The chloride-catalyzed decomposition of ammonium nitrate in nitric acid media at 100°C

.( inorg, nucl. ('hera. Vol. 41. pp. 1583-1588 Pergamon Pres~ Ltd.. 1979. Printed in Great Britain

THE CHLORIDE-CATALYZED DECOMPOSITION OF AMMONIUM NITRATE IN NITRIC ACID MEDIA AT 100°Ct A. D. KELMERS, L. MAYA, D. N. BROWNING:~ and W. DAVIS, Jr. Chemistry Division and Chemical Technology Division, Oak Ridge National Laboratory, Oak Ridge, TN 37830, U.S.A.

(First received 15 November 1978; in revised form 14 March 1979) Abstract--The rate of disappearance of ammonium ion in chloride-catalyzed nitric acid mixtures at 100°C can be expressed by the rate equation d(NH4+)/dt =-k[H+]n[NO3-]°[CI-]qNH4+], where k =(2.430-+0.186)× 10 7 (sec ~), n = 2.426-+0.034, o = 3.096-+0.039, and p = 1.295±0.035. The proposed reaction mechanism is: 2H÷ + C1 + NO~-~HOCI + HNO2 HOCI + N H 4 + ~ N H 2 C I + H + + H.,O NH2CI + HNO2--+ N O 2 N H 2 + H + + CIN O 2 N H 2 --+N20 + H_,O

NHfl + NO3 -+ N20 + 2H20 Fluoride and bromide had no catalytic activity. This is in agreement with the proposed reaction sequence since they do not form stable analogs of hypochlorous acid or chloramine. INTRODUCTION Ammonium nitrate is stable in boiling nitric acid solutions although nitrogen has a formal valence of - 3 in NH4 ÷ and + 5 in NO3 -. The addition of HCI to solutions of ammonium nitrate in boiling nitric acid has been reported[l] to catalyze the decomposition of ammonium nitrate via the reaction HCI

NH4++ NO~

, N20 + 2H20.

(1)

Ammonium nitrate in radioactive waste solutions was destroyed by boiling in 0 . 1 N H C I - 2 . 5 N H N O 3 solutions[2]. We have investigated the rate of disappearance of ammonium ion in HC1-HNO3 solutions and have proposed a possible reaction mechanism and a rate equation. EXPERIMENTAL A four-neck, 500-ml round-bottom flask was equipped with a water-cooled condenser and water-cooled (~ 10°C) sampling burette so that - 10 ml samples of the solution in the flask could be withdrawn by vacuum into the burette and then discharged into ice-cooled brown-glass bottles. Catalyst addition was made via a pressure equalized addition funnel in the third neck. The fourth neck contained a thermometer. Temperature control and measurement utilized chromel-alumel thermocouples inserted in a protective glass tube through the condenser into the solution in the flask. One thermocouple was attached to a Doric Model DS-300 digital indicator and the other to a Leeds and Northrup ~Research sponsored by the Nuclear Power Development Division, U.S. Department of Energy under contract W-7405-eng26 with the Union Carbide Corporation. SGreat Lakes Colleges Association Student Participant.

Electromax 11I temperature controller, which maintained the solution temperature at 100±2"C by means of a flask heating mantle. The flask contents were mixed by a magnetic stirrer. To carry out these experiments, 240 ml of a solution containing HNO3 and NH4NO3 were placed in the flask and 10 ml of an HCI solution was placed in the addition funnel. In some preliminary tests, solutions other than HCI were added to test for possible catalytic activity. In other tests, LiNO3 was added to the flask solution so that the H ÷ and NOC concentrations could be independently varied. The solution in the flask was heated to 100°C and the catalyst solution was then added. After allowing approx. 10 sec for mixing, the first sample (zero time point) was taken. Additional samples were taken at subsequent times, as measured by a digital elapsed-time timer. In several tests an air flow of 10--20cm3/min was introduced through the vapor space of the flask. Samples were stored refrigerated (-4°C) in brown-glass bottles until analyzed. The ammonium ion concentrations were determined by an ion-sensitive electrode as NH3 after the solutions had been made basic with NaOH. Chloride ion concentrations were determined by a potentiometric titration with AgNO~. RESULTS

A series of preliminary tests (not included in the 47 rate experiments described below) with 0.5 M NH4NO~ in 8 M HNO3 gave the following results: (a) without HCI, no analytically significant decrease in the NH4 + concentration occurred after 5400 sec (1.5 hr) at 100°C. (b) NaC! was as efficient a catalyst as HCI. (c) The following showed no catalytic activity: HBr, HF, HI, NaNO2, KCNS, K:Cr207 and Fe(NO3)> (d) At HNO3 concentrations less than 8 M, no loss of CI- from the solution occurred during the reaction, even when air was passed through the flask vapor space.

1583

1584

A. D. KELMERS et al.

A series of 47 experiments was then carried out to measure the rate of disappearance of NH4+. The conditions covered the ranges 1.0-7.0MHNO3, 0.020.30 M HCI, 0-6.0 M LiNO3 and 0.1--0.75 M NH4NO3, all at 100%. Examples of typical data are shown for a series of tests at different HNO3 concentrations (Fig. 1) and for different initial ammonium ion concentrations (Fig. 2). The negligible effect of an air sweep through the reaction flask is shown in Fig. 3. The composition of the gaseous reaction product and its rate of formation was determined by gas chromatographic techniques in one series of experiments with 0.1 M HC1, 6 M HNO3 and 0.5 M NH,oNO3 at 100°C. Helium was passed through the reaction vessel at 300cm3/min and the gas stream was periodically analyzed. Over the entire course of the reaction, N20 was the sole gaseous product. The amount of N20 evolved was compared with the residual NH4+ concentration as a function of time and the values corresponded to those for eqn (1). Thus eqn (1) is believed to be a valid representation of the overall reaction stoichiometry. DISCUSSION Analysis of data The potential validity of a first-order kinetic rate process was examined by plotting fraction of NH4+ remaining, or its reciprocal, on semilogarithmic scale vs time for all 47 rate experiments. With only a few exceptions, the data indicated that NH4+ disappearance was first-order rate. This was verified quantitatively by analyzing data of each run by use of the Marquardt nonlinear least-squares computer program[3, 4]. Results of this analysis and a summary of the pertinent experimental concentrations are listed in Table 1. Analyses of the data are based on the assumption of a constant percentage uncertainty in the analyses for the concentration of NH4+. This is the assumption most consistent with the analytical procedures. On this bases, the standard deviation of fit of data ranges from 0.2527.1% (column 10 of Table 1); values of this parameter equal or exceed 10% in only four experiments. Rate constants, designated as k[(Q) in the equation d(NH4+)/dt = - kf ( Q )(NH4+),

q / : / / ~ +

0.6 0.5

q

jj°." ~

]

I

m~

0.3

~

I

o o

I

0.2

U

D

I

1

2

3 TIME (K sec)

4

5

Fig. 2. Plot of ammoniumion concentrationvs time with 0.1, 0.2, 0.4 and 0.8 M initial NH4NO3.©, 7 M HNO3;O, 5.5 M HNO£ I-1, 4 M HNO3. All experiments with 0.1 M HC1 and 100°C3.

and their standard deviations, plus calculated intercepts [values of (NH4+) at zero time] and their uncertainties, are listed in Table 1. The quantity f(Ci) is some yet-tobe-determined function of the concentrations, or activities, of H +, NO3- and C[-. It should be noted that, excepting the experiment with 5 M HNO3, 0.1 M HCI, 2.0MLiNO3, there is good agreement between the experimental zero-time value of NH4+ concentration and the value (NH4+)o obtained by fitting the data of the individual experiments. Since the values of kf(G) obviously depend on concentrations of the several chemicals present, three plots, Figs. 4-6, were drawn in an effort to estimate the individual contributions of H +, NO3- and CI- to the function f(Ci). Straight lines, in Fig. 4 and 5, and slightlycurved lines, in Fig. 6, discussed below, are based on further nonlinear least-squares analysis. These figures, however, served as a basis for testing the hypotheses that the quantity f ( Q ) is simply a product of concentrations of H +, NO3- and Cl- raised to some powers. A test was made of the model f(C~) = ( H + y ' ( N 0 3 - ) ° ( C I - ) ",

-0.7ii

(2)

I

I

I

I

I

I

I

t

(3)

I

-o.si - 0.6

I

I

I

I

-0.9'

I

-t.0

- 0.8

-1.1 -t.0 +~r-I. 2 -r Z

-t.2

-t.3 -1.4

- 1.4

+,q"1" Z -I.6

-1.5 -L8

-I.6

-2.0

-I.7

I I

1

I

I

L

2

3

4

5

T I M E (K sec)

Fig. 1. Plot of ammonium ion concentration vs time with 4 to 7 M HNO3. O, 4 M, O, 5 M, I1, 6 M and CI,7 M HNO~.All experiments with 0.5 M initial NH4NO3,0.l M HCI and 100°C.

-I.8

\ 0.2 0.4 0.6 0.8 TIME (K sec)

t.o

Fig. 3. Effect of air sweep through the reaction flask. O, no air flow; Q, 20cc/min air flow. Both tests with 7MHNO3, 0.15 M HCI,0.50M initial NH4NO3and 100°C.

1585

The chloride-catalyzed decomposition of ammonium nitrate in nitric acid media at 100°C corresponding to the rate equation d[NH4+l/dt = - k'[H+]"[NO3-]°[CI]P[NH,+].

(4)

P a r a m e t e r s and standard deviations for this model are as follows: k ' = (2.430-+0.186)x 10-7, n = 2.426-+0.034, o =

3.096-+0.039, p = 1.295---0.035 and Std. dev. of fit = 20,5%. T h e 4-parameter model provides a very good fit of the data and s h o w s a stronger d e p e n d e n c e of NH4 + decomposition on nitrate concentration (the e x p o n e n t is 3.0%) then on acidity (the e x p o n e n t is 2.426). Values of k, n, o

Table 1. Least squares analysis and experimental data Analysis of individual runs 105 x No. of points

HNO~ conc. (M)

HCI conc. (M)

LiNO3 conc. (M)

Initial NH4 + conc. (M)

(sec -I)

o-(fit) (%)

4 Param. model, All runs k[(Ci) x 105 Isec i)

7 7 8 8 9 8 8 4

1.0 1.0 1.0 2.0 2.0 2.0 2.0 2.0

0.20 0.30 0.30 0.20 0.20 0.30 0.30 0.30

2.0 2.0 6.0 0 0 0 1.0 5.0

0.729 0.519 0.155 0.245 0.747 0.216 0.460 0.146

0.7988 0.4329 0.1566 0.2462 0.7746 0.2101 0.4690 0.1486

0.160 0.256 3.553 0.295 0.758 0.233 1.202 13.66

0.006 0.026 0.110 0.006 0.022 0.010 0.039 0.378

6.54 17.5 13.3 2.96 4.72 7.20 7.69 3.58

0.277 0.475 4.27 0.250 1.22 0.453 1.8(! 17.0

8 8 9 8

3.0 3.0 3.0 3.0

0.10 0.2(I 0.30 0.3(I

4.0 0 0 4.0

0.733 0.685 0.443 0.198

0.0111 0.0108 0.0042 0.0115

8.741 1.245 2.598 43.92

0.203 0.104 0.113 1.29

2.38 2.35 1.58 10.0

10.8 2.88 4.25 41.7

6 6 6 6 4 8 6 6 9

4.0 4.0 4.0 4.0 4.0 4.0 4.0 4.0 4.0

0.05 0.1(I 0.10 0.10 0.10 0.1(I 0.15 0.25 0.30

0 0 0 0 0 3.0 0 0 3.0

0.4922 0.5126 0.5106 0.4372 0.2040 0.7378 0.5132 0.5097 0.1989

0.0051 0.0034 0.0101 0.0058 0.0037 0.0430 0.0091 0.0030 0.0045

1.364 3.192 2.912 3.242 3.022 23.93 5.928 10.15 76.57

0.346 0.218 1.628 0.440 0.586 1.22 1.176 0.389 0.80

1.43 0.91 2.74 1.86 2.09 9.92 2.44 0.81 4.25

1.57 4.01 3.98 3.78 3.22 21.4 6.99 14.3 79.4

6 5 7 5 6 6 5 5 6 6

5.0 5.0 5.0 5.0 5.0 5.0 5.5 5.5 5.5 5.5

0.05 0.10 0.10 0.15 0.20 0.25 0.10 0.10 0.10 0.10

0.503 0.508 0.602 0.503 0.501 0.502 0.795 0.463 0.197 0.096

0.5141 0.5222 0.3702 0.5044 0.4975 0.5067 0.8172 0.4520 0.2004 0.0936

0.0073 0.0098 0.0625 0.0009 0.0099 0.0076 0.0278 0.0176 0.0023 0.0026

5.524 12.53 17.89 15.74 24.39 30.31 12.93 12.86 12.13 9.35

0.468 0.62 1.39 0.12 1.31 0.98 2.05 2.38 0.75 1.84

1.97 2.59 27.1 0.24 2.73 2.06 4.05 4.50 1.57 3.78

5.01 12.6 34.2 21.8 32.8 44.2 23.9 20.2 17.6 16.6

6 5 6 5 6 6 6

6.0 6.0 6.0 6.0 6.0 6.0 6.0

0 0 0 0 0 0 0

0.486 0.500 0.494 0.500 0.511 0.500 0.420

0.4827 0.5133 0.4965 0.5098 0.5210 0.5002 0.4147

0.0070 0.0113 0.0105 0.0159 0.0068 0.0052 0.00%

4.179 16.59 25.15 26.86 42.37 61.86 80.10

0.478 1.50 1.75 2.37 1.43 1.72 3.84

1.98 2.88 2.91 3.68 1.81 1.43 3.15

3.89 13.0 32.4 32.5 56.4 83.1 109

6 6 6 6 6 6 6 6 5

7.0 7.0 7.0 7.0 7.0 7.0 7.0 7.0 7.0

0 0 0 0 0 0 0 0 0

0.500 0.507 0.103 0.199 0.411 0.754 0.500 0.487 0.488

0.5117 0.5149 0.1012 0.1995 0.4314 0.7806 0.5288 0.4868 0.4831

0.0065 0.0118 0.0029 0.0040 0.0214 0.0257 0.0206 0.0052 0.0093

0.84 1.51 3.75 2.68 6.53 4.33 6.40 1.75 2.94

1.77 3.18 3.86 2.78 6.93 4.61 5.47 1.46 2.23

29.3 73.4 61.9 64.5 70.6 81.2 126 185 251

kf(Ci)

o[kf(Ci)]

(sec b

0.0343 0.0473 0.0105 0.0057 0.0204 0.0088 0.0204 0.0042

0.7251 0.7057 0.4364 0.1898

0.500 0.511 0.500 0.429 0.206 0.752 0.515 0.508 0.200

0 0 2.0 0 0 0 0 0 0 0

0.02 0.05 0.10 0.10 0.15 0.20 0.25 0.05 0.10 0.10 0.10 0.10 0.10 0.15 0.20 0.25

Number of runs Number of points Std. deviation of fit (%)

(NH4+)o] o'[(NH4+)o] (M) (M)

30.01 50.98 48.56 50.24 46.25 48.74 107.6 121.8 130.6

47 299 20.5

1586

A.D. KELMERS et al.

and p of the 4-parameter model were used to draw the lines in Figs. 4-6. These plots show both the goodnessof-fit and the substantial, apparent deviations of data pertaining to a few of the experiments. Values of k[(C) calculated from the 4-parameter model are listed in the last column of Table 1 for comparison with the individual rate constants shown in column 8. The dependence of the rate process is assumed in eqn (4) to be first order with respect to [NH4+]. The validity of this assumption was tested by fitting a major portion of the data to a model in which the ammonium ion concentration exponent was allowed to change. The resulting value was 1.01 with a standard deviation of 0.08. This dependence is so close to unity that the assumption of first-order dependence of rate, eqn (4), on [NH4+] was retained. Theoretically, it should be possible to improve the fit by using the data to calculate parameters of an activitycoefficient equation, such as an extended Debye-Huckei equation. We tested this concept by replacing each concentration in eqn (3) by an activity calculated, for example, as [H +] = y(H+),

10-3

(c_~) O~ ~3 (1.1 Od

T

~._*)

Q

23 0 3.3 @ ~ 4.1 0 ~1

P,I

• ....

I

(J

l )

z 0

p

10-4

P, I

/.

Z tL I-

/

IO-S

(5)

/

/Y

/

I

where y is an activity coefficient calculated by use of the equation In y = - SX/(I)/(1 + I.Sx/(I)) + aI + bI 2 + cP.

(6)

¢_

,~ .,.q I

2 3 4 6 NO~ CONCENTRATION (fill

8

I0

Fig. 5. Effect of NO3- concentration at constant CI- and H+. All lines have slopes of 3.096 and represent the 4-parameter nonlinear least-squaresanalysis. L)

Here I is ionic strength, S is the Debye-Huckel coefficient (1.3829) at 100°(2, and a, b and c are additional parameters. Use of this concept only marginally improved the fit. We concluded that the accuracy of the 4-parameter model and of the data are not sufficient to warrant inclusion of activity coefficient results in this paper.

I--

Z tie

,9 i,-

8 LO r,tw ne O

h

,8

I

2

3

4

5 6 78910

H" CONCENTRATION (M)

Fig. 4. Effect of H+ concentrationat constant CI- and NO3-. All lines have slopes of 2.426 and represent the 4-parameter nonlinear least-squares analysis.

Reaction mechanism Very little information is available on the chemistry of hot HCI-HNO3 mixtures, especially at low HCI concentrations. The first reaction must involve CI- and HNO3 to generate an active chlorine-containing component, which need be present only in trace quantities if it is constantly generated as used during the course of the reaction. The experiments which showed that air flowing through the flask vapor space did not affect the reaction rate constant or remove Cl- eliminates the possibility of a significant concentration of any intermediate with an appreciable vapor pressure over these nitric acid solutions at 100°C. Thus compounds such as chlorine or nitrosyl chloride are unlikely to be the active component. Also, HBr showed no catalytic activity, although Br2 and NOBr are chemically very similar to the corresponding chlorine compounds. Thus the active component should be one in which the chlorine-containing compound is stable and the corresponding bromine-containing compound is much less stable or much less

The chloride-catalyzeddecompositionof ammonium nitrate in nitric acid media at 100*C

o o _o

El 0

l,NOs-) 7 6

A

5

O

4

1587

It is also possible to gain information about the reaction sequence from the final compounds, N20 and H20. The evolution of N20 gas suggests the decomposition of a compound containing two interconnected nitrogen atoms. The decomposition of nitramide

10-5

NH2NO~ ~ N20 + H20

(13)

[,-

has been carefully investigated by Bronsted[6] and others[7, 8]. The reaction is catalyzed by acid and also weakly catalyzed by chloride. An alternative dinitrogen compound is hyponitrous acid, H2N202, since the acidcatalyzed decomposition[9] of H2N202 also leads to N20

0

Z

lq+

H2N202 ~

I0-4

I1£

8

-

.02

£B ,04

,06 .08 .1

.2

.3

.4 ;5

Cl- CONCENTRATION (M) Fig. 6. Effect of CI- concentration at constant NO3- and nearly

constant H ÷. ([H+]=[NOf]+[CI-]).

Slightly curved lines

represent the 4-parameternonlinearleast-squaresanalysis.

reactive. The hypohalous acids meet this criterion[5]. For the reaction X2 + HzO~HOX + H ÷ + X-

(7)

in acid solution the equilibrium constant for C12 is 4 x 10 - 4 , while for Br2 it is 7 x l 0 - 9 . HF also exhibited no catalytic activity. Hypofluorous acid, HOF, is unstable even at ambient temperatures and would not be present in experiments at 100°C. Thus, the initial reactions may be 3H + + 2C1- + NO3 ~C12 + H20 + HN02

(8)

C12 + H20~HOCI + H ÷ + CI

(9)

H + + CI- + HNO3~HOCI + HNO2

(10)

N20 + H20.

Nitramide seems the more likely compound since reaction (18) involves nitrous acid. Nitramide is structurally O2N-NH2 while hyponitrous acid is HON=NOH and it seems more plausible for HNO2 to react with chloramine to form nitramide (the amine end from the NH2C1 and the -NO2 end from HNOz) than for a rearrangement to occur to form hyponitrous acid. However, the distinction between NH2NO2 and H2N202 is not critical to the proposed reaction sequence. Nitramide has been postulated as an intermediate in the decomposition of molten ammonium nitrate [I0]. Consideration of postulated[l 1] mechanisms for the chloride-catalyzed decomposition of molten ammonium nitrate at higher temperatures provides a potential series of intermediate steps between HOCI and NH2NO2. The decomposition at 160--260°C was reported to be greatly accelerated by chloride and the possible role of chloramine, NH2CI, as an intermediate was discussed [ l l]. In acid solution, pH<3, the hypochloride--chioramine equilibrium was reported[12] to yield nitrogen trichloride from the sequence NH4+ + OCI-~NH2C1 + H20

(15)

2NH2CI + H-~NHCI2 + NH4+

(16)

3NHCI2 + H+~NCI3 + NH4+.

(17)

However, the equilibria approaching NC13 were reported to be slow and dichloramine is very unstable. Thus, it is possible that in hot nitric acid solutions containing nitrous acid from equation (10), the proposed reaction of chloramine and nitrous acid NH2CI + HNO2 ~ NO2NH2 + H ÷ + CI

which would yield low but significant concentrations of hypochlorous acid and nitrous acid in solution. Since chloride is not lost from the reaction solution, the concentration of the intermediate Cl2 must be quite low. Likewise, the concentration of HNO2 must be low since no brown fumes characteristic of NO2 from the reaction

(14)

(18)

were observed during the experiments. The concentration of HOCI must also be low since no loss of chlorine as volatile dichlorine oxide via the reaction

may be kineticaily favored over the reaction of chloramine with a proton (reaction 16). The selection of chloramine as an intermediate also is in agreement with the observation that HBr was not catalytic since bromamide, NH2Br, is much less stable in aqueous solutions than chloramine[13], as well as the precursor HOBr being less stable than HOC1. It is, therefore, possible to write a sequence of plausible reactions to represent the chloride-catalyzed decomposition of ammonium nitrate in nitric acid media. The reactions are:

2HOCI ~ 0 2 0 + H20

2H+ +CI + NO3-~HOCI + HNOz

(10)

HOCI + NH4+~NHaCI + H + + H20

(15a)

HN02 + H+ + NO3-~.~-2N02 + H20

occurred. nNC Vol. 41, No. ! I--E

(11)

(12)

1588

A. D. KELMERS et al. NH2CI + HNO2 ~ NO2NH2 + H + + CI-

(18)

NOzNH2--> N20 + H20

(13)

NH4 + +NO3---) N20 + 2H20

(1)

The equilibrium (reaction 10) which generates hypochlorous and nitrous acids must have a very small equilibrium constant, as discussed above, and may also be highly temperature dependent. In the second equilibrium, reaction 15a, hypochlorous acid reacts with ammonium ion to yield chloramine. The chloramine likely is relatively unstable under these conditions. Nitrous acid is known[14, 15] to react very rapidly with other nitrogen containing compounds (hydrazine, amidosulfonic acid and hydrazoic acid) in nitric acid media and thus could react rapidly with transient molecules of chloramine, as shown in reaction 18. This reaction must yield either nitramide or hyponitrous acid, although nitramide appears the more likely, and release the catalyst, CI-, for reuse in reaction 10. It seems unlikely that the reverse of reaction 18 occurs to any significant extent. Finally, the nitramide (or hyponitrous acid) decomposes to nitrous oxide and water. While not unequivocally established, this reaction mechanism is consistent with the known chemical properties of the compounds. The postulated reaction mechanism does not predict the observed third order dependence on nitrate concentration and we are unable to explain that relationship.

Acknowledgements--The authors wish to acknowledge the contributions of D. Y. Valentine, Chemistry Division, ORNL, and J.

A. Haas and R. J. Hydzik, co-op student participants, who performed many experiments; M. R. Bennett, Chemistry Division,ORNL, for gas chromatographic data; and the Analytical Chemistry Division Staff, who performed the many analyses.

REFERENCES 1. J. van R. Smit, Chem. Ind. (London) 2018, 2019 (1964). 2. W. M. Campbell, Nucleonics 14, 94 (1956). 3. D. W. Marquardt, 3. Soc. Ind. Appl. Math. I1,431 (1%3). 4:. D. W. Marquardt, Least-Squares Estimation o/Nonlinear " Parameters (NLIN), Share Program Library, SDA-2093-01 (Rev.), 15 Aug. (1%6). 5. A. J. Downs and C. J. Adams, Comprehensive Inorganic Chemistry (Edited by J. C. Bailor, H. J. Emeleus, R. Nyholm and A. F. Trotman-Dickerson), Vol. 2, p. 1406. Pergamon Press, New York (1973). 6. J. N. Bronsted and K. Pederson, Z. Physik. Chem. 108, 185 (1924). 7. C. A. Marlies and V. K. LaMer, J. Am. Chem. Soc. 57, 1812 (1935). 8. B. Perlmutter-Haymon and M. A. Wolff, Israel J. Chem. 7, 153 (1%9). 9. M. N. Hughes, Quart. Rev. 22, I (1%8). 10. L. Friedman and J. Bigeleisen, J. Chem. Phys. 18, 1325 (1950). I1. G. Guiochan, Ann. Chim. 5, 295 (1960). 12. R. E. Corbett, W. S. Metcalf and F. G. Soper, J. Chem. Soc. 1927 (1953). 13. H. Galal-Gorchev and J. C. Morris, lnorg. Chem. 4, 899 (1%5). 14. J. R. Perrott, G. Stedman and N. Uysal, J. Chem. Soc. (Dalton Trans.) 2058 (1976). 15. J. R. Perrott and G. Stedman, J. Inorg. Nucl. Chem. 39, 325 (1977).