T. HUANGAND J. T.SPENCE
4198
The Oxidation of Hydrazine by Molybdenum( VI)1 by T. Huang and J. T. Spence Chemiatry Department, Utah State University, Logan, Utah 84981
(Received Mag $9,1968)
The oxidation of hydrazine by Mo(V1) has been investigated in the pH range 1.2-3.2. The reaction produces NZin stoichiometricamounts with no detectable formation of NHI. The reaction is first order in each reactant, with a fractional-order H + dependence in the range studied. Evidence for diimide (N2Hz) as an intermediate has been obtained by mass spectrometryand by trapping with unsaturated acids. The results indicate Mo(V1) is acting as a two-electron oxidant, and a mechanism for the reaction involving the intermediate formation of Mo(1V) and NzHzhas been developed.
Molybdenum, in the form of metalloenzymes, is involved in the biological processes of nitrogen fixation2 and nitrate r e d ~ c t i o n . ~As part of a study of the reactions of molybdenum species with possible intermediates of these processes as models for the enzymatic reaction^,^ the oxidation of hydrazine by Mo(V1) has been investigated. Hydrazine is thermodynamically unstable toward both oxidation and reduction in aqueous solution, as indicated by the reduction potentials5
+ 3H+ + NZH5+ 2NH4+ (EO = 1.275 V) 4e- + Nz + 5H+ +N2H5+ (EO = -0.230 V)
2e-
---t
The reduction to NH4+ is generally very slow, however, while the oxidation to I T 2 proceeds at a reasonable rate and has been investigated with a number of oxidizing agents .6 The oxidation of hydrazine by methylene blue, catalyzed by Mo(VI), has been used as a test for hydrazine for some time,’ and recently a preliminary study of the oxidation by NIo(V1) has appeared, but little kinetic data or information concerning the mechanism was given.8
monomeric, at the concentrations used, although this is not certain.$ Kinetics. The order of the reaction was determined by plotting the data using different rate expressions and was found to be second order, first order in Mo(V1) Table I : Quantitative Analysis of Gas Evolved from the Hydrazine-Mo(V1) Reaction at 70” (pH 1.60, [Mo(VI)]a = 5.00 X 10-4 M , [NzHa+]o = 2.41 X M) I O ~ [ N Z H ~ + I 104[Nzl Time, min
104[NzHs *I, M
0
2.41 2.12 1.97 1.81 1.65 1.58 1.53
42.7 74.5 135.5 211.0 253.4 303.0
(reacted),
M
(formed), M
0.29 0.44 0.60 0.76 0.83 0.88
0.33 0.41 0.61 0.73 0.80 0.87
and first order in NzHs+ over a pH range of 1.2-3.2, according to the rate equation
Results Stoichiometry. Preliminary investigation showed the anaerobic reaction between Mo(V1) and hydrazine proceeds a t an easily measurable rate a t 70” in a phosphate buffer at pH 1.6. Qualitative analysis of the reaction mixture for NH, gave negative results, and mass spectrometric analysis of the gas above the solution revealed the presence of Nz only. The rate of disappearance of N2H4 was found to be equal to half the rate of appearance of Mo(V)z, (Figure l), and quantitative analysis of gas formation using a Gilson respirometer showed 1 mol of gas formed for each mole of NzH4 reacted (Table I). These data indicate the stoichiometric equation is 491o(VI)
+ NzH5++Nz + 2n!fo(V)~+ 5H+
Molybdenum(V) is present mainly as a dimer under these conditions, while Mo(V1) is probably mainly The Journal of Physical Chemistry
This expression is easily integrated giving the standard second-order equation (typical results are plotted in (1). Abstracted from the Ph.D. dissertation of T. Huang, Utah State University, 1968. (2) L. E. Mortenson, J. A. Morris, and I. R. Kennedy, Proceedings of the 68th Annual Meeting on Bacteriology, Detroit, Mich., 1968,p 133. (3) A. Nason, “The Enzymes,” Vol. 7, E’. D. Boyer, H. Lardy, and K. Myrback, Ed., Academic Press, New York, N. Y., 1958,p 587. (4) E.P. Guymon and J. T. Spence, J . Phys. Chem., 70,1964 (1966). (5) W.M. Latimer, “Oxidation Potentials,” 2nd ed, PrenticeHall, Inc., Englewood Cliffs, N. J., 1952,p 99. (6) W. C. E. Higginson, “The Oxidation of Hydraaine in Aqueous Solution,” Special Publication No. 10, The Chemical Society, London, 1957. (7) R. Lang, 2. Anal. Chem., 128, 165 (1948). (8) 8. Ostrowetsky and D. Brinon, Compt. Rend., 263, 406 (1966). (9) E. F. Rohwer and J. J. Cruywagen, J . S. African Chem. Inst., 16, 26 (1963).
THEOXIDATION OF HYDRAZINE BY MOLYBDENUM(VI)
4199
2.6
2.0
$
1.6
7
2 X
v
!1 . 0
u"
0.6
f
100
/
I
300
100 200 Time, min.
Figure 1. Oxidation of N&&+by Mo(V1) ( [ N ~ % + ]=o 2.41 X lO-4M, [Mo(VI)]o = 5.00 X 10-4M, pH 1.60, phosphate buffer, p = 0.22 M, 70"): 0, [NzHs+]; A, [Mo(V)z].
Figure 2 ) , and the rate constants, obtained from the slopes of the lines by the least-squares method, are t'abulated in 'Table 11. The second-order plots were linear to -85% reaction and all runs were followed for at least 2 half-lives. Table I1 : Rate Constants'
2.41 2.42 1.91 2.42 1.50 12.50 2.33 2.33 2.33 2.30 2.30 2.35 2.27 2.40
5.00 7.41 4.95 7.46 60.0 5.00 7.46 7.46 7.41 7.46 7.46 7.46 7.46 7.46
*
1.60 1.60 1.60 1.60 1.60 1.60 1.60 1.20 2.00 2.40 2.80 3.20 1.60 1.60
70 70 70 60 60 60 90 90 90 90 90 90 80 90
5.86 5.93 5.83 2.83 2.90 2.80 21.0 19.4 28.0 44.6 44.6 39.0 10.0 17.0"
*
a E, = 14.3 kcal/mol, A S = - 8.4 eu, AF = 16.5 kcal/mol, pH 1.60, p = 0.22,70". Each value is the average of two runs. The average of the deviations for each pair is *1.5% (e.g., 5.86 0.07 1. mol-1 min-1). 1.00 M NaCl added.
*
~~~~
~~
~
~~
It was found that the higher the pH, in the range 1.23.2, the faster is the rate. A plot of log IC2 vs. pH gave a straight line with slope of -0.25, indicating a complex relationship between rate and [H+]. This fractional
200 Time, min.
300
400
600
Figure 2. Second-order kinetic plots ( a = [NzH~+]o, b [Mo(VI)]o, x = [NzHb+]): 0, [Mo(VI)]o 7.46 X M, [NzH6+]0= 2.30 X M, pH 2.80, 90'; 0, [Mo(VI)]o = 7.46 X lo-* M, [NzHsfIo = 2.27 X lO-'M, pH 1.60, 80"; A, [Mo(VI)]o = 7.41 X M, [NzHs+]o= 2.42 X lod4M,pH 1.60, 70'; 0, [Mo(VI)]O= 7.46 X 10-4 M ,[NzH6+]o = 2.42 X lo-' M, pH 1.60, 60'.
dependence on H+ may reflect changes in the state of polymerization of Mo(V1) with pH,IO ionization of a monomeric molybdate species, and/or the involvement of H+ in the rate-controlling step. It has been reported that at M concentrations, Mo(V1) is present mainly as a monomeric dibasic acid with pK1 = 4.00 and pKz = 4.08,9although there is no general agreement on these results. I n any case, the first-order dependence of the rate on the total Mo(V1) concentration indicates the reacting species must be present essentially all in one form, whether it is a monomer or higher polymer, while the H+ dependence cannot be explained until more is known of Mo(V) and Mo(V1) species in solution. As seen in Table 11, only minor variation of rate is observed with a large change in ionic strength at pH 1.60. Therefore, small changes in ionic strength due to small changes of pH are without any appreciable effect on the rates and may be neglected. The activation energy and entropy for the reaction were obtained from an Arrhenius plot at four temperatures from 60 to 90". The free energy of activation was calculated by means of Eyring's equation, and these results are found in Table 11. I n order to obtain information concerning intermediates, an esr study of the reaction was made. Samples taken during the course of the reaction revealed an esr signal, identified as Mo(V) on the basis of its g value (1.94), shape, and hyperfine splittings (naturally occurring Mo contains -25% isotopes with I = 5//2, which give six satellite lines)." (10) J. Aveston, W. Anacker, and J. S. Johnson, Inorg. Chem., 3 , 735 (1964). (11) R. C . Bray, G . Palmer, and H. Beinert, J . Biol. Chem., 239, 2667
(1964).
Volume 7.8, Number 18 November 1968
T. HUANGAKD J. T. SPENCE
4200 By measuring the Mo(V) signal (by double integration of the esr signal) in a control solution containing the same total amount of molybdenum(V) as the samples from the reaction, it was found that the ll'ro(V) monomer is at all times in equilibrium with the dimer, n!to(V), (esr inactive) 11IO(V)Z
2iVIo(V)
No other esr signal was observed at any field strength. Diimide (NzHz)was suspected as an intermediate, since in other reported oxidations of x2H4 giving N2 as the only product this species has been shown to be involveda6 A sample from the reaction solution was found to give a small peak at m/e 30 in the mass spectrometer, corresponding to NzH2. To determine if this might have come from the decomposition of NzH4 in the ionization chamber, a sample of N2H4 under the same conditions was run and no peak at m/e 30 was observed. It was attempted to trap diimide by reaction with unsaturated compounds, since it is known it can reduce olefins.12 A small amount of cis-4-cyclohexene-1,2-dicarboxylic acid was added to the reaction mixture, and after the reaction was essentially complete, the organic compounds from the aqueous solution were extracted by ether and then applied to a silica gel thinlayer plate. The chromatogram was developed with dioxane, dried, and sprayed with brom cresol green. A small amount of cis-1,2-cyclohexane dicarboxylic acid was identified by comparison with Rf values of pure compounds. This arose from the reaction
In order to make the identification of diimide more positive, fumaric acid was used (because of its greater solubility) and the pH was raised to 8.5, since it has been reported that yields in the reduction of olefins by diimide are poor below this pH. A sample was taken from the reaction mixture under these conditions, and the organic components were separated and subjected to analysis in the mass spectrometer. A small but significant peak at m/e 100, most likely due to succinic anhydride, formed by reduction of fumaric acid with diimide and subsequent dehydration under the conditions of separation used, was found. A sample of pure succinic acid, treated in the same way, gave a peak at m/e 100, indicating succinic anhydride is formed under the conditions of separation.
Discussion Previous work using a number of different oxidants indicates the oxidation of hydrazine can proceed by two paths, depending on whether the oxidant is a o n e or a two-electron oxidant The Journal of Physical Chemistry
/ xzH4\ -2e1NzH2
+Nz
These paths are based on evidence from kinetic studies and the use of W-enriched hydrazine.6 Since in the oxidation by Mo(V1) only Nz is formed and N2H2 has been qualitatively detected by mass spectrometry and trapping with unsaturated acids, it appears Mo(V1) is acting as a two-electron oxidant. This conclusion is in agreement with other work on the Rlo(V1)-catalyzed oxidation of oxalate by H2O2I3and the reductions of IZand O2 by Mo(V)2.l4 Furthermore, R/Io(V) monomer is present only at equilibrium levels with ni10(V)~, thus making it unnecessary to involve it as an intermediate. Based on these data, the stoichiometry, and the kinetics, the following mechanism is proposed for the reaction
+ Mo(V1) -% N2H2+ Mo(1V) H + + 2;"\',H2--% Nz+ NzHj+ Mo(1V) + Mo(V1) I_ 2MO(V)
(I)
xZH5+
(11)
k3
(111)
k- a
kr
2Mo(V)
k- 4
(IV)
1Io(V)2
In this mechanism, reaction I is rate determining, reaction 116is fast, the equilibria I11 and IV are quickly established, and both n9o(IV) and N2H2are reactive intermediates. From this mechanism, the disappearance of rVIo(V1) and N2&+ can be written as -d[N2H5+1=
dt
kl[X2H5+][Rilo(VI)]-
k 3 [ n ~ o (1 [n!to(vI) ~~) 1-
~[;\/10(v)( 2 ) 12
Applying the steady-state approximation to XZHZ and Mo(IV), the following equations are readily obtained
-d[n'o(V1)l dt
=
2kl[NzH5+][Mo(VI)]
(4)
(12) E.J. Corey, D. J. Pasto, and W. L. Mock, J . Amer. Chem. SOC., 83, 2957 (1961). (13) P. Saffir and H. Taube, ibid., 82, 13 (1960). (14) E. P. Guymon and J. T. Spenoe, J . Phys. Chem., 71, 1616 (1967).
THEOXIDATIOiN OF HYDRAZINE BY nqOLYBDENUM(V1) Comparing eq 3 with eq 4
-,---.d[NzHS+] -d[Mo(VI)] = dt
*
dt
1:4
(5)
From the ablove derivation, eq 5 is seen to satisfy the stoichiometry of the reaction. The expressions in both eq 3 and 4 are identical with the experimental rate expression, with l/Jc1 = kz', and the proposed mechanism is therefore consistent with the kinetic data. The small amount of Mo(V), as detected by esr, would arise frlom equilibrium IV and should be never more than the equilibrium value, as was observed. Although aquomolybdenum(1V) is thermodynamically unstable, its presence as a reactive intermediate is in agreement with studies of redox reactions of molybdenum(V) and m ~ l y b d e n u m ( V I ) , ' ~ and ~ ' ~ there is evidence that equilibrium I11 is involved in such rea c t i o n ~ . ' ~There ~ ~ ~ are cases, however, where the Mo(V) monomer appears to be the active intermediate,4 and it may be that the biological importance of the metal is due to its ability to react via both states, thus mediating the transfer of either one or two electrons.
Experimental Section Chemicals. lllolybdenum(V1) stock solutions were prepared from reagent grade NazMo04 2Hz0 which had been analyzed to be 99.7% pure.16 Stock solutions of hydrazine sulfate were prepared from Fisher Co. reagents and amalyzed by titration with standard KIO3 s ~ l u t i o n . ' ~Fumaric acid, succinic acid, cis-4-cyclohexene-1,2-dicarboxylic acid, and cis-cyclohexane-1,2dicarboxylic acid were obtained from Eastman Co. and were recrystallized before use. Helium (99.99%) and nitrogen (prepurified, 99.99o/c), used for deaerating solutions, were obtained from Matheson Co. I n all cases, redistilled HzO was used, since traces of Cu2+ or Fe3+ catalyze the autodecomposition of hydrazine. For stoichiometric and kinetic measurements, hydrazine was determined by the spectrophotometric method of Watt and Chrisp,16 which depends on the formation of a yellow color (A, 458 mp) upon the addition of p-dimethylaminobenzaldehyde to a solution of hydrazine in dilute HC1. Molybdenum(V) was measured spectrophotometrically at 289 mp. Since it was necessary to work in the absence of Oz, a specially constructed quartz spectrophotometer cell which could be evacuated, filled with deaerated solution, and sealed was used for molybdenum(V) measurements. After filling, the cell was placed in a thermostated compartment ( A T = :&0.5")on a Becliman Model DU spectrophotometer and measurements were made periodically. For hydrazine analysis, the deaerated reactants were mixed under helium in a flask which was kept in a constant-temperature bath (AT = dz0.1'). Samples were withdrawn with a hypodermic syringe through a rubber diaphragm and were analyzed for NzH4. For mass spectrophotometric gas analysis, a special flask a
4201 was constructed which could be evacuated and connected to the sealed reaction vessel. After the reaction was complete, a stopcock was opened, allowing the gas above the reaction to flow into the flask, and the stopcock was then closed. The flask was placed in Dry Ice-2-propanol in a dewar to freeze out HzO and then was connected immediately to the inlet of the mass spectrometer for analysis. A Gilson respirometer was used for quantitative measurements of Nzformed during the reaction. Nessler's reagent was used to test qualitatively for the presence of NH3. In order to trap diimide, a modified reaction vessel was employed. A 1000-ml three-neck flask equipped with a condensor, gas inlet and outlet, and a thermometer was used. I n a typical run, 600 ml of a solution 0.30 M in Rllo(V1) and 0.06 M in cyclohexene-1,2-dicarboxylic acid was added to the flask and was deaerated with He for 2 hr. Following this, 100 ml of 0.35 M NZH4 in phosphate buffer which had been deaerated with He was added to the flask, and the reaction was allowed to continue for 800 min at 90". The solution was then concentrated to about 100 ml, cooled in an ice bath, and acidified with a few drops of HCl. The solution was extracted with ether and the ether was removed under vacuum, giving white crystals of the mixture of saturated and unsaturated acids. The crystals were dissolved in a small amount of ether and spotted on a thin-layer plate of silica gel. The chromatogram was developed with dioxane as solvent. After drying, the plate was sprayed with brom cresol green, giving yellow spots for the dicarboxylic acids. The saturated and unsaturated acids were identified by comparison of their Rf values with those of pure compounds. The same procedure was used with fumaric acid, except that the succinic acid formed by reduction was identified by mass spectrometry (it was not possible to identify cis-cyclohexane-1,2,-dicarboxylic acid at the small concentrations present by mass spectrometry because of its high melting point). For esr analysis, samples were removed from the reaction with a gas-tight syringe, transferred anaerobically to a quartz esr tube, frozen in liquid nitrogen, and measured in a Varian V-4500-10 X-band esr spectrometer equipped with 100-kc field modulation. For standardization, a solution of K31/lo(CN)s,prepared as described previously,lS was used. The esr derivative spectra of the samples were doubly integrated and compared with the doubly integrated spectrum of the standard to determine the concentration of spins (estimated error of If 10%). The g values were determined by comparison with quinhydrone in alkaline ethanol. (15) A. A. Bergh and G. P. Haight, Jr., Inorg. Chem., 1, 688 (1962). (16) J. T. Spence and G. Kallos, ibid., 2, 710 (1963). (17) G. W. Watt and J. D. Chrisp, Anal. Chem., 24, 2006 (1952). (18) J. T. Spenoe and M. Heydanelr, Inorg. Chem., 6, 1489 (1967).
Volume 72, Number 18 hrovember 1968
4202
A. S. KERTESAND G. MARKOVITS
Acknowledgment. We wish to express our thanks to the U. S. Public Health Service (Grant GM-08347,
National Institute of General Medical Sciences) for financial support.
Activity Coefficients, Aggregation, and Thermodynamics of Tridodecylammonium Salts in Nonpolar Solvents by A. S. Kertesl and G. Markovits Department of Inorganic and Analytical Chemistry, The Hebrew University of Jerusalem, Jerusalem, Israel (Received May 31,1968)
Vapor pressure lowering measurements on benzene, carbon tetrachloride, and cyclohexane solutions of trilaurylamine chloride, bromide, nitrate, perchlorate, and bisulfate have been carried out a t 25, 37, and 50”. The data analyzed via the Gibbs-Duhem relation reveal that the activity coefficients of the salts decrease abruptly with increasing concentration up to about 0.1 m and tend to level off a t higher concentrations. Interpreting the osmometric data in terms of molecular association of the solutes via the Bjerrum relationship, dimers and higher oligomers have been shown to exist in solution even a t the lowest measurable concentration. The formation constant of dimers in benzene increases in the order chloride < bromide < perchlorate < bisulfate. The thermodynamic functions evaluated from the change of the formation constants suggest that the bond energy is not greatly affected by the change in the solvent. The average AH” per dipole bond in a linear oligomer is about 3.3 kcal/mol, while in a cyclic oligomer it is around 4.1 kcal/mol.
The current interest in long-chain aliphatic amine salts in solvent-extraction processes has led us to undertake a thermodynamic study designed t o give information concerning ionic and molecular interactions in solutions of such salts in nonpolar, water-immiscible organic solvents. As a part of that project, we report here the results of osmometric, vapor pressure lowering measurements of several trilaurylammonium salts in benzene, cyclohexane, and carbon tetrachloride at 25, 37, and 50”. Allrylammonium salts, when dissolved in low dielectric constant organic solvents, usually exist in the form of ion pairs.2 Depending upon the nature of the amine salt and its concentration and the nature of the organic solvent employed, the ion pairs may either dissociate or associate into higher aggregates. In nonpolar solvents the ionic association constants for various alkylammonium salts has been shown t o be of the order of lO6-lO’, indicating that the dissociation of the ion pairs is negligible in salt concentrations practical for solvent-extraction processes (>0.01 M ) . 3 On the other hand, electrostatic molecular association of the ion pairs into dimers and higher oligomers in nonpolar solvents is a common phenomenon and is known t o affect metal extraction greatlye2v3 Consequently, for a mass action law treatment of the experimental data, activity coefficients are needed to correct The Journal of Physical Chemistry
for both the nonspecific nonideality of the solutes and for the specific nonideality caused by the aggregation of monomers. The considerable amount of both qualitative and quantitative work on the extent (number) and degree (size, both of aggregated units) of association of various long-chain aliphatic amine salts has recently been re~ i e w e d . ~Diluent ?~ vapor pressure lowering, determined by direct vapor pressure measurements or by isopiestic balancing, has been used by Coleman and Roddy4-6 for investigating departures from ideality in behavior of trioctylammonium sulfate and bisulfate in benzene. Bucher and Diamond,’ Muller and Dia(1) Chemistry Division, Argonne National Laboratory, Argonne, 111. 60439. (2) Y. Marcus and A. 9. Kertes, “Ion-Exchange and Solvent Extraction of Metal Complexes,” Interscience Division, John Wiley & Sons, Inc., London, 1968, Chapter 10. (3) A. S. Kertes in “Recent Advances in Liquid-Liquid Extraction,” C. Hanson, Ed., Pergamon Press Ltd., London, 1968, Chapter 1. (4) Chemical Technology Division Annual Progress Reports, Oak Ridge National Laboratory, Oak Ridge, Tenn., ORNL-3452, 1963; ORNL-3627, 1964, p 206; ORNL-3945, 1966, p 186. (5) C. F. Coleman and J. W. Roddy in “Solvent Extraction Chemistry,’’ D. Dyrssen, J. 0. Liljenzin, and J. Rydberg, Ed., NorthHolland Publishing Co., Amsterdam, The Netherlands, 1967, p 362. (6) C. F. Coleman, 1967, private communication. (7) J. J. Bucher and R. M. Diamond, J . Phys. Chem., 69, 1565 (1965).
