V. K. JINDAL, M. C. AGRAWAL, AND S. P. MUSHRAN
188
Oxidation of Hydrazine by Alkaline Ferricyanide
in Water-methanol Mixtures
V. K. J in d a l , M. C. A g r a w a l , and S. P. M u sh r a n
Department of Chemistry, University of Allahabad, Allahabad, India
(Z. N atu rforsd i. 25
b, 18 8 — 190 [197 0] ; eingegangen am 2. Septem b er 1969)
Kinetics of the oxidation of hydrazine by ferricyanide was investigated in water-methanol mix­
tures using several buffer solutions. The reaction showed first order dependence in both hydrazine
and ferricyanide. The order with respect to hydroxide ion concentration was zero. Increase in con­
centration of methanol had a retarding influence on the rate while the addition of neutral salts
showed a specific ion effect. The energy and entropy of activation were calculated as 12.3
kcals. mole- 1 and —20.8 cals. deg-1 mole-1 respectively. A suitable mechanism has been proposed
which suggests the primary rate determining reaction between N2H4 and Fe(CN)63e. Nitrogen was
found to be the product of the reaction.
Hydrazine and its derivatives have been exten­
sively used as reducing titrants in both acidic and
alkaline m edia.1 The principal product of the oxi­
dation is nitrogen. A little work has been done on
the kinetics of the oxidation of hydrazine in alkaline
solutions 2. In an acidic medium detailed kinetics of
the oxidation of hydrazine by iron (III) has been
studied by H ig g in s o n and WRIGHT3. Velocity of
the oxidation of hydrazine by ferricyanide was stud­
ied by G ilbert 4 at a pH of about 6.
Oxidation of hydrazine by ferricyanide is very
fast in an alkaline medium. In a medium of 2.5 to
5% KOH hydrazine can be directly titrated against
ferricyanide at 70° 5 where the stoichiometry was
found to be 1: 4.0. In presence of methanol, the rate
of the reaction was considerably slowed down. In
the present investigation, the results of the kinetics
of oxidation of hydrazine by alkaline ferricyanide
in 50% methanol have been recorded and sub­
sequently used for the formulation of a suitable
mechanism.
Experimental
Aqueous solutions of hydrazine sulphate are stable
even for a period of twro months and therefore a stock
solution was prepared from a recrystallised sample of
hydrazine sulphate (A. R., B. D. H .). Potassium ferri­
cyanide solution was prepared by dissolving weighed
amount of the AnalaR sample of the reagent. All other
solutions wTere prepared from the reagents of analyti­
cal grade and their concentrations were determined by
appropriate methods.
1 A. B
, J. V
, and J. Z
, Chemist-Analyst 52,
56 [1963].
2 W. C. E. H
, The Chem. Soc. [London], Spl. publi­
cation No. 10, pp. 95.
erka
u l t e r in
ig g in s o n
yka
Several buffer solutions were prepared by mixing
suitable amounts of sodium carbonate and bicarbonate
(0.2 ) and final pH was adjusted in 50% methanol.
The following table gives the pH of the buffers employed
in aqueous and 50% methanolic solutions.
m
CO,20 a + HC03q a
pH in 50%
pH
MeOH
[ml]
[ml]
9.90
4.0
46.0
9.2
10.80
13.0
37.0
9.5
19.5
30.5
9.7
11.20
25.0
25.0
9.9
11.30
30.0
20.0
10.1
11.38
35.5
14.5
10.3
11.45
Table I. Influence of Methanol on pH. a Total volume was
made to 200 mis and 10% of the buffer was used in the
investigations.
pH measurements were made on a Leeds and
Northrup type direct reading pH-meter using glasselectrode. Bidistilled water was used to prepare all the
solutions and diluting where necessary.
The kinetics of the reaction wTere followed by esti­
mating the amount of unreacted ferricyanide colorimetrically using Klett-Summerson Photoelectric Colori­
meter with blue filter No. 42 (transmission 400 —450
m/v). The absorption cell was chilled before adding the
reaction mixture and readings were taken within 10
seconds.
Results
The kinetics of the oxidation of hydrazine bv
ferricyanide was investigated at several concentra­
tions of the oxidising and reducing agent. It was
observed that the rate of disappearance of ferriC. E.
1551.
E. C. G
J. V
3 W.
4
5
H ig g in so n
il b e r t ,
u l t e r in
and
P. W
r ig h t
.
J. chem. Soc. 1955,
Z. physik. Chem. A, 142. 139 [1929].
and J. Z y k a , Chem. Listy 48. 1762 [1954].
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OXIDATION OF HYDRAZINE BY ALKALINE FERRICYANIDE
Expt.
No. a
[Fe(CN)639]
M x 104
1
2
3
4
5
6
7
8
9
10
3.2
3.6
4.0
4.4
4.0
4.0
4.0
4.0
4.0
4.0
[Hydrazine]
M x 103
4.0
4.0
4.0
4.0
1.0
2.0
3.0
4.0
7.0
9.0
k t x103 b
189
k2 c = k1/
[min-4 ]
[Hydrazine]
[/•mole-1 min-1 ]
63.6
66.8
66.0
65.0
38.7
74.4
116
149
235
299
15.9
16.7
16.5
16.2
38.7
37.2
38.6
37.4
33.6
33.3
Table II. Effect of Reactants Concentration. a Methanol = 50 %, pH = 11.3, Temp. = 25° and [NaClOJ = 0.1 m, only
in expts. Nos. 5 to 10. b First order constants in ferricyanide. c Second order rate constants.
cyanide follows first order dependence at all concen­
/cj x 103
pH
/l*2
[min-1 ]
[/•mole-1 min-1 ]
trations of hydrazine (Fig. 1). However, the con­
centration of hydrazine increased the first order con­
9.90
51.6
12.9
10.80
57.6
14.4
stants in ferricyanide almost linearly showing that
11.20
58.5
14.6
the reaction is also of first order in hydrazine (Table
11.30
65.9
16.5
II). The reaction is, therefore, of second order
11.38
66.8
16.7
18.8
11.45
75.3
whose constants k.2 have been calculated by dividing
the pseudo-first order constants in ferricyanide by Table III. Effect of pH. [N2H4] = 4 x 1 0 - 3 m , [Fe(CN)63e]
= 4 x 10- 4 m , Methanol = 50% and Temp. 25°.
the concentration of hydrazine.
It is observed from the pH study that a large
variation in the alkalinity of the reaction mixture
causes a slight increase in the rate constant. It is
therefore concluded that the reaction has an insignifi­
cant pH effect.
Effect o f Neutral Salts
Influence of several neutral salts on the rate of the
reaction was studied and it was observed that though
the addition of different salts usually enhanced the
reaction rate (Table IV), the effect was singularly
specific.
Fig. 1. Plot of log a/a—x vs time, [N2H4] as 1.0, 2.0, 3.0,
4.0, 7.0 and 9.0 x l 0 ~ 3 m in I, II, III, IV, V and VI respec­
tively.
The reactions were studied in presence of buffer
and 0 .1 m NaC104 (Experiments 5 — 10 of Table
II), in order to avoid any variations in the reaction
rate due to ionic strength and pH.
Effect o f pH
The effect of pH was studied between the range
9.9 to 11.45 using several sodium carbonate-bicarbonate buffers (Table II I).
Salt
[^1
[/•mole-1 min-J ]
[M]
None added
0.0
16.5
0.02 KC1
0.02
28.1
0.04 KC1
0.04
46.5
0.03
19.1
0.01 Na,S04
0.06
22.0
0.02 Na,S04
0.03 Na.,S04
0.09
25.6
0.02
18.9
0.02 NaC104
0.04
22.7
0.04 NaC104
0.06
0.06 NaC104
27.1
0.08 NaC104
0.08
30.0
Table IV. Effect of Neutral Salts [N,H4] = 4 x 10 - 3 m ,
[Fe(CN)630] = 4 x 1 0 - 4 m, pH = 11.3, Methanol = 50%
and Temp. 25°.
It is evident from the above table that the accel­
erating effect of neutral salts is not due to a change
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OXIDATION OF HYDRAZINE BY ALKALINE FERRICYANIDE
190
in ionic strength but appears to be a specific salt Cahn and P owell 6, which is further oxidised to
effect. K® ions have more pronounced effect than nitrogen in several subsequent fast steps as follows:
f Qcf
Na® ions while SO420 ions have a retarding in­ N2H3+ Fe(CN)63® —
N2H2+ Fe(CN)640.
(2)
fluence.
2 N2H3 — N2H2+ N2H4.
(3)
Effect of Dielectric Constant of the Medium
Increase in concentration of methanol or a de­
crease in dielectric constant of the medium shows a
marked inhibition on the rate of oxidation of hydra­
zine by ferricyanide (Table V).
Methanol
D o5o
k2
[Z• mole-1 min-1 ]
60.18
27.1
40
55.59
16.5
50
23.0
53.29
55
11.1
50.99
60
46.40
9.0
70
Table V. Effect of Dielectric Constant (Conditions same as in
Table IV).
[%]
Other Effects
Addition of ferrocyanide retards the rate of the
reaction (Table V I). Rise in temperature increases
the rate and the temperature coefficient has been
found to be 1.87. The reaction was studied at several
temperatures (Table V I), wherefrom the energy
and entropy of activation were calculated as 12.3
kcals. mole“ 1 and — 20.8 cals. deg-1 mole-1 re­
spectively.
Mechanism
The independence of the reaction rate on hydrox­
ide ion concentration predicts that N2H4 , rather than
N2H5®, is the main reacting species. As the reaction
shows first order dependence in both hydrazine and
ferricyanide, the primary rate determining step may
be written as follows:
N2H4+ Fe (CN) 63® ^ N2H3+ Fe (CN) 64®. (1)
The formation of N2H3 intermediate, during the
oxidation of hydrazine has also been suggested by
[Fe(C N))64°]
Mx l O 4
[/-mole-1 min-1 ]
k.
Temperature
[°C]
0.0
1.6
3.2
6.4
8.0
16.5
14.6
13.3
12.5
11.3
15
20
25
30
35
6 J. W. C a h n and R. E. P o w ell . J. A m er. chem . Soc. 76,
2568 [1954].
7 M . C. A g r a w a l and S. P . M u s h r a n , J. p h ysic. C hem . 72,
1497 [1968],
N2H2
N2 .
(4)
The stoichiometry of the reaction agrees well with
the above mechanism.
Now as step (1) is slow and rate-determining the
reaction would have first order dependence in both
hydrazine and ferricyanide and due to its reversible
nature it would have significant retarding influence
of ferrocyanide. Our experimental observations are
in accordance with these conclusions (Tables II
and VI).
The rate determining step (1) involves the inter­
action between an uncharged molecule and a nega­
tively charged ion and therefore should correspond
to a positive dielectric effect and a negative entropy
change 7. This has been found to be true from the
kinetic data recorded earlier.
The rate of oxidation of hydrazine by ferricyanide
in an alkaline medium is greatly influenced by the
addition of inert ions but the effect has been found
to be exclusively specific. It was not, therefore, pos­
sible to separate out the influence of ionic strength
on the rate of the reaction, which should have been
insignificant according to proposed mechanism. How­
ever, it has been ascertained that there are specific
ions like Na® and K® which enhance the rate whilst
ions like S 0 42t retard the rate of reaction.
It is, therefore, concluded that during the oxida­
tion of hydrazine by alkaline ferricyanide in 50?o
methanol, neutral hydrazine molecule (N2H4) is
attacked by ferricyanide, due to which the reaction
unlike other oxidations by ferricyanide8’9, shows
zero order dependence in hydroxide ion.
The authors wish to thank Council of Scientific and
Industrial Research, New Delhi and State C.S.I.R..
Lucknow, for fellowships to MCA and VKJ.
k2
[/•mole-1 min- ! ]
8.9
12.7
Table VI. Effect of Ferrocyanide and
Temperature (Conditions same as in
Table IV).
8 U. S. M e h r o t r a , M . C. A g r a w a l , an d S. P. M u s h r a n , J.
physic. Chem. 73. 1996 [1969].
9 M . C. A g r a w a l , V. K. J i n d a l , and S. P. M u s h r a n , J. in­
org. nuclear Chem., in press.
öl '.
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