. This particular problem is further complicated by the possibility of acid hydrolysis products from esters in the original formulation. Additional information is often available from the chromatographic scans. The presence of low-boiling alcohols attributable to ester hydrolysis is evident, in addition to which other esters of uncertain origin are frequently encountered. I n many acrylate latex samples with a discernible sweetish odor, the chromatogram can confirm the
presence of a n alkyl propionate as the source. The amount required to give a very noticeable odor is less than 0.1%. ACKNOWLEDGMENT
The authors are indebted to W. M. Weddell for many ideas on the solution of samples and to H. B. Prindle for invaluable aid and cooperation in sample preparations. LITERATURE CITED
(1) Critchfield, F. E., Funk, G. L., Johnson, J. B., ANAL.CHEX28,76 (1956). ( 2 ) Erley, D. S., Ibid., 29, 1564 (1957). (3) Lacoste, R. J . , Rosenthal, I., Schmit. tinger, C.'H., Ibid., 28, 983 (1956). (4) Nelsen, F. M., Eggertsen, F. T., Holst, J. J., Ibid., 33, 1150 (1961). (5) Rowe, R. G., Furnas, C. C., Bliss, H., Ibzd., 16, 371 (1944). (6) Tweet, O., Miller, W. K., Ibid., 35, 852 (1963). ( 7 ) Zeitler. V. A , . Dow Chemical Co.,
Freeport, Tex., private communication; 195s.
RECEIVED for review February 28, 1964. Accepted May 15, 1964. Southwest Regional Meeting, ACS, Houston, Tex., December 1963.
Formation Constants of Nickel(l1) and Zi nc (II) Complexes of Dithizone and Related Compounds K. S. MATH,' QUINTUS FERNANDO, and HENRY FREISER Department of Chemistry, University o f Arizona, Tucson, Ark.
b The acid dissociation constants of diphenylcarbazone, diphenylthiocarbazone(dithizone), di-o-tolylthiocarbazone, di-p-tolylthiocarbazone, di2,4-dimethylphenylthiocarbazone, di1 -naphthylthiocarbazone, and di2-naphthylthiocarbazone, as well as the formation constants of the 1 to 1 complexes of these compounds with nickel(l1) and zinc(ll), have been determined spectrophotometrically in 50 volume aqueous dioxane at 25" C. Although in the diphenyl-, di-p-tolyl-; and di-2-naphthylthiocarbazones the difference in the formation constants is that which would b e expected with ligands having nitrogen and sulfur as donor atoms, with the remaining thiocarbazones the formation constant of the nickel(l1) complex is much higher than that of the corresponding zinc(ll) complex. These data together with the absorption spectra of the metal complexes suggest that whereas the zinc complexes of dithizone and the substituted derivatives involve nitrogen- and sulfur-metal bonds, the nickel complexes involve only nitrogen-metal bonds.
yo
1762
ANALYTICAL CHEMISTRY
D
[2-phenyl hydrazide of (phenylazo) thioformic acid, dithizone] has been widely used since 1925 (10) as a metal extractant which forms highly colored metal complexes that provide the basis for many spectrophotometric determinations. I n spite of the continuing interest in this reagent, the formation constants of the metal dithizonates necessary for a thorough understanding of its analytical behavior are available only for the copper(I1) ( I S ) and zinc(I1) (16) complexes. ;idditional formation constant data are difficult to obtain because the exceptionally low solubilities of the metal dithizonates preclude the use of the Bjerrum potentiometric technique, and both dithizone and its metal chelates are susceptible to oxidation in aqueous or mixed-solvent media, making the spectrophotometric determination of formation constants more difficult. Even though these difficulties interfere with the determination of the formation constants of many metal dithizonates, it is useful to obtain such values for the nickel(I1) and zinc(I1) complexes. A comparative study of the formation IPHENYLTHIOCARBAZONE
constants of nickel and zinc complexes of a chelating agent has been found useful in evaluating various structural factors that influence chelate stability (1%'). This has been particularly true in dealing R ith sulfur-containing ligands ( 6 ) ,where in contrast to the complexes formed by ligands containing either two nitrogen donor atoms or one nitrogen and one oxygen donor atom, the zinc complex is either almost as stable as or more stable than the corresponding nickel complex. In addition, the comparison of the nickel and zinc formation constants has been shown to be a useful frame of reference within which the influence of steric factors can be measured ( 9 ) . Of course, if ligands have sufficient complexity to provide alternative sites of reaction, it is imperative that both nickel and zinc be shown to react with the same donor atoms, in order to make
'On leave from the Department of Chemistry, Karnatak University, Dharwar, India.
the comparison meaningful. Dithizone is an example of such a ligand.
1-21
-
EXPERIMENTAL
Reagents. Dithizone, obtained f r o m the Fisher Scientific Co., was purified by the method of Cowling and Miller (6). Diphenylcarbazone was also obtained from the Fisher Scientific Co. and purified by the method suggested by Krumholz and Krumholz (20). The di-o-tolylthiocarbazone, dip-tolylthiocarbazone, di-l-naphthylthiocarbazone, di-2-naphthylthiocarbazone, and di-2,4-dimethylphenylthiocarbazone were synthesized by methods that have been described (16). These compounds were purified by precipitating them with ethanol from solutions in chloroform which was freshly distilled and treated with hydroxylamine hydrochloride ( I ) . 'I'he absorption spectra of these compounds were compared with those previously reported (23, 24) (Figure 1). All compounds used as buffer components were of reagent grade purity and included XaC1O4, C1CH2COOH, HCOOH, CH,COOH, SaH2P04,and NaOH. SaCIOl was added to each of the buffer solutions to adjust the ionic strength to a convenient value. A11 buffer solutions below pH 7 were freed from trace metal impurities by extraction with a Cc14 solution of dithizone. Above p H 7 , SaCIOl and S a H 2 P 0 4mere used in the preparation of buffers, and were treated in the same manner. Approsimately O.lM ,solutions of the perchlorates of S i ( I 1 ) and Zn(I1) (obtained from the G. F. Smith Chemical Co.) were prepared and standardized bv E D T A titrations. All solutions were made with deionized water and freshly distilled diosane, purified as described previously (11). Apparatus. A Beckinan Model DB spectrophotometer and a Cary Model 14 spectrophotometer were used for absorbance measurements in the visible region. X Beckman I R 4 spectrophotometer was use'd for obtaining t h e infrared spectra in the region 3500 to 700 em-'. A Beckman Model G p H meter with a glass-saturated calomel electrode pair was standardized with llleckman buffers a t pH 4.00 and 7.00 and used for all p H measurements. Appropriate corrections were made for converting p H meter readings in 50yo aqueous dioxane to hydrogen ion concentrations (29). Determination of Ac.id Dissociation Constants. An aliquot of the reagent in pure dioxane was pipetted into a solution containing the buffer and dioxane, mixed in such a manner as to give an ionic strength of 0 . 1 0 X and a solvent composition of 50 volume % aqueous dioxane in the final solution. The absorbances of these solutions were measured as soon after mixing as possible, and a t appropriate time intervals to permit the extrapolation of absorbance values to zero time (time of mixing). This extrapolation procedure was necessitated by the small
0.8-
400
500
700
600
mp
Figure 1 . Absorption spectra of thiocarbazones in carbon tetrachloride 1. 2. 3. 4. 5. 6.
3.14 3.93 2.21 2.76 1.87 1.20
X X X X X 1 0-6 X
M M M M M M
changes in absorbance resulting from the oxidation of the reagent. It was ascertained that the reagent in pure dioxane did not undergo such changes within the time required for the entire series of measurements. Absorbance measurements were carried out a t wavelengths correspondof the reagent in the 600ing to A,, to 680-mp region, since in this region oxidation products have no appreciable absorption. The values of the extrapolated absorbances were plotted against the p H values of the solutions to give characteristic sigmoid-shaped curves from which the acid dissociation constants were obtained graphically. A plot of p H against log
EXLTL d (-4 - E- L -T ~ .
gave straight lines of unit slope with intercepts equal to pK,. ( E X L and E L are the molar absorbances of the neutral and anionic form of the reagent, respectively, A the measured absorbance,
di-1-naphthyl di-2-naphthyl di-2.4-dimethylphenyl di-p-tolyl diphenyl di-o-tolyl
and T L the total reagent concentration). The results obtained are shown in Table I. Determination of Formation Constants of 1 to 1 Metal Complexes. T h e solutions were mixed in the manner described for the determination of acid dissociation constants, except t h a t a n amount of metal perchlorate solution sufficient t o give a 100 to 1 ratio of metal-reagent (to avoid formation of the 1 to 2 complex) was added before the aliquot of reagent. T h e presence of metal ion increased the sensitivity of the reagent to oxidation, requiring the ube of the extrapolation procedure in every case. Absorbance measurenients were carried out a t wavelengths in the values of the metal neighborhood of A,, complexes (Table I). The variation of these absorbance values with p H provided the basis for the calculation of
Table I. Acid Dissociation Constants of Dithizone and Its Derivatives and Formation Constants of Their 1 to 1 Metal Chelates in 50 Volume 7 0 Aqueous Dioxane
(Ionic strength 0 . 1 ) DiphenylDi-2,4Di-2Di-lthioDi-oDi-pdimethyl naphthyl naphthyl carba- tolylthio- tolylthio- phenylthio- thiocar- thiocar- Diphenylzone carbazone carbazone carbazone bazone bazone carbazone
ki Aa a
5.80 2.5
6.23 3.8
6.40 3.4
6.87 5.2
5.76 2.2
5.23
9.26
, . .
, . .
6.18 500
4.50 500
6.45 510
4.80 510
5.91 500
4.21 520
5.76 530
5.83 660
5.90 700
6.60 690
6.40 700
6.14 690
5.38 750
6.02 530
Wavelength in mfi used for absorbance measurements.
VOL. 36, N O . 9, AUGUST 1 9 6 4
1763
1.2
o.a ui
2 0.4
0
1.21
400
1764
ANALYTICAL CHEMISTRY
500
600
700
A11 the terms in Equation 2 have the same significance as in Equation I, except A , which represents the observed absorbance at 610 mp, where only the neutral form of the reagent absorbs. The formation constants were obtained from straight-line plots as described above. Since in the case of the nickel dithizonates there was considerable averlap of the reagent and the chelate absorption bands, the formation constants of the nickel complexes a t varying metal-ligand ratios could not be studied. When the metal-ligand ratio is unity or less than unity, both 1 to 1 and 1 to 2 complexes are present and the
formation constants from Equation 1,
where k l is the formation constant of the 1 to 1 complex, K , is the previously determined acid dissociation constant, T L and T , are the total reagent and metal ion concentration, respectively, [H'] is the measured hydrogen ion concentration, d is the observed absorbance value corrected by extrapolation to zero time, and C H I , and E . ~ Lare the limiting values of the molar absorbances of the neutral reagent and metal complex, respectively, a t the wavelengths employed. Implicit in Equation 1 are
4 Figure 2.
Absorption spectra of thiocarbazones and their chelates rM/rL
A. B.
>
filtration of the solution. T h e chloroform solution was concentrated a n d the metal dithizonate precipitated b y the addition of CC1,. T h e precipitate was filtered, washed with CCl,, and dried in vacuo over P205.T h e infrared absorption spectra of the solid metal chelates as well as the reagents were obtained by the potassium bromide disk technique. RESULTS AND DISCUSSION
The acid dissociation constants of dithizone and related compounds are shown in Table I. The pK, value of dithizone in dioxane-water is 1.35 units higher than that observed in water [4.46 (S), 4.55 ( I S ) ] , which agrees well with such differences observed for many neutral acids. Many workers have attempted to evaluate the position of the thione-thiol tautomeric equilibrium in dithizone solutions:
loo
Diphenyl Di-o-tolyl C. Di-p-tolyl D. Di-2,4-dimethylphenyI E. Di-2-naphthyl F. Di-1 -naphthyl
rL x
A.
I. 2. 3. 4. 5.
rL x I. 2. 3. 4.
Nickel chelate Neutral reagent iIinc chelate Keagent anion
IO-EM
Thiol
5.26 2.63 2.63 2.63 2.63
Nickel chelate Neutral reagent Copper chelate i!inc chelate Reagent onion
B 3.73 1 .86 1.86 1.86
the assumptions that only a 1 to 1 complex is formed, that no appreciable quantity of uncomplexed reagent is dissociated in the pH range used in the measurements, and that the uncomplexed metal ion concentration [Jf+2] = T.,. Here, as in the determination of acid dissociation constants, the data were treated graphically, and straightline plots whose slopes are unity were used to obtain formation constants. Values of the formation constants are shown in Table I. Values of the formation constants of zinc dithizonates were determined a t metal-ligand ratios less than 100 to 1 (Table 11). At low ratios of metal to ligand, the approximation [A11+2] =Tw is not valid, and Equation 2 was used to calculate the formation constants of zinc dithizonates in solutions which contained a metal-ligand ratio of a t least 4 to 1.
C 6.3 3.8 3.8 3.8
e N = N I
IO-EM
D 2.56 2.56 2.56 2.56
E 6.18 3.71 3.71 3.71
Thione
F 4.11 4.11 4.11 4.11
stepwise formation constants, kl and ICz, were obtained from a plot of the formation function il us. p L a t half integral values of il.
A
The peak heights of two absorption bands in the visible have been used as a measure of the relative concentrations of the two forms and, therefore, of K T (21, 22). Values of K T obtained in this manner (Table I), while only semiquantitative, indicate that the value of K T , approximately 2 for dithizone, does not increase by more than a factor of 2 to 3 for the series of compounds studied here. Kothrol
=
Kaobsd
(3)
I n Equations 3 and 4 all the terms have the same significance as in Equation 1. pL is equal to the negative logarithm of the concentration of the anionic form of dithizone and d is again the measured absorbance at 610 mp, where only the neutral form of the reagent absorbs. Preparation of M e t a l Complexes. A slight excess of the reagent in dioxane was added t o the metal perchlorate solution in water. T h e precipitate obtained was filtered, washed with water, and dissolved in chloroform. Traces of water in the chloroform solution were removed by absorption on a fluted filter paper during
(1
+ KT)
(2)
Table II. Formation Constants of Zinc Dithizonates a t Various Metal-Ligand Ratios in 50 Volume Aqueous Dioxane (Ionic strength 0.1)
70
Ligand concn.,
Metal ion concn., M X 10' M X l o 6 2.28 2.28 2.50 2.34 2.02 2.34 2.60 1.64
VOL. 36,
0.20 1.00 5.00 20.0 50.0 98.6 250 1000
log ki 7.66 7.60 7.20 6.85 6.64 6.54 6.23 6.18
NO. 9, AUGUST 1 9 6 4
log k,k, 14.7 14.4 14.2 , ,
,
, , .
. ,, ... , . .
1765
The observed pK, values of the dithizones are related to the pK, values of the thiol and thione forms by Equations 2, 3, and 4. From Equation 2 it may be seen that the pKathiolvalues are from 0.1 to 0.2 unit lower than the values of pKaobsd, In contrast, the pKnthionevalues may range from 0.5 to 0.6 unit lower than the corresponding pKoobedvalues. I n view of the uncertainty in K T ] the pKothio, values are more reliable than those of pKathione. Since the difference between pKoobsdand pKothiol is small, t’he former is used in the discussion of the effect of substitution in the dithizone molecule on the dissociation of the thiol forms. The acid dissociation constant of dithizone is somewhat larger than those of mercapto compounds, but this might be expected from the electron-withdrawing character of the azo and hydrazo substituents. There is a parallel increase in acid strength observed with diphenylcarbazone, when compared with (ph) enols. Methyl substitution in either the ort’ho or the para position gives rise to the expected increase in pK,. The pK, of the 2,4dimethyl derivative coincides with that which would be predicted on the basis of the additivity of the effects of the separate methyl groups on the pK, of dithizone. The acid dissociation constants for both naphthyl derivatives are somewhat less than that of dithizone. Similar behavior is observed with 1- and 2-naphthol and phenol (19). The formation constants for the metal complexes of dithizone and its derivatives show remarkable differences in substituent influences on the stability of the zinc and the corresponding nickel complexes. I n a number of ligands containing sulfur as a donor atom, the zinc complex is a t least as stable as the nickel complex (5,9, 1 2 ) . I n addition, in many cases where a substituent in a ligand is capable of exerting steric hindrance to chelate formation, zinc complexes tend to be a t least as stable as those of nickel (18). Inasmuch as the dithizones studied here involve one or both of these structural features, a comparison of the nickel and zinc complex stability should be revealing. I n dithizone as well as its p-tolyl and 2-naphthyl anaologs, the relative stabilities of the zinc and nickel complexes are consistent with previous findings (la). The zinc-nickel comparison in the o-tolyl, 2,4-dimethyl, and 1naphthyl analogs, however, is anomalous, in that the nickel complex stability constant is significant’lylarger. Each of these latter groups of analogs has an ortho substituent in the phenyl ring. The only possible effect that these substituents could have. over and above what would be exhibited in the para position, is that of steric hindrance to chelate formation. Since nickel com1766
ANALYTICAL CHEMISTRY
I
I 1600
Figure 3. dithizone
I
I
I
1400
I
1200
I
crn-l
I
Infrared spectra of zinc and nickel chelates of
- - - - Zn(ll) chelate -Nilll) chelate plexes are usually more susceptible to steric hindrance than those of zinc, the reason for the anomaly might lie in the assumption that both of these metals form dithizonates having the same structure-Le., involving identical donor groups. If the structures of the zinc and nickel dithizonates are postulated as I, in which the zinc is tetrahedrally coordinated, and 11, in which the nickel is octahedrally coordinated, the observed stability order can be explained.
from their spin-paired nat,ure cannot be gainsaid. I n the infrared region also, the spectra of the zinc and the copper complexes resemble each other (Figure 3) ( 7 ) . The most noteworthy difference in the nickel dithizonate spectrum is the appearance of a strong band at 1210 cm-1. This band strongly indicates the presence of a thiocarbonyl (>C = S) group in the nickel complex. Because of the variability of the location of the thiocarbonyl band ( 2 ) , it is not possible to assign the 1210 em.-’ band unequivocally in the spectrum of the nickel complex. Again the appearance of such a band may well reflect the influence of the spin-pairing of the nickel. On the basis of this evidence, structures I and I1 seem reasonable for the zinc and nickel dithizonates, respectively. The alternative sixmembered ring structure possible for the nickel complex was rejected, since this structure would have to involve the loss of a proton from the nitrogen I n the case of the zinc complexes, the adjacent to the phenyl ring. The proton absorption spectrum in the visible region, attached to the nit,rogen adjacent to which consists of essentially one band, the > C = S group is much more labile, closely resembles that of the dithizone as shown by the existence of thiono anion. This is true also of the dithizothiol tautomerism. The apparent abnates of most metals, including Hg(I1) sence of steric hindrance in the nickel and Cu(1I) which from x-ray structure complexes of o-tolyl, 2,4-dimethyl, and determinations are known to have a I-naphthyl analogs can be understood, structure similar to I ( 3 , 1 4 ) . The since the aromatic rings in each of absorption spectrum of the nickel these compounds must lie out of the complex, on the other hand, is markedly plane of the chelate ring. From molecdifferent and includes several absorption ular models it is apparent that in the bands (Figure 3). This unique spectrum nickel complex of dithizone itself, the resembles that observed for the complex di [3-methyl-l-phenyl-5-p-tolylformazyl] phenyl ring must lie out of the plane of the chelate ring. Hence any ortho nickel(I1) (I?’), for which a sixsubstituents in the phenyl ring cannot membered chelate ring structure with add to the steric effect already present nitrogen donor atoms, has been adin the dithizone. vanced. The possibility of the inIn contrast, the zinc complexes creased complexity of the absorption exhibit a sensitivity to the presence of spectra in these nickel chelates arising
ortho substituents. A molecular model of zinc dithizonate itself indicates that the free rotation of the phenyl ring is not significantly rest’rii:ted, permitting it to lie in the plane of the chelate ring. The presence of ortho subst.ituents, however, changes this situation. The extent of the steric hindrance with one o-methyl group is such that the formation constant is lowered by approximately one hundredfold. The formation constants of the zinc complexes of the three orthosubstituted dithizones show remarkably little variation with the pK, values of these reagents. This behavior is in sharp contrast with the zinc complexes of the remaining dithizones and of the nickel complexes of all the ligands. This would seem to mean that the electron-releasing alkyl groups in the phenyl rings cannot transmit their effects t o the chelate ring in the zinc complexes, but can to the nickel chelate ring. Since in both zinc and nickel complexes that are hindered, the A aromatic ring is out of plane, t’he B ring must be bet’ter able to transmit an electronic effect to the chelate ring in the nickel complexes. This is not entirely surprising, since the B aromatic ring in the nickel complexes is separated from the chelate ring b y only one atom, whereas in the zinc complexes, the B ring is two atoms removed. The lack of agreement between the value of the formation constant for zinc observed here with that reported on the basis of a solvent extraction method (16) presented a problem. The first series of determinations was carried out a t high metal-ligand ratios (>lo0 to 1) to ensure the formation of the 1 to 1 complex only. Under these conditions, the 1 to 2 complex would be significant only in the unlikely event that kz (the second stepwise forniat’ion constant) would be from 100 to 1000 times greater than ICL. This (does not seem to be the case with zinc dit’hizonate (15) or other zinc complexes with sulfur atom donors (4, 9, 12). When further deteirminations were carried out’ a t lower zinc-ligand ratios (Table 11), the values of k l obtained appeared to depend on the metal concentration in a most unusual manner (Figure 4). The possibilit’y that this variat.ion was an artifact caused by an incorrect assumption in the calculations was carefully considerled. When such
of other complexes without substituents (25-28).
ortho
ACKNOWLEDGMENT
The authors are grateful to the U.S. Atomic Energy Commission for financial assistance. LITERATURE CITED
I
I
-1.0
I
0
I
1
I
1.0
2.0
3.0
log Tm/T
Figure 4. Effect of metal ion concentration on first stepwise formation constant of zinc dithizonate
possibilities as the formation of 2 to 1 metal-ligand or 1 to 2 or higher complexes, were used as a basis for calculation, significantly greater scatter resulted. For example, in the case of the possibility of formation of polynuclear complexes, the apparent variation of kl with metal ion concentration would be expected to show an increase with metal ion concentration rather than the decrease observed. Further, when a 1 to 2 metal-ligand complex was assumed, not only was the scatter of the data much greater, but the klk2 value calculated was also at least a thousand times smaller than that reported previously (15). Although a t this time we cannot explain the apparent variation of kl values with zinc ion concentration, there seems to be little doubt that the values obtained a t high metal-ligand ratio have experimental validity. I n order to evaluate the significance of these formation constants in dithizone reactions of analytical significance, extraction experiments in the presence of both nickel and zinc were carried out with dithizone and its o-tolyl, 1naphthyl, and 2-naphthyl analogs. The observed extraction exchange constants agree reasonably well with those predicted on the basis that the differences in kz values between nickel and zinc would be the same as those seen in the k~ values and that K D , (chelate distribution coefficient) of both complexes id the same. A similar effect was observed with Hg(II), Cu(II), and Zn(I1) complexes of ortho-substituted dithizone analogs. Extraction constants of such complexes were lower than the constants
(l)-Barnbach, K., Burkey, R. E., IND. h N G . CHEM., ANAL. E D . 14, 906-7 (1942). (2) Bellamy, L. J., “Infrared Spectra of Complex Molecules,” p. 356, Wiley, New York, 1958. (3) Bryan, R. F., Knopf, P. M., Proc. Chem. SOC. 1961,203. ( 4 ) Charles, R. G., Freiser, H., J . Am. Chem. SOC. 74, 1385 (1952). ( 5 ) Corsini, A., Fernando, Q., Freiser, H., ANAL.CHEY.35, 1424 (1963). (6) Cowling, H., Miller, E. J., IXD.ENG. CHEM., ANAL. ED. 13, 165 (1941). (7) Dyfverman, A . . Acta Chem. Scand. 17. 1609 (1963). (8) Dyrssen, D.,’Hok, B., Svensk. Kem. Tidskr. 64, 80 (1952). (9) Fernando, Q., Freiser, H., J . Am. Chem. SOC.80, 4928 (1958). (10) Fischer, H.. Wiss. V e r o f m t l . Siemens-Konzern’4, 158 (1925). (11) Freiser, H., Charles, R. G., Johnston, W. D., J . Am. Chem. SOC.74, 1383 (19.52). - _-~ (12) Freiser, H., Fernando, Q., Cheney, G. E . , J . Phys. Chem. 63, 250 (1959). (13) Geiger, R. W.. Sandell. E. B.. Anal. Chim. y4cia 8 . 19i i1953). (14) Harding, hf., J . Chem. SOC.,1958, 4136. (15) Honaker, C. B., Freiser, H., J . Phys. Chem. 66, 127 (1962). (16) Hubbard, D. M., Scott, E. W., J . A m . Chem. SOC.65, 2390 (1953). 17) Irving, H., Gill, J . B., Cross, W. R., J . Chem. SOC.,1960, 2087. 18) Johnston, W .U., Freiser, H., Anal. Chim. ilcta 11, 201 (1956). 19) Kortum, G.) Vogel, W.,Andrussow, K., “Dissociation Constants of Organic Acids in Aaueous Solution.” Butterworths. London. 1961. 20) Krurnholz, ’P., Krumholz, E., Monatsh. 70, 431 (1937). (21) Pel’kis, P. S., Dubenko, R. G., Dokl. Akad. AVauk U S S R 120. 320 (1958). (22) Pel’kis, P. S., Dubenko, R . G., C‘krain. Khim. Z . 23, 748 (19.57). (23) Takei, S., Kato, T., Technol. Repts. Tohoku C‘nz‘v.21, 319 (1957). (24) Ibid., 24, 75 (1959). (25) Ibid., 25, 39 (1961). (26) Ibid., p. 143. (27) Ibid., 26, 19 (1962). (28) Ibid.. D . 35. (29j van’citert, L. G., Haas, C. G., J . Am. Chem. SOC.75, 451 (1953). ~
RECEIVEDfor review March 16, 1964. Accepted May 18, 1964.
VOL. 36,
NO. 9, AUGUST 1 9 6 4
1767
