The reactions of alkoxides with different alkyl halides were measured in the same bottles as used for the free base titrations. Ten ml of the alkyl halide solution were syringed into a capped bottle followed by 1 ml of an alkoxide solution. The mixtures were allowed to react for 1, 3, and 5 minutes. Each bottle was uncapped, ethyl alcohol and indicator were added, and the solution was titrated with standard 0.1N benzoic acid in toluene.
Table 111. Determination of n-BuNa in Presence of Different Alkoxides Carbon sodium Theoretical Found Molar ratio 0.21 0.21 1 n-BuNa :1 t-BuOLi 0.50 0.48 1 n-BuNa:Z t-BuOK 1.00 1.02 1 n-BuNa:Z t-BuOK 0.35 0.35 1 Allylsodium :1 i-PrONa
RESULTS AND DISCUSSION
base, The extent of reaction was measured after one, three. and five minutes. All four alkyl halides reacted with potassium tert-butoxide. However allyl bromide reacted the least, and only negligibly after one minute. None of the other alkoxides showed any significant reaction with the alkyl halides even after five minutes. The carbon-bound sodium content of butylsodium compounds containing different alkoxides was determined using the double titration procedure. Allyl bromide was used to react with the organometallic compound for one minute. The results in Table III demonstrate that accurate assays can be obtained on organometallics in the presence of large amounts of alkoxides. Even in the two mixtures containing potassium tert-butoxide in twice the concentration as butylsodium, accurate determination of the carbon-bound sodium was obtained.
A variety of carbon-bound metal compounds have been accurately analyzed using the double titration outlined in the experimental section. Included are n-butyllithium, s-butyllithium, t-butyllithium, phenyllithium, and n-butylsodium. Table I shows total and free base values obtained on these compounds. Also included in the Table is the manufacturer's assay of the carbon-bound metal concentration. In all cases the carbon-bound metal content was greater than 90 %. Four different alkyl halides were studied in the free base titration. Allyl bromide was found to be best of the four. Gilman has previously done a very thorough investigation of a wide variety of halides with regard to their effectivenessin reacting with different organometallics. Gilman found allyl bromide and 1,2-dibromoethane to be the most effective. Table I shows that higher free base values were normally obtained with 1,3-dibromopropane. This indicated incomplete reaction with the organometallics. The four alkyl halides listed in Table I were reacted individually with different alkoxides. The extent of the reaction is indicated in Table I1 by a decrease in the original titratable
RECEIVED for review July 6, 1970. Accepted September 8, 1970. Paper published by permission of the Firestone Tire and Rubber Company.
Complex Formation in Lead Oxalate Solutions Leon N. Klatt Department of Chemistry, University of Georgia, Athens, Ga. CONFLICTING RESULTS are found in the literature concerning the nature of complexes formed between lead and oxalate. Kolthoff, Perlich, and Weiblen reported the existence of Pb(C20&2- ( I ) . Meites, using solubility and polarographic methods, showed that in neutral solution Pb(Cz0&2- predominates, while in highly alkaline media a basic oxalate, 3 Pb(OH)2.PbC204precipitates (2). Jain, Kumar, and Gaur, based upon polarographic half-wave potential shifts, reported that lead in the presence of excess oxalate forms two complexes, presumably PbC204 and Pb(Cz04)22- (3). Rouseff, in a similar study, obtained the same result (4). In the latter two investigations, the complex or complexes that were proposed in order to explain the experimental data depended upon the data analysis procedure. In a recent article it was shown that this apparent anomaly is probably due to experimental errors in the data (5). Thus, in order to (1) I. M. Kolthoff, R. W. Perlich, and D. Weiblen, J. Phys. Chem.,
46, 561 (1942). (2) T. Meites and L. Meites, J . Amer. Chem. SOC., 73, 1161 (1951). (3) . , D. S. Jain. A. Kumar. and J. N. Gaur. J. Electroanal. Chem.. 17,201 (1968). (4) ~, R. L. Rouseff. M. S. Thesis. Southern Illinois Universitv. _
August 1968. ( 5 ) L. N. Klatt and R. L. Rouseff, ANAL.CHEM., 42,1234 (1970).
I
30601
affirm the nature of these species, the solubility of lead oxalate in the presence of excess oxalate was studied. The following is a report of this investigation. EXPERIMENTAL
Apparatus. All polarographic data were obtained with a Heath polarographic system. pH measurements were obtained with an Orion Model 801 pH meter. Temperature control was =tO.O1 "Cof the desired value. Reagents. Lead oxalate was prepared by mixing equal molar solutions of lead nitrate and sodium oxalate. The precipitate was filtered, washed with water and ethanol, and dried at 110 O C for 18 hr. Assay for lead via EDTA titration and for oxalate via KMn04 titration yielded a purity of 99.98%. All other reagents were reagent grade quality. Procedure. Solutions for the solubility studies were prepared as follows. A specified volume of 0.200M KzCz04 was transferred to a 100-ml volumetric flask, KNOa was added to adjust the ionic strength, 100-200 mg of solid PbC204 was added, and diluted to volume. The solutions were then placed in a constant temperature bath and allowed to equilibrate for 5-10 days with twice daily mixing of the flask contents. No changes in the lead content of the supernatant occurred after 3-4 days. Analysis for oxalate content in the equilibrated supernatant showed that decomposition of oxa-
ANALYTICAL CHEMISTRY, VOL. 42, NO. 14, DECEMBER 1970
* 1837
Table 111. Temperature Dependence of Complex Formation in Lead Oxalate Solutions. T “C K, x 109 PI x 1 0 - 8 pa x 10-5 15.00 1.49 =k 0.03 6.63 d~ 0.23 4.62 i 0.12 5.21 i 0.16 25.00 1.90 f 0.04 6.95 i:0.21 3.59 4 0.11 40.00 4.49 0.13 5.46 i 0.30 a Ionic strength = 0.300.
6 v)
x
5
vi-
Table IV. T ~ e r ~ o ~ y n Dataa a~ic
,
I
0 01
0
I
0 02
0 03
Equilibrium IC,
I
0 04
P1
0 05
[,2Odl
AG kcal/mole 11.9 i 0 . 1 -5.2 I-t 0.1
-7.8
P2
a
Figure I. Molar solubility of PbzC204as a function of oxalate con~entrat~o~
T = 25 “C, I
=
&
0.1
AH kcal/mole 8 . 0 i 1.8 -2.0 rt 2.0
-1.4
+ 1.0
AS cal/deg/mole
-13.1 =IC 6.4 10.9 i: 7.0 21.5 4 3.7
0.300.
T = 25.00 T, I = 0.300
Table 1. Solubility of Lead Oxalate 25 “C Excess Molar solubility, S X 105 oxalate, M I = 0.150 0.300 0.500 1.00 1.50 0 3.52 4.36 5.98 11.2 14.5 0.00050 1.34 1.51 2.10 5.36 7.0 0.00100 1.26 1.55 1.94 3.57 5.3 0.00200 1.34 1.57 1.94 3.69 5.25 0.00300 1.43 1.58 ... 5.40 0.00500 1.69 1.83 2 35 4 76 ... 0.0100 2.05 2.38 3.03 6.65 6.82 0.0150 2.51. 2.77 3.56 8.61 8.0 0.0200 2.86 3.14 4.09 10.8 ... 0.0300 3.69 4.49 5.15 14.8 ... 0.0400 4.44 5.30 6.36 18.9 15.0 0.0500 , . , 6.18 7.55 23.2 18.6 .
Table II. E lonic strength 0 0. I50 0.300 0. So0 1.000 1.500
~
K, X IOs 0.145 i.0.005 1.24 i.0.04 1.90 i:0.04 3.58 f 0.09 12.2 f 3.9 20.4 =IC 2.4
.
I
~ Constants ~ ~ at 25 ~ ”@b
p1 x
(4)
10-3
82.2 i: 3.3 9.92 =IC 0.40 6.95 =IC 0.21 5.00 f 0.17 2.12 =k 0.68 2.15 =IC 0.30
p2 x
~
i
The contribution of PbN03+ to the total solubility of Bb@ 2 0 4 ( s o l i d ) was less than 2 % at the lower ionic strengths studied and was neglected. The data as a function of ionic strength are summarized in Table I. K, is obtained by squaring the molar solubility measured at zero excess oxalate. fil and PZ are then obtained from the intercept and slope of the linear segment of the solubility curve, respectively. The results at various ionic strengths are summarized in Table 11. The values at infinite dilution were obtained by calculating the respective activity coefficients ~using~ the Davies extended form of the Debye-Hiickel equation (6).
10-5
57.8 4 2.1 6.55 4 0.23 5.21 f 0.16 3.21 i 0.08 3.36 =k 1.08 1.25 rt 0.16
late had not occurred. The pH of all equilibrated solutions was between 5 . 5 and 6.5. A portion of the equilibrated supernatant was carefully withdrawn, transferred to a polarographic cell, and the diffusion current of Pb(1I) was determined. The total lead content was determined through comparison with a standard curve which was prepared at the same ionic strength and temperature as that employed in the equilibration. Replicate analysis of the supernatant yielded a precision of approximately 2 % for the total Pb(I1) content. Corrections due to differences between the diffusion current constant of Pb(I1) in mitrate and mixed nitrate-oxalate media were negligible. RESULTS AND DISCUSSION
In Figure 1 is shown the molar solubility of lead oxalate as a function of the concentration of excess oxalate. The molar solubility is described by
- log y N = 0 . 5 2 . ~ ~
1
+
11‘2
The results of a temperature study are summarized in Table 111. Based upon these values, the thermodynamic data in Table IV are obtained for Equilibria 2-4. The solubility product at infinite dilution compares moderately well with 4.8 x obtained by Kolthoff et al. ( I ) . The formation constants obtained at an ionic strength of 1.50 are in excellent agreement with those calculated from the data of Jain et al. (3), PI = 2.09 i. 0.39 X l o 3 and P2 = 1.05 i 0.05 X I O 5 compared with the respective values in Table II. On the basis of comparison of similar complexes, P(PbSO4) = 530 (7), P(NdC204+) = 1.6 X IO’ (a), P(CdC204= ) 8 X IO3(9),and p(CeC2Q4+)= 3.3 X IOe (8)suggest that ‘the values obtained for lead oxalate are reasonable. Comparison of the thermodynamic data of Table IV with other lead species, AH(PbCl+) = -4.4 ksal/mole, AH(PbBr+) = -2.9 kcal/mole, AW(PbN03+) = + O S 7 I
