THE CATALYTIC M I S I - 1 I U ~ IPOINT’ BY M, BERGSTEIS2 AXD MARTI?; KILPATRICK, JR.3
On the basis of the protion theory of chemical reactivity Rice? has predicted that for those reactions catalyzed both by hydrogen and hydroxyl ions in aqueous solutions the point of minimum catalytic activity (maximum stability of the compound) will be practically the same for all reactions and will be on theacid side. In support of this view the work of Karlssonj on ethyl and methyl acetates was cited. Karlson working at 8 j . j j oC. found a minimum for methyl acetate a t pH 4.70 and for ethyl acetate at pH j.10. From the work of Dawson and Powis6 on the autocatalytic reaction between acetone and iodine Rice calculates a minimum at pH 4.7. More recently Bolin7 has reported the following points of maximum stability.
TABLE I T “C 20 2 0
20
80 40
PH
Ethyl ester of formic acid ’, ” ” monochloracetic ,, ” ” aminoacetic Benzamid Acetanilid
4 6j 1 8 3 8
58 6 2
On the basis of these results Bolin criticises the unhydrated ion theory as applied t o hydrolysis. He states that the results for benzamid are not at all in harmony with the theory of Rice, In order to apply the theory one must assume either that the benzamid affects the hydration of the hydrogen and hydroxyl ions differently or that the dissociation conftant of water is changed by the benzamid. I n view of the small solubility of the benzamid and the pH values these appear improbable. E-, 1 9 q - 3 5 Sational Research Felloir Rice J -Am Chem Soc , 45, 2808 f 19231 Iiarlsson Z anorg Chem 119, 69 (1921 Dairson and POXISJ Chem S o ( > .101, 1502 11912 Aolin Z anorg Chem 143, 2 0 1 (1925) 8 Iiarlsson Z anorg Chem , 145, I (19251 4 Euler and L a i r i n . Sr Vet .Aka({ lrkiv f u r Kc~rni 7 , 30 (1920) ~
1617
THE CATALPTI C J l I N I M U J l POIKT
In his earlier work Karlsson reported a distinct shift of the minimum point toward the acid side upon the addition of salts. His later results do not show an appreciable displacement of the minimum point. Hn attributes the discrepancies t o experimental errors in the earlier work. Karlsson's experiments were carried out in buffer solutions. The effect of change of the total equivalent salt concentration on the reaction rate in buffer solutions has been discussed in several recent papers.' From the protion theory as presented by Rice me would expect the addition of salts to shift the minimum point t o the acid side. I n addition the effect of temperature would be to shift the minimum point toward the alkaline side. These considerations would not apply in buffer solution but should apply to the autocatalytic reaction between acetone and iodine in aqueous solution. This reaction has been the subject of a number of studies in acid solution. The reaction in acid solution may be represented by the stochiometric equation CH3COCH3 1 2 H+ +CH3COCHZI H+ I-
+ +
+
+
Since the rate of the reaction is independent of the Concentration of I?, Dawson and Powis gave as the equation expressing the velocity dx d t
=
k(a-x)(c
+ x)
where x is the amount of iodine which has disappeared a t time t , a is the initial concentration of acetone, and c is the initial hydrogen ion concentration, I n the actual experiments the solution was I 11 in acetone and 00 j 11 in iodine, x being negligible in comparison with n . They wrote therefore. dx dt = k a(c x), and integrating, they obtained - kat. In -
+
~
+ C
Extrapolating to zero time they found c to be equal to 2.1 X IO-: (themean of three observations). When a small quantity of hydrochloric acid was added to the solution and this concentration of hydrogen ion 2 x I O - j S subtracted from the initial value obtained byextrapolation, c wasfoundtoequal 1.1xIO-;. They concluded that the catalytic substance whose activity \vas represented by this initial hydrogen ion concentration was actually alkaline in nature. K h e n sodium hydroxide was added to the reaction mixture extrapolation yielded a value for c of z X 1 0 - j . On the whole they could present no satisfactory theory to explain the anomalous catalytic activity of pure water. Their trouble lay in the fact that they neglected the not inappreciable catalytic activity of the hydroxyl ions in pure water. One of us (11. B.) has shown, in an unpublished research, that the rate of the reaction at pH j is approximately twice that at pH 2 . As has been indicated above, neutral water is catalytically active. Therefore, the iodine-acetone reaction starting in a neutral solution will proceed at a rapid rate, liberating HA ion and removing OH- ion until a point is reached where the products of the concentraRronsted and Teeter: J. Phys. Chem , 2 8 . j 7 9 (1924); Bronsted and Icing: J. Am. Chem. Soc., 47, 2 j 2 3 (1925); Iiilpatrick: 48, 2091 (1926).
1618
M . BERGSTEIX A N D M A R T I N ICILPATRICK, J R .
tions of each of them hy their catalj-tic activities are equal. This condition is reached practically instantaneously and cannot be measured because only a minute quantity of iodine has disapreared. At this point the reaction ha< minimum velocity and gradually accelerates because of the liberation of hydrogen ion. K e may consider this point to be the initial point of the reaction. It is actually the catalytic minimum point. Dawson and Powis's reaction velocity equation is subject to the criticism that it does not recognize the esistence of two interdeFendent catalytic SUI)stances at the initial point. If n e ascribe half of the effect that they observed to hydrogen ion, we may assume that catalytically neutral water as a hydrogen ion concentration of IO-:, hut this statement iq. as the following analysis discloses. far from esact. Consider the reaction CH3COCH3 1, Hf OH- ---t C"?COC'H?I IzH+ OHLet a = initial concentration of acetone, c = initial concentration of hydrogen ion, d = concentration of hydrosyl ion at time t , s = mols of I? reacted at time t . k, = catalytic activity of hydrogen ion. k,, = catalytic activity of hydrosyl ion.
+ +
+
+ +
+
+ +
Then ds dt = (a - s) [kH(c s) k,, d]. I t is known that (c 4- a) s). The constants 1 0 - l ~ and we may therefore write ti = 1 0 - l ~ (c k, and k,, are also related. We shall not consider the secondary effect: the formation of iodoform which is also catalyzed by hydrosyl ion in alkaline solution. In catalytically neutral water the rate of reaction due t o hydrogen ion is equal t o the rate due t o hydroxyl ion or k,c = k,, IO-'' c k,, = c2 IO-%, or The equation then takes the form
d
+
=
-dx_ - k ( a - s ) (c
+
x 4dt c+s ) where k = k,. The solution of this equation is C d (c .I2 c2 kt = a+c In a s (a c)' (a c2 ~
+ +
+ +
+ + e')tan-l-
On removing negligible quantities the above equation yields kt,
=
I
-In a
d ( c
+
s)2
a - s
+cz
+K -
When t
=
0, s
=
0. Therefore, K
In cd -, and a a
= -I
c+x+I< C
T H E CATALYTIC 3IINI3IC3I P O I S T
1619
If the initial molar concentration of acetone is about twenty times that of iodine (as is true in the experiments reported), x is negligible in comparison with a and the equation may be simplified to k
=
at
~n
1,'
/(e
+ xi? + c2 2
c-
In order to make a determination of c the following device is resorted to. For large values of t (and therefore of x). c? and z cx are negligible in cornparison with x2. TTe insy write in chis case kat+lnc.\/l=lnx
I1 is now possible to determine c without knowing either k or a. Plotting In s as ordinate against t as abscissa the intercept is equal t o In c . \ / T Thus c may be determined. The experimental details follow. Solutions were made up containing .z"( potassium iodide and approximately .ooj ?tliodine. To this was added a known weight of acetone from a calibrated pipette, the flask was filled to the mark m-ith distilled water and shaken, and the time was noted as zero time for reactions conducted at z j". TITOsamples were removed with a calibrated pipette and titrated against standard sodium thiosulphate using a weight bure tte. The remainder of the solution was transferred through a quick running siphon to glass-stoppered 80 cc. bottles and the bottles were place3 in the thermostat. When the experiments were conducted at temperatures of 30" and above, a time 1 2 minutes after the bottles \{-ere place 1 in the thermostat was noted as zero time. Two 2.; cc. samples were taken from each bottle at regular interval. and the iodine concentrations \?ere determined by titration against sodium thiosulphate. The reactions were stopped by running the solutions into L?, mixture of acetate buffer of approximately pH j an 3 a known weight of thiosulfate. Starch solution was added and the solution titraced t o the disappearance of the blue n-ith sodium thiosulphate form a weight burette. K h e n the experiments were run a t 2 j " and 30' the time noted as final for each sample was the time at which the reaction mixture started to run from the pipette into the titrating flask. At higher temperatures the reactions were stopp2red by placing the reaction vessels in a freezing mixture. The final time noted in such cases was taken as two minutes after the bottles were placed in the freezing mixLure. Check runs were macle without any acetone using only the iodine sollition and the error due to volatilization of iodine from the glassstoppered reaction vessels was found to be negligible. Because, as previously stated the b j drogen ion concentration of the water almost immediatel-y came to approximately IO-^, conductivity water was not required, ordinary distilled water being satisfactory, and any alkalinit) due to potassium iodide could be neglected. The value for the catalytic minimum point in the aboveexperiment was obtained as follows. The drop in titre of the solution in terms of grams of ?;a?S20asolution per 24.92 cc. of sample taken was plotted along the logarithmic axis on semi-logarithmic paper against the time plotted along the ordi-
1620
31. BERGSTEIX AND MARTIN KILPATRICK. J R .
TABLE I1 The Autocatalytic Iodine-Acetone Reaction at z so 2 L of soln. containing 2 7 . 5 4 g. acetone; 2 % K I ,00.;344 A 1 Sa?SgOs
24.92 cc. sample
Time (hours)
Titrc
Time
Titre 12.25
2 4 53 24 16 22 7 7
192.7 196.8 I97 . 4 216.I
21 3 2
220.7
24 66
0
26.j 73.6 119 9 I42 I 16j.9
11.29 10.70
3 . IO I 25
18 69
Catalytic minimum point [H+] 1.57 X
IO+;
PH 4.81
nary axis. The points plotted, with the exception of the earlier ones, lay on a straight line which, of course, intersected the logarithmic axis a t zero time. The point of intersection represented the apparent drop in titre per 24.92 cc. at zero time. 3lultiplying this value by the molarity of the sodium thiosulfate and dividing by the volume of the sample taken gives the apparent initial hydrogen ion concentration which is c d y For experiment 103 =
C d ,
.106 X .005%z44
= 2.274
x
Io-:
24.92
c
= 1.57
x
Io-5
Table I11 summarizes a series of experiments carried out in the same manner. In the experiments at 40' C . and 45' C. the concentration of the acetone was .13; XI, in the others .237 M. 0 . 2 5 KI was present in all experiments.
TABLE I11 Tcmprrat urc
"C.
1Iinimnm point
[H'] X IO-: 1.57 I .30
PI4
1.66 1.57
1.49 I .62 I.j
7
I.32 I.6j 1.56 1.79
1
.74
From Table 111it is evident that any variation in the minimum point with temperature is within the experimental error of the measurements. From data on ester hydrolysis the minimum point can be calculated at 20' (', to
1621
T H E CATALI-TIC' MINIJIUJ1 P O I S T
be a t PH j.47 and a t 40' C. a t PH j.35. The mean value determined for the catalytic minimum point in the above experiments ranging from z jo C . t o 4 j 0 C . , i s [ H + ] = 1 . j 8 X 1 o + o r p H 4.8. Table IT- summarizes: the results in the presence of neutral salts. -411 experiments were carried out at 35' c'.
TABLE IT Salt addctl Molarity
1Iininium point
[H-] X IO-^
P€I
I t is apparent that if there is any neutral salt effect it is within the experiineiitnl error of the measurements. Other Catalytic Minimum Points l17i.is1 studying the saponification of methyl acetate in water found n iniiiirnuiii at pH 5.42 at 24.8' C .
Euler and Laurin found pH 5.10 for ethyl acetate at 69.3' c'. Hudson' studying the mutarotation of glucose, using the data, of Osaka3 a,nd his own calculated the minimuin t o he at pH 4.64at 24.7' c'. More recently Eider4 reporcs the point of minimum catalytic activity for the same reaction at p H .;.o for j . 2 O c. Boglie; in discussing the hydrolysis of collagen noted that the region from pH 4.3 t o 6.0 should he avoided as the hydrolysis is very slow at that concentration of hydrogen ions,
The relation of the miniinuni point to the isoelectric points of protein? has heen discussed by Thomas and Kelly.6 l l a n y other examples of phenomena which involve consideration of a minilnilin point in the neighborhood of p H j may he found in the fields of industrial and biological chemistry. Onc of us (11.B.) wishes t o express his thanks to Dr. F. 0 . Rice of Johns Hopkins L-niversity for directing his work in this investigation. 1
d
2 1:
IYijs, Z. physik. Chcxm.. 11. 492 (1893). Hudson: J. Am. Chem. SOP.,29, 1571 ' 1 9 0 7 1 ;31, 1136 r1909!. Osaka, Z. physik. Chrm., 35. 702 /1900). Euler: Z. anorg. Chem., 146, ~j (19251. Bogue: J. I n d . En$. Chem., 15, I I j1 (1923 1 . Thomas and Kelly: J. Ind. En-. Cliem.. 15, I 148 (1923'1.
1622
11. BERGSTEIS .1ND MARTIN KILPATRI CK, J R .
Summary I. The catalytic ininimum point for the reaction between iodine and acetone h2s been found to he a t pH 4.8.
This value is unaffected within the limits of experimental error by 2. either temperature or neutral salts. 3. The results do not support the unhydrated ion theory of catalysis as proposed by Rice. S e 1 c . I’ork, x. Y . Rolf?more, X d .
