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Kinetics of the oxidation of hydrazine by chromium(VI)
L inorg, nucL Chem., 1976, Vol. 38, pp. 549-552. Pergamon Press. Printed in Great Britain
KINETICS OF THE OXIDATION OF HYDRAZINE BY CHROMIUM(VI) KALYAN K. SEN GUPTA,* SHIPRA SEN GUPTA and HRIDAY RANJAN CHATTERJEE Department of Chemistry, Jadavpur University, Calcutta-32, India
(First received 1August 1974;in revisedform 10December 1974) Abstract--Kinetic studies of the oxidation of hydrazine by chromic acid were made under various experimental conditions. The enthalpy and entropy of activation have been calculated to be 10.5-+0.5 kcal mole-1 and -24.5 +-1.0 (e.u.) respectively and an attempt has been made to correlate activation parameters with those of the chromic acid oxidations of some phosphorus(Ill) and arsenic(Ill) compounds. The data suggest that oxidation follows formation of a complex between N2H5+ and the reacting species of chromium(Vl) in the rate determining step. The decomposition of the complex by a fast step gives the reaction products. Various mechanistic possibilities have been suggested.
chromium(VI)) and the products have negligible absorptions. Figure 1 shows the plot of the log of the optical density of the reaction mixture against time with substrate in excess and illustrates first order dependence of rate with respect to oxidant. The slopes obtained from these plots give different values of k, which were reproducible to within -+3%.
INTRODUCTION
THE reaction of hydrazine with several oxidising agents such as manganese(VII)[1], vanadium(V)[2], cobalt(III) [3], manganese(III) [4,5], iron(III)[5, 6] and copper(II)[7] have been studied. Several other oxidants like perdisulphate [8], thallium(III) [9] and cerium(IV) [10] have been used to study the kinetics of the oxidation of hydrazine. It was found that in almost all cases the oxidation product consists of a mixture of nitrogen and ammonia especially when one electron transfer metal ions are used as oxidants. Recently it has been suggested that[l l, 12] hydrazine can be quantitatively oxidized to nitrogen by chromium(VI) and a potentiometric method for the estimation of hydrazine by chromium(VI) has been developed. The reaction has been represented stoichiometrically by the equation,
c~
4 HCrO4- + 3 N2H5+ + 13 H + + 8 H 2 0 ~ 4 [Cr(H20)6] +3 +3N2.
._1
They have suggested that nitrogen is formed by the decomposition of a complex or complexes [12] formed by the interaction of hydrazine with chromium(VI). However, the details of the mechanism of the hydrazine-chromium(VI) reaction are not yet understood. The present investigation was therefore undertaken to elucidate the reaction mechanism of the oxidation of hydrazine by chromic acid. -u.~ F
EXPERIMENTAL
0
Reagents. B.D.H. (AnalaR) and E. Merck (G.R.) grade reagents were used throughout. A stock hydrazine solution of strength 10 2M was prepared and stored in a refrigerator. The materials were standardised by the usual procedures. Ethanol was purified twice by distillation. Kinetic measurements. The rate of decrease of gross chromium(VI) was followed spectrophotometrically in a Beckman DU model spectrophotometer using a cell of path length 1 or 0.1 cm depending upon the concentration of chromium(VI) used. Sodium acetate-acetic acid buffer was used throughout and ionic strength was held constant by the addition of sodium perchlorate where necessary. The temperature of the cell compartment of the spectrophotometer was kept constant by circulating water from a thermostat. The reactants previously equilibriated to bath temperature were mixed externally and immediately transferred to the cell and optical densities were noted at 350 nm at definite time intervals. At this wavelength, all the reactants (except *To whom all correspondence should be sent.
I
4
i
I
8
t,
i
I
12
3
I
16
I
20
rain
Fig. I. Pseudo first order plots at 39-5°C and at constant chromium(VI) concentrations of 3.30 x 10-" M. 1, 2.50 × 10-~ M; 2, 2.25 × 10-3 M; 3, 2.0 x 10-3 M; 4, 1.5 × 10-3 M; 5, 1.0 x 10-3 M. 1-5 denote substrate concentrations. RESULTS
Effect of reactant concentrations. The first order rate constants have been calculated at various initial chromium(VI) concentrations and at I = 1.0 M. The rate is not directly proportional to gross chromium(VI) (Table 1). This suggests that HCrO4- is the reacting species [13] and the values of k~[Cr(VI)]/[HCrO4-] at various chromium(VI) concentrations have been calculated, k~oris seen to be constant. The details of the calculations of HCrO4concentrations have been described elsewhere[14,15]. Rates were also measured in the
549
K. K. S. GUPTAet al.
550
Table 1. Effect of oxidant on the reaction rate. [N2H5÷]= 2.0 × 10-3 M, pH = 5.48, Temp. = 33.5°C,I = 1.0 M No. [Cr(VI)]xl04M
k, × 103 (sec-~) [HCrO,-]×10"M
0.833 1.666 3.332 4.168
2 3 4
2.03 1.96 1.85 1.74
k~o~× 103 (sec ~)
1-613 3.161 6.064 7.454
1.049 1.033 1.016 0.995
concentration ranges 1.0 x 10-3-3.0× 10-3M substrate and 3.30 × 10-4 M oxidant respectively. The plots of kl vs various substrate concentrations indicate that the order with respect to substrate is also unity (Fig. 2).
4.0
3.0
× 2.0 2.
DISCUSSION The acid catalyzed oxidation of hydrazine by chromium(VI) may take place in different ways. In aqueous solution hydrazine exists[19] as both hydrazinium (+1) ion, N2Hs* and hydrazinium (+2) ion, N2H62+. The equilibrium constants of eqns (1) and (2) are 8.5 × 10-7 and 9.0 × 10-t6 respectively at 25°C.
1.0
I I'0 [Substrate],
I 2"0 x [0 3 M
I 30
Fig. 2. Variationof pseudo first order rate constants with substrate concentrations at different temperatures. [Chromium(VI)]= 3-30x 10-3 M; pH = 5.48; 1,30°C;2, 35'5°C;3, 39.5°C;4, 44'3°C.
Effect o f p H o n the rate. The reactions were studied at various hydrogen ion concentrations adjusted by the addition of buffer, but the ionic strength of each reaction mixture was adjusted to 0.20 M. The results have been recorded in Table 2. The plot of log k~ vs pH is linear with a slope which is significantly less than one (0.4) unlike previous studies [14--16]. The apparent order (0.4) in H ÷ is much more likely to be consistent with HCrO4- being the important oxidant because at this pH, CrO42- is substantially protonated, so that [HCrO4-] would be dependent upon a low power of [H +] as observed. H2CrO4 is present in near zero amount; thus if H~CrO4 were involved, [H2CrO4] would be virtually directly proportional to [H+] [HCrO4-], i.e. the reaction would be first order in H +. A similar explanation may also apply to the noninvolvement of N2H62+. Table 2. Effect of pH on the reaction rate. [N2Hs+]= 2.0 x 10-~ M, [Cr(VI)]= 3.33 × 10-4 M, I = 0.20 M, Temp. = 39°C pH k~× 103(sec-~)
Effect of salt and solvent on the rate. The effect of addition of a salt such as sodium perchlorate on the rate has been studied. However, when [salt] < 10-5 M, the salt had no influence on the rate but above this value the salt had a retarding influence and the retardation is significant when [salt] > 10-~ M. The dielectric constant was varied by the addition of ethanol (v/v) to the reaction mixture. It will be noted that ethanol does not react with chromium(VI) at the low hydrogen ion concentration of ~10-6M. The rate increases as the dielectric constant fails. The results have been recorded in Table 3. Influence of temperature on the rate. The specific reaction rates have been calculated from the relation, k,p = k~/[substrate]. The values are 0.839, 1.01, 1.35 and 1.89 (lmole -1 sec -~) at 30, 35.5, 39.5 and 44.3°C respectively. The heat of activation has been calculated to be 10.5 ---0.5 (kcal mole -l) from the gradient of log ksp vs 1] T plot. The other activation parameters like entropy and free energy of activation have been computed to be -24.5-+ 1.0 (e.u.) and 18.1 -+0.6 kcal mole -~ respectively. The details of the calculations of other activation parameters have been described earlier[14]. It will be interesting to compare activation parameters of the present reaction with those for the oxidations of some phosphorus(llI)[14, 15] and arsenic(Ill) compounds[16] by chromic acid (Table 4). The heats as well as entropies of activation are not very different suggesting that some similar mechanism involving complex formation may be operative.
6.49 6.10 0~768 1.35
5.60 5.48 5.18 5.00 2.23 2.39 2.60 3.45
N2H4 + H20 ~ N2H5+ + OH-
(1)
N~H5+ + H:O ~ NEH6z* + OH-.
(2)
The second equilibrium is of little importance at this low acidity. Consequently, N:H5 + may be considered as the reacting species in the reaction of hydrazine with chromium(VI) leading to the formation of nitrogen. ESR measurement has indicated[17,18] that protonated amines are attacked by electrophiles at a point furthest from the site of protonation and the electrophile would attack the unprotonated N-atom in N2H~+. The electron transfer from hydrazine to chromium(VI) possibly takes place through the formation of complex between N2H5+ and HCrO4-. If the equilibrium constant for a 1:1 complex is high, a plot of l/kt vs I/[N2Hs*] would give a straight line making an intercept and the rate expression would be -d[Cr(VI)] dt
kK[Cr(VI)I[N2H,+] 1 + K[N2H5+]
(3)
The rate expression, on the contrary, at constant hydrogen ion concentration is given by -d[Cr(VI)] dt
k,p [Cr(VI)][N:Hs+].
(4)
The formation of 1:1 complex appears to be the rate
Kinetics of the oxidation of hydrazine by chromium(VI)
551
Table 3. Effect of salt and solvent on the reaction rate. [N~Hd] = 2.0 × 10-~ M, [Cr(VI)] = 3.33 × 10-" M, Temp. = 31°C, pH = 5.48 [NaCIO.] in M k~ x 103 (sec-') % Ethanol (v/v) kt x 103 (sec ')
0 1.58 0 1.58
0-12 1.45 5 2.78
0.20 1.15 10 2.82
Table 4. Activation parameters of the oxidation reactions of nitrogen(Ill), phosphorus(Ill) and arsenic(III) compounds by chromium(Vl) No. Reducingsubstrate
AH* (kcal mole ')
-AS + (e.u.)
10.5 10.4 10.3 9' 1 9"4
24.5 39,3 40.4 43"5 39"9
0.40 0.60 0.959 0.728 15 3.30
0.80 0.703 20 3.38
1.0 0.650 ---
Scheme 3
N2H5+ + Cr(IV)~ N2H3' + Cr(III) + 2H*
(12)
N:H3. + Cr(VI) ~ N2H3÷ + Cr(V)
(13)
Cr(IV) + Cr(VI)~ 2Cr(V) 1 2 3 4
N2H5+ H3PO3TM PhPOzH2 ~5
5
~'AS20316
~ As20316
a ~ Low chromium(Vl) and high arsenic(III) concentrations. b --*High chromium(Vl) and low arsenic(Ill) concentrations. determining step, for which a small equilibrium constant is to be expected. The complex decomposes by a fast step to give a protonated diimide which is known to be a powerful reducing agent[19]. This is subsequently oxidized either by chromium(VI) or chromium(V) to nitrogen. The formation of 1:2 or 2:1 complex as N2Hs+: Cr(VI) seems unlikely in view of the dependence of the rate on the substrate and reactive chromium(VI) species. However, even if these complexes are formed, the concentrations must be negligible. 0
H
H
-O--~r--OH + H_N_N+ I I 0
H
H
s~ow O H H
II I I -O--Cr-- .N.--N*--H + H20 O H O H H II I I .... -O_Cr~N_N+_H ~HCrO3- + NH NH2+. O
(5)
(6)
H
Assuming that all subsequent reaction steps are fast, three alternative schemes may be proposed 2 (HCrO4- + N~H~+ ,
slow
"X + H20)
2 (X-~ NzH3+ + Cr(IV)).
(11)
3 (N,,H3+ + Cr(V) ~ N,_+ Cr(III) + 3H+).
(5)
(10)
The various redox couples (relevant at p H = 0 ) mentioned below should enable us to isolate the reaction scheme which is involved in the reaction. Redox couple Cr(IV) ~ Cr(III) Cr(V) -* Cr(III) Cr(VI) ~ Cr(III) N2H4 ~ N2
>
E ° (V) 1.75 1.75 1.36 0'23.
Chromium(IV) from the decomposition of the complex might be converted to chromium(III) by reacting rapidly with NEH5+ producing a free radical according to step (12) or become involved in reaction with chromium(VI) producing chromium(V) according to step (11). In addition to these reactions, chromium(IV) may also disproportionate to give chromium(V) and chromium(III) according to step (8). The absence of free radicals (failed to polymerise acrylamide) indicates that the reaction of chromium(IV) with N2Hd is unlikely. However, the redox potential of N2H4~ N2 is so small that chromium(IV) would prefer to react with N2H5+ before it reacts with chromium(VI). Consequently Schemes 1 and 2 appear to be less likely. Again, it has been suggested by Higginson et a/.[2,3] that vanadium(V) behaves mainly as a two electron transfer oxidant where nitrogen is the main product of oxidation of hydrazine. The oxidation of the same substrate by vanadium(V) to ammonia and nitrogen is insignificant and takes place only when metal ions behave as one electron transfer oxidants[3-6]. The absence of ammonium ions in the solution further indicates that chromium(VI) behaves as a two electron oxidant and not as one electron oxidant. Acknowledgements--Thanks are due to Prof. M. N. Das for
(6)
encouragement during the progress of the work. Thanks are also due to the referee for constructive suggestions.
(7) (8)
REFERENCES 1. A. W. Browne and F. F. Shetterly, J. Am. Chem. Soc. 31,221 (1909). 2. W. C. E. Higginsonand D. Sutton, Jr. Chem. Soc. 14020953). 3. W. C. E. Higginson, D. Sutton and P. Wright, 3. Chem. Soc. 1380 (1953). 4. W. A. Waters and J. S. Littler, Oxidation in Organic Chemist~ (Edited by K. B. Wilberg), Part A, p. 196. Academic Press, New York (1965). 5. E.J. Cuy and W. C. Bray, J. Am. Chem. Soc. 46,1786 (1924). 6. E. J. Cuy, 3. Am. Chem. Soc. 46, 1310 0924). 7. H. T. S. Britton and M. Konigstein, J. Chem. Soc. 673 (1940).
Scheme 1
2 (N2H3+ + Cr(VI)-* N2 + Cr(IV) + 3H÷) 4 Cr(IV)-* 2Cr(V) + 2Cr(III) N2Hd + Cr(V) --*N2H3+ + Cr(III) + 2H* N2H3+ + Cr(V)~ N2 + Cr(lII) + 3H÷.
(9) (10)
Scheme 2
2 (Cr(lV) + Cr(VI) ~ 2 Cr(V)) NzH,+ + Cr(V)-* N2H3 + + Cr(IIl) + 2H+ 3 (N~H~+ + Cr(V) ~ N2 + Cr(IlI) + 3H+).
(1 l) (9) (10)
552
K. K. S. GUPTA et al.
8. A. P. Bhargava, R. Swaroop and Y. K. Gupta, J. Chem. Soc. A, 2183 (1970). 9. K. S. Gupta and Y. K. Gupta, Indian J. Chem. 11, 1285 (1973). 10. K. K. Sen Gupta and K. Chakravarti, J. Proc. Inst. Chem. 36, 97 (1964). 11. S. Syamsunder and T. K. S. Murthy, Indian J. Chem. 11,669 (1973). 12. M. T. Beck and D. A. Durham, 3. Inorg. Nucl. Chem. 32, 1971 (1970). 13. K. B. Wiberg and T. J. Mill, J. Am. Chem. Soc. 80, 3022 (1958). 14. K. K. Sen Gupta, J. K. Chakladar and A. K. Chatterjee, J. Inorg. Nucl. Chem. 35, 901 (1973).
15. K. K. Sen ¢3upta and J. K. Chakladar, 3. Chem. Soc. Perkin II, 929 (1973). 16. K. K. Sen Gupta and J. K. Chakladar, J. C. S. Dalton trans. 226 (1974). 17. W. T. Dixon, R. O. C. Norman and A. L. Buley, J. Chem. Soc. 3625 (1964). 18. W. T. Dixon, R. O. C. Norman, J. Chem. Soc. 90, 1619 (1968). 19. W. L. Jolly, The Inorganic Chemistry of Nitrogen p. 61. Benjamin, New York (1964).
KINETICS OF THE OXIDATION OF HYDRAZINE BY CHROMIUM(VI) KALYAN K. SEN GUPTA,* SHIPRA SEN GUPTA and HRIDAY RANJAN CHATTERJEE Department of Chemistry, Jadavpur University, Calcutta-32, India
(First received 1August 1974;in revisedform 10December 1974) Abstract--Kinetic studies of the oxidation of hydrazine by chromic acid were made under various experimental conditions. The enthalpy and entropy of activation have been calculated to be 10.5-+0.5 kcal mole-1 and -24.5 +-1.0 (e.u.) respectively and an attempt has been made to correlate activation parameters with those of the chromic acid oxidations of some phosphorus(Ill) and arsenic(Ill) compounds. The data suggest that oxidation follows formation of a complex between N2H5+ and the reacting species of chromium(Vl) in the rate determining step. The decomposition of the complex by a fast step gives the reaction products. Various mechanistic possibilities have been suggested.
chromium(VI)) and the products have negligible absorptions. Figure 1 shows the plot of the log of the optical density of the reaction mixture against time with substrate in excess and illustrates first order dependence of rate with respect to oxidant. The slopes obtained from these plots give different values of k, which were reproducible to within -+3%.
INTRODUCTION
THE reaction of hydrazine with several oxidising agents such as manganese(VII)[1], vanadium(V)[2], cobalt(III) [3], manganese(III) [4,5], iron(III)[5, 6] and copper(II)[7] have been studied. Several other oxidants like perdisulphate [8], thallium(III) [9] and cerium(IV) [10] have been used to study the kinetics of the oxidation of hydrazine. It was found that in almost all cases the oxidation product consists of a mixture of nitrogen and ammonia especially when one electron transfer metal ions are used as oxidants. Recently it has been suggested that[l l, 12] hydrazine can be quantitatively oxidized to nitrogen by chromium(VI) and a potentiometric method for the estimation of hydrazine by chromium(VI) has been developed. The reaction has been represented stoichiometrically by the equation,
c~
4 HCrO4- + 3 N2H5+ + 13 H + + 8 H 2 0 ~ 4 [Cr(H20)6] +3 +3N2.
._1
They have suggested that nitrogen is formed by the decomposition of a complex or complexes [12] formed by the interaction of hydrazine with chromium(VI). However, the details of the mechanism of the hydrazine-chromium(VI) reaction are not yet understood. The present investigation was therefore undertaken to elucidate the reaction mechanism of the oxidation of hydrazine by chromic acid. -u.~ F
EXPERIMENTAL
0
Reagents. B.D.H. (AnalaR) and E. Merck (G.R.) grade reagents were used throughout. A stock hydrazine solution of strength 10 2M was prepared and stored in a refrigerator. The materials were standardised by the usual procedures. Ethanol was purified twice by distillation. Kinetic measurements. The rate of decrease of gross chromium(VI) was followed spectrophotometrically in a Beckman DU model spectrophotometer using a cell of path length 1 or 0.1 cm depending upon the concentration of chromium(VI) used. Sodium acetate-acetic acid buffer was used throughout and ionic strength was held constant by the addition of sodium perchlorate where necessary. The temperature of the cell compartment of the spectrophotometer was kept constant by circulating water from a thermostat. The reactants previously equilibriated to bath temperature were mixed externally and immediately transferred to the cell and optical densities were noted at 350 nm at definite time intervals. At this wavelength, all the reactants (except *To whom all correspondence should be sent.
I
4
i
I
8
t,
i
I
12
3
I
16
I
20
rain
Fig. I. Pseudo first order plots at 39-5°C and at constant chromium(VI) concentrations of 3.30 x 10-" M. 1, 2.50 × 10-~ M; 2, 2.25 × 10-3 M; 3, 2.0 x 10-3 M; 4, 1.5 × 10-3 M; 5, 1.0 x 10-3 M. 1-5 denote substrate concentrations. RESULTS
Effect of reactant concentrations. The first order rate constants have been calculated at various initial chromium(VI) concentrations and at I = 1.0 M. The rate is not directly proportional to gross chromium(VI) (Table 1). This suggests that HCrO4- is the reacting species [13] and the values of k~[Cr(VI)]/[HCrO4-] at various chromium(VI) concentrations have been calculated, k~oris seen to be constant. The details of the calculations of HCrO4concentrations have been described elsewhere[14,15]. Rates were also measured in the
549
K. K. S. GUPTAet al.
550
Table 1. Effect of oxidant on the reaction rate. [N2H5÷]= 2.0 × 10-3 M, pH = 5.48, Temp. = 33.5°C,I = 1.0 M No. [Cr(VI)]xl04M
k, × 103 (sec-~) [HCrO,-]×10"M
0.833 1.666 3.332 4.168
2 3 4
2.03 1.96 1.85 1.74
k~o~× 103 (sec ~)
1-613 3.161 6.064 7.454
1.049 1.033 1.016 0.995
concentration ranges 1.0 x 10-3-3.0× 10-3M substrate and 3.30 × 10-4 M oxidant respectively. The plots of kl vs various substrate concentrations indicate that the order with respect to substrate is also unity (Fig. 2).
4.0
3.0
× 2.0 2.
DISCUSSION The acid catalyzed oxidation of hydrazine by chromium(VI) may take place in different ways. In aqueous solution hydrazine exists[19] as both hydrazinium (+1) ion, N2Hs* and hydrazinium (+2) ion, N2H62+. The equilibrium constants of eqns (1) and (2) are 8.5 × 10-7 and 9.0 × 10-t6 respectively at 25°C.
1.0
I I'0 [Substrate],
I 2"0 x [0 3 M
I 30
Fig. 2. Variationof pseudo first order rate constants with substrate concentrations at different temperatures. [Chromium(VI)]= 3-30x 10-3 M; pH = 5.48; 1,30°C;2, 35'5°C;3, 39.5°C;4, 44'3°C.
Effect o f p H o n the rate. The reactions were studied at various hydrogen ion concentrations adjusted by the addition of buffer, but the ionic strength of each reaction mixture was adjusted to 0.20 M. The results have been recorded in Table 2. The plot of log k~ vs pH is linear with a slope which is significantly less than one (0.4) unlike previous studies [14--16]. The apparent order (0.4) in H ÷ is much more likely to be consistent with HCrO4- being the important oxidant because at this pH, CrO42- is substantially protonated, so that [HCrO4-] would be dependent upon a low power of [H +] as observed. H2CrO4 is present in near zero amount; thus if H~CrO4 were involved, [H2CrO4] would be virtually directly proportional to [H+] [HCrO4-], i.e. the reaction would be first order in H +. A similar explanation may also apply to the noninvolvement of N2H62+. Table 2. Effect of pH on the reaction rate. [N2Hs+]= 2.0 x 10-~ M, [Cr(VI)]= 3.33 × 10-4 M, I = 0.20 M, Temp. = 39°C pH k~× 103(sec-~)
Effect of salt and solvent on the rate. The effect of addition of a salt such as sodium perchlorate on the rate has been studied. However, when [salt] < 10-5 M, the salt had no influence on the rate but above this value the salt had a retarding influence and the retardation is significant when [salt] > 10-~ M. The dielectric constant was varied by the addition of ethanol (v/v) to the reaction mixture. It will be noted that ethanol does not react with chromium(VI) at the low hydrogen ion concentration of ~10-6M. The rate increases as the dielectric constant fails. The results have been recorded in Table 3. Influence of temperature on the rate. The specific reaction rates have been calculated from the relation, k,p = k~/[substrate]. The values are 0.839, 1.01, 1.35 and 1.89 (lmole -1 sec -~) at 30, 35.5, 39.5 and 44.3°C respectively. The heat of activation has been calculated to be 10.5 ---0.5 (kcal mole -l) from the gradient of log ksp vs 1] T plot. The other activation parameters like entropy and free energy of activation have been computed to be -24.5-+ 1.0 (e.u.) and 18.1 -+0.6 kcal mole -~ respectively. The details of the calculations of other activation parameters have been described earlier[14]. It will be interesting to compare activation parameters of the present reaction with those for the oxidations of some phosphorus(llI)[14, 15] and arsenic(Ill) compounds[16] by chromic acid (Table 4). The heats as well as entropies of activation are not very different suggesting that some similar mechanism involving complex formation may be operative.
6.49 6.10 0~768 1.35
5.60 5.48 5.18 5.00 2.23 2.39 2.60 3.45
N2H4 + H20 ~ N2H5+ + OH-
(1)
N~H5+ + H:O ~ NEH6z* + OH-.
(2)
The second equilibrium is of little importance at this low acidity. Consequently, N:H5 + may be considered as the reacting species in the reaction of hydrazine with chromium(VI) leading to the formation of nitrogen. ESR measurement has indicated[17,18] that protonated amines are attacked by electrophiles at a point furthest from the site of protonation and the electrophile would attack the unprotonated N-atom in N2H~+. The electron transfer from hydrazine to chromium(VI) possibly takes place through the formation of complex between N2H5+ and HCrO4-. If the equilibrium constant for a 1:1 complex is high, a plot of l/kt vs I/[N2Hs*] would give a straight line making an intercept and the rate expression would be -d[Cr(VI)] dt
kK[Cr(VI)I[N2H,+] 1 + K[N2H5+]
(3)
The rate expression, on the contrary, at constant hydrogen ion concentration is given by -d[Cr(VI)] dt
k,p [Cr(VI)][N:Hs+].
(4)
The formation of 1:1 complex appears to be the rate
Kinetics of the oxidation of hydrazine by chromium(VI)
551
Table 3. Effect of salt and solvent on the reaction rate. [N~Hd] = 2.0 × 10-~ M, [Cr(VI)] = 3.33 × 10-" M, Temp. = 31°C, pH = 5.48 [NaCIO.] in M k~ x 103 (sec-') % Ethanol (v/v) kt x 103 (sec ')
0 1.58 0 1.58
0-12 1.45 5 2.78
0.20 1.15 10 2.82
Table 4. Activation parameters of the oxidation reactions of nitrogen(Ill), phosphorus(Ill) and arsenic(III) compounds by chromium(Vl) No. Reducingsubstrate
AH* (kcal mole ')
-AS + (e.u.)
10.5 10.4 10.3 9' 1 9"4
24.5 39,3 40.4 43"5 39"9
0.40 0.60 0.959 0.728 15 3.30
0.80 0.703 20 3.38
1.0 0.650 ---
Scheme 3
N2H5+ + Cr(IV)~ N2H3' + Cr(III) + 2H*
(12)
N:H3. + Cr(VI) ~ N2H3÷ + Cr(V)
(13)
Cr(IV) + Cr(VI)~ 2Cr(V) 1 2 3 4
N2H5+ H3PO3TM PhPOzH2 ~5
5
~'AS20316
~ As20316
a ~ Low chromium(Vl) and high arsenic(III) concentrations. b --*High chromium(Vl) and low arsenic(Ill) concentrations. determining step, for which a small equilibrium constant is to be expected. The complex decomposes by a fast step to give a protonated diimide which is known to be a powerful reducing agent[19]. This is subsequently oxidized either by chromium(VI) or chromium(V) to nitrogen. The formation of 1:2 or 2:1 complex as N2Hs+: Cr(VI) seems unlikely in view of the dependence of the rate on the substrate and reactive chromium(VI) species. However, even if these complexes are formed, the concentrations must be negligible. 0
H
H
-O--~r--OH + H_N_N+ I I 0
H
H
s~ow O H H
II I I -O--Cr-- .N.--N*--H + H20 O H O H H II I I .... -O_Cr~N_N+_H ~HCrO3- + NH NH2+. O
(5)
(6)
H
Assuming that all subsequent reaction steps are fast, three alternative schemes may be proposed 2 (HCrO4- + N~H~+ ,
slow
"X + H20)
2 (X-~ NzH3+ + Cr(IV)).
(11)
3 (N,,H3+ + Cr(V) ~ N,_+ Cr(III) + 3H+).
(5)
(10)
The various redox couples (relevant at p H = 0 ) mentioned below should enable us to isolate the reaction scheme which is involved in the reaction. Redox couple Cr(IV) ~ Cr(III) Cr(V) -* Cr(III) Cr(VI) ~ Cr(III) N2H4 ~ N2
>
E ° (V) 1.75 1.75 1.36 0'23.
Chromium(IV) from the decomposition of the complex might be converted to chromium(III) by reacting rapidly with NEH5+ producing a free radical according to step (12) or become involved in reaction with chromium(VI) producing chromium(V) according to step (11). In addition to these reactions, chromium(IV) may also disproportionate to give chromium(V) and chromium(III) according to step (8). The absence of free radicals (failed to polymerise acrylamide) indicates that the reaction of chromium(IV) with N2Hd is unlikely. However, the redox potential of N2H4~ N2 is so small that chromium(IV) would prefer to react with N2H5+ before it reacts with chromium(VI). Consequently Schemes 1 and 2 appear to be less likely. Again, it has been suggested by Higginson et a/.[2,3] that vanadium(V) behaves mainly as a two electron transfer oxidant where nitrogen is the main product of oxidation of hydrazine. The oxidation of the same substrate by vanadium(V) to ammonia and nitrogen is insignificant and takes place only when metal ions behave as one electron transfer oxidants[3-6]. The absence of ammonium ions in the solution further indicates that chromium(VI) behaves as a two electron oxidant and not as one electron oxidant. Acknowledgements--Thanks are due to Prof. M. N. Das for
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encouragement during the progress of the work. Thanks are also due to the referee for constructive suggestions.
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REFERENCES 1. A. W. Browne and F. F. Shetterly, J. Am. Chem. Soc. 31,221 (1909). 2. W. C. E. Higginsonand D. Sutton, Jr. Chem. Soc. 14020953). 3. W. C. E. Higginson, D. Sutton and P. Wright, 3. Chem. Soc. 1380 (1953). 4. W. A. Waters and J. S. Littler, Oxidation in Organic Chemist~ (Edited by K. B. Wilberg), Part A, p. 196. Academic Press, New York (1965). 5. E.J. Cuy and W. C. Bray, J. Am. Chem. Soc. 46,1786 (1924). 6. E. J. Cuy, 3. Am. Chem. Soc. 46, 1310 0924). 7. H. T. S. Britton and M. Konigstein, J. Chem. Soc. 673 (1940).
Scheme 1
2 (N2H3+ + Cr(VI)-* N2 + Cr(IV) + 3H÷) 4 Cr(IV)-* 2Cr(V) + 2Cr(III) N2Hd + Cr(V) --*N2H3+ + Cr(III) + 2H* N2H3+ + Cr(V)~ N2 + Cr(lII) + 3H÷.
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Scheme 2
2 (Cr(lV) + Cr(VI) ~ 2 Cr(V)) NzH,+ + Cr(V)-* N2H3 + + Cr(IIl) + 2H+ 3 (N~H~+ + Cr(V) ~ N2 + Cr(IlI) + 3H+).
(1 l) (9) (10)
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K. K. S. GUPTA et al.
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