Table 11. Rate Constants at 25 “C for the Reaction: MLi Dz- --t M D z + iL (or MLiDz) Nickel Zinc
+
Ligand Acetate Thiocyanate Mercaptoacetate Oxalate Tartrate
+
1 .o 2.5 14.0 1.0 1 .o
,.. -0
4 0.23
25.0 10.0 7.0 1 .o 1 .o
... -0 -0 -0 -0
cyanate, forming a stronger zinc complex than does acetate, to a lesser extent. In contrast the thiocyanate complex of nickel, probably more stable than that with acetate, reacts more rapidly than either the hydrated or monoacetate nickel ions. Again, although mercaptoacetate forms much stronger complexes with nickel and zinc than does thiocyanate, the former set react more rapidly than do the hydrated metal ions. In the cases of the 1 : 2 complexes, all of which react much more slowly with dithizone than do the hydrated metal ions, there is a possibility that in addition to the deterrent effect of the added negative charge, the coordination sphere may be (almost) completely filled by the auxiliary ligand further hindering reaction. One might speculate that because the l :2 nickel complex has observable reactivity whereas the zinc does not, perhaps in these tartrate complexes zinc has a coordination number of 4 and nickel, 6. At this point it can be seen that much useful information can be obtained about the kinetic as well as equilibrium aspects of auxiliary complex formation but that meaningful generalizations will have to await the accumulation of information about many more systems. Useful analytical application of these kinetic studies follows from the different effect these auxiliary ligands have on various metal reactions. Thus, one can further increase the already significant difference in the rates of formation of zinc and nickel dithizonates by another 25 times simply by extracting from acetate containing solutions. This further change significantly improves a kinetically based procedure for the separation of nickel and zinc (7) and indicates an approach which will surely contribute significantly to the more widespread use of differential kinetic methods of analysis.
nickel complexes, Hammes ( 6 ) found the 1 :1 Ni-glycine complex released a water molecule 13 times faster than did the aqueous NiZ+. Even the 1 :2 complex was faster (3.3 fold) than the aqueous ion. Hammes explained the observed kinetic behavior on the basis of the reduced attraction for water in the primary hydration sphere resulting from the reduction of the net positive charge on the nickel ion. Although the rate of water release can be expected to be increased by the lowered net charge on the metal ion complexed by an anionic ligand, the case of reaction with another negatively charged species should also decrease. In the reactions with dithizonate reported here, it would seem that the first factor is of greater influence than the second in the cases of ZnOAc+, ZnCNS+, NiCNS+, Zn(SCHZCOO), and Ni(SCH.2COO), that the two factors are balanced in the case of NiOAc+, as well as in the 1 :1 Ni and Zn complexes of lactate, oxalate, and tartrate, and that the second factor outweighs the first in each of the 1 :2 complexes studied (Table 11). Quite obviously, a simple electrostatic explanation of these observations is inadequate. With zinc, both acetate and thiocyanate accelerate the reaction with dithizone; the thio-
RECEIVED for review May 12, 1969. Accepted July 11, 1969. This work was supported by a grant from the National Science Foundation.
(6) G. H. Hammes and V. I. Steinfeld, J. Amer. Chem. SOC.,84, 4639 (1969).
(7) B. E. McClellan, Murray, State University, private communication, 1968.
Determination of Water in Crystalline Ammonium Perchlorate by High- Resoht ion Proton Magnetic Resonance A. G . Keenan and Robert F. Siegmund Department of Chemistry, University of Miami, Coral Gables, Fla. 33124
THEUSE OF proton magnetic resonance spectra for the quantitative determination of water is well known ( 1 ) . Detailed methods have been developed ( 2 ) to cover a wide variety of solids and liquids. The method offers a rapid and nondestructive means of analysis which compares favorably ( 3 ) with classical methods. However, high resolution NMR does not appear to have been applied successfully to the accurate determination of low water contents in crystalline solids. The present note describes a procedure for ammonium perchlorate which gives high precision and an accurately linear calibration curve at low water contents. The method depends (1) J. A. Pople, W. G. Schneider, and H. J. Bernstein, “HighResolution Nuclear Magnetic Resonance,” McGraw-Hill, New York, N. Y., 1959, p 458. (2) W. L. Rollwitz, Humidity Moisture Papers Intern. Symp., Washington, D. C., 4,149 (1963) (pub. 1965). (3) S. V. Kuznetsov, Ve.rtn. Sel‘skohor. Nauki Vses. Akad. Sel’shokhoz. Nauk, 6,117(1961). 1880
on the proton exchange between ammonium ion and water which occurs when the perchlorate is dissolved in the high dielectric nonaqueous solvent, N,N-dimethylformamide (DMF). Up to 0.5 wt % water, the precision of determination is *0.018% absolute as measured by the mean deviation of points from the calibration curve. Up to 8% water, the 2 The observed radical difference in precision is i ~ 0 . 0 5 %. behavior between ammonium perchlorate and nitrate is suggested as a basis for a method of measuring relative base strengths of anions in nonleveling solvents. Previous studies of ammonium salts by NMR have been concerned primarily with additions of small amounts of salt to an aqueous solution. It has been found (4-6) that the (4) R. A. Ogg, Jr., Discussions Faraday Scx., 17,215 (1954). (5) H. Benoit, H. Ottavi, and J. Pommer, Compt. Rend., 256, 359 (1963). (6) H. Ottavi, Ann. Phys. (Paris),1, 5 (1966).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
6ot
% HzO
Figure 2. Shift of NMR peak for NH4+ as a function of water content from 0 to 100%
1
I
1
I
2
I
3
I
4
I
5
I
6
1
7
1
8
Yo H 2 0
Figure 1. Calibration curve showing width of NMR peak for NH4+at half-height as a function of water content in the range 0 to 817,. AP: ammonium perchlorate. A N : ammonium nitrate
determined. Peak heights and peak half-widths--i.e., peak widths at half-height-were measured on the arbitrary scale of the chart paper and integrated peak area were also recorded. Chemical shifts (As) were measured to within +0.01 ppm. All data reported are averages of three spectra. RESULTS AND DISCUSSION
resonance absorption of NH4+, a triplet in acid media, collapses to a single absorption line in basic media because the chemical exchange between the ammonium ion and H 2 0 becomes so rapid as to produce an averaging effect. Early studies (7) showed that increasing the concentration of the ammonium salt, in an aqueous solution, shifted the position of the H 2 0absorption. The spectrum of equimolar amounts of NH3 and H20 showed a single resonance line intermediate between the resonance absorptions of pure H 2 0and ammonia (4). Again this was attributed to a rapid exchange in the acid-base reactions of the two components. EXPERIMENTAL
Recrystallized Fisher Certified ammonium perchlorate was dried in an oven at 85 "C for 7 days and stored in a desiccator over silica gel. Calibration solutions were prepared by adding weighed amounts of distilled water to 5-gram samples of the perchlorate and dissolving in weighed 15-1111 portions of Fisher Certified N,N-dimethylformamide. The D M F contains 0 . 0 2 z H20 according to the label. Double distillation, rejecting the first and last 25% each time, was tried but is unnecessary because it does not change the precision of the results. All weighings were carried out to 0.1 mg. The concentration of ammonium perchlorate in D M F was 26.12 + 0.046z:. The proton magnetic resonance spectra of the solutions were measured on a Perkin-Elmer Hitachi Model R-20 high resolution spectrometer, operating at 60 MHz with the Samples at the normal operating temperature of 34 "C. Tetramethylsilane (TMS) was added to all solutions as an internal standard. The spectra were first swept over 600 cps to locate all peaks and then more accurately over a 300-cps range enand D M F singlet peaks. The compassing the H20, "*+, latter was done under various sensitivities depending on the accuracy required for the range of water concentrations being (7) H. S. Gutowsky and S. Fujiaara, J. Chern. Phys., 22, 1782 (1954)
The spectrum of pure DMF at 34 OC contains two closely spaced peaks at 2.82 6 and 2.98 6 relative to TMS, due to hindered internal rotation of the methyl groups ( I ) , and a singlet at 8.05 6 due to the lone hydrogen. When water alone is added, a new peak appears at 4.49 6 and all four peaks remain unchanged in position up to water concentrations of at least 7 5 z , indicating no interaction between H20 and DMF. Similarly, when dry ammonium perchlorate is added to DMF, a peak at 7.38 appears due to the NH4+ and all peaks remain at the same position up to a concentration of at least 35 perchlorate. When, however, the H 2 0and ammonium perchlorate are present together, the NH4+ peak broadens and, with increasing water content, shifts downfield toward the pure H 2 0peak. This presumably is due to proton exchange between NH4+and H 2 0 . In principle any of three parameters, namely the peak shift, the peak broadening, or the peak area, might be used for determination of water. For ammonium perchlorate, as shown in Figure 1, the peak width at half-height provides a sensitive measure of water content with an accurately linear calibration curve from 0 4 % H 2 0 in the perchlorate. The average deviation of the points from the curve is 0 , 0 5 2 z absolute. For lower water contents, the spectra can be run at higher sensitivities. For example, in the range of water contents from 0-0.5 an accurately linear calibration curve with a mean deviation of only 0.018 absolute was obtained. At water contents appreciably higher than about lox, the peak half-width gets difficult to measure accurately because of broadening, and the best parameter to use is peak shift downfield from the D M F peak at 8.05 6. The calibration curve is not linear, but as shown in Figure 2, useful sensitivity can be obtained to about 50z H 2 0 in the perchlorate. The shift is continuous with decreasing sensitivity up to 100% or pure water. The third possible parameter, the peak area, was found not to be comparable in precision with either of the other two parameters in any composition range for this system.
z
z,
z
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
1881
A limited amount of work was also done with ammonium nitrate. As shown in Figure 1, the sensitivity of peak halfwidth to changes in water content is very much less than for the perchlorate. This is presumably due to the different base strengths of the anions in DMF. This affects the proton exchange equilibrium as shown by the equations
+ HzO %NH3 + H30+ H30+ + Nos- % HNOa + Hz0 Ha0+ + clod- % HC104 + H20 NH4+
(1)
(2) (3)
For a given water content, Reactions 2 and 3 will come to equilibrium at different concentrations which will in turn influence the exchange rate in Reaction 1. The results with ammonium nitrate show that although the method depends on proton exchange with the ammonium ion, independent calibration curves will be required if the procedure is to extend to salts with various anions.
The observed sensitivity of peak half-width to the type of anion appears also to provide a basis for developing a new and powerful method for measuring relative base strengths of anions in nonleveling solvents. Perchlorate is known to be a weaker base than nitrate. It gives a broader peak (Figure 1) which indicates a faster exchange rate in Reaction 1. This in turn results from the fact that Reaction 3 lies farther to the left than Reaction 2. Thus, either the slope or the ordinate at some convenient abscissa in Figure 1 is a parameter for measuring relative base strengths. The method can obviously be generalized to a series of ammonium salts of various anions dissolved in nonaqueous solvents of sufficiently high dielectric constant to ensure adequate solubility and dissociation.
RECEIVED for review June 9, 1969. Accepted September 8, 1969. This work was supported by the Office of Naval Research, Power Program, under Contract Nonr-4008(07).
Thiocyanate Induced Adsorption of Zin,c Ion at the Mercury Electrode George Lauer and R. A. Osteryoungl Science Center, North American Rockwell Corp., Thousand Oaks, Calg. 91360
ANIONinduced adsorption of various white metal ions at the mercury electrode has been reported in a number of recent papers ( I ) . We report here on the thiocyanate induced adsorption of zinc, The adsorptive behavior of zinc ion is distinctly different from the previously studied metals in a number of ways. Two qualitative theories rationalizing the induction effect of anions on the adsorption phenomenon have been presented ( I , 2). The data which we present here fit neither completely and appear to give credence to portions of both theories. We also report on the coadsorption behavior of zinc and cadmium. We believe that this is the first quantitative study of competitive adsorption to be reported in the literature.
COMPUTER
1
11'
1 1-L
REFERENCE
1 7
COUNTER ELECTRODE
INDICATOR ELECTRODE
I
1 AN+bOG DIGITAL CONVERTER
Figure 1. Block diagram of instrumental arrangement. A l , A2, and A3 are analog devices, Type 201, operational amplifiers
EXPERIMENTAL
Solutions were made up using triply distilled water and A. R. grade chemicals without further purification. The computer system has been described previously (3). A block diagram of the interconnections is shown in Figure 1. All the data reported here are for 25 "C maintained by circulating water from a temperature bath through a jacketed cell. Measurement of Adsorbed Species. The quantity of adsorbed species at a given potential was measured using the technique of chronocoulometry (4). A Kemula-type hanging mercury drop electrode was used throughout. This elec1 Present address, Department of Chemistry, Colorado State University, Fort Collins, Colo.
(1) F. C. Anson and D. J. Barclay, ANAL. CHEM., 40, 1791 (1968). (2) G . W. O'Dom and R. W. Murray, J. Electroanal. Chem., 16, 327 (1968). (3) G. Lauer and R. A. Osteryoung, ANAL. CHEM., 40 (lo), 30A (1968). (4) D. H. Christie, R. A . Osteryoung, and F. C. Anson, J . Electroanal. Chem., 13, 236 (1967). 1882
0
trode was modified by fitting a 3-inch diameter head with markings at 18" intervals to the micrometer screw; this modification greatly increases the precision with which one can extrude a drop of a given volume. In most cases an electrode area of 0.032 cm2 was used. The chronocoulometric technique requires that the electrode be initially at a potential Ei where no faradaic current flows-i.e., well anodic of the polarographic wave. The potential is then stepped to a point where the faradaic current is controlled solely by diffusion-Le., well on the polarographic diffusion plateau. This potential is maintained for a period of time, T , and then stepped back to the original potential. The integral of the current, Q,is measured as a function of time. During the forward step, 0 < t < T , Q is given by
Q = *nFD"2CoAt u 2 + Q d l + nFp
6
(1)
where Qdl is the charge required for the double layer,
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
Q ~=I JE;Cdiff(E) dE
(2)
