THE JOURNAL OF
PHYSICAL CHEMISTRY (Registered in
VOLUME65
U. 6. Patent 0 5 c e )
(0Copyright, 1961, by the American Chemical Society)
DECEMBER 6, 1961
SUMBER 11
KINETICS OF THE OXIDATION OF MOLTBDEXJRI(V) BY HYDROXYLAMINE* BY G. P. HAIGHT, JR., AND ALICECARROLL SWIFT Department of Chemistry, Swarthmore College, Swarthmore, Pennsylvania Recezued J u l y 2.2, 1960
The kinetics of the oxidation of Mo(V) by hydroxylamine have been studied in hydrochloric acid. In concentrated (12.0 M )acid there is no reaction. In 3.0 M acid there is essentially complete reaction. 2Mo(V) XH80H+ = 2Mo(VI) "I+. The rate law for the reaction is d[Mo(V)]/dt = ~ [ M O ( V ) ] [ N H ~ ~ H + ]in[ H solutions + ] ~ 3.0 molar in chloride ion. The rate constant, k, is 4.0 X 10-6 a t 25' and 1.20 X lo-' a t 35" if concentrations are expressed in moles per liter and time in seconds. The energy of activation is 23 kcal. and the entropy of activation is 4.8 e.u. A side reaction which might have interfered with colorimetric analysis of Mo(V) was observed. A color forming reaction between Mo(V1) and hydroxylamine was examined but found not to be the side reaction in question. The side reaction apparently involves complex formation between hydroxylamine and Mo(V).
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Introduction A study of the molybdate-catalyzed reduction of hydroxylamine carried out in this Laboratory suggested to one of the authors the need for a study of the direct oxidation of Mo(V) by hydroxylamine. This reaction was postulated as a possible key step in the catalytic process. The rate of the reaction is too slow and its activation energy too high for it to be directly involved in the catalysis. However, the kinetics and mechanism of the reaction have proved to be of sufficient interest to constitute the subject of a separate report. We have examined the stoichiometry and kinetics of the oxidation of Mo(V) by hydroxylamine in 3.0 molar hydrochloric acid, examined the nature of an observed side reaction, and deduced a possible mechanism for the reaction. Experimental Reagent grade materials were used without further purification. Hydroxylamine solutions were prepared by dissolving hydroxylamine hydrochloride in hydrochloric acid so that the final concentration of chloride ion was 3.0 molar. Sodium molybdate was dissolved in 3 .O M hydrochloric acid, reduced to Mo(V) with mercury, and stored over mercury in an atmosphere of nitrogen. Nitrogen, scrubbed with chromous sulfate solution, was bubbled through reagent flasks containing Mo(V) whenever they were opened for sampling. Hydrochloric acid solutions were standardized with sodium carbonate, using methyl orange as indicator. Hydroxylamine solutions were standardized by the method of (1) Presented at Delaware Valley Regional Meeting of %heAmerican Chemical Society, Philadelphia, 1960.
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Bray, Simpson and McICenzie.2 ?do(V) solutions were standardized by diluting 40.fold in dilute sulfuric acid and titrating with permanganate and also by measuring their absorbance a t 410 mp with a Beckman D.U. Spectrophotometer. The rate of the reaction of Mo(V) with hydroxylamine was determined by observing the decrease in absorbance due to Mo(V) a t 410 mp. A secondary reaction appeared during the latter half of runs containing high excesses of hydroxylamine which produced an interfering yellow color. Therefore, a second analytical method was employed. Five ml. samples of reaction mixture were withdrawn and quenched in 200 ml. of a solution of ferric sulfate in dilute sulfuric acid. The solution was boiled, cooled and then titrated with standard potassium permanganate. The procedure is essentially that of Bray, Simpeon and hlcICenzie for hydroxylamine. However, in this case both the Mo(V) and hydroxylamine reduce ferric ions to ferrous ions. The permanganate titer is thus a measure of the sum of XIo(V) and hydroxylamine present. Two equivalents of ferrous ion are produced per mole of hydroxylamine and one per gram atom of Mo(V) present. The method was useful only when neither &lo( V) nor hydroxylamine was present in excess. The two analytical methods gave identical rate laws with rate constants that agreed within experimental errors. Only initial portions of colorimetric runs involving excess hydroxylamine were used. Qualitative observations were made on the oxidation of Mo(V) and the reduction of Mo(V1) by hydroxylamine in concentrated (12 M) hydrochloric acid by observing the change in absorbance of solutione due to grren,s monomeric4 Mo(V) a t 720 mfi. (2) Wm. C.Bray, M. E. Sirni)son and A. A. McIiennie, J. Am. Chem. Soc., 41, 1362 (1919). (3) C. F. Hiakey and V. W. Bleloche, ibid., 62, 1819 (1940); 63,
964 (1941). (4) L. Sacconi and R. Cini, ibid., 7 6 , 4239 (1954).
1921
1922
G. P. HAIGHT, JR.,.4ND ALICECARROLL SWIFT
Results Stoichiometry of the Oxidation of Mo(V) by Hydroxylamine.-Tables of oxidation potentials5 and other potentiometric data6 indicate that hydroxylamine is thermodynamically capable of reducing Mo(V1) as well as oxidizing Mo(V). It was imperative, therefore, to determine whether the oxidation of Mo(V) could be studied without interference from the reduction of Mo(V1) or not. Mixtures of Mo(V) and hydroxylamine in 3.0 M hydrochloric acid reacted as follows: a. If Mo(V) were present in excess, absorbance at 410 mp decreased until the number of moles of Mo(V) consumed was equal to twice the number of moles of hydroxylamine initially present. b. If hydroxylamine were present in excess, the absorbance of solutions a t 410 mp decreased until 50 to 80% of the Mo(V) had disappeared, then slowly increased due to a secondary reaction. The total l t o ( V )plus hydroxylamine concentration continued to decrease, however, until the reducing power of the solution toward ferric ions corresponded to a hydroxylamine value which would be left if all the Mo(V) had been consumed along with one mole of hydroxylamine for each two moles of Mo(V). It thus appeared that in the time needed to observe the reaction, the only significant oxidationreduction process to occur was the reaction 2Mo(V) + NHzOH 4 2Mo(VI) NHa (I) Moreover, solutions of n'Io(V1) and hydroxylamine in 3.0 M hydrochloric acid and in water showed no change in reducing power toward ferric ions over a period of five days despite the fact that complicated color changes were observed. Thus, except for the secondary color forming reaction which does not involve oxidation-reduction, reaction I appears to be the only process occurring during the time of observation under the conditions employed. Kinetics of the Oxidation of MOW) by H y droxylamine.-Runs which were analyzed by decreasing absorbance a t 410 mp gave the following results. a. For runs in which [Mo(V)] = 2[NH20H], the change in absorbance with time followed the law
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where A is the measured absorbance, E is the molar absorbance of Mo(V) in 3.0 M hydrochloric acid at 410 mp, and kobs is the observed bimolecular rate constant. b. When a high excess of hydroxylamine was used, the change in absorbance with time initially followed the law dA /dt = k0bs14 [NH2OH] / E
(2)
I n case 2, kobs was lower than in case 1. This was later found to be caused by the decrease in hydrogen ion concentration resulting from the use of high hydroxylamine concentrations while main( 5 ) W. M. Latimer, "Oxidation Potentials," 2nd Ed., PrenticeHall, Inc.. New York, N. Y., 1952. ( 6 ) N. H. Furman and K. Murray, J. Bm. Chem. SOC.,68, 1689 (1936)
1'01. 65
taining a constant chloride ion concentration of 3.0. Although the results in each case agreed with the rate law d[Mo(V)l/dt = kobs[Mo(V)I [PU'H~OHI
(3)
the change in absorption spectrum a t the end of runs containing a large excess of hydroxylamine led us to use the titrimetric procedure outlined in the Experimental portion of the paper. The results obtained by this procedure were analyzed as follows. Since both hydroxylamine and Mo(V) reduce ferric ion to ferrous ion, the concentration of reducing equivalents ( T ) is given by r = [Mo(V)]
+ P[NHZOH]
For the case in which Mo[ (V)] find that
=
drjdt = 2 d[Mo(V)]/dt =
2 [NH20H],we
e2 4
( 4)
if the rate law given by equation 3 is correct. For the case in which [Mo(V)] # 2[ KH*OH],we define b
2[SH2OH]
-
[Mo(V)]
where b is a constant during any given run, but may vary from run to run. The quantity b may be either positive or negative. The rate law of equation 3 requires that aridt
= kobs(T
- b)2
4
( 5)
The integrated forms of equations 4 and 5 give straight line plots for the functions of r us. time in the cases where they apply. The values of kobs calculated from the slopes of the lines are in agreement with the values obtained spectrophotometrically within limits of experimental error. Effect of Changing Hydrogen Ion Concentration on kob,.-In order to account for lorn results in kobs obtained with high concentrations of hydroxylamine, the hydrogen ion concentration was varied between 1.5 and 3.0 molar by substitution of lithium chloride for hydrochloric acid. NOfurther variation in the medium of reaction was attempted. Mo(V) species are very sensitive to changes in hydrogen ion or chloride ion concentrations.2 Over the narrow range of hydrogen ion concentration employed kobs was found to be directly proportional to the square of the hydrogen ion concentration. Thus the rate law finally obtained is d [ M ~ ( V ) l / d t= ~ [ ~ ~ o ( ~ ~ ) ] [ N I ~ z O H(6) ][H+]~
Effect of Changing Temperature on Reaction Rates.-Results obtained a t 25 and 35" using each analytical method are summarized in Tables I and 11. Temperature control was somewhat better for runs analyzed by the titration procedure, the results of which are to be preferred. The activation energy for the reaction as determined from the results at two temperatures is 23.0 kcal. and the entropy of activation is 4.8 e.u. The activation energy was checked by performing a run a t 19.5'. us. The rate constant found was 1.G X 1.50 X 10+ calculated. Effect of Reaction Products on Reaction Rates.Seither sodium molybdate nor ammonium ion in concentrations comparable to those of r\lo(V)
Nov., 1961
KINETICSOF
THE
OXIDATION OF MOLYBDENUM(V) BY HYDROXYLAMINE
1923
TABLE I RESULTSOF RUNSAT 25' Run number
[Mo(V)lo,* [NHa+OH]o, mole/l.
rnole/l.
[H+
mole)?
k X 106, Method 1.P mole-3 of sec.-1 analysisb
0.010 0.1 2.90 3.64 A .010 .25 2.75 3.39 A .010 .05 2.95 4.05 A .010 .005 2.99 3.18 A .017 .433 2.57 3.48 -4 .036 .030 2.97 4.00 T .050 .025 2.95 3.78 T .040 .020 2.96 4.08 T .040 .020 2.96 4.13 T k = :3.75 f 0.30 X 10-6 1.3 mole-3 sec.-l IC = 4.00 f 0.20 X 10-6 for runs 6-9 a Measured as though Mo(V) were monomeric. b A for light absorbance (initial slopes where "%OH wm in excess); T for titration (most reliable). T control poor 0.3 to0.5' below 25".
1 2 3 4 5 6 7 8 9
1.8 1.6
1.4 1.2 6
:
-n, 1.0
w
2
0.8
0.6
TABLE I1 RESULTSOF RUNSAT 35" Run
i I o ~ , V ) l o , ~[KHzOHIo,
['mole/l.
moleil.
[H+]a, mole/l.
k X IO', 1.8 mole-3 Analysis sec. -1
1 0.040 0.020 2.96 1.16 T 2 ,050 .025 2.50 1.23 T 3 .050 .025 2.00 1.21 T 4 .050 .025 1.50 1.35 T 5 ,010 .250 2.75 1.20 A 6 ,020 .lo0 2.90 1.14 A 7 .010 .050 2.95 1.39 A 8 .020 .020 2.98 1.06 A 9 .010 .010 2.99 1.11 A a See Table I. 12 = 1.20 2= 0.19 X 10-*l.3mol.-3 see.-'. E, = 23 kcal., S* = 4.8e.u.
0.4
0.2
400 450 500 Wave length (mp). Fig. 1.-Absorption curves for 0.01 M Mo(V) in 3.0 M hydrochloric acid (A), and for 0.05 M Mo(V) plus 0.25 M hydroxylamine in 3.0 M hydrochloric acid after 23 hours (B). 350
the reaction under study must be such as to weaken the K-0 bond since the activation energy of the reaction is only 23.0 kcal. We therefore suggest a mechanism involving an equilibrium step in which hydroxylamine displaces an hydroxide ion (or water) in the coordination sphere of dimeric Mo(V). The complex formed then breaks up with the aid of hydrogen ions to form Mo(V1) and ammonia. Comparing this mechanism with
or hydroxylamine produced any measurable effect on the rate of reaction. Mechanism of Oxidation of Mo(V) by Hydroxylamine. Nature of Reactants in 3.0 M Hydrochloric Acid.-Hydroxylamine is certainly completely protonat,ed to give hydroxylammonium ion in 3.0 M acid. The species of Mo(V) is not known. It is a di:imagnetic dimer3 probably containing K oxygen or hydroxide bridges between the two (NoV)2-OH "BOH' ( MoV)2-O-NH3 HzO molybdenum at'oms. Compounds prepared by (11) James and Wardlaw7 from dilute hydrochloric acid k' indicate that there are probably two chloride ions (MoY-ONHa 2Hf +(Mo"')-OH~ NH3 (111) per molybdenum atom in the complex species. the empirical rate law, it is evident that the The rest of the octahedral coordination sphere probably is filled out with water molecules and empirical rate constant k is equal to the product of hydroxide ions. However, a mechanism that is K and k', which are the equilibrium constant and reasonable can be proposed without exact knowl- rate constant for the first and second steps in the mechanism, respectively. The ammonia and Moedge of the structure of the Mo(V) dimer. (VI) species resulting from the break up of the actiFormation of the Activated Complex.-Hydroxylamine, like water and hydrogen peroxide, vated complex will react rapidly with the solvent to probably can coordinate through its oxygen atom achieve their final equilibrium forms- ammonium with molybdenum in its higher oxidation states. ion and some unidentified chloromolybdate species. Such coordination to an atom of high positive The fact that neither ammonium ion nor molybcharge should serve to weaken the N-0 bond. T h e date ion affects the reaction rate indicates that single bond energy for N-0 is estimated to be over these final equilibrium forms must be produced in 30 kcal. per mole from t'he N-N and 0-0 single steps subsequent to the break up of the activated bond energies and the electronegat,ivity values complex. The role indicated for hydrogen ion is given by Pauling.8 The activated complex for only one of many possible and indistinguishable roles. ( 7 ) R . G. J8.mes and W. Wardlaw, J . Chem. SOC.,2145 (1927). Reactions of Mo(V) and Mo(V1) with H y ( 8 ) L. Pauling, "The Nature of the Chemical Bond," 3rd Ed., Cordroxylamine in 12 M Hydrochloric Acid.-For ne11 University Press. Ithaca. N. P., 1960.
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G. P. HAIGHT,JIZ.,A N D ALICECARROLL SWIFT
1924
Vol. 65
Mo(V1) and hydroxylamine is not responsib!e for the interference in the kinetic studies. The failure of the reaction of Mo(V1) and hydroxylamine to account for the color-giving curve B in Fig. 1, forced us to consider the possible formation of a more intensely colored Mo (V) species than that normally observed in 3.0 M hydrochloric acid. The form of the absorption curve B in Fig. 1 is very similar to that observed for hlo(V) in 4.0 M hydrochloric acid.3 It appears that hydroxylamine may have a similar effect to hydrochloric acid in displacing water or hydroxide from the coordination sphere of Mo(V). We believe M00C16-NHaOH+ 1 7 N o O C ~ O N H+~ H + + C1- that a strong enhancement of the peak at 420 mp occurs when hydrochloric acid causes breaking of (IV an oxygen bridge in the Mo(V) dimer while leavfar to the left by the high concentrations of hy- ing a second bridge intact. If hydroxylamine drogen ion and chloride ion. I n addition it is could accomplish the same effect as hydrochloric probably easier for dimeric Mo(V) to undergo acid, but a t a much slower rate, the species giving direct divalent oxidation by hydroxylamine than curve B in Fig. 1 would be accounted for. This for monomeric Mo(V) to undergo univalent oxi- explanation is very speculative on present evidence, dation with the production of radicals such as being offered merely as an hypothesis to explain NH3+or "2. the similar effects of hydroxylamine and hydroNature of the Side Reaction Interfering with chloric acid on the spectrum of Mo(V). This Colorimetric Detection of Mo(V) .-Figure 1 shows explanation would allow a required distinction the change in absorption spectrum accompanying between reactions leading to the species giving the appearance of an interfering yellow color dur- curve B and reaction I1 leading to formation of the ing the latter portion of runs containing excess activated complex. I n reaction I1 hydroxylamine hydroxylamine. The Mo(V) concentration at replaces a peripheral oxygen in the Mo(V) dimer, the time of observation of curve B would be less while in the react)ion giving the species responsible than 1.0 millimolar according to the observed re- for curve B of Fig. 1 hydroxylamine becomes action rates. bound a t a site left vacant by the breaking of an Two possible origins of species giving rise to oxygen bridge, the latter process being very slow. I n conclusion, the reaction between Mo(V) and curve 13 were investigated qualitatively. The fact that Mo(V) was being destroyed by oxidation sug- hydroxylamine to give Mo(V1) and ammonia gested that the new color might result from a re- (reaction I) has been shown to proceed quantiaction between Mo(V1) and hydroxylamine. Such tatively though at a very slow rate in 3.0 ICI hya reaction producing a yellow color was indeed drochloric acid. No reaction occurs in 12.0 M found to occur. It was too slow to cause inter- hydrochloric acid. The kinetics of reaction I and ference with most of the kinetic observations and a possible mechanism have been described. A involved no oxidation of hydroxylamine or reduc- side reaction producing a change in the Mo(V) tion of lCIo(V1) for several days. In addition, spectrum was observed and shown not to be the the spectrum of the yellow color produced showed same as a color forming reaction between Mo(V1) only general absorption increasing toward the ultra- and hydroxylamine. Acknowledgment.-The authors wish to thank violet, with no peak at 420 mp as observed in curve B. Furthermore, color formation between Mo(V1) the Faculty Research Committee of Swarthmore and hydroxylamine is inhibited by high concentra- College for a grant, and the Office of Ordnance Retions of hydroxylamine rather than enhanced. search, U. s. Army, for use of the spectrophoIt is clear that this color forming reaction between tometer.
purposes of comparison reactions were attempted in concentrated hydrochloric acid where mixtures of Mo(V) and hydroxylamine showed no decrease in absorbance due to Mo(V) a t 720 mp over a period of one week. On the other hand mixtures of hydroxylamine and Mo(V1) in concentrated hydrochloric acid slowly became green indicating formation of Mo(V). The latter reaction was by no means quantitative, however, even after standing for several weeks. In concentrated hydrochloric acid hfo(V) probably is the ion MoOCls--. An equilibrium similar to reaction (IV) would be forced
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