Mechanism of the periodate oxidation of ethane-1,2-diamine,NNprime;-dimethylethane-1,2-diamine, and 2-aminoethanol

摘要

1980 39 Mechanism of the Periodate Oxidation of Ethane-I ,2-diamine, NN'-Di- methylethane-I ,2-diamine, and 2-Aminoethanol By LAszl6 Maros," lbolya Moln5r-Per1, Enik6 Schissel, and Vilmos Szerdahelyi, Institute of Inorganic and Analytical Chemistry, L. Eotvos University, 1088 Budapest, Hungary The kinetics of the oxidation by periodate of ethane-l,2-diamine, NN'-dimethylethane-l,2-diamine, and 2-amino- ethanol have been studied at pH 4-1 2 and 293.2 K. The reaction rates are highest at pH 9-1 1. It is assumed that periodate monoanion reacts with the unprotonated amines. The reactions of 1,2-diamines with periodate are catalysed by acids and bases. The oxidation of 2-aminoethanol is catalysed by bases, and at high concentration of catalysts the rate seems to approach a limiting value.The pH dependence of the rate of oxidation of these compounds can be explained in terms of multi-stage reactions similar to those in the oxidation of 1,2-diols by periodate. THE periodate oxidation of 1,2-diols involves the form- ation of a cyclic periodate ester of the diol which decom- poses to pr0ducts.l The periodate ester is formed by the electrophilic attack of periodate,2 and the products are formed by the rupture of the cyclic ester mono-anion, probably in its dehydrated form.3 Simple 1,2- diols, c.g. ethane-1,2-diol, form a cyclic ester (C,) with periodate (Per.) rapidly and reversibly, which slowly decomposes to products 374 reaction (l),and the kinetic form of the reaction is given by equation (2),where the kPer.+ diol h' C, -products (1)Slow k' = kKdiol/(l + Kiliol) diol is in large excess over periodate, and k' is the observed pseudo-first-order rate constant with respect to periodate. Highly substituted 1,2-diols, e.g. pinacol, have second-order kinetics in the periodate oxidation over a wide range of reagent concentration. This may be due to the fact that Kdiol 1, and the second- order rate constant k" equals kK. The kinetics of the oxidation of pinacol indicate the occurrence of general acid-base cataly~is.~.~It has been shown by Buist, Bunton and his co-workers that the catalysed step is the formation of the cyclic intermediate, and they suggested a mechanism which is probably common to all oxidations of 1,2-diols by periodate, vix.the formation of a periodate monoester (Cl), followed by ring-closure to a cyclic ester, then decomposition of the monoanion of the cyclic ester to products reaction (3). k kz kPer. + diol 4C, ===C, +products (3)k-t k-a We will show that several features of the kinetics of periodate oxidation of 1,2-diamines and 2-aminoethanol resemble closely those observed for the oxidat ion of pinacol, and a plausible mechanism can be given by an acylation-type reaction of periodate with the amines forming steywise a cyclic intermediate which decomposes to products. Dahlgren and his co-workers have studied the kinetics of periodate oxidation of 2-aminoethanol and 2-(methylamino)ethanol over the pH range 3.6-6.4 and temperature range 273--323 K, and found second-order kinetics for the reactions.The pH-dependence of the rate is in agreement with the assumption that either the periodate monoanion reacts with the unprotonated 2-aminoethanol or the periodate dianion reacts with the protonated compound.76 The assumption that the dehydrated periodate monoanion, 104--,is the reactive species leads to a negative enthalpy of activation.'" Second-order kinetics and a catalytic effect of buffers were observed in the oxidation of cyclic and substituted a-amino-alcohols by peri~date.~?~ The oxidation of ethane-1,2-diamine by periodate yields formaldehyde and ammonia.lOill" The anomalous decrease of the reaction rate (if periodate is in excess) is due to the reaction of formaldehyde, or the imine-like primary oxidation product, with ethane-1,2-diamine. The oxidation of ethane-1,2-diamine and NN'-dimethyl- ethane-1,2-diamine becomes fast and stoicheiometric in the presence of a many-fold excess of propane-1,3-dianiine.llb EXPERIMENTAL 1VateriaZs.--Ethane-l,2-diamine, NN'-dimethylethane-1,2-dianiine and 2-aminoethanol (Fluka) were vacuum-distilled and tested for purity.Sodium periodate was recrystallized. All other chemicals were of analytical grade. Kinetics.-Reaction rates were measured in solution by following the decrease of periodate concentration with time. Titrimetric and spectrophotometric methods were used. For titrinietric measurements (in the pH-range 4-7) aliquot portions of the solutions were added to acidic potassium iodide solution and titrated with 0.02~sodium thiosulphate. A Unicam SP 700 spectrophotometer with water-jacketted cell holders, and 0.1-, 1-, and 4-cm silica cells was used.All reactions were carried out at 293.2 K, and solutions were previously therinostatted overnight. In the kinetic experiments, the comparison cell was filled with amine solution containing all components except periodate. The amine and periodate solutions were mixed in a volu- metric flask and a portion was added into the reaction cell. In most cases recording began after 20-25 s. The lowest ratio of diamine : periodate was 20 : 1 at the beginning of the kinetic runs. For 2-aminoethanol this ratio was 10 : 1.The lowest initial periodate concentration was 1 x 10-5~for the spectrophotometric measurements. Perchloric acid and sodium hydroxide were used to make the solutions acidic or alkaline, respectively. The ionic 40 strength was maintained constant by sodium perchlorate. The pH of the solution was measured during the runs and was constant within f0.02 pH unit. 2-4 parallel runs were carried out. Errors in Kobs. for titrimetric measurement were f5, but for the fast reactions followed spectrophotonletrically errors were f20. RESULTS AND DISCUSSION The oxidation of 2-aminoethanol by periodate follows second-order kinetics, first-order in each reactant. Self-catalysis by the 1,2-diamine occurs in the oxidation of ethane-1,2-diamine and NN'-dirnethylethane-lJ2-di-amine, the reactions being first order in periodate.The reactions are catalysed by buffers. The observed second-order rate constants, kohs.,were extrapolated to zero buffer concentrations, and, in the self-catalysed reactions of the 1,2-diamines, to zero 1,2- diamine concentrations (ki). In equations (4) and (5), Rate = kobs.Per. ~Amine (4)T hobs. = k," + kiAmine~+ kiBufferT (5) Per.jT, Amine~, and BufferIT are the corresponding total concentrations. The regular changes with pH and the concentration of 1,2-diamines are in good agreement with the assumption that the periodate monoanion reacts with the unproton- ated 1,2-diamines in general acid-base catalysed reactions. The pseudo-second-order rate constants of the re-actions of periodate monoanion and unprotonated 1,2- diamines, klo4, were Calculated according to equation (6), where I?, and I?3 are the second and third acid dissoci- ation constants of periodic acid, Ka1 and K,2 the first and second acid dissociation constants of the diprotonated l,2-diamines,f and the f values are the activity co-efficients of the ions calculated according to the Davies equation.15 For the oxidations of lJ2-diamines the rate constants kg show a broad maximum in the pH range 9-11, and the values of the pseudo-second-order rate constant extrapolated to zero buffer and 1,2-diamine concen-trations, k,",g, have a minimum in the pH range 8-9 (Tables 1 and 2).We assume that the reactions of periodate monoanion with unprotonated 1,2-diamines are catalysed by water, hydronium, and hydroxide ions equation (7).kyod = k3OH2O + Ko'H30++ kp,-OH- (7) Table 3 shows the pH-dependence of the second-order t The dissociation constants for periodic acid l2 at 293.2 K are: 5.43 x 10-9 mol dm-3 and I?, 6.3 x mol dm-3; for the diprotonated ethane-l,2-diamine l3 Kal 1.0 x lo-' mol dm-3 and Ka28.2 x lo-" mol dm-3; and for the diprotonated NN'-dimethylethane-l,2-diaminel4 Kal 5.97 x lo-* mol dmP3 and Ka2 6.16 x lo-" mol dm-3. J.C.S. Perkin I1 rate constants kg and k,",~calculated for the periodate oxidation of 2-arninoethanol.* Acetate, phosphate, ammonia, and borate buffers have been used to maintain the pH constant and to examine TABLE1 Second-order rate constants extrapolated to zero buffer concentrations for the periodate oxidation of ethane- 1,2-diamine at I 0.300 mol dm-3.(Units for all rate constants are dm3 mol-1 s-1) 6.07 6.20 7.28 8.10igki 1.03 1.20 10.7 22 10-2kio$ 33.5 21.8 2.8 1.0 8.45 8.60 8.82 9.22igki 31 32 34 67 10-2kamp;~ 1.0 1.0 1.0 1.8 9.84 10.31 10.80 11.35igki 82 82 90 35 1 0-2k;'o) 2.7 4.4 11.5 18.7 11.90 12.46 12.82!$ki 20 7.0 3.0 1 0-2ki* 72 255 540 the catalytic effect of the buffers. A catalytic effect was not found for borate. In the oxidations of 1,2-diamines the reaction rate increases linearly with increase in buffer concentrat ion. The self-catalysis of lJ2-diamines was expressed in TABLE2 Second-order rate constants extrapolated to zero buffer concentrations for the periodate oxidation of NN'-dimethylethane-l,2-diamine.(Units for all rate coii-stants are dm3mol-l s-l) pH (I0.300 mol 4.52 4.72 4.99 5.07 5.42 dm-3)103~; 1.00 1.85 3.80 4.20 8.60 10-3~;* 1010 735 435 355 105 pH (I0.300 mol 6.13 6.55 d m-3) 103~; 44.0 155 lO-3ky04 28 15.4 pH (I0.100 mol 8.08 8.86 9.32 9.80 10.00 dm-3) k; 4.4 8.7 8.7 10.0 10.0 10-3k;~ 2.8 3.0 2.8 3.7 3.9 pH (I0.010 mol 10.26 10.80 11.08 11.66 11.93 dm-3) 15.6 10.2 10.3 9.7 9.5 10-3kyO4 4.7 6.9 12.0 51 140 PH 12.20 12.83 13.06 I/mol dm-3 0.020 0.060 0.100 ki 6.0 1.o 0.45 10-3k;~ 230 890 1 520 terms of the observed overall third-order catalytic constant k: in equation (5).Furthermore we have calculated the pseudo-t hird-order catalytic const ant kf$, for the oxidation of unprotonated 1,2-diamine by periodate monoanion catalysed by the monoprotonated 1,2-diamineJ AH+, according to equation (8). The variation of the rate constants ki; and k;$' * The acid dissociation constant of protonated 2-amino-ethanol is K, 2.26 x 10-lomol dm-3 at 293.2 K. with pH for the oxidation of ethane-1,2-diamine is shown in Table 4. At pH 7 no measurable self-catalytic effect for ethane-l,2-diamine was found, and the pH-dependence TABLE3 Second-order rate constants extrapolated to zero buffer concentration for the oxidation of 2-aminoethanol by periodate at I 0.300 mol dm-3. (Units for all rate constants are dm3 mol-l s-l) PH 4.92 5.23 5.65 6.27 6.93 7.65 k," 0.066 0.098 0.23 0.60 3.5 13.7 10-3k;04 3.9 3.3 2.8 2.0 2.7 3.0 8.15 8.53 9.38 10.34 10.53 10.84;? 33 48 51 27 15 8.3 10-3k;oll 4.2 5.0 6.0 11.0 9.0 10.2 PH 11.30 11.58 11.90 11.93 11.95 k: 3.4 1.6 0.59 0.64 0.52 10-3k;,~ 15 18 21 21 22 of kf$' suggests the monoprotonated and unproton- ated ethane-1,2-diamine to be the effective catalysts.Using equation (9), graphical evaluation yielded the values of kff' and kfi. (9) The pH-dependence of the pseudo-third-order cataly- tic constant for ammonium, kFa+, is shown in Table TABLE 4 Third-order self-catalytic constants for the oxidation of ethane-1,S-diamine by periodate at I 0.300 niol d1-11-~ (Units for all rate constants are din6 niol-2 s-') PH 7.28 8.10 8.45 8.60 8.82 I AH+ 2.1 0.2 0.09AH+ 4 68 70 81 111 tamp;4kAHt104 3 3.0 2.5 3.0 3.3 PH 9.48 9.84 10.34 10.80 A0/ AH+ 0.177 0.415 1.22 3.79 k" 111 113 62 23 1amp;4kAH+ 3.7 5.3 7.3 13.8104 6 for the oxidation of ethane-lJ2-diamine.The values of kyp and kyp' were evaluated graphically. For the oxidation of NN'-dimethylethane-l,2-diamine the values of the catalytic constants k$, kf?', kyp, and kTF4' were evaluated from the data in Table 6. Self-catalysis was not detectable for the oxidation of 2-aminoethanol, probably owing to the very low con- centrations needed for the rate measurements. Plots TABLE5 Pseudo-third-order catalytic constants (in dm6 mol-2 s-1) for ammonium in the oxidation of ethane-l,2-diamine by periodate; I 0.300 mol dm-3 PH 8.64 9.02 9.30 9.45 9.65 10.00 "Hi/ 12.9 3.4 1.8 1.26 0.79 0.35 "H3I 10-4kyst 0.75 1.15 1.43 2.0 3.4 8.8 of kiob against acetate concentration are linear, but plots against liydrogenphosphate (Figure I), ammonia (Figure 2), and hydroxide concentration show marked curvatures.The slopes of the linear regions of the curves are independent of buffer ratio; i.e., only the 15i +2 m e I I 1 5 10 1O2HP0~-lmoldm'3 1FIGURE Hydrogenphosphatecatalysis in the oxidation of 2-aminoethanol by.perioclate: (A) pH 8.15, (B) pH 7.65, (C) pH 6.93. The zeros of successive curves are displaced by one unit on the concentration axis bases are effective catalysts in the oxidation of 2-aminoethanol. However, the small increase in kamp; with decrease in pH in the yH range 4.92-6.27 (see Table 3) may be due to catalysis by hydronium ion.The catalytic constants for the oxidation of 2-amino-ethanol were evaluated according to the treatment used by Ruist, Bunton, and their co-workers for 1,2-diol oxidat ions. 30 0Yo 20 + ms -0 10. d--7--'-7----1-I I 12345 10NH3l/mol drr~'~ FIGURE Ammonia catalysis in the oxidation of 2-aminoethanol2 by periodate: (A) pH 10.34, (R) pH 9.38, (C) pH 8.53. The zeros of successive curves are displaced by one unit on the concentration axis The catalytic constants for the reactions are sum-marized in Table 7. The exponents of the Brijnsted equation, and a, were calculated with the statistical corrections p and q used by Bell and Evans.17 The @-values calculated for the OH-and H,O catalysis are : for ethane- ,2-diarnine oxidat ion, 0.33; for Nh7'-dimet hyle t hane- 1,2-diamine oxidation, 0.31 ; and for 2-aminoethanol oxidation, 0.32.TABLE6 Pseudo-third-order catalytic constants (in dma mol-2 s-l) for the oxidation of NN'-dimethylethane- 1,2-diamine by periodate (I0.100 mol dm-3) 8.08-9.32 9.80 10.80 * FfSk AH'14 652 9 42 9.00 9.45 10.00 ~~SkNHlflU* 1.o 4.2 12.5 * I = 0.010 mol dm-3. The a-values calculated for the H,O+ and H20 catalysis are: for ethane-1,2-diamine oxidation, 0.51 ; and for NN'-dimethylethane-l,diamine oxidation, 0.49. The acid-base catalysis can be explained in terms of a reaction in which the periodate and dianiine react reversibly to form a monoamide, C,, which cyclises to a diamide, C,, which in turn decomposes rapidly to yro- ducts, according to reaction (3).The general acid-base catalysis can be attributed to the catalysis of the cyclization of the monoanion, Ci, TABLE7 Catalytic constants for the periodate oxidations of 1,2-diamines and 2-aminoetlianol. (Units for all constants are dms mol-2 s-l) "'-Dimethyl-Ethane-1,2-ethane-1,2-2-Amino-diamine dianiine ethanol kyio 1.8 50 108a kOH-1.0 x 106 1.5 x 107 1.6 x 107" 10kO+ 2.7 x lo9 2.7 x 1OIo 1.0 x lo8 kr;6HJ 2.8 x 104 4 x 105 2.6 x lo5" kNIT4+ 6 x lo3 5 x 104 k$ 2.6 x 104 1.2 x 106 k#+ 3.0 x 104 6 x lo5 kAc0-1.5 x 10410 kAcOR 5 x 106 k#'O.'-1.4 x loBa a Catalytic constants were calculated from equations (13)- (15)9 formed rapidly and reversibly by the electropliilic attack of periodate monoanion on the unprotonated amino-group.Cyclization is rate-determining in the oxidation of 1,2-diamines over the whole range of concentration of reactants and catalysts used in the kinetic study. Bufer Catalysis in the Oxidation of 2-Aminoethano1.-In the oxidation of 2-aminoethanol, when strong bases are used as catalysts, cyclization is accelerated so much that another step of the reaction becomes partly or wholly rate-limiting. We have only observed second- order kinetics for the oxidations, and therefore we assume that the concentrations of the intermediates C, and C, are low, and for reactions (10)-(12) the steady-state approximation is applicable (AE = 2-aminoethanol ; A and B = conjugated acid-base pair.) J.C.S.Perkin I1 This leads to equation (13a) for the pseudo-second- order rate constant klo$, where K, = kl/k-,, Kf= kF/kaz, and Kf = k!/k". k k-i Per.-+ AE A Ci (10) kBCi + B amp;Ci-+ A;kh Ci- ks- products (11) Ci-+ A kA C; k-+ B; Cg +products (12) kS w104 = 1k/, + 1/K,H BI+ 1/~(KlK2-Bl/") + K,K!k$k-B/(k?3B+ k-)} (134 If k!,B amp; k-, equation (13a) can be simplified into equation (13b) for k,,$, and equation (14) is obtained for the limiting pseudo-second-order rate constant, kloJ(lim.)-1/k1amp; = 1/K1k?PI + 1/klo$(lim.) (13W 1/k~o~(~im.)= l/k, + 1/{(K1K2"k2-B1/A1) +fvGwk-) (14) The following treatment was used to evaluate the catalytic constants for bases in the oxidation of 2-aminoethanol : at constant pH, when base-catalysed cyclization predominates, the catalytic effect of water and hydroxide ion cannot be neglected at low concen- trations of B; the kinetic results can be described by equation (15), where kyo, is the pseudo-second-order rate constant at zero buffer concentrations.1/amp;$ = 1/(K,@BI + klO~)+ l/klO$(lim.) (15) Results were obtained by successive approximations : at high base concentrations equation (13b) was used to obtain values for kl@h(,ima)by extrapolating a plot of l/klo$ against 1/B to l/B = 0. The results at low buffer concentrations were treated according to equation (l6), a transformation of equation (15), to give the first approximation to values of K,",$ and Klkf.The second approximation to kloS(lim.) was found by plotting l/klo$ against l/(KIKfB + k;o$). For hydrogenphosphate catalysis (Figure 1) the broken lines are fitted by equation (15) with Klk2HP0:-1.4 x lo6 dm6 moF2 s-l, and at pH 6.93, 7.65, and 8.15 with kiM 2.7 x lo3, 3.0 x lo3, and 4.2 x lo3 dm3 mol-l s-l, and with klo$(lim.) 9.6 x lo3, 1.5 x lo4, and 1.9 x lo4 dm3 mol-l s-l, respectively. For ammonia catalysis (Figure 2) the broken lines are fitted by equation (15) with K1kfHs2.6 x 105 dms moF2 s-l and at pH 9.38 and 10.34 with kamp; 6.0 x lo3 and 1.1 x lo4 dm3 mol-l s-l and with klW(lim.1 3.4 x lo4 and 4.3 x lo4dm3 mol-l s-l, respectively. At pH 8.53 k,",* is 5.1 x lo3 dm3 mol-l s-l, and kloamp;lim. is ca.1.4 x lo4dm3 mol-l s-l. The change of l/klwclima) with the buffer ratio A/B according to equation (14) can be fitted by successive approximation. This yields, for hydrogenphosphate 1980 2catalysis, k2-K1KHPoaa-5.0 x lo3 dm3 mol-l s-l,k-KIK~Po~a-KH~Po~-1.0 x 104 dm3 mol-1 s-l, and k, 2.2 x lo4 dm3 molt1 s-l; and for ammonia catalysis, k2-K1KgH32 x lo5 dm3 mol-1 s-l, k-K1KfH3KFH4+ca. 1 x lo3 dm3 mob1 s-l and k, 4.3 x lo4dm3 mol-l s-l. From these data we obtain, for k2-/k-KfZPo4-0.5, and for k2-/k-KfHd+ca. 200. The equilibrium constant Kk of equation (12) can be expressed as Kf = K,/K:;, where K:; is the second acid dissociation constant of the cyclic intermediate C,, and K, the acid dissociation constant of the conjugate acid of the base catalyst.ls Therefore,ls a value euro;or K:;k2-/k-which can be calculated, from the hydrogen- phosphate catalysis data, is 3.0 x mol dm-3 and from the ammonia catalysis data is ca.2 x mol dm-3.* Hydroxide, Water, and Hydroniuna Catalysis in the Oxidation of 2-Aminoethanol.-To evaluate the catalytic constants for hydroxide and water we used equation (17) for kyo$ (see Table 3) which was fitted by successive approximations where the terms KOH-and K' are defined 1/kTo$ = l/(KoH-COH-I + K'I + 1/kPOamp;im.) (17) as shown below. In the pH range 10.5-12 we obtain for kamp;(lim.) 2.5 x lo4 dm3 mol-l s-,; for K' 6 x lo3 dm3 mol-1 s-l; and for KOH-1.6 x 107 dmG rnol-, s-l.In comparing these constants with the constants for hydro- genphosphate catalysis we assume kio$(lim.)= k,, KOH-= K,k;"-, and K' = K,k~oH20;i.e., equ-ation (17) becomes equation (17a). l/kamp; = l/(K,kfH-OH- + KlkpoH,O) + 1/k, (17a) In the pH range 7-9.5 the best fit of equation (17) was found with kamp;,~(~~~.)6 x lo3 dm3 mol-l s-l, with K' 2.3 x lo3 dm3 mol-l s-l, and with KOH-9 x lo9 dm6 mol-2 s-l. In this pH range hydroxide catalysis is negligible and we assume that water-catalysed cycliz- ation and hydronium ion-catalysed ring-opening is the partly rate-determining step. In this case the l/k, term of equation (13a) can be neglected, the ' limiting ' rate constant k:o$(lim.) is K,kFaoH,O and the transform- ation of KOH-yields for k2-KlKpo a value of 2.0 x loW6dm3 mol-l s-l, i.e., equation (17) becomes equation (17b), where 1/H30+ = OH-/K,, K, being the ionic product of water.l/kamp;4 = 1/{(k2-K1K~soH20/H~O+)+ I'} + l/Klk~~oH,O (17b) From the values of K,kF and k2-K,KF for the hydrogen phosphate, ammonia, and water catalysis we calculate for the acid-catalysed ring-opening k?F4+/k2-= 1.3, kHiP04-/k2-= 2.7 + lo2,and k?i0+/k2-= 5.5 x lo7 dm3 mol-l, respectively. The Bronsted plot for ring-opening is necessarily related to the corresponding * The error in this value can be large. This is largely due to the error in k-K,KNH3KY"4+, which depends strongly upon the klOd(lim.)value calculated from the values measured at pH 8.53.We have calculated the value of klo~(li,,)from 4 measured values (shown by C in Figure 2) and we use this only as an estimated valne. plot for cyclization. The slope, a', equals 1 -p, i.e. 0.68. The term K' in equation (17b) is equal to k-K,KpO-KF0+ if k?;oH,O k-see equation (13a), and leads to K:;k2-/k- 1.4 x 10-8 mol dm-3, which is in good agreement with the values calculated for the hydrogen- phosphate catalysis. K' = K2K~20k~0H20 if k?io-H,O k-, but this assumption need not be necessary (see later). In the pH range 6.0-4.9 kf0+ increases linearly with increasing hydronium ion concentration (see Table 4). We assume for a hydronium ion-catalysed reaction the sequence in reactions (18).K1Per.-+ AE Ci f kF20 Ci + H30+ kH$ "+ H20 bsol;kz:?+-C; + H30+; Cz k;products (18) We assume that the constant for hydronium ion cataly- sis, kFGo+ = Klkf3O'. In order to estimate the order of magnitude of the rate constants, we assume all second-order rate constants for protonation by H30+ to be at the diffusion-controlled limit,ls vix. ca. 5 x 1O1O dm3 mol-l s-l. This can be done for kF30f (the rate constant for protonation of Ci-to C; by H30+)and for kF30f (the rate constant for protonation of Ci to C: by H,O+) because they are ' normal ' acid-base reactions : simple proton transfers.18 This leads (from K,kFyO' = 1.0 x lo8 dm6 moP2 s-l) to a value of 2 x dm3 mol-l for K,.In the water-catalysed reaction K' cannot be equal to KlK~20k~o+H,0(k!!;oH20 k-) because this leads to (k~30'/k!30')= 0.4, and kHio' (the rate constant for ring-opening of Ci-to Ci by H30+) cannot exceed kF30+,i.e., the diffusion controlled limit. The assumption K' = k-K1K~20K~30+seems to be valid, and we obtain k?;'H,O/k-= 2.5, as a lower limit, if kyio' = k!!go'. The latter assumption seems reasonable, and we will use this to provide estimates for other constants. The third-order catalytic constant for hydroxide ion, KlkgH-, is 1.6 x lo7 dm6 molk2 s-l, and the estimated value of k;=-becomes 8 x lo9 dm3 mol-l s-l, which is near the diffusion-controlled limit reported for simple proton-transfer reactions of hydroxide with a negative ion (6 x lo9 dm3 mol-l st1).19 This fact is surprising, and makes it possible to assume that the intermediate Ci is not an open intermediate monoanion, but a cyclic intermediate zwitterion.If so, the assumption kpo+ = k!io' must be valid, and the equilibrium constant Kf must represent an equilibrium of simple proton-transfer react ions. According to this, the estimated rate constant for the decomposition of the intermediate Ci-to products, kz-, becomes 9 x lo2 s-l from the value of (k?j0'/kz-) = 5.5 x 107 dm3 mol-l. Furthermore from the estimated value of KpO= 1 x lopG,the second acid dissociation constant of the intermediate, K:.y = 1.6 x mol dm-3, can be calculated (KpoH20f2- = Ktg). We assume that in the equilibria (19) KfSodoes not differ from l/KSIJo+, and so we obtain an estimated value for k-of 1 x lo6s-l (K:iK2-/k-= 1.4 x 10-8 mol dm-3).The estimated rate constants are summarized in Table 8. TABLE8 Estimated rate constants for reactions involving the intermediates Ci, Cy, Ci-, and C; in the oxidation by periodate of 2-aminoethanol Reaction Per.-+ AE --t Ci C; +Per-+ AE Ci + H,O --+-Ci-+ H,O+ Ci + OH-+ C,Z-+ HZO Ci---+ Products Ci + H30++Cp + HSO Ct-+ H30+--t C, + H,O C; + H,O+ +Ct + H,O Ci-+ H,O+ +C, + H,O C, + H,O +Ci-+ H30+ C; --t Products We note that in the periodate oxidation of propane- 1,2-di01~~ is 5.2 x lo2 dm3 mol-l s-l, while in the k, oxidation of 2-aminoethanol K, is 2.5 x lo4dm3 mol-l s-l.The catalytic constant for hydroxide (K,k,OH-) is 1.1 x 107 dm6 moP s-l and this value is almost the J.C.S. Perkin I1 According to the estimated acid dissociation constant, the amide-ester intermediate is an acid about 1 000-fold stronger than periodic acid. Probably this is why the catalytic effect of weak acids was not detectable. The estimated second-order catalytic const ants for ammonia, hydrogenphosphate and acetate are kFEs 1.3 x lo8, @Pop*-7 x lo8, and kkco-7.5 x lo6 dm3 mol-l s-l, respectively. The second-order catalytic constant for ammonia and hydrogen phosphate is only about one order of magnitude lower than k2H-, and it has been shown20 that for strong bases such catalytic constants can approach the diffusion-controlled limit.In this case the two proton transfers follow a stepwise mechan- ism. The catalytic constants for water, acetate, ammo- nia, and hydroxide lie on the Bronsted plot, but the hydrogenphosphate catalytic constant is higher than expected by about an order of magnitude. This can be due either to bifunctional catalysis for the two proton transfers for the conversion of Ci into Ca, or to the fact that the stepwise proton transfers are involved, in which case must change from zero to one in the sequence NH,, HPOi-, and AcO-, and the catalytic constant for water must be abnormally high.20y21 For the general acid-base catalysed periodate oxid- ation of 1,2-diamines we suggest the mechanism in Scheme 2. We have made a detailed kinetic analysis on the basis of the mechanism in Scheme 2 as follows.The third-order catalytic constant for acid, kamp; = K,K:k; and k,B/ke2 = 1. The third-order catalytic constant for base, k:, = KIK*-KFkf and kf/kt2 = 1. We assume that all second- order catalytic constants for hydronium ion are 5 x 1O1O ( c;-CH2-NH CH2-NHCy-NH2I 1.1+ IOL bsol;2 CHZ-OH k-1 *I--' k2-'IOkH2-I04H2-Products CH2-0 '+ CH2-0'I k*_2IAI H ( CYl CCi 1 SCHEME1 same as that for the oxidation of 2-aminoethanol dm3 mol-1 s-l, the second-order catalytic constant for (K,kiH-1.6 x lo7dm6 moP s-l). hydroxide, kiH-, is 6 x lo9, and kfH-is 1 x 1O1O dm3 Mechartism.-The mechanism in Scheme 1 is one mol-1 s-1, and water can act both as an acid and a base.which is in harmony with the experimental facts already In the oxidation of ethane-1,2-diamine this leads to discussed for the oxidation by periodate of 2-amino- estimated acid dissociation constants for the acid Cp, ethanol. Kfi, of 3.7 x lo-* mol dm-3 and for the acid C;*, 1980 K:;*, of 3.5 x mol dm-3, with K, 5.4 x dm3 mol-1 and K* 3.1 x loW3.In the oxidation of NNrsquo;-dimethylethane-1,2-diamineall these constants have similar orders of magnitude except K, which is estimated to be 0.54 dm3 mol-l. In the oxidation of 1,2-diamines the third-order k1_10; -1-1 i jH2-NHlo,n;K* Fast Products~ CHZ-NH I H SCHEME2 catalytic constants for ammonia, 1,e-diamine, and monoprotonated 1,2-diamine are of the same order of magnitude and only 20-30 times lower than the hydroxide catalytic constant.The order of magnitude of the estimated acid dissociation constant K:;* (10-6 mol dm-3) makes it possible to assume that the monoprotonated 1,e-diamines act as bases, and the second-order rate constant k; can assume the value of the diffusion controlled limit. 8/2001 Received, 17th November, 19781 REFERENCES C. A. Bunton, in lsquo; Oxidation in Organic Chemistry,rsquo; Part A, ed. K. B. Wiberg, Academic Press, New York, 1965, ch. 6; G. J. Buist, Kinetics of Oxidations by Periodate,rsquo; in lsquo; Com-prehensive Chemical Kinetics,rsquo; eds. C. H. Bamford and C. F. H. rsquo;llsquo;ipper, Amsterdam, 1972, vol. 6, p. 435. C. A. Bunton and V. J. Shiner, jun., J. Chem. SOC., 1960, 1593. G.J. Buist and C. A. Bunton. J. Chem. SOC.,1954, 1406; 1957, 4580. F. R. Duke, J. Amer. Chem. SOC., 1947, 69, 3054; F. R. Duke and V. C. Bulgrin, ibid., 1954, 76,3803. P. Zuman, J. Sicher, J. Krupicka, and M. Svoboda, Nature, 1956, 178,1407. (a)G. J. Buist. C. A. Bunton, and J. Lomas, J. Chem. SOC. (B),1966, 1094, 1099; (b) G. J. Buist, C. A. Bunton, and W. C. P. Hipperson, ibid., 1971, 2218; (c) G. J. Buist and C. A. Bunton, ibid, 1971, 2117. (a)G. Dahlgren and J. M. Hodsdon, J. Phys. Chem., 1964,68,416; (b)G.Dahlgren and E. M. Rand, ibid., 1967, 71,1955.* G.E. McCasland and D. A. Smith, J.Amer. Chem. SOC.,1951,73,5164. J. Kovar, J. Jary, and J. Blaha, Coll. Czech. Chem. Comm., 1963, 28,2199. 10 P.Fleury, J. Courtois, and M. Grandchamp, B.ull. SOC.chim. France, 1949, 88. 11 (a) L. Maros, I. Molnir-Perl, and E. Schulek, Acta Chim. Hung., 1962, 30, 119; (b) L. Maros, I. Molnbr-Perl, and M. Molnar, ibid., in the press. 12 G. J. Buist, W. C. P. Hipperson, and J. D. Lewis, J. Chem. SOC. (A), 1969, 307. l3 C. R. Bertsch, W. C. Fernelius, and B. P. Block, J. Phys.Chem., 1958, 62,444; (b) G. H. McIntyre, jun., B. P. Block, and W. C. Fernelius, J. Amer. Chem. SOC., 1959, 81,529. 14 F. Basolo, R. K. Murmann, and Y. T. Chen, J.Amer. Chem. SOC.,1953, 75,1478; F. Basolo and R. K. Murmann, ibid., 1954, 76,211. 1s C. W. Davies, Ion Association,rsquo; Butterworths, London, 1962, p. 41. l6 R. G. Bates and G. D. Pinching, J. Res. Nut. Bur. Stand., 1951, 56, No. 5. l7 R. P. Bell and P. G. Evans, Proc. Roy. SOC., 1966, A291,297. 18 R.P. Bell, lsquo; The Proton in Chemistry,rsquo; 2nd edn.. Chapmann and Hall, London, 1973. lo M. Eigen, W. Kruse, G. Maass, and L. DeMaeyer, Progr.Reaction Kinetics, 1964, 2,308. 20 W. P. Jencks and M. Sayer, Faraday Symp. Chem. SOC.,1975,10,41, and references therein. 2rsquo; R.E. Rarnett, Accounts Chrm. Res., 1973, 6, 41.