3830
l'ol. 77
HENRYG . KUIVILA A N D ROBERT A. ~YII,ES [CONTRIBUTION FROM THE
DEPARTMENT O F CHEMISTRY,
UNIVERSITY O F N E W HALIPSHIRE]
Electrophilic Displacement Reactions. VII. Catalysis by Chelating Agents in the Reaction between Hydrogen Peroxide and Benzeneboronic A ~ i d ' , ~ , ~ Rs HENRYG. KUIVILAAND ROBERT A.WILES RECEIVED MARCH28, 1955 Xineteen chelating agents have been tested for catalytic activity on the reaction between benzeneboronic acid and hydrogen peroxide. Free oxalic and malonic acids have no effect. Malonate and/or acid malonate retard the reaction. Of five vicinal diols and two 1,3-diols only one, pinacol, has :I catalytic effect, Among ten hydroxy acids the order of catalytic activity is : benzilic > I-hrdroxycyclohesanecarboxylic > a-hydroxyisobutyric > salicylic > tartaric > mandelic > lactic > glycolic > hydracrylic. A detailed study with a-hydroxyisobutyric acid indicates that i t functions as a catalyst in a PHindependent reaction. It also ties up boronic acid in an unreactive complex anion; the behavior of citric acid is similar.
Benzeneboronic acid reacts with hydrogen peroxide to yield phenol and boric acid. ;Imong the interesting characteristics of this reaction' and of the reactions of areneboronic acids with bromine6 and iodine6 is the fact that they are catalyzed by certain chelating agents. The present paper describes a survey of the effects of nineteen chelating agents on the peroxide-benzeneboronic acid reaction and a detailed examination of the nature of the catalysis by a-hydroxyisobutyric acid. General Survey The preceding study3on the kinetics of the reaction was carried out in solutions a t ionic strength 0.2. I n order to provide more latitude in varying concentrations of ionic species the present work was carried out at ionic strength 0.5. It was therefore necessary to prepare new pH L I S . log k curves for the reactions first and second order in boronic acid. The results are summarized in equations 1 and 2. TTalues of k , indicate a positive salt log k i = -6.87 + pH ( p = 0.5) (1) log k ! = -580
+ fiH
(p =
0.5)
(21
TABLE I SUMMARY OF THE EFFECTS OF MALOSIC ASD OXALICACIDS h-0.
Concn., of moles/l. runs
Acid
Sone
Jlalonic
0 05 10 30 50 Ovalic 05 Initial concn.
(1) For t h e preceding paper in this series see H . G. Kuivlla, 77, 4014 (lg55).
THIS
JOCRXAL,
(2) From t h e 1.1.S. Thesis of R. A . Wiles, August, 1954. (3) Supported in p a r t b y t h e Office of S a v a l Research. (4) H. G. Kuivila, THISJ O U R X A L , 76. 870 11954). (5) H.G.Kuivila and E. J. Soboczenski. ibid.,76, 2675 !1954). (6) H. G. Kuivila and R. ?if. Williams, ibid.. 76, 2679 (1954).
103k, C ~ H ~ B ( O H ) I 1. , mole-' mole~/l.~ sec.-l
1.45 1.00 2.00 1.00 3.21 2.00 4 1 .95-2.04 1.Oc-5.00* 1 3.24 2.00 1 3.19 2.00 1 3.21 2.00 1 1.42-1.56 1.00-5.00" b Range. Mean value.
0.87 0.83 1.OO 0.92"
,98 .82 .65 88"
col and cis-indane-1,%diol have previously been shown t'o form water-stable esters with benzeneboronic acid.7 I t was observed in this work that 2,2TABLE I1 EFFECTS O F DIOLS (Diols, 0.05 X ; acetate buffers, PH 4.72.) SUMMARY OF CATALYTIC
lo?Diol
Sone
effect; the data for k? and k 3 (pH-independent reaction) are essentially the same as those a t p 0.20. The chief source of error was the pH measurement which a t best gave values precise to 0.03 unit, cor- Ethylene glycol responding to $yoin the value of the specific rate constants. Dicarboxylic Acids.-Malonic and oxalic acids were chosen as the simplest dicarboxylic acids Pinacol which might have some effect due t o their chelating abilities. The data obtained with these acids are presented in Table I. It is evident t h a t the free Catechol acids have no significant effect on the rate of the reaction. In solutions containing 55% malonate and 45% acid malonate (pH 3.2) a pronounced de- cis-Cyclohexane-1,2crease in rate occurs as the concentration of buffer diolb is increased. In half-molar malonate the rate is decreased by 35%. The simplest interpretation is cis-Indane-1,2-diol that the malonate or acid malonate and boronic acid react to form a complex which reacts very slowly if a t all with hydrogen peroxide. Propane-l,3-diol Diols.-The specific rate constants in the presence of five vicinal diols and two 1,3-diols are summarized in Table 11. Of these diols catechol, pina-
*
p H range
[CaHaBa(OH)2]
0.40 1.00 3.00 5.00 1.00 3.00 5,OO
103kobsd., 1. mole-] sec. - 1
8.39 8.98 9.96 10.94 8.7 11.0 13.1
0.400 10.6 ,700 11.3 1.00 11.4 1.00 9.2 3.00 11.1 5.00 13.3 1.00 8,5 3.00 10.8 5.00 12.8 0.04 9.0 .07 7.9 .I0 8.4 1.00 8.9 3.00 10.9 5.00 12.8 2,2-Dimethyl0.40 8.6 propane-1,3-diolc .70 9.1 1.00 9.2 a Initial concn., moles/l. b 0.001 111.
lo%, 1. mole-' set.-'
lo%, 1. mole-2 sec. --2
8.00
98
7.6
110
10.2
130
8.1
104
7.4
118
8.0
98
8.3 0.01 dl.
100
(7) H. G. Kuivila, A. H. Keough and E. J . Soboczenski, J . Org. Chetn., 19,780 (1954).
TABLE I11 SUMMARY OF EFFECTS OF HYDROXY ACIDS R1-CCOOH, 0.050 114 /I Rz OH No. of R'
4881
CATALYTIC ACTIVITYOF CHELATING AGENTS
Sept. 20, 1955
Rz
runs
9H range
4 1
1.96-2.07 1.86 1.71 2.42-2.53 2.38-2.45 2.31-2.45 2.23 2.40 2.26-2.46 2.22 2.07 2.21 2.10 (2.63) (2.58) (2.62) (2.73) (2.73) (2.73) 2.08 2.03 1.96 2.01 2.50-2.60 2.71-2.88
[CsHsB(OH)zl"
lOakot>sd., 10Zkog 1. mole-' sec. -1 1. mole-' rec.
mean
mean
-'
167 1.00-5.00 84 149 150 1.00 131 1.oo 1 263 2.5 2.07 1 .00-5.00 4 158 8.7 4 1.00-5.00 74.4 1.00-5.00 38.0 4 71.5 2.00 72.3 1 80.0 160 2.00 1 120 1.00-5.00 4 61 35 1 1.00 18.3 35 1 3.00 18.3 38 19.7 1 4.00 38 1 5.00 19.9 51 .4h 25.8 1 1.00 54.4h 1 3.00 27.3 59 .gh 30.0 4.00 1 1350h 1 0.400 135 1610h 1 .700 161 1560h 1 1.00 156 46f CHOHCOOH H 1 23.4 1 .oo 57' 1 3.00 27.7 61' 1 4.00 31.4 63' 1 32.4 5.00 59 Salicylic acid 3 1.00-5.00 3.80 0.2 1.00-5.00 3-Hydroxypropanoic acid 4 0.95 " lo2 X initial concn., moles/l. Hydroxy acid, 0.10 M. Hydroxy acid, 0.20 M. Solvent, 50% methanol by volume. Hydroxy acid, 0.10 JI. kl, = 40 1. mole-' set.-'; kZc = 460 1. sec.-2. k, = ( k o b a d - 8.0 X 10-4)/[cat.]. kunoataiyeed = (1.40 =?= 0.02) X 10-4 1. mole-' sec.-l in the pH range 2.44-2.76 (measured by Mr. A41bert G. Armour). Hydroxy acid, 0.0050 M . CHzCOOH
CHzCOOH
dimethylpropane-1,3-diolreacts in a similar manner. Because of the relatively low solubility of the esters, i t was necessary to use low concentrations of diol or boronic acid or both to avoid precipitation. cis-Indane-1,2-diol had to be used in especially low concentration. Pinacol is the only one of the diols which shows unmistakable catalysis of both of the reactions. The effects of the others, with ethylene glycol and cis-cyclohexane-1,2-diolas possible exceptions, are negligible. It is thus apparent that formation of a chelate with the boronic acid is not of itself a guarantee of catalysis whether the boron is present in a five- or six-membered ring. The indanediol is a catalyst for the iodinolysis of p methoxybenzeneboronic acid6 but not for the present reactions. This may be due to the difference in mechanism in the two reactions. Hydroxy Acids.-Eight a-hydroxy acids and two P-hydroxy acids were examined in unbuffered solutions and all but one proved to be catalysts. Results are summarized in Table 111. Runs were usually carried out with four different concentrations of boronic acid. When the rate increased with boronic acid concentration data for individual runs are given; otherwise the mean value of the rate constant for the set of runs is given. The order of effectiveness of the free hydroxy acids in catalyzing the reaction first order in boronic acid is benzilic > citric > 1-hydroxycyclohexanecarboxylic>a-
hydroxyisobutyric > salicylic > tartaric > mandelic >lactic > glycolic > hydracrylic. Thus, replacement of the a-hydrogens of glycolic acid by methyl, pentamethylene, phenyl, carboxymethyl or hydroxycarboxymethyl results in greater catalytic activity. Mandelic acid was studied in water and in SOY0 aqueous methanol. This provided a link for examining benzilic acid which is insoluble in water and only slightly soluble in the aqueous methanol. Each acid catalyzes reactions first and second order in boronic acid, benzilic acid being far more effective in each case. The reactions proceed faster in the 50% methanol than in water. Tartaric acid is a slightly better catalyst than mandelic acid for both reactions. The catalytic rate constant for citric acid decreases as its concentration increases. This implies that a significant fraction of the boronic acid forms a complex a t the lowest citric acid concentration, 0.05 M . Then as this is increased to 0.10 ill and further to 0.20 M , the concentration of the complex increases, but not linearly as would be necessary for constancy in kc. An alternative explanation (vide post) is that an unreactive complex is formed decreasing the actual concentration of free boronic acid with increasing citric acid concentration. Consequently the rate constants are low because no correction for this has been made in their computation.
HENRYG. KUIVILAAND ROBERTA. WILES
4832
Citric and a-Hydroxyisobutyric Acids, Effect of pH Citric Acid.--A number of runs, in addition to those reported in Table 111, involving variation in pH were made with citric acid. The data thus obtained are presented in Table IV. A striking pattern is observed. At low pH catalysis of the reaction first order in boronic acid is very important. EFFECTO F pH #H
TABLE IV CITRICACID" CATALYSIS
ON
102 kohsd. [CsHsB. 1. mole-' sec. - 1 (OHhIt
10'h Obsd. Calcd.6
103k~ Obsd. Ca1cd.c
27.5 28.1 28.2 28.0 28.3 0.96 .. 2.2 29.6 12.8 12.4 12.5 12.6 5.6 .. 63 12.6 21.0 20.3 21.7 24.5 27.3 188 436 16 2 38 30.6 49.0 66.0 72.0 118 1410 01 37 120 96.0 a Total citrate 0.05 M i n all runs. * Initial concn., moles,' At mean value of pH range. 1.
3 3 3 3 3 4 4 4 4 5 5 5 5 5 5 5 5 6 6
22 19 12 11 02 61 62 66 62 45 45 44 42 48 45 92 90 00
1.00 2.00 3.00 4.00 5.00 1.00 3.00 4 00 5.00 0.50 1.00 2.00 3.00 4.00 5 00 1.00 2.00 3.00 ..i 00
-1s the PH increases catalysis becomes less important. Above pH 5 the over-all reaction rate is lower than in the absence of citrate, i.e., the citric acid is retarding the reaction. Thus a t pH 6.0 with 0.01 M boronic acid the observed rate constant is 0.0490 1. mole-' sec.-l whereas in the absence of citrate the rate constant a t the same pH is 0.134 1. mole-' sec.-l. Citric acid does not catalyze the reaction second order in boronic acid a t arlv $H in
Vol. 77
the range observed but i t obviously retards this reaction decreasing its observed specific rate constant to less than 8.5% of its normal value a t pH 6.0. Samples of citric acid alone and in the presence of equimolar benzeneboronic acid were titrated with sodium hydroxide. Two distinct pH vs. titer curves were obtained. Below 1.2 equivalents of base the boronic acid decreases the pH; above this point it increases the PH. Inflections in the presence of the boronic acid occur a t 1.2,3.0 and 4.0 equivalents of base. The last inflection indicates that the complex which is formed breaks down completely a t high values of the PH. Because citric acid is tribasic i t was decided to use the simpler a-hydroxyisobutyric acid for more detailed examination of the a-hydroxy acid catalysis. a-Hydroxyisobutyric Acid. Complex Formation. -The catalytic behavior of a-hydroxyisobutyric acid is similar in all respects to that of citric acid as can be deduced from the data of Table V. Curves of pH vs. titer are depicted in Fig. I, for benzeneboronic acid, a-hydroxyisobutyric acid and a combination of the two using sodium hydroxide as titrant. As manifested by these plots a species more acidic than either of those introduced is formed in the mixed solution. Furthermore, the second inflection in this curve indicates that the species breaks down to yield benzeneboronic acid and, presumably, a-hydroxyisobutyrate above pH 6. These titration data can be treated in a manner similar to that applied by Vermaas8 to solutions of boric acid and glycols. The following equilibria are assumed in any solution prepared from benzenebor(CH8)zCCOOH
I
+ HzO
KHA
(CH3)zCCOO-
+ H30+
OH
OH KB
CsH5B(OH)z -I- 2Hz0
+ H30'
C&%(OH)3
OH (COH)
[3)2
20 30 40 A-aOH, ml. Fig. 1.-Titration curves of 50 ml. of solutions with 0.495 fl NaOH; each solution 0.50 M in r\'aAT03: I, 0.200 JI C6H5B(OH)z,p 0.50 throughout; 11, 0.187 M a-hydroxyisobutyric acid; p 0.51 at start and 0.50 a t end of titration; 111, 0.187 M a-hydroxyisobutyric acid, 0.200 M CcHsB(OH)$; p 0.51 at start and 0,49 a t end of titration.
4-SaOH
+ H20
K4
10
OH
II
CeH5-B-OH
+ (CH3)2CCOO- + X a + I
OH
OH (BOH)
(A-)
(8) N. Vermaas, Rec. Irov. rhim., 51, 67 (1932).
4833
CATALYTIC ACTIVITYOF CHELATING AGENTS
Sept. 20, 1955
TABLE V DATAON THE EVALUATION OF BH
100 B,b moles/l.
100 kobs.
THE
100 kcor.
CATALYTIC CONSTANT FOR a-HYDROXYISOBUTYRIC ACID" 100 kunoat.
100 HA,C
100 Ak
4.28 0.08 3.66 4.36 1.00 2.45 4.35 .08 3.92 4.43 2.00 2.40 3.85 .08 3.93 4.00 3.70 2.32 .os 4.13 3.42 4.21 5.00 2.31 8.66 .08 7.23 8.74 2.00 2.23 20.8 .08 20.9 2.00 16.0 2.10 .09 3.89 2.90 3.98 3.00 2.96 2.84 4.00 3.16 4.16 .09 4.07 .09 3.77 3.86 5.00 3.11 2.83 .12 3.05 1.93 3.21 3.57 1.00 .11 3.21 3.32 3.00 2.28 3.40 .ll 3.21 2.40 3.32 3.33 4.00 3.32 .I1 3.21 5.00 2.54 3.30 1.78 .19 1.OO 1.20 1.97 4.10 .20 1.87 3.00 1.40 2.08 3.99 .20 2.13 2.31 4.00 1.51 3.91 19 2.18 5.00 1.56 2.36 3.81 a Total hvdroxv acid 0.05 M exceDt where noted. Boronic acid. Total, 0.26 M . .
moles/l.
e
10aKa
kc
0.92 .95 .85 .92 .92 1.08 1.05 1.11 1.03 1.03 1.08 1.09 1.09 1.15 1.15 1.30 4.5 1.34 Total, 0.10 M . a-Hydroxyisobutyric acid. 4.65 4.60 4.52 4.51 9.41d 19. 3 i e 3.69 3.66 3.65 2.97 2.96 2.95 2.95 1.54 1.64 1.64 1.64
9
100 B, moles/l.
0.84 1.77 3.66 4.63 1.65 1.53 2.19 3.04 4.02 0.60 2.06 2.89 3.82 0.61 2.02 2.61 3.31
14.5 11.3 9.8 8.2 13.2 12.6 11.0 12.8 9.6 6.0 6.2 6.1 5.2 3.1 3.0 4.0
onic acid, a-hydroxyisobutyric acid and sodium hy- obtained with a-hydroxyisobutyric acid a t different values of p H and initial boronic acid concentradroxide. At any given point in the titration the following tions. It will be assumed that the catalytic reacare known quantities: K H Aand K B from previous tion involves free boronic acid and free hydroxy acid titrations, K,, initial concentrations of boronic acid or their kinetic equivalent, species C for example. and hydroxy acid, concentrations of sodium ions The reaction rate is therefore independent of $H. and hydronium ions. From these, nine equations Species COH- then is simply the product of a side in the following ten unknowns can be assembled?: reaction which ties up boronic acid and hydroxy K1, Ka, K3, K4, B, BOH-, HA, A-, C, COH-. acid, thereby decreasing their effective concentrations. In order to treat the data properly the Above pH 8 the concentrations of HA, C and H30' are small enough to be neglected when they are concentration of free boronic acid must be known added to or subtracted from other terms. With and can be computed. Fortunately the fraction of these approximations i t is possible to evaluate K4. free acid changed very little (less than 3%) in any From the titration curve a t 0.50 ionic strength KB of the runs because the initial concentration of perM. This fraction a t 25" is 1.9 X From data at one-half pH unit oxide was always about 3 X intervals between pH 8.1 and 10.4 the value of K4, was therefore calculated for the initial concentration of boronic acid and then used throughout the run (5 0 0.2) X IO3, is obtained. I t can be shown that K3 = K B K H A ' K ~ K in ~ . the ~ calculation of a corrected specific rate conUsing K B , 1.5 X K,, and the above stant, kcor,. Using kl and kz from equations 1 and value for K4 we compute K3, 6.0 X lop3. Alterna- 2 and the corrected boronic acid concentration the tively K , can be estimated from data a t low pH rate constant for the uncatalyzed reaction, kuncat. where BOH- and HO- are present in very low was estimated and the difference, k,,, - k,,,,t concentration and can be neglected. It is also as- Ak, measured the catalytic effect of the hydroxy sumed that the concentration of C is small. When acid. Dividing Ak by the concentration of free athis is done the values of K 8 drift downward from hydroxyisobutyric acid then gave the specific rate a t p H 2 to 4.9 X IOF3 a t p H 4. This is ronstants, k,, for the catalyzed reaction. 6.5 X seen to be good agreement when it is realized that an Pertinent data in the p H range 2.45-4.10 are error of 0.03 unit in the pH measurement leads to gathered in Table V. Values of K , were evaluated .t7% in value of K?. The value of K H Ais subject for each run. Lacking the internal consistency proto the same error and is used in evaluating K,. vided in a single titration, these vary somewhat .lbove pH 4 small errors in stoichiometric concen- more widely than those previously cited. Xevertrations of HA or NaOH loom very large in the theless these were used in estimating the concenvalues of K 3 . trations of the various species. Values of k, drift In principle it should be possible to evaluate K1 upward with pH but not enough to invalidate the and K Balso, but their values are extremely sensi- assumptions outlined above. Beyond PH 4.1 attive to variation in K4. The best that can be de- tempted calculations were fruitless because of the duced from attempts a t their estimation is that K , great sensitivity of the quantities to be calculated is very small and Kz is very large. to small errors in those measured. Mechanism of a-Hydroxyisobutric Acid CataIn view of the above discussion the transition lysis.-Now that the equilibria portrayed above state complex for the a-hydroxyisobutyric acidhave been shown to be consistent with the titration catalyzed reaction must be made up from a moledata we may proceed to examine the kinetic results cule each of boronic acid, peroxide and hydroxy
*
4834
Vol. 77
HENRYG. KUIVILA AND LAWRENCE E. BENJAMIN
acid. The chelating tendency of the catalyst must play an important role in making it effective. One possible intermediate is represented by structure I .
0
0 I
a
H-0
\
T'
I f.--(j
CsHs--R
I
II
A, >c-o0
--+
R-0
~, )c=o
'CR2 111
'CK? IV
boron--oxygen bond formation cannot occur simultaneously. This path is attractive in that the slow step involves the loss of opposite charges. This process should be favored by a decrease in polarity of the solvent and, being a rate process, the effect should overbalance the opposing effect on the equilibrium I e111. Changing from water to 50'; methanol does in fact increase the rate of the mrtndelic acid-catalyzed reaction. This same solvent change has the opposite effect on the rate of the uncatalyzed reaction (compare the first three data in Table I and that in Table IIT, footnote h) for which the
OH
\'I
Experimental
Transfer of H a to Oa aiid cleavage of the peroxide bond to give I1 would occur in the slow step. -411 alternate path from the same intermediate could involve the processes I I11 + IV with the last step rate determining. The conformational possibilities for I are such that the proton transfer and
CBH-R
+ HOOH & HOH + CsHsBOOH +HO- + CsHoBO' I
2
OH
+
1O-O
C6Hs-B
Ct,H,B(OH
OH.
a i l
rate determining process 1. -t VI appears to be reasonable
Materials.--Berizerieboronic acid was prepared froiii butyl borate and phenylmagnesiuni bromide by the niethotl of Bean ant1 Jol~nsoii.~The samples, used :is the anhydride, Iiad neutral equivalents of 103.9 f 0.5. 1-1-Iydro\.yc~clohesarircurljo\:~lic acid \ v i i s graciously prw vided b y Ilr. R . E. Lyle. The ol-~i!.dros!.isobut!-ric acid had ni.11. 78.0~-79.0". Anhydrous piiiacol \vas prepared from the crude hydrate.lL' 8-Hydroxypropanoic acid x i s prepared by warming :t solution of 0-propiolactone on the steam-bath. The resulting solution was assal-ed by titration xvith sodium hydroxide arid used as such. All other reagents were C.P. or reagent grade where available and used x-ithout further purification, or were prepafed bj; conventional methods and had melting points agreelng with those given in the literature. Rate Measurements.-These ivere inade ;LS described prcviously3 using the colorimetric method of Eisenberg" for determining the conceiitrations of uiireactecl peroxide. The optical density of the pertitanate \\':is diminished in the presence of oxalic acid so that L: separate calibration curve \vas prepared in this case. -411 kinetic runs nere made at 24.93 =k0.02" a t ionic strength 0.3.). Titrations.-All pH measurements were made with a Deckman Model H-2 meter and were usually reproducible to 0.0:3 ,bH unit. IVhen ionic strengths were adjusted with sodium perchlorate difficulties were encountered in the titrations after some time probably because of the formation of potassium perchlorate at the calomel electrode. When sodium nitrate was used in adjusting ionic strengths the difficulty disappeared. Sodium perchlorate was used in all the kinetic runs. ~
f9) I;. I
