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Hypochlorite oxidation of cyanate under mildly alkaline conditions
Wat. Res. Vol. 21, No. 6, pp. 677-682, 1987 Printed in Great Britain. All rights reserved
0043-1354/87 $3.00+0.00 Copyright © 1987 Pergamon Journals Ltd
H Y P O C H L O R I T E O X I D A T I O N OF C Y A N A T E UNDER MILDLY ALKALINE CONDITIONS W. K. TEO and T. C. TAN Department of Chemical Engineering, National University of Singapore, Kent Ridge, Singapore 0511 (Received January 1986) Abstract--Kinetic study was carried out on the hypochlorite oxidation of cyanate, over the pH range of 8.6-10.0 and temperature range from 27 to 50°C. The decomposition reaction was found to be first order each with respect to cyanate, hypochlorite and hydrogen ion concentration. The experimental results indicate that the decomposition reaction may probably follow a series reaction in which the cyanate first hydrolyses to ammonium which then decomposes in alkaline hypochlorite solution to nitrogen. Under the range of pH and hypochlorite/cyanate molar ratio examined in this study, the stoichiometric ratio of hypochlorite to cyanate for the decomposition reaction is found to be approx. 1.5. The conversion of cyanate to nitrate is negligible compared with that to nitrogen. This study suggests that complete treatment of cyanide waste is best carried out in two stages at room temperature and in the presence of excess hypochlorite. The first stage involves the hypochlorite oxidation of cyanide to cyanate at pH 11 and the second stage involves the conversion of the cyanate to nitrogen and other stable inert products at a reduced pH of about 9. Key words---cyanate oxidation, cyanides, hypochlorite, kinetics
adverse effect on the natural environment. It is therefore desirable to further convert the cyanate to other stable inert end products before the effluent is discharged. The reaction of chlorine or hypochlorite on cyanate has received relatively little attention compared with that on cyanide. This reaction is apparently rather complex and its stoichiometric equation is generally approximated by equation (1)
NOMENCLATURE a = b= C = E = h= ki= -r~ R T t
= = = =
Order of reaction with respect to cyanate Order of reaction with respect to hypochlorite Concentration (moll -l) Activation energy (Jmol -~) Order of reaction with respect to hydrogen ion Reaction constant with respect to reactant i (12tool -2 s-i) Rate of decomposition of reactant i (mol 1-I s -1) Universal gas constant (J mol -t K -t) Reaction temperature (K) Reaction time (s)
2 C N O - + 3OC1- + H20 = 2HCO~- + 3C!- + N 2.
(1)
Subscripts A = cyanate B = hypochlorite H = hydrogen ion i = reactant A or B 0 = initial condition at t = 0.
This reaction has been studied in some details by Lister (1956) at high p H in the presence of sodium hydroxide. Measurable rates of cyanate oxidation were obtained under these conditions and at elevated temperatures from 50 to 75°C. However the relative amounts of cyanate and hypochlorite decomposed were found to deviate from the stoichiometric ratio given by equation (1). This discrepancy was attributed to a simultaneous parallel reaction as shown by equation (2).
INTRODUCTION The destruction of cyanides by chemical, electrochemical and biochemical methods generally results in the formation of cyanate. Depending on the process conditions, the cyanate may further undergo decomposition to other inert end products such as nitrogen and carbonates. Price et al. (1947) and Eden et al. (1950) showed experimentally that in the presence of excess chlorine or hypochlorite, cyanide oxidized rapidly to cyanogen chloride which then hydrolysed to give cyanate. Since cyanate is very much less toxic than cyanide and for economic reasons, industrial treatment of cyanide waste is usually carried out up to the cyanate stage. However cyanate is chemically unstable and may have an
C N O - + 4OC1- + H20 = NO~- + 4C1- + HCO~- + H ÷.
(2)
Lister showed that the oxidation reaction is first order with respect to cyanate and to hypochlorite concentration and inversely proportional to the hydroxide concentration. Kieszkowski (1968) and Teo et al. (1985) experimentally examined the conditions for the complete destruction of cyanide and cyanate. Teo et al. (1985) reported that at r o o m temperatures cyanate remained stable at high p H even in the 677
678
W . K . TEO and T. C. TAN
presence of excess hypochlorite. However at p H 9.5 and lower, the oxidation reaction was reasonably fast. This paper describes the experimental study on the kinetics of the hypochlorite oxidation of cyanate under conditions of practical interests, in particular, cyanate concentration level, pH and temperature. EXPERIMENTAL
Materials and analyses Analytical reagent grade chemicals, except for potassium cyanate, were used for the experimental study. Both potassium cyanate and cyanate produced by the hypochlorite oxidation of cyanide at high pH were used for the present study. Cyanide stock solution were prepared in 0.1 N NaOH solutions and standardized against silver nitrate standard solution with 5-(4-Dimethylamino-benzyliden)-rhodanin as indicator according to the modified Liebig titration method given by Standard Methods (APHA, 1980). Stock solutions of cyanate were prepared also in 0.1 N NaOH solution and stored in the refrigerator at 4°C. The analytical method developed by Shaw and Bordeaux (1955) and modified by Teo et al. (1985) for the analysis of ammonium and cyanate was used to standardize the cyanate stock solution and to monitor the progress of the cyanate oxidation reaction. The anlaytical procedure involved the acid hydrolysis of cyanate to ammonium ion which was then analysed by spectrophotometric measurement using Nessler's reagent. Chlorine was found to interfere with the analysis. Since the reaction product normally contained residual chlorine, the sample was first treated with sodium sulphite to remove all the chlorine before carrying out the acid hydrolysis. Concentrations of nitrate in the product mixtures were determined based on the reduction of nitrate to ammonia using Devarda's alloy in strong alkaline solution as described in Vogel (1978). Fresh stock solution of hypochlorite was prepared for each experimental run using sodium hypochlorite. The available chlorine contents of the stock solution and the reaction product were determined by the arsenite method which is described in details by Vogel (1978). Procedure Kinetic study on the decomposition of cyanate in the presence of hypochlorite was carried out in a 1 litre laboratory batch reactor maintained at a desired constant temperature in a thermostatically controlled water bath. The pH of the reaction mixture was continuously monitored by the pH electrode inserted into the reaction vessel. For the experiments using sodium cyanide as the starting material, the unbuffered cyanide solution at a pH of about 11 was quantitatively converted to cyanate in a known excess quantity of hypochlorite. Complete conversion to cyanate was usually obtained in less than about 10 min. The resulting cyanate solution was then buffered and made up to give the desired pH and solution volume. For those experiments using sodium cyanate solution directly, the solution was buffered and equilibrated at the given pH prior to the addition of hypochlorite. The kinetics of the reaction was studied by following the concentration change of the cyanate and hypochlorite as the reaction progressed. For some experimental runs, the amount of nitrogen gas evolved was measured continuously by a gas burette connected to the reaction flask. RESULTS AND DISCUSSION
Stability and oxidation o f cyanate The stability and oxidation of cyanate in hypo-
chlorite solution were experimentally examined under various operating conditions. Figures l and 2 show some typical experimental concentration-time curves of the various reactants and products in the oxidation process using sodium cyanide as the starting reactant. At p H 11 and above and in the presence of excess hypochlorite, cyanide was rapidly and quantitatively oxidized to cyanate. At these high pH, the cyanate formed remained fairly stable for several hours with negligible or no further decomposition even in the presence of excess hypochlorite as shown in Fig. 1 for the case of p H = 11. Measurable rates of the cyanate decomposition were only obtained at pH below about 10.5. Figure 2 shows the concentration-time plots for both reactants (cyanide and hypochlorite) and reaction products (cyanate and nitrogen gas) at p H = 9.1 and 27°C. An induction period generally prevailed prior to the liberation of nitrogen. This i n d u c t i o n time was observed to increase with increasing pH of the reaction mixture. The amounts of nitrogen gas liberated and the hypochlorite consumed during the oxidation of cyanate were found to differ from the values given by equation (I) for a corresponding amount of cyanate decomposed. The amounts of hypochlorite consumed were generally 5-12% higher while the nitrogen yield was generally lower than those predicted by equation (1). These may be attributed partly to the liberation of chlorine gas especially at higher temperature and in the presence of excess hypochlorite and partly to the parallel reaction leading to the formation of nitrate according to the reaction given by equation (2). Lister (1956) observed
? 16o
+--
0
4
x
6
3
x Z .c_
2
--O-O--O
•
8 ~_1
e-8
©'5
8 0
910~
10
+
t
130
cl
t
210
0
Time (min)
Fig. 1. Hypochlorite oxidation of cyanide at pH = l l.O and 27°C (CN- A, OCI- O, CNO [~, N 2 ©).
_J
16 b x
? o3 x z 2 ._c ¢
8
~
'~
o 0
10
50
90 Time (rain)
130
1200
Fig. 2. Hypochlorite oxidation of cyanide at pH = 9.1 and 27°C (CN- &, OC1 O, CNO- 15], N 2 ©, N O ; V).
Hypochlorite oxidation of cyanate Table 1, Effect of pH and hypochlorite/cyanatemolar ratio on the formation of nitrate at 27°C (OCI-/CNO-)0= 1.78 pH 8.7 9.2 9.5 10.0 10.5 NO~- x 103moll-I 0.52 0.45 0.32 0.20 0.15 pH = 9.15 (OCI-/CNO-)0 1.54 1.78 2.10 2.40 2.73 NO{ x 103moll 1 0.48 0.45 0.38 0.42 0.41 that the formation of nitrate was more prevalent in the presence of large excess of hypochlorite. The effects of pH and hypochlorite concentration on the formation of nitrate at 27°C were examined in this study and the results are tabulated in Table 1. Table 1 shows that reaction (2) leading to the formation of nitrate was more favourable at low pH and the effect of hypochlorite concentration was only marginal. Under the experimental conditions used in this study, the extent of decomposition of cyanate by reaction (2) would be relatively small compared with that observed by Lister (1956) at 65°C and in strong sodium hydroxide solutions. Most of Lister's (1956) experiments were carried out with a reactant ratio of (OC1-/CNO-) lower than the stoichiometric ratio given by reaction (1). No attempt to identify the presence of other products besides nitrogen and nitrate was indicated. In our present study, ammonium ion was always detected to be present in the product when the reactant ratio used was lower than the stoichiometric ratio. However, in the presence of excess hypochlorite, no ammonium ion was detected. Lister (1955) and Teo et al. (1979) reported that in the absence of hypochlorite, cyanate hydrolysed to ammonium ion very rapidly under acidic conditions but the rate became increasingly slower as pH increased to neutral and to mild alkaline conditions. These observations suggest that the hypochlorite oxidation of cyanate may probably proceed according to a series reaction by which the cyanate first hydrolyses according to reaction (3) to ammonium ion which is then oxidised by the hypochlorite to nitrogen gas according to reaction (4). C N O - + 2H20 + H ÷
OCI-
NH~ + HCO~-
(3)
2NH~- + 3OC1- + 2OH- = N 2 + 3C1- + H:O. (4) These would give an overall reaction as shown by equation (1). The significant increase in the rate of decomposition of cyanate in the presence of hypochlorite may be attributed to reaction (4) and/or to the catalytic effect of hypochlorite on reaction (3). The increase, especially at pH below 10, may also be brought about by the presence of hypochlorous acid, which is a much stronger oxidising agent than hypochlorite. Hypochlorous acid co-exists with hypochlorite according to the dissociation reaction as shown in equation (5). HOCI = H + + OCI-.
(5)
679
At 25°C, the dissociation constant is given as 3.4 x 10 -s (Cotton and Wilkinson, 1972). Lower pH favours the presence of hypochlorous acid which would therefore enhance the oxidation rate of cyanate. These experimental observations on the stability and oxidation of cyanate under various pH and (OC1-/CNO-) reactant ratios suggest that a complete treatment of cyanide waste to stable and inert end products is best carried out in two stages both in the presence of excess hypochlorite. The first stage involves the hypochlorite oxidation of cyanide to cyanate at pH about 11 and the second stage at a reduced pH of ,,-9 whereby the cyanate could be favourably oxidized to nitrogen and other stable products. In view of this proposed treatment scheme, subsequent kinetic study of the hypochlorite oxidation of cyanate was carried out under mild alkaline conditions at pH below 10.
Kinetics of hypochlorite oxidation of cyanate Effect of hypochlorite and cyanate concentrations. The kinetics of the oxidation of cyanate at pH = 9.15 and 27°C were examined over the range of initial molar ratio of (OCI-/CNO-) from 1.10 to 2.70. The ratio was varied by varying the hypochlorite concentration while maintaining the cyanate concentration at 7.14 x 10 -3 moll 1. Figures 3 and 4 show the concentration-time curves at various initial (OC1-/CNO-) molar ratios for cyanate and hypochlorite respectively. The decomposition rates of both the reactants were found to increase with increase in the initial reactant concentration ratio. Since the ratio of the experimental rate of decomposition of hypochlorite to that of cyanate was observed to differ by as much as 12% of the stoichiometric value, differential method of analysis of the kinetic data was preferred to the integral method. This would then avoid the necessity of having to choose a suitable average stoichiometric ratio of the reactants for the computation. The rates of decomposition of cyanate and hypochlorite were determined from the respective concentration-time curves given in Figs 3 and 4 at various fractional conversions of cyanate. The corresponding hypochlorite concentrations were also
8
7o
o<2 300
600 Time (s)
900
Fig. 3. Cyanate concentration-time curves at pH = 9.15 and 27°C and various initial reactants ratios [Cao/C^o= 1.10 (&), 1.54 (11), 1.78 (O), 2.10 (A) 2.40 (I-q), 2.73 (O)].
680
W.K. TEO and T. C. TAN Time (s) 240
120
360
480
16 14 12
~d~\ \ i
..alO ? o
\
o ~o
12 ~
o
~ 8
o
8.
o'2
x
I
0
300
I
600
Time (s)
g00
Fig. 4. Hypochlorite concentration-time curves at pH = 9.15 and 27°C and various initial reactants ratios [CBo/CAo= 1.10 (A), 1.54 (B), 1.78 (@), 2.10 (A) 2.40 (I-1), 2.73 (C))].
determined. Since the initial cyanate concentrations were the same at 7.14x 10-3moll -l for all the experimental runs, the corresponding cyanate concentration at a given fractional conversion of cyanate will be the same for all the runs. Figures 5 and 6 show the logarithmic plots of rates of decomposition of cyanate and hypochlorite against hypochlorite concentration for various constant fractional conversions of cyanate or constant cyanate concentrations. The plots are all found to be almost parallel straight lines giving a slope of approx. 1.0. The results therefore show that the rate of decomposition is first order with respect to hypochlorite concentration. The general kinetic expression may be written as ai bi h - ri = k, CACBC~
Figure 7 was then plotted accordingly for both the cyanate and hypochlorite. Parallel linear plots are obtained with a slope equal to 1.0 (a, = 1.0) for both reactants. The results therefore show that the reaction proceeds with first order kinetics with respect to cyanate and to hypochlorite. These were further confirmed by the graphical plots shown in Fig. 8 where the rates of decomposition were plotted against CACB. The plots are essentially linear giving a slope of 0.65 1mol-l s - 1 for - r A vs CACB and 0.3851mol-ls -I for - r e vs CAC~. This gives a ratio of --rB/--rA equal to 1.69 which is about 12.6% higher than the stoichiometric ratio given by equation (1). Effect o f p H . As discussed previously, the rate of cyanate decomposition increased with decreasing pH. At pH above about 10, a drastic reduction in the reaction rate was observed and the reaction practically ceased to occur at pH beyond 11. The significant enhancement of reaction rate at lower pH can most probably be attributed to the presence of HOCI. For this reason, only the data obtained at pH below 10 will be used for the kinetic analysis. The experimental data obtained for various pH were treated in the same manner as the preceeding section. The plots of ( - r , ) against CACB at various pH 10.C 8.° 6.0
i = reactant A or B - r i = rate of decomposition of the reactant i k, = reaction constant with respect to the reactant i CA, Ca, Ca = concentration of cyanate, hypochlorite and hydrogen ion respectively
i
2.0 o
1.C x 0.8
°.2
(6)
where
o/
4.0
i
,
681o
2'0
C B x l o "3 (M/L)
Fig. 5. ( - r ~ - C ~ plots at pH = 9.15, 27°C and at constant CA[CA x 103mo11-1;7.14 (O), 5.71 (@), 5.00 (A) 4.29 (V), 3.57 (UI)].
10.0 8.0
and a, b, h = order of reaction.
9 2.0 x
Since k~, CA and CH are constant, the plots in Figs 5 and 6 are described by
I 0.! 0.1
l o g ( - ri) = log(k, C~ C~) + b~log CB
(7)
0.4 / a J l L I 210 I 4 6 8 10 40 CB x 10 "~ (M/L)
where b~ is found to be 1.0 for both A and B. Rearranging equation (6) in the form shown by equation (8),
Fig. 6. ( - - r B ) - C B plots at p H = 9.15, 27°C and at constant
log(-r~/CB) = Iog(k~C~) + a, log CA.
3.57 (D)I.
(8)
C A [C A x l03 mol l-i; 7.14 (O), 5.71 (O), 5.00 (/k) 4.29 (V),
Hypochlorite oxidation of cyanate
b,
x
4.0
, ~ l~ ,
2.0 1.0
//o / .
0.8
~2- o.8 '~"
2.59 x 10 t3 and 3.80 x 1013 for the cyanate and hypochlorite respectively giving thus a ratio of --rB/--rA of 1.47, very close to the stoichiometric ratio of 1.5 given by reaction (1). This implies that on the overall average, the stoichiometry for the reaction is well approximated by equation (1). For the range of parameters examined in this study, reaction (1) therefore predominates over reaction (2) thus supporting the earlier observations and deductions. The overall rate expression for the decomposition reaction can now be written as
10.0 8.0 6.0
0
,A/I/ ~1 /
°
0.4
I
[
I
I
3
4
5
6 7 8910
I
681
I I
CA x l 0 -3 (M/L)
Fig. 7. (--rA/CB) and (--rs/CB) vs cA plots at pH=9.15, 27°C [(--rA/CB) O, (--rB/CB) I-q].
13 /-27030"X --rA=2.585 X lO e x p ~ - - - - ~ - ) C A C a C H \
(11)
/
and 10
- % = 1.5 (--rA) where the gas constant R = 8.314 J tool -l K -l. The orders of reaction with respect to the three reactants agree with those proposed by Lister (1956) in his study of the same reaction but at very high pH and moderately high temperature at 65°C.
6
b
4!
CONCLUSION
X
S
The decomposition of cyanate to its end products at room temperature was found to be negligible at pH
I
I
I
2o
4o
eo
I
I
6.0
I
80 lOO12O14O
4.0
t"~...("-- x 1 0 "e ( M ' / L ' ) vA
/
,.-- o
"~ 2.0
Fig. 8. (--rA) and (--ra) vs CACe plots at pH=9.15, 27°C [CBo/CAo= 1.10 (A), 1.54 (ll), 1.78 (O), 2.10 (A), 2.40 (fl)].
1.0 0.8
" ~ o.e (.~ 0.4
are linear. The values of k,CH or ~i are given by the slopes of the lines. Logarithmic plots of log ~biagainst the hydrogen ion concentration for both the reactants were shown in Fig. 9 and described by equation (9).
0.2 I
4
I
I
I
410
I
6 8 10
20
60
C H x l O -~° (M/L)
log ~ = log k~ + h~log C a.
(9)
The results show that the reaction is first order with respect to hydrogen ion. The reaction constants for cyanate and hypochlorite were found to be 5.85 x 109 and 9.50 x 10912 m o l - : s -l respectively giving a ratio of - rB/- rA equal to 1.62.
and (--rB/CACB) vs CH plots at 27°C [(--rA/CACs) 0, (--rB/CACB) 1"7].
(--rA/CACB)
8.0 6.0 4.0 ^
Effect of temperature
2.0
The reaction constants, k~, for the decomposition of cyanate and hypochlorite, were similarly obtained for the different reaction temperatures. For the effect of temperature on the reaction constants, Fig. 10 is plotted according to the Arrhenius expression lnk i=lnk~-
Fig. 9.
~
5 v
\ \ o -.
\ 1.0
o
os
\
._ 0 . 8
\
0.4
~-~ 0.2
I 2.9
for the reactants. Both the plots are parallel giving an activation energy of 27.0 kJ tool -I for the decomposition reaction. The frequency factors were
t 3.1
I
I 3.3
I
I 3.5
l I T x 10 -a (K -~)
Fig. 10. Reaction constant-temperature Arrhenius plot (CNO- C), OCI- ~).
682
W.K. "lEo and T. C. TAN
above 11 even in the presence of excess hypochlorite. However at pH ~< 9.5, the rates of decomposition were observed to be reasonably fast. The experimental results suggest that the decomposition most probably follows a series reaction C N O - + 2H20 = NH~ + CO 22NH~- + 3OC1- + 2 O H - = N 2 + 3C1- + H20. The increase in the decomposition rate in the presence of excess hypochlorite may be attributed to (a) catalytic effect of hypochlorite on the hydrolysis of cyanate to ammonium ion and (b) the removal of ammonium by a subsequent reaction to nitrogen and other products. The increase in the rate at lower pH in the alkaline region is mainly due to the increase in the concentration of a more powerful oxidising agent, hypochlorous acid at pH below 10. The experimental results yield a rate ratio of hypochlorite to cyanate of approx. 1.5. This supports the general acceptance of the approximate stoichiometric equation for the reaction, 2CNO
+ 3OC1- + H20 = 2HCO~- + 3C1- + N2
and also substantiates that under the present range of experimental conditions, the extent of decomposition to nitrate is small. The overall rate expression for the decomposition of cyanate is given by --rA(mOl 1-1 s -1) = 2.59 × 1013 {--27030\ × exp~)CACBCH
over the temperature range of 27-50°C and pH range of 8.6-9.5. The activation energy for the reaction is 27.0 kJ mol- 1. It may therefore be concluded that a complete cyanide treatment process at room temperature is
best carried out in two stages; both stages in the presence of excess hypochlorite. The first stage involves the conversion of cyanide to cyanate at pH 11 and the second stage from cyanate to inert and stable end products at pH 9. Acknowledgements--The authors gratefully acknowledge Ong Seng Eng and Yau Tai Yin for their assistance in part of the experimental work. REFERENCES
APHA (1980) Standard Methods for the Examination of Water and Wastewater, 15th edition. American Public Health Association, Washington, D.C. Cotton F. A. and Wilkinson G. (1972) Advanced Inorganic Chemistry, p. 569. Wiley, New York. Eden G. E., Hampson B. L. and Wheatland A. B. (1950) Destruction of cyanide in wastewaters by chlorination. J. Soc. chem. Ind. 69, 244-249. Eden G. E. and Wheatland A. B. (1950) Effect of temperature and pressure of hypochlorite on the rate of hydrolysis of cyanogen chloride in alkaline solution. J. Soc. chem. Ind. 69, 166-169. Kieszkowski M. (1968) Investigations on oxidation of cyanides in waste water by means of sodium hypochlorite. Pr. Inst. Mech. precyz., XV, 62-71. Lister M. W. (1955) Some observations on cyanic acid and cyanates. Can. J. Chem. 33, 426-440. Lister M. W. (1956) The reaction between cyanate and hypochlorite. Can. J. Chem. 34, 489-501. Price C. C., Larson T. E., Beck K. M., Harrington F. C., Smith L. C. and Stephanoff I. (1947) Hydrolysis and chlorinolysis of cyanogen chloride. J. Am. chem. Soc. 69, 1640-1644. Shaw W. H. R. and Bordeaux J. (1955) Semimicro method for the determination of cyanate ion in the presence of interfering substances. Analyt. Chem. 27, 136-138. Teo W. K. and Ng C. S. (1979) Optimum conditions in cyanide destruction using hypochlorite. Chemsia Conference, 16-19 January 1979, Singapore. Teo W. K., Ng C. S. and Ho W. Y. (1985) Cyanate determination in cyanate waste. Asia Chem. Conference, 8-I1 April 1985, Singapore. Vogel A. I. (1978) Quantitative Inorganic Analysis, 3rd edition, pp. 382-383. Longman, London.
0043-1354/87 $3.00+0.00 Copyright © 1987 Pergamon Journals Ltd
H Y P O C H L O R I T E O X I D A T I O N OF C Y A N A T E UNDER MILDLY ALKALINE CONDITIONS W. K. TEO and T. C. TAN Department of Chemical Engineering, National University of Singapore, Kent Ridge, Singapore 0511 (Received January 1986) Abstract--Kinetic study was carried out on the hypochlorite oxidation of cyanate, over the pH range of 8.6-10.0 and temperature range from 27 to 50°C. The decomposition reaction was found to be first order each with respect to cyanate, hypochlorite and hydrogen ion concentration. The experimental results indicate that the decomposition reaction may probably follow a series reaction in which the cyanate first hydrolyses to ammonium which then decomposes in alkaline hypochlorite solution to nitrogen. Under the range of pH and hypochlorite/cyanate molar ratio examined in this study, the stoichiometric ratio of hypochlorite to cyanate for the decomposition reaction is found to be approx. 1.5. The conversion of cyanate to nitrate is negligible compared with that to nitrogen. This study suggests that complete treatment of cyanide waste is best carried out in two stages at room temperature and in the presence of excess hypochlorite. The first stage involves the hypochlorite oxidation of cyanide to cyanate at pH 11 and the second stage involves the conversion of the cyanate to nitrogen and other stable inert products at a reduced pH of about 9. Key words---cyanate oxidation, cyanides, hypochlorite, kinetics
adverse effect on the natural environment. It is therefore desirable to further convert the cyanate to other stable inert end products before the effluent is discharged. The reaction of chlorine or hypochlorite on cyanate has received relatively little attention compared with that on cyanide. This reaction is apparently rather complex and its stoichiometric equation is generally approximated by equation (1)
NOMENCLATURE a = b= C = E = h= ki= -r~ R T t
= = = =
Order of reaction with respect to cyanate Order of reaction with respect to hypochlorite Concentration (moll -l) Activation energy (Jmol -~) Order of reaction with respect to hydrogen ion Reaction constant with respect to reactant i (12tool -2 s-i) Rate of decomposition of reactant i (mol 1-I s -1) Universal gas constant (J mol -t K -t) Reaction temperature (K) Reaction time (s)
2 C N O - + 3OC1- + H20 = 2HCO~- + 3C!- + N 2.
(1)
Subscripts A = cyanate B = hypochlorite H = hydrogen ion i = reactant A or B 0 = initial condition at t = 0.
This reaction has been studied in some details by Lister (1956) at high p H in the presence of sodium hydroxide. Measurable rates of cyanate oxidation were obtained under these conditions and at elevated temperatures from 50 to 75°C. However the relative amounts of cyanate and hypochlorite decomposed were found to deviate from the stoichiometric ratio given by equation (1). This discrepancy was attributed to a simultaneous parallel reaction as shown by equation (2).
INTRODUCTION The destruction of cyanides by chemical, electrochemical and biochemical methods generally results in the formation of cyanate. Depending on the process conditions, the cyanate may further undergo decomposition to other inert end products such as nitrogen and carbonates. Price et al. (1947) and Eden et al. (1950) showed experimentally that in the presence of excess chlorine or hypochlorite, cyanide oxidized rapidly to cyanogen chloride which then hydrolysed to give cyanate. Since cyanate is very much less toxic than cyanide and for economic reasons, industrial treatment of cyanide waste is usually carried out up to the cyanate stage. However cyanate is chemically unstable and may have an
C N O - + 4OC1- + H20 = NO~- + 4C1- + HCO~- + H ÷.
(2)
Lister showed that the oxidation reaction is first order with respect to cyanate and to hypochlorite concentration and inversely proportional to the hydroxide concentration. Kieszkowski (1968) and Teo et al. (1985) experimentally examined the conditions for the complete destruction of cyanide and cyanate. Teo et al. (1985) reported that at r o o m temperatures cyanate remained stable at high p H even in the 677
678
W . K . TEO and T. C. TAN
presence of excess hypochlorite. However at p H 9.5 and lower, the oxidation reaction was reasonably fast. This paper describes the experimental study on the kinetics of the hypochlorite oxidation of cyanate under conditions of practical interests, in particular, cyanate concentration level, pH and temperature. EXPERIMENTAL
Materials and analyses Analytical reagent grade chemicals, except for potassium cyanate, were used for the experimental study. Both potassium cyanate and cyanate produced by the hypochlorite oxidation of cyanide at high pH were used for the present study. Cyanide stock solution were prepared in 0.1 N NaOH solutions and standardized against silver nitrate standard solution with 5-(4-Dimethylamino-benzyliden)-rhodanin as indicator according to the modified Liebig titration method given by Standard Methods (APHA, 1980). Stock solutions of cyanate were prepared also in 0.1 N NaOH solution and stored in the refrigerator at 4°C. The analytical method developed by Shaw and Bordeaux (1955) and modified by Teo et al. (1985) for the analysis of ammonium and cyanate was used to standardize the cyanate stock solution and to monitor the progress of the cyanate oxidation reaction. The anlaytical procedure involved the acid hydrolysis of cyanate to ammonium ion which was then analysed by spectrophotometric measurement using Nessler's reagent. Chlorine was found to interfere with the analysis. Since the reaction product normally contained residual chlorine, the sample was first treated with sodium sulphite to remove all the chlorine before carrying out the acid hydrolysis. Concentrations of nitrate in the product mixtures were determined based on the reduction of nitrate to ammonia using Devarda's alloy in strong alkaline solution as described in Vogel (1978). Fresh stock solution of hypochlorite was prepared for each experimental run using sodium hypochlorite. The available chlorine contents of the stock solution and the reaction product were determined by the arsenite method which is described in details by Vogel (1978). Procedure Kinetic study on the decomposition of cyanate in the presence of hypochlorite was carried out in a 1 litre laboratory batch reactor maintained at a desired constant temperature in a thermostatically controlled water bath. The pH of the reaction mixture was continuously monitored by the pH electrode inserted into the reaction vessel. For the experiments using sodium cyanide as the starting material, the unbuffered cyanide solution at a pH of about 11 was quantitatively converted to cyanate in a known excess quantity of hypochlorite. Complete conversion to cyanate was usually obtained in less than about 10 min. The resulting cyanate solution was then buffered and made up to give the desired pH and solution volume. For those experiments using sodium cyanate solution directly, the solution was buffered and equilibrated at the given pH prior to the addition of hypochlorite. The kinetics of the reaction was studied by following the concentration change of the cyanate and hypochlorite as the reaction progressed. For some experimental runs, the amount of nitrogen gas evolved was measured continuously by a gas burette connected to the reaction flask. RESULTS AND DISCUSSION
Stability and oxidation o f cyanate The stability and oxidation of cyanate in hypo-
chlorite solution were experimentally examined under various operating conditions. Figures l and 2 show some typical experimental concentration-time curves of the various reactants and products in the oxidation process using sodium cyanide as the starting reactant. At p H 11 and above and in the presence of excess hypochlorite, cyanide was rapidly and quantitatively oxidized to cyanate. At these high pH, the cyanate formed remained fairly stable for several hours with negligible or no further decomposition even in the presence of excess hypochlorite as shown in Fig. 1 for the case of p H = 11. Measurable rates of the cyanate decomposition were only obtained at pH below about 10.5. Figure 2 shows the concentration-time plots for both reactants (cyanide and hypochlorite) and reaction products (cyanate and nitrogen gas) at p H = 9.1 and 27°C. An induction period generally prevailed prior to the liberation of nitrogen. This i n d u c t i o n time was observed to increase with increasing pH of the reaction mixture. The amounts of nitrogen gas liberated and the hypochlorite consumed during the oxidation of cyanate were found to differ from the values given by equation (I) for a corresponding amount of cyanate decomposed. The amounts of hypochlorite consumed were generally 5-12% higher while the nitrogen yield was generally lower than those predicted by equation (1). These may be attributed partly to the liberation of chlorine gas especially at higher temperature and in the presence of excess hypochlorite and partly to the parallel reaction leading to the formation of nitrate according to the reaction given by equation (2). Lister (1956) observed
? 16o
+--
0
4
x
6
3
x Z .c_
2
--O-O--O
•
8 ~_1
e-8
©'5
8 0
910~
10
+
t
130
cl
t
210
0
Time (min)
Fig. 1. Hypochlorite oxidation of cyanide at pH = l l.O and 27°C (CN- A, OCI- O, CNO [~, N 2 ©).
_J
16 b x
? o3 x z 2 ._c ¢
8
~
'~
o 0
10
50
90 Time (rain)
130
1200
Fig. 2. Hypochlorite oxidation of cyanide at pH = 9.1 and 27°C (CN- &, OC1 O, CNO- 15], N 2 ©, N O ; V).
Hypochlorite oxidation of cyanate Table 1, Effect of pH and hypochlorite/cyanatemolar ratio on the formation of nitrate at 27°C (OCI-/CNO-)0= 1.78 pH 8.7 9.2 9.5 10.0 10.5 NO~- x 103moll-I 0.52 0.45 0.32 0.20 0.15 pH = 9.15 (OCI-/CNO-)0 1.54 1.78 2.10 2.40 2.73 NO{ x 103moll 1 0.48 0.45 0.38 0.42 0.41 that the formation of nitrate was more prevalent in the presence of large excess of hypochlorite. The effects of pH and hypochlorite concentration on the formation of nitrate at 27°C were examined in this study and the results are tabulated in Table 1. Table 1 shows that reaction (2) leading to the formation of nitrate was more favourable at low pH and the effect of hypochlorite concentration was only marginal. Under the experimental conditions used in this study, the extent of decomposition of cyanate by reaction (2) would be relatively small compared with that observed by Lister (1956) at 65°C and in strong sodium hydroxide solutions. Most of Lister's (1956) experiments were carried out with a reactant ratio of (OC1-/CNO-) lower than the stoichiometric ratio given by reaction (1). No attempt to identify the presence of other products besides nitrogen and nitrate was indicated. In our present study, ammonium ion was always detected to be present in the product when the reactant ratio used was lower than the stoichiometric ratio. However, in the presence of excess hypochlorite, no ammonium ion was detected. Lister (1955) and Teo et al. (1979) reported that in the absence of hypochlorite, cyanate hydrolysed to ammonium ion very rapidly under acidic conditions but the rate became increasingly slower as pH increased to neutral and to mild alkaline conditions. These observations suggest that the hypochlorite oxidation of cyanate may probably proceed according to a series reaction by which the cyanate first hydrolyses according to reaction (3) to ammonium ion which is then oxidised by the hypochlorite to nitrogen gas according to reaction (4). C N O - + 2H20 + H ÷
OCI-
NH~ + HCO~-
(3)
2NH~- + 3OC1- + 2OH- = N 2 + 3C1- + H:O. (4) These would give an overall reaction as shown by equation (1). The significant increase in the rate of decomposition of cyanate in the presence of hypochlorite may be attributed to reaction (4) and/or to the catalytic effect of hypochlorite on reaction (3). The increase, especially at pH below 10, may also be brought about by the presence of hypochlorous acid, which is a much stronger oxidising agent than hypochlorite. Hypochlorous acid co-exists with hypochlorite according to the dissociation reaction as shown in equation (5). HOCI = H + + OCI-.
(5)
679
At 25°C, the dissociation constant is given as 3.4 x 10 -s (Cotton and Wilkinson, 1972). Lower pH favours the presence of hypochlorous acid which would therefore enhance the oxidation rate of cyanate. These experimental observations on the stability and oxidation of cyanate under various pH and (OC1-/CNO-) reactant ratios suggest that a complete treatment of cyanide waste to stable and inert end products is best carried out in two stages both in the presence of excess hypochlorite. The first stage involves the hypochlorite oxidation of cyanide to cyanate at pH about 11 and the second stage at a reduced pH of ,,-9 whereby the cyanate could be favourably oxidized to nitrogen and other stable products. In view of this proposed treatment scheme, subsequent kinetic study of the hypochlorite oxidation of cyanate was carried out under mild alkaline conditions at pH below 10.
Kinetics of hypochlorite oxidation of cyanate Effect of hypochlorite and cyanate concentrations. The kinetics of the oxidation of cyanate at pH = 9.15 and 27°C were examined over the range of initial molar ratio of (OCI-/CNO-) from 1.10 to 2.70. The ratio was varied by varying the hypochlorite concentration while maintaining the cyanate concentration at 7.14 x 10 -3 moll 1. Figures 3 and 4 show the concentration-time curves at various initial (OC1-/CNO-) molar ratios for cyanate and hypochlorite respectively. The decomposition rates of both the reactants were found to increase with increase in the initial reactant concentration ratio. Since the ratio of the experimental rate of decomposition of hypochlorite to that of cyanate was observed to differ by as much as 12% of the stoichiometric value, differential method of analysis of the kinetic data was preferred to the integral method. This would then avoid the necessity of having to choose a suitable average stoichiometric ratio of the reactants for the computation. The rates of decomposition of cyanate and hypochlorite were determined from the respective concentration-time curves given in Figs 3 and 4 at various fractional conversions of cyanate. The corresponding hypochlorite concentrations were also
8
7o
o<2 300
600 Time (s)
900
Fig. 3. Cyanate concentration-time curves at pH = 9.15 and 27°C and various initial reactants ratios [Cao/C^o= 1.10 (&), 1.54 (11), 1.78 (O), 2.10 (A) 2.40 (I-q), 2.73 (O)].
680
W.K. TEO and T. C. TAN Time (s) 240
120
360
480
16 14 12
~d~\ \ i
..alO ? o
\
o ~o
12 ~
o
~ 8
o
8.
o'2
x
I
0
300
I
600
Time (s)
g00
Fig. 4. Hypochlorite concentration-time curves at pH = 9.15 and 27°C and various initial reactants ratios [CBo/CAo= 1.10 (A), 1.54 (B), 1.78 (@), 2.10 (A) 2.40 (I-1), 2.73 (C))].
determined. Since the initial cyanate concentrations were the same at 7.14x 10-3moll -l for all the experimental runs, the corresponding cyanate concentration at a given fractional conversion of cyanate will be the same for all the runs. Figures 5 and 6 show the logarithmic plots of rates of decomposition of cyanate and hypochlorite against hypochlorite concentration for various constant fractional conversions of cyanate or constant cyanate concentrations. The plots are all found to be almost parallel straight lines giving a slope of approx. 1.0. The results therefore show that the rate of decomposition is first order with respect to hypochlorite concentration. The general kinetic expression may be written as ai bi h - ri = k, CACBC~
Figure 7 was then plotted accordingly for both the cyanate and hypochlorite. Parallel linear plots are obtained with a slope equal to 1.0 (a, = 1.0) for both reactants. The results therefore show that the reaction proceeds with first order kinetics with respect to cyanate and to hypochlorite. These were further confirmed by the graphical plots shown in Fig. 8 where the rates of decomposition were plotted against CACB. The plots are essentially linear giving a slope of 0.65 1mol-l s - 1 for - r A vs CACB and 0.3851mol-ls -I for - r e vs CAC~. This gives a ratio of --rB/--rA equal to 1.69 which is about 12.6% higher than the stoichiometric ratio given by equation (1). Effect o f p H . As discussed previously, the rate of cyanate decomposition increased with decreasing pH. At pH above about 10, a drastic reduction in the reaction rate was observed and the reaction practically ceased to occur at pH beyond 11. The significant enhancement of reaction rate at lower pH can most probably be attributed to the presence of HOCI. For this reason, only the data obtained at pH below 10 will be used for the kinetic analysis. The experimental data obtained for various pH were treated in the same manner as the preceeding section. The plots of ( - r , ) against CACB at various pH 10.C 8.° 6.0
i = reactant A or B - r i = rate of decomposition of the reactant i k, = reaction constant with respect to the reactant i CA, Ca, Ca = concentration of cyanate, hypochlorite and hydrogen ion respectively
i
2.0 o
1.C x 0.8
°.2
(6)
where
o/
4.0
i
,
681o
2'0
C B x l o "3 (M/L)
Fig. 5. ( - r ~ - C ~ plots at pH = 9.15, 27°C and at constant CA[CA x 103mo11-1;7.14 (O), 5.71 (@), 5.00 (A) 4.29 (V), 3.57 (UI)].
10.0 8.0
and a, b, h = order of reaction.
9 2.0 x
Since k~, CA and CH are constant, the plots in Figs 5 and 6 are described by
I 0.! 0.1
l o g ( - ri) = log(k, C~ C~) + b~log CB
(7)
0.4 / a J l L I 210 I 4 6 8 10 40 CB x 10 "~ (M/L)
where b~ is found to be 1.0 for both A and B. Rearranging equation (6) in the form shown by equation (8),
Fig. 6. ( - - r B ) - C B plots at p H = 9.15, 27°C and at constant
log(-r~/CB) = Iog(k~C~) + a, log CA.
3.57 (D)I.
(8)
C A [C A x l03 mol l-i; 7.14 (O), 5.71 (O), 5.00 (/k) 4.29 (V),
Hypochlorite oxidation of cyanate
b,
x
4.0
, ~ l~ ,
2.0 1.0
//o / .
0.8
~2- o.8 '~"
2.59 x 10 t3 and 3.80 x 1013 for the cyanate and hypochlorite respectively giving thus a ratio of --rB/--rA of 1.47, very close to the stoichiometric ratio of 1.5 given by reaction (1). This implies that on the overall average, the stoichiometry for the reaction is well approximated by equation (1). For the range of parameters examined in this study, reaction (1) therefore predominates over reaction (2) thus supporting the earlier observations and deductions. The overall rate expression for the decomposition reaction can now be written as
10.0 8.0 6.0
0
,A/I/ ~1 /
°
0.4
I
[
I
I
3
4
5
6 7 8910
I
681
I I
CA x l 0 -3 (M/L)
Fig. 7. (--rA/CB) and (--rs/CB) vs cA plots at pH=9.15, 27°C [(--rA/CB) O, (--rB/CB) I-q].
13 /-27030"X --rA=2.585 X lO e x p ~ - - - - ~ - ) C A C a C H \
(11)
/
and 10
- % = 1.5 (--rA) where the gas constant R = 8.314 J tool -l K -l. The orders of reaction with respect to the three reactants agree with those proposed by Lister (1956) in his study of the same reaction but at very high pH and moderately high temperature at 65°C.
6
b
4!
CONCLUSION
X
S
The decomposition of cyanate to its end products at room temperature was found to be negligible at pH
I
I
I
2o
4o
eo
I
I
6.0
I
80 lOO12O14O
4.0
t"~...("-- x 1 0 "e ( M ' / L ' ) vA
/
,.-- o
"~ 2.0
Fig. 8. (--rA) and (--ra) vs CACe plots at pH=9.15, 27°C [CBo/CAo= 1.10 (A), 1.54 (ll), 1.78 (O), 2.10 (A), 2.40 (fl)].
1.0 0.8
" ~ o.e (.~ 0.4
are linear. The values of k,CH or ~i are given by the slopes of the lines. Logarithmic plots of log ~biagainst the hydrogen ion concentration for both the reactants were shown in Fig. 9 and described by equation (9).
0.2 I
4
I
I
I
410
I
6 8 10
20
60
C H x l O -~° (M/L)
log ~ = log k~ + h~log C a.
(9)
The results show that the reaction is first order with respect to hydrogen ion. The reaction constants for cyanate and hypochlorite were found to be 5.85 x 109 and 9.50 x 10912 m o l - : s -l respectively giving a ratio of - rB/- rA equal to 1.62.
and (--rB/CACB) vs CH plots at 27°C [(--rA/CACs) 0, (--rB/CACB) 1"7].
(--rA/CACB)
8.0 6.0 4.0 ^
Effect of temperature
2.0
The reaction constants, k~, for the decomposition of cyanate and hypochlorite, were similarly obtained for the different reaction temperatures. For the effect of temperature on the reaction constants, Fig. 10 is plotted according to the Arrhenius expression lnk i=lnk~-
Fig. 9.
~
5 v
\ \ o -.
\ 1.0
o
os
\
._ 0 . 8
\
0.4
~-~ 0.2
I 2.9
for the reactants. Both the plots are parallel giving an activation energy of 27.0 kJ tool -I for the decomposition reaction. The frequency factors were
t 3.1
I
I 3.3
I
I 3.5
l I T x 10 -a (K -~)
Fig. 10. Reaction constant-temperature Arrhenius plot (CNO- C), OCI- ~).
682
W.K. "lEo and T. C. TAN
above 11 even in the presence of excess hypochlorite. However at pH ~< 9.5, the rates of decomposition were observed to be reasonably fast. The experimental results suggest that the decomposition most probably follows a series reaction C N O - + 2H20 = NH~ + CO 22NH~- + 3OC1- + 2 O H - = N 2 + 3C1- + H20. The increase in the decomposition rate in the presence of excess hypochlorite may be attributed to (a) catalytic effect of hypochlorite on the hydrolysis of cyanate to ammonium ion and (b) the removal of ammonium by a subsequent reaction to nitrogen and other products. The increase in the rate at lower pH in the alkaline region is mainly due to the increase in the concentration of a more powerful oxidising agent, hypochlorous acid at pH below 10. The experimental results yield a rate ratio of hypochlorite to cyanate of approx. 1.5. This supports the general acceptance of the approximate stoichiometric equation for the reaction, 2CNO
+ 3OC1- + H20 = 2HCO~- + 3C1- + N2
and also substantiates that under the present range of experimental conditions, the extent of decomposition to nitrate is small. The overall rate expression for the decomposition of cyanate is given by --rA(mOl 1-1 s -1) = 2.59 × 1013 {--27030\ × exp~)CACBCH
over the temperature range of 27-50°C and pH range of 8.6-9.5. The activation energy for the reaction is 27.0 kJ mol- 1. It may therefore be concluded that a complete cyanide treatment process at room temperature is
best carried out in two stages; both stages in the presence of excess hypochlorite. The first stage involves the conversion of cyanide to cyanate at pH 11 and the second stage from cyanate to inert and stable end products at pH 9. Acknowledgements--The authors gratefully acknowledge Ong Seng Eng and Yau Tai Yin for their assistance in part of the experimental work. REFERENCES
APHA (1980) Standard Methods for the Examination of Water and Wastewater, 15th edition. American Public Health Association, Washington, D.C. Cotton F. A. and Wilkinson G. (1972) Advanced Inorganic Chemistry, p. 569. Wiley, New York. Eden G. E., Hampson B. L. and Wheatland A. B. (1950) Destruction of cyanide in wastewaters by chlorination. J. Soc. chem. Ind. 69, 244-249. Eden G. E. and Wheatland A. B. (1950) Effect of temperature and pressure of hypochlorite on the rate of hydrolysis of cyanogen chloride in alkaline solution. J. Soc. chem. Ind. 69, 166-169. Kieszkowski M. (1968) Investigations on oxidation of cyanides in waste water by means of sodium hypochlorite. Pr. Inst. Mech. precyz., XV, 62-71. Lister M. W. (1955) Some observations on cyanic acid and cyanates. Can. J. Chem. 33, 426-440. Lister M. W. (1956) The reaction between cyanate and hypochlorite. Can. J. Chem. 34, 489-501. Price C. C., Larson T. E., Beck K. M., Harrington F. C., Smith L. C. and Stephanoff I. (1947) Hydrolysis and chlorinolysis of cyanogen chloride. J. Am. chem. Soc. 69, 1640-1644. Shaw W. H. R. and Bordeaux J. (1955) Semimicro method for the determination of cyanate ion in the presence of interfering substances. Analyt. Chem. 27, 136-138. Teo W. K. and Ng C. S. (1979) Optimum conditions in cyanide destruction using hypochlorite. Chemsia Conference, 16-19 January 1979, Singapore. Teo W. K., Ng C. S. and Ho W. Y. (1985) Cyanate determination in cyanate waste. Asia Chem. Conference, 8-I1 April 1985, Singapore. Vogel A. I. (1978) Quantitative Inorganic Analysis, 3rd edition, pp. 382-383. Longman, London.