T H E SYSTEM AZIDE-IODINE-"HIOCYA.NATE
1151
substance at the surface of the mercury drop. This decrease in concentration is caused by the incompleteness of the anodic process occurring during the quiescent period of the voltage wave. REFERENCES ( 1 ) D E L A H APY:, J. Phys. & Colloid Chem. 59, 1279 (1949). (2) D E L A H APY:, J . Phys. & Colloid Chem. 64, 402 (1950). (3) D E L A H AP: Y , J. Phys. & Colloid Chem. 64, 630 (1950). (4)DELAHAY, P., AND PERKISS, G . :J . Phys. &Colloid Chem. 66, 586 (1951). Y , AND STIEHL, G. L.: J. Phys. & Colloid Chem. 66, 570 (1951). (5)D E L A H A P., (6) KOLTHOFF, I. M., AND LINQANE,J. J . : Polarography, p. 45. Interscience Publishers, Inc., Kew York (1941). (7) RAXDLES, J. E . B . : Trans. Faraday SOC.44, 322 (1918). (8) RANDLES, J. E. B.: Trans. Faraday SOC.44, 327 (1948). (9) SEVCIK, A , : Collection Czechoslov. Chem. Commun. 13, 349 (1918). (10)WEIDEMANN, S . : Thesis, University of Bern (Switzerland), 1947.
REACTIONS IN T H E SYSTEM AZIDE-IODINE-THIOCYANATE PASCHOAL SENISE
Departamento de Qulmica da Faculdade de Filosofia, Citncias e Letras da Fniversidade de Sdo Paulo, S a Paulo, Brazil Received J u l y IS, 1960 INTRODCCTION
I t was found by Raschig (12) that no reaction can be observed when aqueous solutions of sodium azide and iodine (as KI.12) are brought together but that an immediate and vigorous evolution of nitrogen occurs if very small amounts of sodium thiosulfate or sodium sulfide are added t o the system. The same effect was further observed when either sodium thiosulfate or sodium sulfide was replaced by small quantities of other inorganic and organic sulfur-containing compounds, including thiocyanates, and very sensitive tests for the detection of these compounds have been introduced and developed, especially by Feigl (3, 4, 5, 6, 11). Such tests have found wide application in the analysis of a great variety of materials (6). Measurements of the nitrogen evolved were carried out by Friedmann (7), using different compounds of sulfur as catalyst. Mechanisms have been advanced for the different cases of this catalyzed reaction by Raschig (12), Feigl (5, 6), Friedmann (7), Browne and Hoe1 (l), and Weiw (15),but very few quantitative data were available before the appearance of the paper of Dodd and Griffith (2) on the systems Iz-S~Oy--XT and 12-S40a--NY. For the first time kinetic studies were made and though the mechanisms proposed for these two special cases, which were studied in connection with the direct reactions I z S Z O ~and - Iz-S~OF-, are-as stated by the authors-in some respects speculative, abundant experimental data and very
1152
P.4SCHO.kL SENISE
accurate measurements over wide ranges of conditions support the author's views. A kinetic study of the uncatalyzed reaction between sodium azide and bromine in aqueous solution and studies of the systems Brz-NT-S?O;- and Br2-N;-S40i'- have been published by Griffith and Irving (8). The papers of Dodd, Griffith, and Irving reached the author when he was carrying out research on the reaction between sodium azide and iodine with various sulfur compounds as catalysts, including sodium thiosulfate and sodium tetrathionate. As these last two syptems had been extensively studied by Griffith and Dodd, it was decided to confine the first publication on the subject to the system KT-I2-SCK-, especially because of the considerable number of experiments that had already been carried out. Recently the papers of Hofman-Bang (9) and Lgvtrup (10) were encountered. Hofman-Bang studied the catalysis of the iodine-azide reaction by tetrathionate and Lbvtrup studied its catalysis by cptine. The purpose of this investigation was to study the system Sa-I?-SCS- under conditions that could be considered similar to those ordinarily met with in the analytical tests when sodium azide and iodine are used to detect thiocyanate ion. Thus, almost every experiment was carried out in solutions having excess initial concentrations of azide ion and iodine xith respect to thiocyanate ion and so, with the exception of a few special cases, the iodine was not completely decolorized before or at the ending of the evolution of nitrogen. The main part of the esperimental n-ork consisted in determining the rate of evolution of nitrogen and observing the dependence of this rate on several factors. The factors considered nere: ( 1 ) the concentration of thiocyanate ion, (2) the concentration of azide ion, ( 3 ) the concentration of iodine, (4) the concentration of iodide ion, ( 5 ) the pH of the solution, and (6) the temperature. The following were also determined: (a) the total volume of nitrogen evolved under a variety of conditions relative to the above six factors; ( b ) the total iodine consumed in correspondence with the cases mentioned in ( a ) ; (c) the rate of consumption of iodine in some cases. The reaction between thiocyanate ion and iodine (excess iodine) was also studied under different conditions. Simple kinetic relations have been xyorked out with the data provided by the experiments for both the N;-I?-SCS- and the SCK--IZ systems. EXPERIMENTAL
Velocity of reaction The rate of evolution of nitrogen was measured using a differential manometer of the Barcroft type. The U-tube had a total length of 80 cm. and an average cross-section of 3.37 mm? The volumes of the vessels, including the space above the meniscus in the corresponding side arm of the apparatus, \yere found to be 39.246 mm.3 and 39.260 mm.8, respectively. The apparatus constant for the conversion of readings in millimeters to volume of gas in cubic millimeters a t N.T.P. v a s found to be 7.23 at 20°C.
1153
THE SYSTEM AZIDE-IODINE-THIOCYAN.4TE
Experiments were always carried out in a buffered medium and in a thermostat with continuous shaking. The total volume of solution in both the reaction and the blank vessels was 3.0 ml. Into the reaction flask were introduced 0.5 ml. of an aqueous solution of sodium azide, 1.0 ml. of iodine solution (KI.I2), 1.2 ml. of a suitable buffer solution, and 0.3 ml. of the buffered solution of potassium thiocyanate. The blank flask contained the same volumes of sodium azide and iodine solutions, 1.5 ml. of the buffer, and no potassium thiocyanate. These conditions are exactly the same as adopted by Friedmann ( 7 ) . Except when influences of temperature or of pH on the rate of gas evolution were studied, the reaction was allvays followed a t 20°C. and the buffer solution used was a 1:l mixture of 0.2 M acetic acid and 0.2 M sodium acetate (pH = 4.6). The reaction solutions so prepared showed a pH of about 4.9. In the majority of the experiments the reaction was followed over a period of 2 hr. and the manometric readings were generally taken every 5 min. A great number of experiments have been carried out under a variety of conditions, but only a few examples are given in this report. The values for the initial concentrations in all tables or graphs are given in gram-moles or gram-ions per liter. [ZIP] and [XI-] represent total iodine and total iodide concentrations, respectively. In the K I .IOsolution used, the iodine was always estimated by titration with thiosulfate and the iodide titrated by oxidation to iodine monochloride with potassium iodate in strong hydrochloric acid solution.
Effect of variation of different factors on the rate of evolution of nitrogen ( I ) Variation of [SCN-1: It was found that the volume of nitrogen produced over definite intervals of time is a function of the concentration of thiocyanate ion. The evolution of nitrogen increases with increasing initial [SCN-] in the range tested. In many rases, provided small intervals of time are considered as shown in figure 1, the plot is linear. (2) Variation of [NJ: An increase in the initial concentration of azide ion determines an increase in the evolution of nitrogen, the plot of mm! of nitrogen against initial [XJ being approximately linear over relatively small intervals of time and not too wide concentration ranges. One example is shown in figure 2. ( 3 ) Variation of [ZI,]: The influence of the increase of total initial iodine concentration on the rate of evolution of nitrogen is shown in the graph of figure 3. ( 4 ) Variation of [XI-]: Figure 4 shows how the evolution of nitrogen is decreased by increasing the initial total concentration of iodide ion. The decrease is very m a 3 :d a t relatively low concentrations of iodide ion but small at high values of [XI-]. It was found that many of these curves are very regular and of the hyperbolic type according to the general equation y = ax-*, y being the volume of nitrogen evolved and z the initial [XI-]. The following equations were calculated for the curves of figure 4, corresponding to 10. 20, and 30 min.: y = 1.832-2 3 9 ;
y = 1,79x-? 7 3 ;
y = 2.39z-O
76
120'
150
-
1.o
2 .o
3.0 In; ti a1 [S C N-7 x 10'
4.0
FIG.1. Initial concentrations: IS;] = 5.00 X 10-2; [z12] = 2.22 x 10-2; [XI-] = 1.23 x
x
FIG.2. Initial concentrations: [XI21 10-6.
-
4.47 X lo-*; [XI-]
1154
-
1.60 X W1;[SCK-] = 9.67
60'
200
1%
N
5z
-r sc
C 1.0
6
1.4
l.b
Initial
x
2.2
2.4
2.b
E IJXIO'
FIG.3 Initial concentrations: IN;] = 5.00 X 10-2: @I-] =. 6.03 X IO-*; [SCN-] = 1.w 10-4. 10(
71
\
\30' N
5 z
5c
25
0
Initial I Z I ~ X I O FIG.4. Initial Concentrations: [N;] = 5.00 X 10-2; [212] = 5.04 x 10-2; [SCN-] = 1.94 10-4.
1155
x
1156
PASCHOAL SENISE
I n order to exclude other possible interpretations of the r81e played by iodide ion, a few experiments have been performed by replacing partially the potassium iodide by potassium chloride. No perceptible influence of the potassium chloride was observed. In addition, some experiments were carried out in the presence of varying amounts of another electrolyte, sodium nitrate being chosen. The same results were obtained. This shows that-in the range tested-the iwrease of ionic strength does not affect significantly the rate of evolution of nitrogen. I n the following two cases, for instance, the molar concentration of sodium nitrate in the reaction solution was varied from 0.2 to 0.5, the initial concentrations of the other reactants being as follow: (a) [Nil = 6.04 X 10-8 [SCK-] = 1 94 x lo-' (b) [N;] = 5.00 X 10-2 [SCX-] = 1.94 X lo-'
'
1
~
[XI21 = 2.91 X 1G+ [XI-] = 5.35 x 10-2 [XI21 = 8.35 X [XI-] = 0.399
' ,
[HAC] = 5 . 0 X [ S s A c ] = 5.0 X IO-* [HAC] = 5.0 X lo-* [XsAc] = 5.0 X lo-*
The data presented under paragraphs (3) and (4) may be considered from a different point of view. Instead of [ZI,] and [CI-] one can use the concentrations of free initial iodine and iodide ion, meaning for free initial concentrations the values obtained for [I21 and [I-] from the triiodide equilibrium:
[Id[I-l [IF1
=
1.2
x
10-3 (20oc.)
In fact, plotting the volume of nitrogen produced in definite intervals of time against [I*]/[I-] or [Iz]/[I-]~it can be seen that the rate of evolution increases with the increase of these quotients. A number of experiments were carried out using a progressively diluted iodine solution. Thus, the value of [212]/[21-] was kept constant in the whole series. It was found that the rate of evolution increased with decreasing [CI,]. However, if [I21 and [I-] are calculated, it can be seen that the ratio [I2]/[I-], and consequently [12]/[I-]?, increases with dilution, thus explaining the increase of the rate of nitrogen evolution (see table 1). Collecting the data of many experiments performed under a variety of conditions with respect to initial [ZI?] and [ZI-1, graphs can be constructed which are of the type shown in figure 5, if the evolution of nitrogen is plotted against [12]/[I-l2. The rate of evolution increases with increasing [I?]/[I-]' regardless of the initial values of [ZIZ]and [ZI-1. (5) Variation of p H : The effect of pH changes on the rate of evolution of nitrogen was studied through a number of experiments carried out with acetic acid-sodium acetate buffers with pH ranging from 3.4 to 5.9 and Na:HPOaKH2POa buffers, with pH ranging from 5.5 to 8.0. The buffers were carefully prepared, so that every mixture had a constant value for [HC2H302] and [SaC2H30?]and for [Na2HP04]and [KH2POa],respectively, and about the same value for the ionic strength in the reaction solutions. For control purposes phosphate buffers were replaced in some samples by
1157
THE SYSTEM AZIDE-IODINE-THIOCYANATE
borate buffers and practically no difference was observed in the results obtained. The initial pH of the reaction solution was always checked in a Cambridge pH meter.' TBBLE 1
mm3cztlTP
3 35 99
4 2s 2 09
3 73 s3
0 344
II
'1
52 8
102
14
1 ii i) 1
1
i
~
127
'1
308
4 46
1 08 0.520
I]
7'33
0 397
]
8 99
I
1 41
mm3atlTP
~
'
512
582
Figure 6 is representative of the results obtained. The evolution of nitrogen reaches a maximum-for a given interval of time-at pH values of about 5.7-5.9. The limits of the reaction according to the pH may thus be established for each of the cases studied. 1
Cambridge Instrument Company, London, England.
1158
PASCHOAL SENISE
It is to be noted that in the range of pH below 6.0, even when the reaction proceeds very slowly, the evolution of nitrogen can be followed over a wide period of time and the reaction may have a long duration. On the contrary, in the region of high pH-from the maximum on-the evolution of nitrogen ceases completely after a relatively short length of time, this interval of time being shorter with increasing pH. In figure 6 the plots for pH 7.4 in curve I11 and pH 7.0 in curve IV are final values for the evolution of nitrogen.'In such cases further production of gas was not observed by changing the pH of the medium or increasing the azide-ion concentration. If, however, an experiment carried out at low pH-say, 3.04.0-is interrupted and the pH slightly increased, the rate of formation of nitrogen also increases.
The above-mentioned difference indicates a more rapid consumption of the catalyst (SCN-) by iodine a t higher pH than about 6.0. This assertion is also illustrated by curve I in figure 6: the usual solutions of iodine, thiocyanate, and buffer were kept in a thermostat for 120 min. and then the sodium azide solution was added. Results obtained for the first 10 min. in terms of nitrogen evolved are shown in curve I in correspondence with those given in curve I1 and obtained with exactly the same reaction mixtures-in normal conditionsafter the same initial period of 10 min. (6) Variation of temperature: Determinations of the rate of gas evolution were made in some cases a t 15", 20", 25", and 30°C. The increase of the rate with increasing temperature nil1 be discussed later (page 1162).
Determination of the total nitrogen evolved Determinations of nitrogen were made by allowing the reaction to proceed to completion, Le., until the readings in the Barcroft, manometer became constant
1159
THE SYSTEM AZIDE-IODINE-THIOCYANATE
regardless of the length of time elapsed. The final readings were always checked by increasing the azide-ion concentration and sometimes the iodine concentrstion also. According to the concentrations of the various reactants some experiments had a very long duration (sometimes a few days). Table 2 summarizes the results of part of the experiments carried out under different conditions. Inspection of the data shows that in spite of the lower value of [SCN-] in comparison with concentrations of azide ion, iodine, and TABLE 2 INITIAL CONCENTUTIONS
.-x 10'
rN3 .
5.00 5.00 5.00 5.00 5.00 5.00 5.00 2.50 10.0 23.7
[XI11
x
108
0.82 1.41 2.58 3.17 2.58 2.58 2.58 2.58 2.58 2.58
[XI-] x 10'
[SCN-I X 104
6.03 6.03 6.03 6.03 12.1 18.1 24.1 6.03 6.03 6.03
1.94 1.94 1.94 1.94 1.94 1.94 1.94 1.94 1.94 1.94
1.80 1.72 1.64 1.58 2.15 2.34 2.62 1.18 1.75 2.62
93.0
88.9 84.6 81.5 111 121 135 60.4 90.2 135
* R is the ratio of moles of azide ion decomposed to moles of thiocyanate ion consumed. TABLE 3 EPPECI OF INCREASING
[SCX-] . . . . . . . . . . . . . . . . . . . . . . . . . [S,].. . . . . . . . . . . . . . . . . . . . . . . . . .
Oh' THE PATE OF EYOLDTIOS OF
??I
Increase Increase [ZI*].. . . . . . . . . . . . . . . . . . . . . . . . . . . Increase [ZI-].. . . . . . . . . . . . . . . . . . . . . . . . . Decrease [I*]/!I-]P.. . . . . . . . . . . . . . . . . . . . . . . . Increase Temperature . . . . . . . . . . . . . . . . .Increase pH. . . . . . . . . . . . . . . . . . . . . . . . . . . . A maximum was observed
o s THE
~ 0 t PRODUCTION h ~
or Ns
Increase Increase Decrease Increase Decrease Decrease A maximum was observed
iodide ion, the variation of these excess concentrations affects the final result in terms of [NJ consumed. I t can be noticed that with increasing initial [ZI,]the total production of nitrogen decreases slightly. This decrease seems to be linear in the range tested, The increase of [ZI-]determines the increase of the total nitrogen evolved. It was also found that the volume of gas decreases with increasing initial [12]/[1-]* regardless of the values of the initial [XI,] and [ZI-]. The total nitrogen evolved also decreases with increasing temperature. I t is to be noted that in each series of the foregoing examples a linear relationship holds between the initial [SCN-] and the total nitrogen evolved, in the range tested. The total production of nitrogen seems to be affected by the variation of pH in a similar manner as the rate of evolution. The effect of variation
1160
PASCHOAL SENISE
of the different factors studied on the rate of evolution and on the total volume of gas produced is summarized in table 3.
Determination of the iodine consumed Though titrations with N/20 arsenite of the excess iodine have been done in almost every experiment mentioned in the preceding section, the accuracy of this method was not very satisfactory for the estimation of the total iodine consumed in the reaction process. Much better results were obtained by titration with N/20 thiosulfate after acidification to pH about 2 with dilute sulfuric acid and by colorimetric measurements. Results obtained by the last two procedures show a satisfactory agreement. I n the colorimetric process iodine was determined by extraction with chloroform (0.1 or 0.2 ml. of the reaction mixture was extracted) and measurements were made in a Klett-Summerson photoelectric colorimeter, readings being taken against pure chloroform, using standard tubes and filter K-S n. 50. The law of Lambert and Beer was perfectly obeyed in the range tested and no interference was observed.? Results obtained for the total iodine consumption together with the values found for the total nitrogen evolved agree satisfactorily with the following stoichiometric equations, vhich are assumed to be representative of the net reactions taking place: 3x2 + 21+ 312 + 4H;O -+SO;- + 61- + C S - + 8H11 (acid medium) 3x2 + 212N; + I* ‘I (alkaline medium) SCK- + 412 + 4H20 SOT- + 71- + ICX + 8H+(
ax; + I*
---t
SCN-
-+
---f
The experimental results found for the iodine consumption were compared with data calculated on the basis of the corresponding values obtained for nitrogen. Considering that these “calculated” values involve the experimental errors of the nitrogen determination, results can be considered satisfactory for the purposes of this work. Comparison of results obtained by the t n o different methods show deviations that vary from 0.5 to 2.5 per cent for a consumption of iodine of about 1.5 X to 3.9 X 10-3 gram-mole per liter. Qualitative spot tests for the detection of sulfate ion were carried out in the cases mentioned above with barium rhodizonate, as indicated by Feigl (6). The tests were made after complete decolorization of excess iodine by arsenite, follo\ved by slight acidification with dilute acetic acid. Cyanide ion was detected by the copper acetate-benzidine acetate test, carried out after complete consumption of the sodium azide (to avoid any interference of hydrogen azide) and decolorization of the excess iodine by thiosulfate.
* This procedure was suggested by Dr. Rubens SalomB Pereira from the Department of Zoology, Laboratory of General and Animal Physiology of this University, who kindly allowed the author the use of his colorimeter and personally assisted in this part of the work. His kind cooperation is gratefully acknowledged.
THE SYSTEM AZIDE-IODINE-THIOCYANATE
1161
The presence of hydrogen ion in unbuffered samples was also proved. The determination of the rate of consumption of iodine in the course of the reaction was also carried out in some cases, as will be seen later (page 1163). The reaction was stopped at the desired interval of time through the addition of dilute sulfuric acid, and the iodine was titrated with N / 2 0 thiosulfate.
Thiocyanate ion and azide ion do not react in the absence of iodine Because of the statement of Hofman-Bang (9), who observed evolution of nitrogen when large amounts of tetrathionate and sodium azide were brought together, some experiments were carried out as follows: ( a ) The usual amounts of potassium thiocyanate and sodium azide solutions and buffer were mixed in the reaction vessel of the Barcroft apparatus and kept in a thermostat, a t 20°C., for definite intervals of time, a maximum of 26 hr. being reached. KO formation of gas could be observed. The iodine solution waa then added and the evolution of gas measured. The results agreed with those of the corresponding experiments as carried out in the normal way. ( b ) Five-tenths of a milliliter of a 1.8 M solution of sodium azide was treated in the Barcroft apparatus with 1.5 ml. of a 0.5 M solution potassium thiocyanate and kept in a thermostat a t 20°C. over a period of 46 hr. No evolution of gas was detected. ( c ) One milliliter of a 1.8 M buffered solution of sodium azide was treated in the same apparatus with 2.0 ml. of a 0.9 M buffered solution of potassium thiocyanate (the pH of the mixture being 5.6). The solution was kept at 20°C. during 18 hr., but no evolution of gas was observed.
Kinetic considerations The experimental data reported in the foregoing sections have been used for simple kinetic calculations. I t was found that the graphs of log ( a - 2) against t are linear in the majority of the experiments, a being the initial azide ion and z the amount reacted in time t , provided not too large an interval of time is considered. Linearity holds over the time the reaction was followed (120 min.) only when evolution of gas proceeds very slowly (figure 7 , curves I and 11).When more rapid evolution occurs, deviations from linearity are observed and straight lines are obtained only for shorter periods of time, the retardation being more marked the more rapid the reaction (figure 7 , curves 111, IV, and V). The rate of evolution of nitrogen can thus be considered as unimolecular with respect to azide ion. Values for k,,, were calculated from the simple relation
'k
2303 a - x1 = L log a - xz t~ - t1
for the interval of time in which linearity is observed in' the graphs mentioned. Table 4 gives the values of kuni calculated from the experimental data in some tested cases.
1162
PASCHOAL SENISE
0.71
0.6
06
n
X
8
v
3
-F
0.6
06
0.6. 10
30
20
40
60
50
70
80
90
100
110
120
TIME-MIN.
IN;]
-
FIG.7. Initial concentrations
5.00 X I : [E111 = 8.35 X 11: [ZLI = 5.04 X 111: 12121 = 6.81 X IV: 12121 = 2.22 X v: [ZM = 2.82 x
10-* lo-*; [XI-] = 3.99 X lWa; [XI-] = 2.36 X lWa; [XI-] 6.03 X lWz; [XI-] = 1.23 X 10-y 121-1 = 1.21 x
-
lW1; [SCN-] = 1.94 X 1O-I; [SCN-] = 1.94 X lO+; [SCN-] = 9.67 X lW1; [SCN-] = 1.94 X 10-1; [SCN-I = 1.94 x
10-6 10-4
l(r4 10-4
The following two examples show the effect of the temperature on the rate of reaction: ( 1 ) Initial concentrations: [Nil = 5.00
x 10-2;
[&]
=
5.44 X l e 2 ;
[XI-] = 2.78 X le';[SCN-] = 1.94 X 10-4 15OC.: k = 3.70 X lW4 min.-' 25°C.: k = 12.4 X le4 rnin.-l
kw/kw
3.35
2OOC.: k = 7.02 X lW4 rnin-l 3OOC.: k = 20.8 X lO-' rnin.-' k a y / k r o O = 2.97
1103
THE SYSTEM AZIDE-IODINE-THIOCYANATE
(2) Initial concentrations: [K;] = 1.00 x 10-1; [ZIJ = 5.44 [ZI-]= 2.78 X 10-l; [SCS-] = 9.67 X 1W6
x
15OC.: 12 = 2.40 X lo-' min.-l 25OC.: k = 8.35 X lW4 min.-l
2OOC.: 12 = 4.55 X min.-l 30°C.: k = 14.8 X lW4 min.-'
kw/kjs* = 3.47
kaoO/ktp
1W2;
3.25
The graphs of log k against 1 / T are linear, in agreement with the Arrheniua equation:
-AHa 1 log k = 2.303R T
+
The values calculated for AHa, the apparent activation energy, are 20.6 kcal. and 21.1 kcal., respectively. TABLE 4 ISITIAL COSCENIPAIIOh'S
w;1
x 10'
PI*1 x 10'
IrI-1
x
10
[SCN-I X 10' min.-l
2.50 2.50 2.50 5.00 5.00 5.00 5.00 5.00 5.00 5.00 5.00 5.00 5.00 5.00 5.00 7.50 7.50 10.0 10.0 10.0
4.47 4.47 4.47 0.681 0.681 0.681 1.36 2.82 2.22 2.22 2.22 2.58 5.04 5.04 5.04 2.22 2.22 2.22 2.22 3.71
1.60 1.60 1.60 0.605 0.605 0.605 1.21 1.21 1.23 1.23 1.23 1.81 2.36 2.63 3.61 1.23 1.23 1.23 1.23 1.68
0.967 1.94 2.90 0.484
0.967 1.94 1.94 1.94
0.967 1.94 2.90 1.94 1.94 1.94 1.94 1.94 0.967
0.967 1.94 0.967
9.29 18.1 27.2 5.83 12.4 24.2 8.45 25.2 9.84 19.4 27.9 7.00 7.34 5.32 1.84 20.1 10.7 10.5 20.8 8.35
When the reaction is very rapid it seems without significance to present pseudo-unimolecular constants that would fall markedly in the course of the reaction. This fall, as well as the deviation from linearity in the examples mentioned above, can be ascribed to the considerable changes in the concentrations of the various reactants. In many of these cases of rapid evolution of gas, the consumption of iodine, a t definite intervals of time, was determined by titration in the manner described on page 1160, and thus the amount of thiocyanate ion consumed in the course of
1164
PASCHOAL SENISE
the reaction could be calculated with the help of the stoichiometric equations given on page 1160. In addition, the variation of the ratio [12]/1-]2 was also calculated over the time the reaction was followed. With these data, attempts were made to work out more precise kinetic relations that could express the overall rate of reaction, but though many experiments were performed, the accuracy of the results was not considered satisfactory for this purpose. However, in a few cases it was found that the rate of gas evolution could be expressed approximately as ANz/At
k[([N;]
- Z)([IZ]/[I-]~- y ) ( [ S C N - ] - z)]
z,y , and z being the decrease of [N,], [Iz]/[I-]2,and [SCN-1, respectively, in time t. I t is to be noted that in such cases the ratio [12]/[1-]2can be replaced in the equation by the quotient [12]/[1-] and a linear relationship is even so obtained. Two examples are given below, with mean values of k calculated from the plot of log [([NJ - r)([I2]/[I-I2 - y ) ( [ S C N - ] - z)] against t, which was found to approach linearity. (1) Initial concentrations: [NJ = 5.00 X [XI2] = 2.56 X 1W2; [ZI-] = 6.05 X le2; [SCK-] = 1.94 X pH = 4.9;T = 30°C. k = 8.5 X 1W2 rnin.-l (9) Initial concentrations: [NJ = 6.04 X IO+; [ X Z ] = 2.91 X lo+; [ZI-] = 5.35 X pH = 4.9; T = 30°C. [SC-U-] = 1.94 X k = 2.7 X 10-1 min.-'
T h e reaction between thiocyanate i o n and iodine The behavior of thiocyanate ion in the presence of excess iodine (KI.12) in buffered acetic acid-sodium acetate solution of pH 4.6 was followed and the rate of consumption of thiocyanate ion determined. The quantitative estimation of thiocyanate ion in the course of the reaction was made by an indirect method with the aid of the reaction with azide ion. In the reaction vessel of the Barcroft apparatus, solutions of iodine, buffer, and potassium thiocyanate were mixed in exactly the same volumes as used in the study of the N;-12-SCN- system, Le., 1.0 ml. of iodine solution, 1.2 ml. of buffer, and 0.3 ml. of buffered potassium thiocyanate solution. After being kept in a thermostat during a definite time (several hours, owing to the slow reaction rate) the sodium azide solution was added (0.5 ml.). The evolution of nitrogen was then measured as usual. As the consumption of iodine was found to be very small in the majority of the experiments, a graphical method could be adopted to evaluate the amount of thiocyanate ion which had reacted with iodine before the addition of azide ion. With data obtained from adequate examples of the system N~-Iz-SCN-, graphs were constructed by plotting volume of nitrogen evolved as a function of initial [SCN-J for given intervals of time (as already shown, for instance, in figures 2 and 3). From these graphs the concentration of thiocyanate ion at the
1165
THE SYSTEM AZIDE-IODINE-THIOCYANATE
moment the sodium azide was added could be found with acceptable accuracy, in the samples in which a known [SCS-] had been previously treated with excess iodine. Check experiments were carried out by preparing solutions of sodium azide, iodine, and potassium thiocyanate with exactly the same [SCN-] as was found by the graphical method. Results were confirmed with an accuracy of 1 to 5 per cent according to the different cases tested. Mean values were generally used for calculations. This method has provided reproducible and more accurate results than the attempts made to determine the iodine consumed by titration or by colorimetric ~neasurements.~ TABLE 5 1niti:il [ S C S - ] = 2.32 X lo-' kuni ( S C W ) [ZIll
x
102
1x1-1
x
7 7 7 7 7 10 14 21
21 24
x
10'
102
nin.+
0.708 1.69 2.50 3.81 3.10 3.10
3.10 3.10
I I
I
24 21 24* 8* 5'
7*
3.43 4.25 7.12 19.2 10.9" 4.25' 2.29" 0.928"
The decrease of [SCS-] was studied as a function of time. The plot of log ( a - z) against t , a being the initial [SCK-] and z the amount reacted in time t , shows in every case studied a linear relationship. Thus a unimolecular constant can be calculated for the consumption of thiocyanate ion. The rate of consumption of thiocyanate ion was also studied in relation ts different factors. It was found that increasing of [ZI-] decreases the rate of reaction (see also table 5). The graph of various values of k,,, (or [SCK-] consumed) against the initial concentrations of [ZI-] gives a curve of the hyperbolic type, since by plotting log X: against log [ZI-]a straight line is obtained. The following general mathematical equation can thus be written y = ax-b
y being k,,,
(or concentration of thiocyanate ion consumed) and z the initial [ZI-1. The curve obtained with the data marked with an asterisk in table 5 obeys the equation :
Y
=
2.95
x
1 0 - ~ 2 5
' A micro method for the determination of thiocyanate, based on this reaction, was developed and presented at the First International Microchemical Congress, Graz, Austria, July, 1950. See Senise, P. Mkrochemie 36, 210 (1951).
1166
PASCHOAL SENISE
It is to be noted that an expression of the same type was found in the system Nr-Iz-SCN-, relating the evolution of nitrogen (or consumption of azide ion) to the [ZI-] (page 1153). It was also found that values of kuni increase with increasing [XIz] (table 5) and with increasing [Iz]/[I-]~regardless of the initial [Z12]and [XI-]. Some experiments of this series were repeated by adding enough sodium nitrate to make the reacting solution 0.2 M with respect to this salt. The results mere practically the same as obtained before. The increase of ionic strength has thus no effect on the rate of consumption of thiocyanate ion by iodine. The temperature coefficients were also determined by following the reaction a t various temperatures. An example is given below: Initial concentrations: (SCN-] = 2.32 X [ZrIi] = 3.10 X [ZI-] = 7.21 X kuni X lOr
IEKPEPATUPE
‘C.
min.-i
15 20 25 30
0.608 1.09 2.07 3.65
The plot of log k against 1/T is linear, AHa being 20.8 kcal. It is to be noted that this calculated value of the apparent activation energy is very nearly that obtained in the study of the system K - I Z S C N - (page 1163). The effect of pH is evident from figure 6 (considered on page 1157). The marked increase of the rate of reaction with increasing pH was to be expected from the data already known (13, 14). The formation o€ sulfate mas qualitatively proved in separate tests even a t low pH ranges. -4number of experiments were carried out with potassium thiocyanate and excess iodine a t pH 4.6 (buffered medium) in the presence of barium chloride. The formation of barium sulfate, which was easily observed in cases of low iodide concentration, occurred only very slowly when large amounts of potassium iodide were present. However, if sodium azide was added to the samples, the precipitation of barium sulfate was considerably accelerated. An experiment was performed with very large amounts of potassium iodide. The solution with barium chloride a t pH 4.6 was kept at room temperature over a period of 7 months. Only extremely slight turbidity could be observed. Sodium azide was then added; a slow evolution of gas began and after 1 day the white precipitate formed was identified as barium sulfate. DISCUSSION
The equation ANz/At-= k[([NJ
- Z)([IZ]/[I-]~- 1/)([SCN-] - z)]
THE SYSTEM AZIDE-IODINE-TRIOCYANATE
1167
given on page 1164 is assumed to be a probable expreasion of the overall rate of reaction in the system N;-I,-SCN-. Such an equation would summarize the experimental findings and give a logical explanation of the results obtained in the study of the rate of the reaction. In fact, it appears clear that the deviations from linearity observed when the evolution of nitrogen was plotted against the initial [SCN-] or the initial [NJ are to be mainly ascribed to the changes in the concentrations of the factors that are not considered in those graphs ([NJ and [I2]/[I-I2 in the first example; [SCN-] and [I2]/[I-]* in the second) and which become the more marked the larger the interval of time considered or the more rapid the course of the reaction. This explanation is also true for the deviation from linearity in the graphs of log (a - 5) against t as pointed out when the values of kuniwith respect to azide ion were given (see page 1161). The influence of the concentrations of iodine and iodide is much better understood if the [I2]/[I-I2 ratio is used instead of [XI31 and [ZI-]. [I2]/[I-I2 is directly related to the oxidation-reduction potential of the iodine-iodide system and so it can be considered as a meamre of the “active iodine” in the solution. Thus, it could be stated that the rate of evolution, in the range tested, is a direct function of the concentrations of azide ion and thiocyanate ion and of the oxidation-reduction potential of the iodine-iodide system a t the moment considered. The results of the determination of the total nitrogen evolved and of the corresponding iodine consumed agree with the equations (given on page 1160) assumed as representative of the net reactions taking place. The results obtained in the study of the reaction of thiocyanate ion with excess iodine show-as has been reported (page 1166)-that the influence of the initial and [ZI-] on the unimolecular rate of consumption of thiocyanate icn is very similar to that observed in the system XT,-12SC9-. The calculated temperature coefficients are also very close and the same was found for the apparent activation energy. Yo effect on the rate of reaction in both systems was observed by the increase of ionic strength. Xot only does the thiocyanate ion catalyze the consumption of azide ion by iodine, but also and a t the same time the oxidation of thiocyanate ion by iodine is catalyzed by azide ion. The observation that large amounts of sodium azide are oxidized to nitrogen whereas only very small quantities of thiocyanate ion are consumed and that by starting from a definite concentration of thiocyanate ion the production of nitrogen can be varied by the variation of the excess concentrations of iodine, iodide ion, or azide ion (see table 2) supports the idea khat the mechanism of consumption of sodium azide involves a reaction chain. Though further work seems to be necessary to support a detailed mechanism for the reactions in the system X~-ITSCX-, a consideration of our findings, taken together, points out the main lines of the reaction mechanism. In order to help follow the argument a general and simplified scheme of the assumed mechanism is tentatively given on page 1168. There is no intention to indicate in the scheme the real number of reaction steps or the true formulation of the intermediates.
1168
PASCHOAL SENISE
As the conditions that favor the increase of the rate of oxidation of thiocyanate ion by iodine leading to sulfate ion also favor the rate of evolution of nitrogen in the system NT-I&CIX-, it can be assumed that the initial stepor steps-is the reaction between thiocyanate ion and iodine (I), giving an intermediate product (SCNI). The oxidation would proceed further (reactions I1 and 111), and also a competing reaction between the intermediate and azide ion would occur (IIa) with production of nitrogen and regeneration of thiocyanate IIb, would be broken by the comion (IIb). This reaction chain, I IIa peting reaction (IIc) between the assumed intermediate (SCNI(KT)) and iodine, leading to sulfate ion (reaction 111).
+
I + Is c--I SCNI + 1+ I/-.+ 'I SCNI(12) SCNI 4
+
SCN-
+N;'-=+
SCNI(NT)
+N;/-+
IIb
SCN-
+ I- + 3N2
----( +IA2'. SCNI(In) +
+
N;
A mechanism with various steps taking place consecutively and simultaneously could thus be assumed. The variation in the velocity of these different steps due to different conditions of concentration, temperature, or pH would favor in a higher or lower degree the production of nitrogen.4 As the increase of ionic strength does not affect the rate of reaction, it could be assumed that the rate-determining reaction step involves an ion and a neutral molecule (reaction I in the scheme). I n the light of these ideas a satisfactory interpretation of the experimental results summarized in table 3 is possible. The effect of the increase of [SCX-] on the rate of evolution of nitrogen and on the total production of gas is obvious, and the same can be said of the influence of increasing [NJ. The increase of [ZIz], as has been shown (page 1166),favors the rate of reaction between thiocyanate ion and iodine. The assumed intermediate (SCNI) will be formed more rapidly, its concentration will be increased, and so the reaction with azide will be accelerated. But the competing reactioq (11) must be more accelerated than reaction IIa and consequently a smaller amount of the intermediate (SCNI) will be able to react with azide ion, the total produc-
'
It is to be noted that the explanation given above is analogous to the detailed mechanism proposedbyDoddandGriffith(2) forthesystems N;-IrS,O; - andS',- I:!-S?O;-and also agrees fundamentally writh the earlier hypothesis of Raschig (12), adopted by Feigl in various cases (6).
THE SYSTEM AZIDE-IODINE-THIOCYANATE
1169
tion of nitrogen decreasing slightly. On the contrary, when [XI-] is increased the reaction between thiocyanate ion and iodine is retarded, determining a slower formation of the intermediate (SCNI) and consequently a decrease in the rate of evolution of nitrogen; however, reaction I1 must be more retarded than reaction IIa, thus allowing a larger amount of SCKI to react with azide ion. The reaction in this case may have a long duration because of its slowness but a larger volume of gas is obtained. The increase of [I2]/[1-]2 representing the increase of the oxidation-reduction potential corresponds to the raising of [ZI2] or lowering of [ZI-1. The influence of temperature promoting the increase of the rate of reaction between thiocyanate ion and iodine is comparable to the effect of increasing [XI,]. The effect of variation of pH finds also s n explanation. The oxidation of thiocyanate ion by iodine takes place with production of hydrogen ion and so this reaction must be favored by raising the pH. This is why even at a pH of about 7, in buffered solutions, the amount of thiocyanate ion consumed by oxidation to sulfate ion (through reactions I I1 111) may be relatively large and in the alkaline region when also the oxidation-reduction potential of the iodine solution is highly increased the formation of sulfate is extremely rapid and no reaction can take place with azide ion at high pH values. In the region of low pH, as described on page 1158, the phenomenon is different. The lower velocity of evolution and smaller amount of gas produced are to be ascribed not only to the extremely slow reaction of thiocyanate ion with iodine, but also to the decrease of the effective [S;] owing to the increase of the ratio [HNJ/[KJ, the reaction occurring only with the free azide ion, as has also been pointed out by Dodd and Griffith (2) and by Hofman-Bang (9).
+ +
SUMMARY
The iodine-azide reaction catalyzed by thiocyanate has been studied. The rate of erolution of nitrogen in the aqueous system NT-L-SCX- was measured under different conditions of concentration, pH, and temperature. The rate of consumption of thiocyanate ion in the system thiocyanate ioniodine was also determined. The main lines of a mechanism nere discussed which explain satisfactorily the oxidation of azide ion to nitrogen catalyzed by thiocyanate ion and the oxidation of thiocyanate ion to sulfate ion accelerated by azide ion. The primary step is considered to be the reaction between thiocyanate ion and iodine, giving an intermediate compound. The writer wishes to express his most sincere thanks to Professor Heinrich Rheinboldt, Director of this Department, for mggesting the problem and for his valuable guidance and interest during the course of this work. He is also greatly indebted to Professor H. Zocher, Ministerio da Agricultura, Laborat6rio da Produ@o Mineral, Rio de Janeiro, for his criticisms and suggestions and to Professor F. Feigl, of the same laboratory, for helpful discussions.
llrn
M . H. KURBATOV, G. B . WOOD, AND J. D. KURBATOV
REFEREKCES (1) BROWNE, A. W., AND HOEL,A. B . : J. Am. Chem. SOC.44,2106 (1922). (2) DODD,G . , A N D GRIFFITH, R . 0.:Trans. Faraday SOC.46,546 (1949). (3) FEIGL,F . : Z. anal. Chem. 74, 369 (1928). (4) FEIGL,F . : Mikrochemie 16, 1 (1934). ( 5 ) FEIGL,F . , AND CHARQAFF, E . : 2. anal. Chem. 74, 376 (1928). (6) FEIGL,F . : Qualitative Analysis by Spot Tests, 3rd edition (1946);Chemistry of Specific Selective and Sensitive Reactions (1949). (7) FRIEDYANN, E.:J. prakt. Chem. 146 (II), 179 (1936). (8) GRIFFITH, R . O., AND IRVINGR . : Trans. Faraday SOC.46, 563 (1949). , Acta Chem. Scand. 3, 872 (1949). (9)H o r 3 r a s - B a x ~K.: (10) L@VTRUP, S.:Compt. rend. trav. lab. Carlsberg, SBr. chim. 27, 63 (1949). (11) METZ,L.:2. anal. Chem. 76, 348 (1929). (12) RASCHIG, F . : Ber. 48, 2088 (1915). (13) RUPP,E.,A X D SCIED,A , : Ber. 36, 2191 (1902). A , : Ber. 36. 2766 (1902). (14) THIEL, (15) WEISS,J.: Trans. Faraday SOC.43, 119 (1947).
ISOTHERMAL ADSORPTION OF COBA4LTFROM DILUTE SOLUTIONS hl. H. KURBATOV, GWENDOLYK B . WOOD,
AND
J. D. KCRBATOV
Department of Chemistry, The Ohio Slate University, Columbus 10, Ohio Receibed July 17, lQ60 INTRODUCTION
I n a series of investigations, carried out in this laboratory (4, 5 , 6, 7, 9) and elsewhere (2, 3), concerning the adsorption of divalent cations on hydrous ferric oxide in the presence of ammonia and ammonium chloride, certain general relations have been found to hold when there was no complex ammonia ion formed in solution which interfered with adsorption. Some of the ions which have been studied are magnesium, calcium, strontium, barium, radium, manganese, cobalt, nickel, copper, and zinc. The general relations which have been observed are as follows: Increased adsorbent, adsorbate, and pH increase adsorption, while increased ammonium chloride decreases adsorption a t constant volume. The influence of changes in quantities of adsorbent and adsorbate has been presented in various types of isotherms. However, no single expression has been proposed for quantitatively relating pH, salt concentration, amount adsorbed, and quantity of adsorbent. The following is a study of the adsorption of cobalt on hydrous ferric oxide. In this investigation low concentrations of cobalt to 10" M or less) were used. Also, the quantities of adsorbent were large compared to those of the adsorbate: l e 5 gram-atom of iron to lo-* gram-atom of cobalt. Under these conditions and others to be described it was possible to evaluate the variables in a mass law expression relating concentrations of ions in solution and the quantity of adsorbent and cobalt adsorbed.
