Proc. Indian Acad. Sci. (Chem. Sci.), Vol. 91, Number 5, October 1982, pp. 371-375.
~) Printed in India.
Kinetics of TI(III) oxidation of hydroxylamine hydrochloride in aqueous sulphuric acid
V A N G A L U R S S R 1 N I V A S A N * and N V E N K A T A S U B R A M A N I A N Vivekananda College, Mylapore, Madras 600 004, India
MS reoeived 22 December 1981 ;revised 25 June t982
Abstract. The kinetics of TI(tID oxidatiort of hyd,roxylamine hydrochloride in 1"0 M~ H~SO4 at a fixed chloride concentration has been investigated to make a formal comparison of the observed rate with that of h.ydrazine sulphate. The reaction exhibits total second order kinetics--first order in each reactant. The rate of the reaction depends inversely on the first power of [H +] and [CV] suggesting the possible reactive species as TIOH 2+. To account for the stoiehiometry of the reaction [TI(HI)]: [H,,NOH.HCI)= t'5 : t and tire products of the reaction, HNO~ and N.,O, two rcaction schemes have been proposed.
Keywords. TI(III)-HzNOH.HCI reaction ; T1OH :E -active spccics;
Ks = 3"3 ;< 10 -5 IVI~ 2
1. Introduction
T h o u g h the kinetics of TI([[[) oxidation of hydrazine sulphate has been investi- gated in aqueous sulphuric acid by Srinivasan a n d V e n k a t a s u b r a m a n i a n (1970 ; 1977), the kinetics of Tl([[i) oxidation of hydroxylamine has not been investi- gated so far. This p a p e r has, therefore, been undertaken with a view to c o m p a r e the reactivity towards these two similar reducing agents.
2. Results and discussion
2.1. Dependence o f rate on Tl(III) attd H.,NOH. HCl concentrations
The kinetics of Tl(II1) oxidation of hydroxylamine hydroehloride were studied in I M H , S Q at 30 ~ C. F r o m the rate dependence on the concentration of TI(III) and H ~ N O H . H C 1 (table l) it has been found that the reaction exhibits total second order kinetics--first order with respect to [Tl(lll)] and first order with respect to [HaN . O H . H C I ] . As T I ( I I I ) is k n o w n to complex with C I - (Halver- son et al 1956) total concentration of CI- has been adjusted to be 0"025 M by adding necessary a m o u n t of NaCl.
* To whom correspondence should be made.
371 P. (A)--t
372 Vangalul S Srinivasan and N Venkatasubramanial~
Table 1.* Dependence of rate on [TI([[I[)] ann [I-I~NOI-I.~C|].
[KffSO,] = 0' t0 ~t ; [aaSO4] = I'0 M ; Temp. : 30 ~ C
[TI([II)] M "< l0 :~ [IfaNOH.HCl] 1V[ • 10~ ka x l0 t litres mole -~ see~X
2"1 0"50 4"2
2"0 1"00 4"1
2"0 1"50 3"9
2"0 2"00 4" 1
3"0 2-00 4"0
4"0 2-0O 4"1
* Total e:lloritte ion collc:ntration is ketJt at 0'025 M by adding required amount of NaCI ill each ease.
2.2. Dependence on acid concentration
The reaction rate is decreased by increasing sulphuric acid concentration from 0"50 M to 2"0 M only (table 2). The plot o f tog k.,, versus log [acid] is linear with a slope equal to --1 and this is similar to what was observed in the TI(III) oxidation of hydrazine sulphate (Srinivasan and Venkatasubramanian 1970).
2.3. Effect o f added chloride ion y
With increasing concentration of chloride ions rate decreases (table 3) and a plot of log Ir verstts log [CI-] is linear with a slope very nearly equal to --1.
2.4. Temperature influence
The effect of temperature on the reaction rate was studied between 30 ~ and 50~ (table 4) and from the plot of log lq versus 1/T, energy of activation has been calculated. The thermodynamic parameters derived therefrom are summa- rised in table 4.
2.5. A comparative rate picture of hydrazine and hydroxylamine kz • 10 ~ (N~Ha) = 52 litre mole -t sec -1 and /q • 10 ~- ( N H . , O H . H C I ) = 0.158 1 mole -1 sec -t at 30~ under identical conditions) clearly indicates that hydroxylamine gets oxidised about 3.5 x 10" times slower than hydrazine sulphate. This is probably due to the lower reducing property o f hydroxylamine wherein the removal of hydrogen from nitrogen may be difficult due to the - I effect of the - O H group attached to the nitrogen.
3. Mechanism of the oxidation of hydroxylamine by TI ([II) 3.1. Nature o f active TI(III) species
In the presence of added chloride ion, TI([II) forms a complex, TICI 2 ~, which may get hydroiysed under the reaction conditions, according to the equilibrium
Kinetics o f T1 (HI) oxidation
Table 2.* 1),~p~.udc~t~c of r~dc oll acid t:onc, cn(ratio)).
[Ho.NOH.HCI] ~ 5"0 • 1 0 - a M ; [TI(III)] = 2 " 0 "< 10-:~M ; Temp. : 3 0 ~
i
[HISO,] M ka X 10* litres mole-X see --a
0"50 7"1 1"00 4"2 1"50 3"2
2"00 1-86
* Total ionic strength is m a i n t a i n e d at 2"0 by adding suitable amount of KHSO4.
373
Table 3. Effect of added chloride i o n o n the rate of oxidation.
[H~,NOH.HCI] ---- 5 ' 0 • 10-3M ; [H..SO,] = I ' 0 M ;
[TI(III)] = 2"0 • 10 -a M ;~ [Na.oSO,] = 0 ' 2 0 M ; Temp. : 30 ~ C
10 '~ • [CI-] M k. z • 1041itres mole -1 see-1
5"0 15.8
10-0 6-7
20 4-0
25 2.0
Table 4. Temperature dependence a n d thermodynamic parameters.
[ H ~ N O H . H C I ] = 5"0 • 10 -'~ M ; [H~SO,] = l ' 0 M ; (TI(III)) = 2"0 x 10 -'~ M ; [Na,.SO~] = 0" 20 M
kz • 10 '~ litres mole -a see -1
30 ~ 40 ~ 50 ~
I" 60 5" 1 36
E a kcal[mole 30
/ X H * kcals/mole 29
/ k S * eal/mole/deg + 2 5
374 Vangalur S Srinivasan and N Venkataxubramanian
TICI~+ + H~O ~-TIOH 2+ + H + + C1- (1)
The hydrolytic constant,
]r
= [ T I O H ~+ ] [H + ] [CI-][TICI2~ ] 9 (2)
Therefore,
[TIOH2+] = ~,
[TICI 2+][H,] [Cl] (3)
Under the reaction conditions, the observed inverse first order dependence on [H+] and [CI-], suggests the probable reactive species as TIOH ~"+.
From the rate expression,
d iT1 (III)] = kob~ IT1 (III)] [H~NOH.HCI] (4)
dt it follows that
- - d [TI (liD] k /r [TICI2+]
dt - ''" [-H U] ~ [Hydroxylamine hydrochloride] (5) whence
kJs [T1CI 2+]
ko,,~
=
-=[H~] [CI-J (6)It follows from (6) that a plot of ko,,~ versus 1/[CI-] should be linear, at a given [H § ], passing through the origin, with the slope equal to k2 Kh/[H +] where kz is the specific rate, in the absence of CI-, obtained by extrapolation method.
From the slope, Kh, evaluated comes to 3.3 • 10 -5 M 2 which is very close to the hydrolytic constant reported for TICI-"+ (Basolo and Pearson 1967).
3.2. The oxidation products, in the hydroxylamine--Tl(IlI) reaction, are N,O and HNO2 (Wiberg 1965) and the reaction has stoichiometry of [TI(III)] : [H2N OH.HCI] = 1.5 : 1. A mechanism consistent with the experimentally observed rate-law
- - d [TI(III)]
dt = k2 [TRill)] [H=NOH.HCI]
could be the following :
H~N-O-H + H* ,H'N-O,*-H.
H H H
TI 3§ ,H-N_6_ H StOW )TI" * N--O " H 2 H +
H I ~ H t
H ~ . ~ H
H20:r~ N=~_ H,....~ f a s t _ H / U . N = U - .Tt§ § H TI3§
H " - ~ _ N = ~ _ H f a s t ~- O = N _ O _ H §
2H §
H / SCHEME I
Kinetics o f 7"I(11I) oxidatiol~ 375 According to scheme 1, [TI(III)] : [Hydroxylamine] = 2 : 1.
N~O on the other hand may be picturised as follows : I-I~ ,
H \ N - O - H - 9 N - O - H
/ / !
H H H
The formation of
H'-N.(~
SlOW *TL 3 . 9 - H - 9 H - N = O H * T I * § 2H +
H H
H
-N=OH
9 f a s t , H N = O - + H "H-,,.N = 0 fast
N = OH - I ~ = ~ - = H - N - O H - - - - ~ ' H 2 0 * N 2 0
SCHEME II
The stoichiometry according to scheme 2 is 1 : 1. As the experimentally observed stoichiometry is [TI(I[I)] : [H~NOH.HCI] = 1.5 : 1, it will be consistent with both the above schemes. Since the products of the reaction are also reducing agents, there may be more uptake of oxidant during stoichiometric studies and hence stoichiometry tended to vary with time and acidity of the medium.
4. Experimental
Thallic oxide (BDH) was used as such and this had 99% purity as evidenced by iodometric estimation of its solution in sulphuric acid. Sulphuric acid was standardised after suitable dilution using standard carbonate free, N a O H solution.
Hydroxylamine hydrochloride (BDn) was used after two recrystallisations. The course of the reaction was followed by pipetting out 5 ml aliquots of the reaction mixture at various intervals and quenching in iodate free K I solution and esti- mating the iodine liberated using standard thiosulphate to a starch end point.
The specific rate (k2) has been evaluated using the integerated rate equation and velocity constants are reproducible within + 5 ~ .
References
Basolo F and Pearson R G 1967 Mechanism o f Inorganic reactions (New York : John Wiley) Halverson H N and Halpern J 1956 J. Am. Chem. Soc. 78 5562
Srinivasan V S and Venkatasubramanian N 1970 Curr. ScL 39 p. 254 Srinivasan V S and Venkatasubramanian N 1977 Indian J. Chem. A15 791
Wiberg K B 1965 Oxidation in Organic chemistry (New York and London : Academic Press) p. 27
