6234 [CONTRIBUTION FROM
THE
DEPARTMEST O F CHEMISTRY O F LEBAXON \-ALLEY COLLEGE]
Kinetics of the Hydrolysis of Cyanic Acid
The rate of hydrolysis of cyanic acid has been determined by measuring the rate of addition of standard hydrochloric acid (L~/K,)(HOCX)2(OH-). The to maintain a constant pH. The rate equation may be expressed as RO = kl(H0CN) reaction intermediates are carbamic acid and a dimer of cyanic acid. S o salt effect mas found, but the conditions of the experiment may have minimized a salt effect. Studies were made a t 20, 25 and 35'. The free energy of ionization of cyanic acid has been determined as 5.0 kcal./mole.
+
The hydrolysis of cyanic acid in acid solution to give CO: and NH4+ has been long recognized. However, kinetic studies made on the reaction are incomplete.2
Experimental Materials.-Solid potassium cyanate reacts with moisture in the air t o form potassium carbonate and ammonia. Bottles of the reagent which have been on the shelf for some time smell strongly of ammonia when opened. .addition of B a + + t o an aqueous solution of the reagent forms a precipitate of barium carbonate. X-Ray analysis of such samples shows a large % of carbonate. However, newly purchased potassium cyanate does not indicate any carbonate upon X-ray analysis, does not smell of ammonia, and no precipitate is formed upon addition of B a + + to a fresh aqueous solution. I n preliminary experiments three specimens of cyanate were used: (1) C.P. potassium cyanate without any preliminary treatment; ( 2 ) C.P. potassium cyanate recrystallized from alcohol and water; ( 3 ) potassium cyanate prepared from urea and potassium ~ a r b o n a t e . ~.a11 three samples were stored in a desiccator. The reactions of all appeared identical and so studies were carried out on the C.P. sample without a n y further treatment. All other reagents used were C.P. grade used without any further purification. Distilled water was used throughout. Measurements of p H were made with a Beckman Model G PH meter with glass-calomel electrodes. The meter was standardized against commercial buffers of pH 4.0. Analytical Method.-The stoichiometry of the reaction is shown by the equation OCN-
+ 2H' + HZ0
= COz
+ NH?'
T h e ratio of H + t o O C S - was checked by adding a weighed sample of potassium cyanate to a n excess of standard acid. The solutions were allowed t o stand overnight. The excess acid was titrated with standard base. Results are given in Table I. Two moles of acid is used for each mole of cyanate, corresponding t o the stoichiometry above. The reaction is complete as written and the formation of urea from XH4+and OCN- does not occur rapidly enough under the conditions of the reaction to change the stoichiometry.
TABLE I RATIOOF H + TO O C S - USEDDURISG HYDROLYSIS Moles OCN -
x
103
3.25 2.15 3.60 4.67 3.76
Moles H + X I O 3
6.56 4.30 7.20 9.11 7.42
M o l e s H-/mole OCS-
2.02 2.00 2.00 1.95 1 98
(1) Chemistry Department, University of Xew Hampshire, D u r h a m , New Hampshire. (2) (a) 0. and I. Masson, Z. p h y s i k . Chem., T O , 290 (1910); (b) E. A. Werner, J . Chem. Soc., 113,81 (1918); (c) W. R. Fearon a n d G. C. Dockeray, Biochem. J., 'do, 13 (1926); (d) P. A. H. W y a t t and H. L. Kornberg, Trans. Faraday Soc., 48, 454 (1952). (3) "Inorganic Syntheses," Vol. 11, McGraw-Hill. I n c . , N e w Tiork, ?; Y , , 1940. p. Si:.
Kjeldahl determinations were run upon hydrolyzed samples of cyanate t o determine the nitrogen content. The O C i V to K ratio does not indicate anything concerning the reaction since total nitrogen is determined. The theoretical amounts of X were recovered. The first experiments were followed by the evolution of COa from the system. These were abandoned for two reasons: the reaction is so fast t h a t i t was felt t h a t the limiting factor might be the evolution of the COn and not the hydrolysis, and since there are two moles of H + used for each mole of O C S - i t was diflicult t o maintain a constant pH, even in a buffered system. I t is important that a constant pH be maintained because of the critical nature of the p H upon the reaction. httempts to follow the reaction by the precipitation of the O C S - as .%gOCN were abandoned because of the solubility of the .%gOCN and the pH change. Some runs were made by titrating to determine the change in acidity of the sample. In a buffered system the total change is small and the procedure is subject to much error. A . modification of a procedure of Pearson and Piette? \?-as finally used. Standard HCl solution was added to the reaction vessel at a rate such t h a t thc pH was maintained constant. The rate of addition of the HC1 can be used to determine the rate of the reaction, as shown below. The standard HC1 was added from either a 50-1111. buret or a 5-ml. buret. The volumes of acid used were between 19 and 40 ml. with the 50-ml. buret "and about 4 nil. with the 5-ml. buret, although a few runs were outside of these limits. The normality of the acids varied from 0.01 to 0.5 ,Vdepending upon the amount of cyanate used. -111 reactions were in vessels immersed in a water-bath. The temperature was controlled by a mercury thermal regulator and was constant to within 0.03". The electrodes uf the pH meter were immersed in the reaction vessel. -4 glass stirrer in the reaction vessel ensured complete and rapid stirring during the reaction. Procedure .-The K O C S was added to the reaction vessel containing water a t the temperature of the reaction. IVith rapid stirring the sample would dissolve in 15 to 20 seconds. Standard HCI was then added rapidly to the vessel to bring the pH of the solution t o the desired value. This addition took from 5 to 10 seconds. The HC1 was then added a t :I rate necessar>- to keep the pH constant. Buret readings were taken as a function of time. .After the reaction had proceeded through approximately two half-lifes the temperature Tras raised to 45-50' to complete the reaction. The total amount of HCI used enabled a calculation of thc original concentration of cyanate to be made. In sonie experiments Treighed quantities of dried Sac1 were dissolved in the mater before the K O C S was added to study the salt effect on the reaction.
Results At a constant pH the rate of addition of the HCl is proportional to the rate of hydrolysis. The proportionality constant is found as follows: Let: .lft = moles of HCl added a t time t M , = moles of HCI added a t completion of reaction Mo = moles of HCl added to reach desired pH = initial moles of HOCX Vt = total vol. of HCl added a t time f = ( H +I (OCS -) K, (HOC~%)(4) R . G . I'e.irwn a n d I.. 1%.Piette, TIIISJouRxAr., 76, 3087 1103-0
KINETICSOF
Dec. 20, 1956
IXITIAL
4.0 pH
CQ
0.18 0.54 1.09 1.35 1.97 2.11 2.73 3.56 4.00 4.77 4.80 6.25 7.17
Vw
0.9 3.0 6.3 8.0 13.2 14.6 19.1 24.2 28.4 36.8 37.0 55.0 66.7
19.0 19.3 21.4 38.6 26.5 38.0 48.9 17.4 19.5 23.3 23.4 30.5 35.1
RQ
Co
0.18 0.54 1.03 1.07 1.09 1.35 1.97 2.10 5.70 7.69 9.53
0.18
COSCEKTRATION O F CYANATE' Temperature 35' 4.2 4.3 CO Ro Vm Co Ro Vm
0.8 2.5 4.9 5.0 5.5 6.9 11.2 11.8 42.7 67.0 88.5
v w
19.0 19.3 22.0 23.0 21.4 38.6 26.5 45.0 27.8 37.5 46.5
Temperature 28' CO
9H
v w
4.0
0.5 1.6 3.3 7.6 6.4 8.1
Ro
42
4.1
0.5 1.3 2.9 6.6
0.4 1.2 2.2 5.2
0.18 0.54 1.09 1.35 1.97 2.11 4.14 6.30 6.78 7.17
co
4.51 5.54 5.62 6.50 6.91 7.60 8.40
0.7 2.1 4.3 6.5 9.3 9.3 22.8 40.3 42.0 47.4
Vm
22.0 27.0 27.4 31.7 33.7 37.2 41.0
0.5 19.0 0.18 1.7 19.3 0.54 21.4 1.09 3.7 38.6 1.35 4.7 26.5 1.97 7.9 38.0 4.26 21.7 20.2 5.75 31.0 30.7 10.6 81.6 33.1 12.1 109 35.1
Rn 4.0
4.1
Since -Ifm/?is the total cyanate present a t time 0 ??? -KO
( H +)
Ro
19.0 19.3 21.4 38.6 26.5 20.8 28.0 51.6 59.0
0.18 0.54 1.09 1.35 1.97 2.11 4.45 6.35 6.89 7.17
0.4 1.4 3.1 4.1 6.3 7.0 19.6 30.5 31.7 35.2
co
V,
V m
19.0 19.3 21.4 38.6 26.5 38.0 21.7 31.0 33.6 35.1
Temperature 20'
45.3 18.1 46.0 47.6 42.1 28.8 31.2
40
0.7 2.8 6.8 6.8 14.8 25.1 25.8
From (6) and (9)
(1)
;If0
d(H0CS)
+ d(OCX-)
(2)
The ratio of the decrease in HOCN to OCX- can be determined from the ionization constant K, Kd ( H O C S ) = d(OCK-) -~ IH+)
(3)
Since two moles of HC1 is used for each mole of OCN- that disappears but only one for each mole of HOCN -d(JIt
0.21 0.85 2.14 2.22 4.32 7.14 7.72
RO 3.8
- Mo
2
The HC1 that is added reacts in two ways: (1) to combine with OCX- to form HOCN; ( 2 ) to neutralize the NHBformed in the hydrolysis. For each mole of NH4+ formed a total of one mole of cyanic acid and cyanate ion has disappeared. -d(SH4+)
21.2 27.1
4.2
4.4 Co
0.5 1.8 25.3 5.1 30.5 5.0 33.9 11.0 41.0 38.0 30.4 20.4 12.5 42.8 21.0 a Rates are given as (moles liter-' min.-l) X lo*and concentrations in (moles liter-' X lo2). Vm is total volume of HCI used in reaction. Delivered from 5-ml. buret.
0.49 0.97 2.01 2.01 2.19 4.04
3.8ih 18.5 19.0 19.6 9.8 10.7 19.7
6235
HYDROLYSIS OF CYANIC ACID
TABLE I1 RATEAS A FUXCTION O F INITIAL 4.1
Ro
THE
- M o ) = 2Kc dMnocs $- d.lfHoch ( H +) iu +\
(5)
Starting with 100 ml. of solution a t time 0, a t time t
The rate of the reaction can be determined by Vt) V S . t and obtaingraphing ( M , - M t ) / ( l O O ing the slope of the curve. A sample plot is given in Fig. 1. The values of K , used were obtained from (1) and are given below. Data for the initial rates obtained are given in Table 11. The initial rates determined in this way were found to agree with rates determined from the product of the initial concentration of cyanate and the initial first-order rate constant determined graphically. The log of RO was plotted vs. log CO, the total initial concentration of cyanate. A line with a slight curvature, and a slope changing from 1 to 1.3 was obtained. This suggested that the reaction might be first order in cyanate. However, firstorder plots deviated from a straight line and the first-order rate constants determined from the data indicated a linear dependence on the initial concentration with a positive intercept. This dependence might be explainable if the reaction were second order. Second-order plots deviated from a straight line and the rate constants calculated on the basis of second-order dependence again showed a dependence on initial concentration. Because of this and the positive intercept the second-order reaction was eliminated. It is possible that the dependence is not on initial cyanate concentration but upon ionic strength. To change the ionic strength independently of the cyanate concentration, NaCl was added to some solutions before the start of the reactions. Results
+
ALEXANDER K. ANELL
6238 30
VOl. 7s
I-
Initial cyanate coIicn., trioles/l. X 102. Fig. 2.-Initial rate of hydrolysis as a function of total cyanate concentration: I, PH 4.00, 28"; 11, PH 3.80, 20"; 111, PH 4.00, 20"
TABLE 111 woix IsrrIsL R A ~ E EFFECTOF I o s r c STRENGTII Temperature 3.5' 'l'atal cyanate moles/l.
PH
4, iJiJ 4.10
4.20 4.30
x 102 1.09 1.09 1.09 1.09
Ionic strengths X 10'
4.00 4.00 4.00 4.00
4.40
1,09
1.00
4.10
1 , 07
9.13 3 . 2 5.52
Experimental r a t e X 102
6.9
5.5 4.3 3 .7 3.1
KL6
&2 5.4 4 5 4.0 3.3 5.3
Ku
29.i
25.6 '1 2 19.2
15.8 5.0 82 4.9 4.10 1.03 5.1 42 2.10 4.10 11.8 11.5 39 Cala Ionic strength increased by addition trf S a C l . culated using K = aco bco2, assuming dependence on initial From Fig. 2, 3, or 4 assuming cyanate concentration. dependence on ionic strength.
+
Initial cyanate concn., Inoles/l. X 102. Fig. 3.-Initial rate of hydrolysis as a futiction of total cyanate concentration: I, pH 4.10; 11, PH 4.30; 111, PH 4.40; all temperatures 35".
The experimental data cannot be explained by either a simple rate equation or a salt effect. The linear dependence of the first-order rate constant upon initial cyanate concentration suggests the form: k = a bco for the rate constant or RO = aco + bco2 (11)
+
for the initial rate. The data were treated by curve fitting methodss and found to fit an equation of this form. The values of the parameters a and b were determined from the data by the method of least squares. The agreement with the experimental data is shown in Figs. 2 , 3 and 4. Values of a and b determined are given in Table IV. The significance of the parameters a and b may be investigated fxrther. The initial concentration of HOCN, d l o ,is related to the initial concentration of total cyanate (0 by
Initial cyanate concn., moles/l. X lo3. Fig. 4.-Initial rate a5 J. function of total cyanate concetitr.itiott: I, pH 4 10, 2 8 O , 11, PH 4 00, 35'; 111, PH 4 2 0 , .'EJ~;I\', p H 4 20, 28'
Solving for cl)and substituting in (11) (5) I. S.a n d E. S . Sokolnikoff, "Higher Mathematics f o r Engineers a n d Physicists," 2nd ed., hIcGraw-Hill Book Co., Tnc.. New I'ork, S . \i 1941, C h a p t e r XII.
+)
+ K,
b [(H +)
A-
K,j2 - l " Z
(13)
Dec. 20, 1956
K I N E T I C S O F THE
TABLE IV VALUESOF PARAMETERS IN Ro = aco Temp, 'C.
fiH
35.0 35.0 35.0 35.0 35.0 28.0
4.00 4.10 4.20 4.30 4.40 4.00 4.10 4.20 3.80 4.00
28.0 28.0 20.0 20.0
a X 102, mh-1
+ bc2
b, 1. moles-'
5.0 4.4 3.7 3.2 2.7 3.0 2.7 2.1 2.9 2.2
HYDROLYSIS OF C Y A N I C
mh-1
0.59 .50 .40 .42
.31 .32 .30
constants are found to be: 6.4 a t 20', 9.0 a t 28" and 15.5 a t 35" for kl X lo2 in min.-l and 0.59 a t 20°, 2.9 a t 28' and 5.3 a t 35' for ka X lo4in (moles/ 1.)-l min.-l. Temperature Dependence.-If log k is plotted vs. 1/T in the conventional plot for activation energy the curves given in Fig. 6 are obtained. An explanation of the deviation from the normal linear function is given in the discussion.
+ kz(H+)"(HOCN)'
x
(H+)m+
f
kl = a[(H+)
+ K O ] = A and b[(H+) + KO]* = B
( H + ) " + 'kz
30
-
20
-
5
(14)
From a comparison of (13) and (14) it is seen that 1
40
-
50
.25 .08 .07
If the assumption is made that the reacting forni is cyanic acid the rate equation should have the form Ro = ki(H+)"(HOCN)
6237
.ACID
L
(15)
3
N
A plot of log A or B in (15) against pH should give a straight line the slope of which will give the order with respect to (H+) and the intercept the rate constant for each step. This plot is shown in Fig. 5. The values of K O used are given below under the discussion of ionization constants for HOCN.
2 X
325
330 3.35 3.40 3.45 ( i p ) x 103. Fig. 6.-Temperature dependence of rate constants: 6, unimolecular rate constant; 0 , bimolecular rate constant.
10 9 8 m.
0
7
6 X Q 5
8
B
4
X v 3
2
1 3.7
3.9
4.1
4.3
4.5
PH . Fig. 5.-Dependence of uni- and bimolecular rate constants upon PH: unimolecular, ( A ) 35'; (B) 28"; ( C ) 20"; bimolecular, ( D ) 35'; (E) 28'; (F) 20'. Piumerical value X 3 plotted in F on ordinate.
The slopes of all curves in Fig. 5 are 1 within experimental error. The unimolecular reaction is 0 order with respect t o (H+) and the bimolecular reaction -1 order with respect to (H+) or first order with respect to (OH-). Values of the rate
Ionization Constant of H0CN.-The ionization constant can be determined from (I). The values increase with ionic strength, although the experimental error is too large to determine the exact dependence. However, in the experiments with the added salt the constant is larger than without the salt, helping to confirm the absence of the salt effect upon the rate. The best values of the ionization constants calculated a t zero ionic strength appear to be: 2.1 X a t 35', 2.0 X a t 28' and 1.9 X at 20'. The heat of dissociation and free energy of dissociation are 1.3and 5.0 kcal. 'mole. Previously reported values are 2.2 X for KOa t 2506 and 5.0 kcal. 'mole for AF.' Discussion of Results The following mechanism is proposed for the hydrolysis
+
HC O C N - z HOCN _c H S C O HNCO HzO +NHqCOOH SHzCOOH +NHI CO, H++ XHi+ NH3 2HNCO +NH?COOCN NHZCOOCN OH-+ NHzCOOH OCN-
+
+
+
+
+
(ii) R.Taufel, C. Wagner and H. Dunwald,
fast slow fast fast fast
(1) (2) (3) (4)
slow
(6)
2.Elekliochem., 34, 115
(1928)
(7) J C .AlcGowan. Che!rziati.y a i d
(5)
1 i i d z i ~ t i . yti32 (1948).
6238
ARTHURVEIS AND
The rate-determining steps are the formation of the carbamic acid. In the bimolecular mechanism this is from the reaction of a dimer of cyanic acid with an OH-. The existence of a dimer of cyanic acid has been proposed previously to explain reactions of the acid.* Werner and Gray, on the other hand, do not feel that there is evidence for the existence of the dimer.g It is felt that the rate equation obtained strongly suggests the existence of the dimer, even though it may be in very low concentration. A dimer of phenyl isocyanate is known to exist.1° Whether the form of the dimer is NH2COXCO or NHZCOOCN cannot be distinguished from this work. The dimer is attacked by an OH- t o form
an intermediate of the form
[
xH2C--OCS
]
which
XH then dissociates to carbamic acid and OCK- in the rate-determining step. Whether cyanic acid is HOCN or HNCO or a mixture of the two has not been completely determined. Infrared spectra of the vapor and liquid appear to show that the acid is HNCO with less than 0.2% HOCN present.” In aqueous solution the tautomerism would be more probable than in the vapor or liquid. Werner and FearonI2 suggest that the mixture is an equilibrium between the two (8) (a) T. L. Davis a n d K. C. Blanchard, THIS JOURNAL. 51, 1806 (1929); (b) H. W. Blohm a n d E . I. Becker. Chem. Reus., 51, 471 (1952); (c) M. Linhard, 2. unorg. allgem. Chem., 236,200 (1938). (9) A. E . A. Werner a n d J. G r a y , S c i . Proc. Roy D u b l i n S o c . , 24, 300 (1947). (IO) C . J . Brown, J . Chem. Soc., 2931 (1955). (11) E . Eyster, R. Gillette a n d L. Brockway, THIS J O U R N A L , 62, 2326 (1940). (12) E . A. Werner a n d W. R. Fearon, J . Chem. Soc., 117, I356 (1920).
[COSTRIBUTION FROM THE
JEROME
COHEN
Vol. ‘78
forms and use the formation of the polymers of the acid as procf. The formation of HNCO is favored as the temperature increases. Evidence supporting the tautomerism is obtained from the temperature dependence of the rate constants (Fig. 6 ) . If it is assumed that only one form of the tautomer can react by a given mechanism, the total concentration of acid used to determine K should be replaced by the concentration of the reacting form. If Kt is the constant for the = HOCK HKCO and (HNCO) tautomerism, &lo would be [K,,’(l K,)]A,. This ratio would be included in the rate constant and a linear temperature dependence would not exist. It is believed that the HXCO form is the one that reacts because of steric considerations. The mechanism proposed is a molecular one and there should be no primary salt effect. Because of the equilibrium which exists between the cyanic acid molecule and the ions, a secondary salt effect might be expected.13 Since the concentration of OCN- and HOCN is of the order of 100 times the concentration of H + the increased dissociation of HOCN because of the added salt will be very small and will not appreciably change the concentration of un-ionized acid. Under these conditions the salt effect will be very small and less than the experimental error. Acknowledgment.-The author is indebted to Drs. J. K. Dixon and G. L. Simard for helpful discussions during the early part of this work, and to Dr. Henry Kuivila who first suggested the dimer as an intermediate in the reaction.
+
+
(13) A A Frost and R. G Pearson, “Kinetics and Mechanism,” John Xt’iley and Sons, I n c , S e w York, N Y , 1923, p 200
XSSVILLE, PA.
RESEARCH DIVISION, -4RMOUR & CO ]
A Non-random Disaggregation of Intact Skin Collagen BY ARTHURVEIS AND
JEROME
COHEX
RECEIVED MAY7 , 1956 Studies of the disaggregation of intact bovine hide collagen show the presence of lateral polypeptide chain cross-linkages of different reactivity and, probably, different distribution along the fibrillar long axis. Under the acid-extraction conditions there appear t o be at least two classes of such cross-linkages and the order in ahich these are broken determines the molecular weight distribution of the solubilized portion. If the extractions are carried out under PH and temperature conditions where peptide bond hydrolysis is not significant then discrete soluble fibrillar fragments are obtained which may retain the gross features of the peptide chain configuration present in the native tissue. The onset of peptide bond hydrolysis is readily recognizable.
Various studies of the conversion of collagen to gelatin have had two major objectives: (1) to throw some light on the structure of intact collagen; ( 2 ) to explain the variations in the stability of the collagens from different Ames’ work’ led him to propose two possible models for collagen, a multi-chain structure in which the peptide chains are held together by non-peptide covalent cross-
linkages of some sort and a single chain model in which the peptide chains are too large to be water soluble, but become so as certain particularly weak peptide linkages are broken. In spite of additional data from other quite different points of view6-’ one coiild not chose between these alternatives. Our studies of the niild thermal solubilization of collagen in the acid pH range2 showed, however,
W.hf. Ames, J. Sor. Chem. Ind., 66, 979 (1947); J . Sci. Food Agric., 3, 454 (1952). (2) A . Veis a n d J. Cohen, THISJ O U R N A L , 76, 2476 (1954); 77,
( 5 ) P. hl. Cowan, S. McGavin and A. C. T. N o r t h , A’atwc, 176, 1062 (1925). (6) a’.If.Ames, J . Sci. Food A g r i c . , 3, 579 (1952). ( 7 ) R. D. Harkness, A. X I . hlarko, H. M. Muir a n d A. l\-euberger. “ N a t u r e a n d Structure of Collagen,” Academic Press, New York. N. Y.. 1953,pp. 208-212.
(1) aid
23G4 (1955). (3) A . Veis, D. S . Egpenberaer and J. Cohen, ibid., 77, 2368 (1925). (4) K. H. Gustavson, J . Amev. Leather Chem. Assoc., 60, 239 (1955).
