4014
HENRY G. KEIVILA
V O l . 77
[COSTRIBUTION FROM THE DEPARTMENT O F CHEMISTRY, USIVERSITY O F N E W HAMPSHIRE]
Electrophilic Displacement Reactions. VI. Catalysis by Strong Acids in the Reaction between Hydrogen Peroxide and Benzeneboronic BY HENRYG. KUIVILA RECEIVED MARCH28, 1935 The kinetics of the reaction between hydrogen peroxide and benzeneboronic acid has been investigated using aqueous solutions of sulfuric, perchloric and phosphoric acids as solvents. The rate in each case increases more rapidly than the stoichiometric concentration of acid. HOcorrelates the data. for phosphoric acid; log CLHX X auZ0 correlates the data for perchloric acid. So simple correlation of the data for sulfuric acid has been found.
An earlier investigation of the kinetics of the reaction between benzeneboronic acid and hydrogen peroxide which yields phenol and boric acid as products dealt with the characteristics of the reaction in the pH range 2 to 6.3 Two ionic reactions were clearly distinguished in this range: each is first order in peroxide; one is first order and the other is second order in boronic acid; both have specific rate constants which increase linearly with the first power of hydroxide ion concentration. I n addition the data indicated the appearance of a pH-independent reaction which is detectable between p H 3 and p H 1.5. This paper is a report on the kinetics of the reaction beginning a t p H 3 and extending to about 9 molar solutions of sulfuric, perchloric and phosphoric acids. Results and Discussion Perchloric Acid.-The data for solutions of perchloric acid are assembled in Table I. It is apparent from the first four items in the table that a minimum extending over about one-half fiH unit exists in the observed specific rate constant. 1. mole-' set.-' is Therefore the value 7.7 X taken as the value for an uncafalyzed reaction. The specific rate constant ka for the acid-catalyzed reaction is then the difference between kobs and 7.7 X
10-4. TABLE I KINETICDATAIN
0 007; 0030 .007> .0100 ,200 ,750 2 469 3 73 4.12 5.93
8. 15
AQUEOUS
2.625 0.781 2.32" ,771 ,773 2 . l.ja 2.00= ,793 0.70 1.56 0.0:' 3.94 16.0 -0.88 31.9 - 1 44 -1.84 44.0 137 -2.58 605 -4.31
.... .
.
I
.
....
.... 0.79 3.17 I:.? 31.3 43.2 130 60.4
..... ..... .....
...
.....
.. .. ..
3.102 2.4913 1.818 1,507
-0.808 -0.222 f0.507 0.9.58 1.0-41 1,707 "333
2.31 2.28 2.34 2.47 2.40 2..i7 2.50
...
1.3GI 0.8fjO 0.219
+
log k, = coust. log af*zoX UHCIO, (1) h log-log plot (Fig. 2) gives a line with a slope 0.95 This type of correlation in strongly acid media
appears to be unique and resembles the termolecular mechanism for general acid and base catalysis originally suggested by Lowryj and more recently favored by Swain.6 In the present case the transition state complex must be built up from one molecule each of boronic acid, perchloric acid, hydrogen peroxide and water. Xs one possible example reactions 2 and 3 could lead to structure I from which assembly I1 could be formed uia the transition state complex (reaction 4). Two molecules of water are needed in I : that demanded by equation (1) and that formed in reaction ( 2 ) . This can be seen by adding reactions (2) and (3). CsHjB(O€f)2
PERCHLORIC ACID
... ...
acid concentration leads to a seven hundred fiftyfold increase in rate constant. Thus the latter is not a simple function of acid concentration. Furthermore i t is not a simple function of Ifo4which varies over five units while log k, varies over less than three units. This is also shown in Fig. 1 which represents a summary of all of the data on acid dependence of the reaction over a range of more than ten powers of ten in acidity. The last two columns of Table I demonstrate that k, can be correlated with the product of the activities of perchloric acid and water, ey. 1
CsHsB
0011 C J I;B( o I ~
(,,,
0013
fast
+ HOOH
+ IIOH
(2)
fast
+ H C ~ O I$- 2IIOH I-I
(3)
..
Measured. * L. P. Hamniett and A . J . Deyrup, THISJ O U R K A L , 54, 2721 (1932); L. P. Harnmett and hl. A . Paul, ibid., 5 6 , 827 (1931). Units 1. mole-' set.-'. kobs -0.00077 = k,. e J . N. Pearce and A . F. Nelson, ibid., 5 5 , 3075 (1933). ' A = log k , - log C Z H ~ OX arrc~o,; mean value = -2.42.
Items in the table beyond the fourth reveal several significant facts. First, a forty-fold increase in ( I ) F o r preceding paper in this series see XI. G. Kuivila and R. hI. Williams, THIS J O U R N A L , 76, 2679 (1954). ( 2 ) Substantial support of this work by t h e Office of Kava1 Research is gratelully acknowledged. (3) H. G. Kuivila, T H I S J O U R N A L , 76, S70 (1934).
c611j-B-oII
CtiHj-0
I
I
H
I I /O\ c1
\
'so-I€
I
H
I
B-OII
I1
/
€10 Ha0+
r
X proton transfer from the perchloric acid molecule to a peroxidic oxygen facilitates ionic cleavage a t ( 4 ) L. P. H a m m e t t and A. J. Deyrup, ibid., 5 4 , 2721 (1932); L. P. H a m m e t t , "Physical Organic Chemistry," LlcGraw-Hill Book Co., New York, . ' h Y.,1940, p. 267ff. ( 5 ) T. hf. L o a r y and I . J. Faulknet, J . C / w o . Soc., 127, 28833 (1925). (6) ( a ) C . G. Swain and J. F. B r o w n . TEIS J O U R N A I . , 74, 2.i33 ( 1 9 5 2 ) . and preceding papers in t h a t series; ( b ) i b d , 7 4 , 2588 (19.52).
4015
RE.XTIONBETWEEN HYDROGEN PEROXIDE AND BENZENEBORONIC ACID
Aug. 5, 1955 I
t
\ "20,
t Fig. 1.-Dependence of velocity constants on acidity in water and strong acids. The points on the broken line are k l ; all others are corrected for k , . values of kabs = k ,
+
b leaving an electron deficient oxygen to which the phenyl group migrates. Water molecule q is in a position to expedite the migration by supplying a pair of electrons to fill the sp2orbital vacated by the departing phenide group. \Vater molecule r is in a position to solvate the proton thus released. Details such as the extent of covalent bonding between q and boron or the degree of participation by phenide in promoting the cleavage of the peroxide link must be unspecified a t this point. Phosphoric Acid.-Kinetic runs were made in phosphoric acid solutions ranging in concentration from 7 to 05yo. The data are presented in Table 11. TABLE I1 KIXETICDATAIS AQUEOUS PHOSPHORIC ACID
~~
-I
0
Haa 0.99 0.28 -0.61 -1.29 -1.98
100
100
kob8.h
ksb
0.255 0.332 1.40 1.32 12.1 12.0 49.7 49.8 214 214
Ad
ka
aH80
2.593 1.60 -0.20 1 . 8 7 9 1 . 6 0 +0.50 1.43 0.921 1 . 5 3 2.11 0 . 3 0 4 1.59 1.65 -0.330 2.81
BA
+P +
a E. Heilbronner and S. Weber, Helv. Chim. Acta, 32, Units, 1. mole-' set.-'. e K. L. Elmore, C. 1513 (1949). M . Mason and J . H. Christensen, THISJOURNAL, 68. 2528 (1946). - A = HO log k.; mean A = 1.59. a A = log U A ~ ~ O ~ / U-log H , ~ k,; mean A = 2.40.
+
Here again it is clear that the rate does not vary in a simple manner with the concentration of phosphoric acid since a thirteen-fold increase in its concentration produces an eight hundred forty-fold increase in specific rate constant. However, the rate constant is now correlated accurately by Ho (columns 2, 5 and 6). This appears to be the first recorded example of a reaction whose rate constant is correlated by one function in solutions of one acid and by a distinctly different function in solutions of another acid. Dependence on HO demands a transition state composed of a molecule of boronic acid, one of peroxide and a proton. This might be similar to that formed from I with water molecule r omitted and the perchloric acid molecule replaced by a proton. 111. Thus the combination, HX HzO in aque-
+
-= __ KIII
H30'
fast
A' 2.39 2.38 2.35 2.41 2.48
'ancio, ) ,
ous perchloric acid plays the same role as H + in aqueous phosphoric acid. In the latter case we may write the following equations where BA is boronic acid, P is peroxide and I11 is the intermediate
-a
--logo
(aHzO
of rate constants in perchloric and sulfuric acids.
Fig. 2.-Correlation
log HiPOc, moles/ I. 0.726 2.267 5.19 6.87 9.67
2
I
log
111
slow R
CeHr-B---OHz I OH
+ H20
(5)
t
products I11
which is converted to products through a transition state X +. Now, the rate expression is
or UH+ fBA fP
k. = kKrri ___ fx+
(8 )
By d e f i n i t i ~ nho ,~ = aH;fB/fBH+
where f B and f m + are the activity coefficients of a base and its conjugate acid. The proportionality between k, and ho implies that f B q f P / f X + is affected by the medium in the same way as . f B l f B K +, a reasonable implication for a transition state consisting of an activated complex of BA, P and a proton. From columns 7 and 8 of Table I1 it is evident that the rate constants are also proportional to the ratio of the activity of phosphoric acid to the activity of water.? An intermediate like I V can accommodate this fact. The difference in mechanism between phosphoric and perchloric acids can be at(7) Ho must therefote be a linear function of the log of this ratio. This obtains also for hydrochloric and perchloric acids but not for
sulfuric add: H.G.Ruivila, J . Phys. Chrm., in press.
30 1(i
EARLEVANHEYNINGEN
tributed to the fact that much un-ionized acid is present in the solutions of the former but not in solutions of the latter. The passage of IV through the transition state to Iris an example of bifunc-
I;i ,
O/o\\\H, '0 i CtjH:--B----O +-- ?-OH 6 1I 611 11.
0
C6H6-0
t
13- 0- -I'--OII
oIr
OH
TABLE I11 KISETICDATAIX AQUEOUS SULFURIC ACID
0.193 0 886 1.13 2 'ti 3 64 1 ; 38 8
