Metal ion catalysis of phosphoryl hydrolysis from thiolphosphate monoesters

BIOINORGANIC CElZMISTRY, 1,107-131 (1972)

Metal

Ion Catalysis

TL:-l-L-e-L~*.% ~~~lu~pm_mp~la~~

S. J. BENKOVIC*

of Phosphoryl

107

Hydrolysis

from

mK~~~~~C~~~ ~~xumut;cILtxcT

and E. Ri. MILLERt

Department.of Chemisf+y, The PennsyluaniuState UniversiLy, Uniuerdy Park, Pennsyhni5, U.S.A.

ABSTRACT The catalytic effect of metal ions on the hydrolysis of ethyl-Z-phosphate (la), cysteamine-S-phosphate (lb),and 2-carboxyethyl-S-phosphate (IC) has been investigated. Fe(II1) and Al(III) markedly inhibit the rate of hydrolysis of all three thiolphosphates, decreasing kOb by cu. 102 at pH 3; Z&II) and M.n(II) are ahuost without effect on the hydrolysis rates of lb and lc; and C&II) in leas than substrate concentration increaczs xb, by >lO* for lb and lc but has no effect on la. The requirement of a suitably located hgand for catalysis is obvious. The mechanism of the Cu(II)-catalyzed hydrolysis of lb has been examined in detail and inchrdes the following pertinent characteristics: (1) ratedetermining formation of two hydrolytically labile complexes from monoanion (NHrR-SPOSH-) and/or dianion (NIL-R-SPOs*) and Cu(Ka)sf+; (2) product inhibition by cysteamine because of the formation of a catalytically inactive cysteamin+Cu(II) complex; and (3) incursion of an oxidative Cu(II)-catalysed conversion of cysteamine to cystamine at pH >7, which relieves product inhibition. The synthetic phosphorylating potential of this reaction in acetonitrile-alcohol solvent mixtures indicates its utility in the synthesis of simple phosphate monoesters. A rationale for the effectiveaess of Cu(II) catalysis in these systems is offered.

Introduction Thiolphosphate monoesters (1) have been utilized as chemical phosphorylating agents under oxidative and nonoxidative conditions to synthesize both P-N and P-O bonds [l, 27. *National Institutes of Health Career Development Awardee, AXred P. Sloan Fellow, 196%70_ t National Institutes of Health Predoctoral Fellow, 1967-71. Copyright @ 1972 by American E&evier Publishing Company, Inc.

108

AT?TDE-

S. J_ BENEOVIC

OH

ii

R--S-P<

(1)

OH

l3ew-e

of the lability

la:

R = -CH&&

lb:

R = -CH,CH,NH,

lc:

R = -CH&H&OOH

of the P-S

bond

of

thiolphOSPhate

monoeSter

~ono~o~

(EtSPOa-, tlr_ = 29 min, 350) [3], several authors have speculated that such species may participate iu biological phosphorylations, pSkUkW~Y as envnebound intermediates [4]_ Although tbioIphosphate monoesters per se have not been detected in living organisms, enzyme preparations, such as alkaliue phosphatase [5], which catalyze the hydrolysis of these monoesters through P-S bond cleavage have been reported [S, 7, S]. In addition, a monothiolphosphate hexose transferase extracted from E. coli apparently can catalyze the transfer of the phosphoryl moiety from the parent phosphorothioic acid to hexose c9]. A possible rationale for the absence of observations of thiolphosphate intermediates in biological phosp%oryl transfers is the increased sensitivity of the P-S bond to hydrolysis in the presence of various metal ions [5] and other functional groups [lo], superimposed on their spontaneous hydrolysis. This apparent catalysti by metal ions has not been quantitated. For these reaSOns we have chosen to study a series of thiolphosphate monoesters (la-lc) as

possible models for enzymic thiolphosphate intermediates in an effort to elucidate the eifects of metal ion interactions on the transfer of the phosphoryl moiety in this system_

Au melting points are uncorrected. Microanalyses

for carbon, nitrogen, and sulfur were performed by Midwest &Iicrolab, and those for orthophosphate by the method of Martin and Doty [ll] as modified by Jencks and Gilchrist [12]. Infrared spectra were obtained with either a Perkin-Eimer 137 spectrophotometer or a Beckman IR5A infrared spectrophotometer. Descending paper chromatograms mere run on Schleicher and Schuell orange ribbon 5S9c paper, employing ~3 solvent the mixture 0.1 b!f KZCOrabs EtOH 3.5 : 6.5, followed by development with Hanes-Isherwood spray [13 3. Materials Metal n&ate solutions (Baker, reagent grade) were standardized by the appropriate EDTA procedure of Flaschka Cl41 or Schwarzenbach [15] before

METAL

ION

CATALYSIS

OF THIOLPHOSPHATE

HYDROLYSIS

109

utilization. Methanol (Baker, reagent grade), Hz 180 (50/ Biorad) , and twicedistilled deionized water mere employed as solvents for the kinetic runs. Buffer materials and solvents used in synthetic procedures were reagent grade (Baker, Eastman, Mallinckrodt, Matheson, Coleman, and Bell). Recrystallization of cystamine dihydrochloride (AIdrich) from methanol did not change its melting point, 223” Dt. [16] mp 212-212.5” (corr.) 3. Descending paper chromatography on Whatman No. 4 (n-butanol-acetic acid-water 2: 1: 1) showed the material to be homogeneous (RI: 0.53, 25O). Development w-as with ninhydrin reagent (0.1 g ninhydrin-40 ml abs EtOH-10 ml pyridine), followed by incubation at SO0 for 3-5 min. Cysteamine (%aminoethanethiol, Columbia) was sublimed (70”, 1 mm) immediately before use. 2,2’-Biquinoline (Aldrich) was used without further purification. The desired thiolphosphate monoesters were obtained by the methods of Akerfeldt [17-191, and the molecular formufas below quote his experimentally determined compositions.

Ethyl-S-phosphate

(dilithium

salt)

Ir: (KBr) 3.33 (sharp), 3.38 (m), 3.45 (sharp), 6.S9 (m), 7.25 (sharp), 7.90 (m), 8.95 (broad), 9.35 (broad), and 10.0 p (broad). Anal. Calc. for GH~POJ_&(H~O)O 5: C, 14.7; H, 3.7; P, 19.0; S, 19.6. Found: C, 15.1; H, 3.2; P, 19.1; S, 19.4. 2-Carboxyethyl-S-phosphate Ir: (Nujol)

(mixed lithiumsodium

9.35

(broad),

and

AnxzZ_Calc. for (PO~CH&H&OO)Li,_snT~_s(CH,0H)~_~(H~O)~.5: H, 2.7; P, 13.5; S, 13.9.

C,

17.3;

6.30

(m),

6.38

(shoulder),

S.90

salt) (broad),

9.98 p (broad).

Found: C, 17.2; H, 2.3; P, 13.3; S, 14.3.

Cysteamine-S-phosphate

(monosodium

saU)

Ir: (IS&) 3.254.00 (broad multiplet), 4.85 (w), 6.16 (m), 6.25 (m), 6.49 (m), 6.92 (sharp), 7.10 (sharp), 7.32 (w), 7.60 (w), S-01 (w), W-9.2 (broad multiplet), 9.60 (m), 10.26 (s), 11.10 p (m). Anal. Calc. for NHpCH&H&POJHNa: C, 13.5; H, 3.9; N, 7-S; P, 1’7.3. Found: C, 13.6; H, 4.0; N, 7-S; P, 17.0. The monosodium compound was obtained as the free acid by the procedure described by Akerfeldt [20], mp 15&156O, [ht. [20] mp 156” (corr.:]- Conversion to the dicyclohexylammonium salt was performed by dissolving the free acid (4.1 g, 2.6 x 10-Z moles) in a minimum amount of distilled water (uz. 100 ml), folIowed by dropwise addition of freshly distilled cyclohexylamine (Abbott) until the pH of the solution was M. 9. The desired salt was precipitated

110

S_ .T. BENI(OvIc

AND

E.

M_ MILLER.

from solution by adding-200 ml absolute ethanol and 400 ml anhydrous ether. Thorough washiig with 100 ml absolute ethanol, followed by 50 ml anhydrous ether and clryirg for 24 hr under vacuum at room temperature, yielded 3.7 g (1.01 X 10-4 moles) dicyclohexylammonium crysteamine-S-phosphate, mp 257259*. Descending paper &xomatography showed the material to be homogeneous (I&: 0.71, 25*). i‘r: !?i ujal) 4.59 (broad), 6.05 (shoulder), 6.19 (m) , 6.42 (m), 7.11 (w), 7.92 (broad), 8.8-9.7 (broad muhiplet), 10.29 (m), 10.41 (m), and 11.02 p (m).

This compound, used as a reagent in the quantitative detection of tbiols, was prepared by the method of Akerfeldt [Zl], mp 185-186”, fit. [Zl] mp 1900]. To avoid decomposition, the compound was stored at 0” in the dark. Apparatus The instrumentation used in this study has previously been described [22]. Most kinetic runs p :re carried out in Kimax screw-cap tubes (No. 9447-A6) v&h Teflon-lined caps (No. 94?7-B3), maintained at constant temperature (&O-l> by a circulating xvater bath. Kinetics The hydrolytic rates of ethyl-S-phosphate (9.4 X lo-4 Al), cysteamine% phosphate (8.9 X lo-4 M), and 2-carboxyethyl-Sphosphate (8.3 X lO+ M) were determined by monitoring the release of orthophosphate by the method of Martin and Doty [ll] as modified by Jencks [127. Pseudo-first-order kinetics were observed to at least three ha&lives. The observed first-order rate constants for the hydrolyses were calculated from the slopes of plots of Iog,[OD,/OD, - OD,] against time. Buffers employed were HNOJ (pH <2.5), form&e (OZK, pH 2.8-3.81, acetate (0.2&f, pH 3.~5.4), succinate (O.O67iW, pH 5X-6.3), ltnd Tris (OS&f, pH 7.0-8.9), allat p = 0.2 @NO,). No accelerakry buPer effects were noted (0.03-0.2OM). The pE of the kinetic runs was measured at 35” (glass electrode) upon initiation and after completion of the experiment; those exhibiting pH drift greater than f0.05 unit were discarded. The release of orthophosphate in the presence of the metal ions employed (5.0 X 106-2.4 X 1W2 M) was also followed by the modified method of Martin and Doty. The analys& was not affected by the presence of metal ions. AlI runs were maintained at p = 0.2 with KNOa, with the pH range limitedby

h4ETAL

ION

CATALYSIS

OF THIOLPHOSPHATE

HYDROLYSIS

111

metal ion hydrolysis employing the buffers listed above. Buffer effects in the Cu(II)-catalyzed hydrolysis of cysteamine$-phosphate n-ere negligible with the exception of Tris (0.01-0.2OM), which inhibited the rate of hydrolysis with increasing buffer concentration. Buffer effects in the presence of other metal ions were not investigated. Slight precipitation’ (pH 2-5) occurred at the end of the Cu(II)-catalyzed hydrolysis of 2-carboxyethyl-Sphosphate and the Fe (III)-catalyzed (pH 1.8) hydrolysis of all three substrates, but did not lead to sign&ant deviations in the ilrst-order release of orthophosphate. Anaerobic hydrolysis of cysteamineS-phosphate was carried out by first degassing the buffer solution (25 ml) with pure nitroger? (1 hr) , which n-as then piped under positive nitrogen pressure to a stoppered two-neck, pear-shaped flask (50 ml) containing the compound. Rubber septums were used at all connecting joints. The system was vented by placing a hypodermic needle through one of the septums at the top of the reaction vessel, with a piece of tygon tubing attached, and the end was immersed in water to insure against any backflow of air into the system as samples were extracted_ The entire system was flushed with nitrogen before the start of each experiment and during the kinetic run to assure exclusion of oxygen. SampIes (1 ml) of the reaction mixture mere removed through the septum, using a gastight syringe while maintaining the solution under constant nitrogen atmosphere. The rapid Cu (II) -catalyzed hydrolysis of cysteamine-S-phosphate was found to be inhibited by increasing concentrations of cysteamine (a product of hydrolysis) at pH <7 (see “Results”). For this reason, the rate of hydrobsis for this reaction was calculated from data obtained for the -first15-200/oof hydrolysis, necessitating a rapid and accurate method of extracting samples from the hydrolysis solution. The following initial rate technique was devised for this purpose. The buffer solution was equilibrated at room temperature (25 3~ lo), transferred to a tube containing the substrate (cu. 4.0 X 10-j g), shaken vigorously for 10 set, and then transferred to a 2%ml calibrated buret, The total elapsed time n-as ca. 20 sec. Samples (1 ml) mere removed at 5-10 set intervals and quenched in 4 ml HN03 buffer (pH l-l), which replaced the 4 ml of water prescribed by the modified Martin and Doty procedure. Control experiments employing known concentrations of orthophosphate showed that quenching in this acid medium had no effect on the accuracy of the analysis. Products Hydrolysis of ethyl-S-phosphate and 2-carboxyethyl-Sphosphate yielded, quantitatively, orthophosphate and thiol over the pH range investigated under aerobic conditions. Thiol was determined by the method of Akerfeldt [217, using CMNP 9s a trapping reagent. The less than 50yo recovery of cysteamine (M&CH&H$3H) during the spontaneous and metal ion-catalized hydrolysis

METAL

ION

CATALYSIS

OF THIOLPHOSPHATE

HYDROLYSIS

113

evaporated to less than 1 ml. After addition of 5 ml absolute ethanol and 20 ml anhydrous ether, the compound precipitated as a white gel. Thorough washing with alcohol and recrystallization from water-ethanol yielded, after drying, 12.5 mg (570/o) ethyl phosphate (2 Li+), RI: O-SO, 26O; ir: (I(Br) 3.35 (sharp), 3.40 (shoulder), 6.7-6.9 (broad multiplet), S.99 (broad), 9.30 (broad), 9.SO (broad), 10.55 (broad), 11.40 (s), 12.5-12.9 P (broad multiplet). The phosphate monoesters, methyl and phenyl, were also characterized, using descending paper chromatography1 Since no pyrophosphate was detected, the analysis in Ref. 26 could be applied.

ResliItS The pH-rate profiles for ethyLSphosphate (:a), Cystesmine-S-phosphate (lb), ad 2carboxyethylSphosphate (lc) were determined; they resemble the profiles obtained by Akerfeldt [20] and by Dittmer et al. [2S] for S-phosphate monoesters. The observation of a pH-rate maximum (pH 34) recalls other phosphate esters for which the monoanion is considered the most reactive species [29, 30, 317. No reinterpretation of the conclusions of these other authors is necessary. Our main intent \-as to establish reference pH-rate profiles under our experimental conditions. The effects of metal ions on the observed rate constants for hydrolysis of la, lb, and lc are shown in Table 1. Co(II), Ni(II), Cs(II), and Mg(I1) (2.4 X lo-’ M) do not catalyze the reaction (pH l-5). Zn(I1) (2.4 X lo-’ M) has a two acceleratory effect on lb (above pH 3) and a 1.4fold effect on lc at pH 5. Similarly, &In(II) (2.4 X lO_’ Ji) enhances the hydrolysis of lc (above pH 4) by 15fold. Note that lb and lc have additional chelating groups available. In contrast, the presence of Al(II1) (1.6 X lO_’ M) leads to an inhibition of the rate of hydrolysis of la, lb, and lc (above pH 1.5), an effect that becomes progressively stronger as pH increases (pH 3, twofold; pH 5, twenty- to thirtyfold). Likewise Fe(II1) has an inhibitory effect on the rate of hydrolysis of la, lb, and lc_ At pH 1.8 the observed rate constant for spontaneous hydrolysis of la, lb, and lcisca.SO-140 times greater than in the presence of 1.6 X lo-” M Fe(II1). The data reveal that relatively small Cu(I1) concentrations (M[Cu(II)]:M[substrate], 1:9) greatly enhance the hydrolysis of lb and 1~ (above pH 4) but have very little catalytic effect on la. With a Cu(II) concentration only 10% that of substrate, the rate constant for hydrolysis of lb is increased by a factor of > 102 (pH 4-5) and by a factor of > 103 ($I 6-S) The pH-r&e profile for the Cu(I1) -catalyzed hydrolysis of cystenmine-% phosphate, ali rates obtained using initial rate measurements (see “Experimental”), ia shown in Fig. 1. The observed rate of hydrolysis in TJ?S buffer (pH ~7) w~ obtained by extrapolation to zero buffer concentration. The overall rate of hydrolysis over the pH region investigated, 4-8, may he expressed

S. J. BENKOVIC

114

AND

E. M. MII;I;ER

TABLSl Meet of MetaI Ions on the Observed Rata of H~Iy& of Ethyl-S-Phosphate (la), Cssteamine-s_Phospha~(lb),and !LCarboxyethyLS-Pho8phat.e (1~)at 35" (p = 0.2)a Relative Hydrolysis Rates Metal Ion

!Gic(II)

lkf

PH

la

lb

1.3 >lXlW

lc

2:2

2.4 X 1.0 x 2.0 x 4.0 x 1.0 x

lo-' lo-' lo-' lo-' lvi

1.2-1.4 4.0 4.0 4.4 7.0

1.0 1.0

1.1

>lXlW

2.4 X 2.4 X 2.4 X 2.4 X

lO+ lO-= 10-t

1.5-i-1.9 2.8-3.0 4.0 4.9-5.0

1.0 1.1

1.0 2.3

0.7

2.3

1-o 1.0 1.0 1.4

1.3-1.5 3.0 4.9

1.0 1.0 0.8

1.0 1.0 0.9

1.3 1.5 3.0 3.6 4.9 5.3

1.0

lo-’

Cobalt@)

2.4 X 101 2.4 X lo-2 2.4 X lo*

l!?ickel(lJJ

2.4 X 2.4 x 2.4 x 2.4 x 2.4 x 2.4 X

1F 10-= lO+ 10-t lO-' lo"

2.4 4.4 >lXl@

1.0 1.0 1.0 0.8

0.9 1.0 0.9 1.5

2.4 X 101 2.4 X IO-= 2.4 X lo-2

1.2-1.3 3.0 4.9

1.0 1.0 0.8

2.4 X l(r 2.4 X 103 2.4 X 10-f

1.2 I_5 3.0

1.0 1.0

2.4 X 1W

3.7 5.3

1.0

1.0 1.0

1.0 1.0

1.0 1.0

2.4 x lO+ 2.4 X

lO+

2.4 x 10-t

2-4 x lo-, 2.4 X 101 1.6 x 1.6 X 1.6 X 1.6 x 1.6 x

1.5 3.1 5.0 5.3

101 1.3 lo-, 1.5 lO_' 1.9 1F 101 i4.;:.9

1.6 x 101

1.0

1.8

0.8

0.8 0.9 0.9

0.7 0.4 0.05

0.5 0.4

0.4 0.03

0.007

0.62

0.016

o The concentrations of thiolphosphates empIoyed were BS foUowB: lh9.35 lb, 8.92 X lo-‘M; slid lc, 8.32 X 10-M.

x lo-'M;

MEZ%L ION CATALYSIS OF THIOLPHOSPHATIZ HYDROLYSIS

c, Ok

I

I

I

I

I

,\P

2

3

4

5

6

, 7

8

115

I

9

PH

Figure1, The p&r&e profjksfar the apantan~w (0) and Cu@)-cataIyzed (I) hydmlyais at29,p =O.Z ThesoLidlinesaretheoreticslcurvescalculstRd fro~valueslistedin Table2

ofcyatesmineSphosp~i.e

in terms

of the following kinetic equation:

+ ~D~*R~PO~~-I[~(H~O)~~+J,

(1)

where R = [CH&, Aa is a first-order rate coostaot associated with the spontanetrus hydrolysis of the neutral species, and k~ and kD are second-order rate constants for hexaaquo Cu(I1) -catalyzed hydrolysis of the monoanionic and dianionic species, respectively.3 Equation 1 may be transformed into an expression for the experimentally determined pseudo-fkst-order rate constant, kpbg,by introducing the mass balance relationships: ET = cN&t_R_sPO3H_7

+ m3-R+PO?-]

+ FJHTR-sP0;H-J + ~rR-SPO?-I],

C+r = [Cu(H,O)Z+]

+ ~cU(H,O)~OH+],

(2) (3)

where ET and Cm are the stoichiometric concentrations of cysteti~ phosphate and Cu(II), respectively. Partitioning tn additional species is not

11s‘

S.

J. BEXKOI’IC AND E. M. MILTZR TABLE 2

ate and Eqoi.likiumConstantsfor the Cu@I-Cata&zed Hyd&sis of cysteaminGPhoaphate(r= O.%W") Kinetic Psrameter Ku”

k,_ hilb K&b

-

K;a b K,=.= K,d

K6 Kf The aDDmt

co~st.antsfor acid diskation,

Numerical Value

(2.64* 1_10)1o%w-~r&l-1 (5.51zk 3_33)1O~M-1 lx&-* (1.78 * 0.15)10-1 (2.44f 1.19)10-' (5-00 f (5.01 f

140)10-” 020)10-7

3.5 x 10' (1.6 f 0.6)10-‘ =3 x 10-T K,, Kt, KS, and K,; formstion, Kr a& Ka; and

aTheses vaIues were obtainedtbrougbiterative calcuhtione employing a nonlinearleast Squares computer programon the IBM 360 computer facility. bRef. 20. =Ref. 43, p. 59. dRef_ 42. ‘HII = WIG insofaraS!kBcrasoOpic equ&br;Um cons&~.& msy be ueed as BPapproGma~On for ticroacopic(groupsnot interacting)values.

ski@cant in the PH range investigated.Employing the apparent acid WCiation and tautomeriaation constants as defined in Table 2 and neglecting the koterm &&ive to the kx and ko terms, it, can be shown that, since 21= kObSCErl,

(4) whereaH = activity of the hydrotiu ion BSmemed by the glass electrodekpection of Fig. I reveaIsthat V~UCB of kobe ~&~&ted from Es- 4, Utikiag

METAt

ION CATALYSIS OF THIOLPHOSPHATE

HYDROLYSIS

117

the numerical assignments of the parameters listed in Table 2, are in satisfactory agreement with the experimental observations. Below pH 4 the release of orthophosphate deviated from first-order kinetics. At PH 4 and 7, Ctc~ was varied from 5.0 X 106 to 1.0 x lO+ M at constant cysteamin~phosphate concentration (8.92 x l(r M). Logssithmic plots of kobsversus CUT were linear with slopes of unity witbin experimental error. The cataIy=d hydrolysis, thus, is fkst order with respect to C;, over the range of catalyst concentration investigated. At CzcT > 1.0 X UP M the appearance of orthophosphate becomes too fast to folIow by the present method of analysis, and the inquiry as to whether a change in the kinetic order of Cur occurs at high concentrations remains unanswered. Measurement of the rate of Cu(II)-catalysed hydrolysis of c&,teamine-% phosphate was carried out under both aerobic and anaerobic condittons. In the pH region 4-6, no change in the rate of orthophosphate release n-as observed (Fig. 2), although thiol rather than disuhide was the major product under anaerobic conditions (see “Ekperimental”). Only trace amounts ( < lOTo) of Cu(1) were detected under the latter conditions. At pH >6.5, however, the rate of hydrolysis was drastically reduced in the absence of oxygen (Fig. 3). In addition, at pH 7, formation of Cu(1) (>50%) and disulfide (30 + 5%) was detected under anaerobic conditions_ Hydrogen peroxide, which could form during oxidation of the sulfhydryl group of cysteamine to disulfide (cystamine) by Cu(II), was not found (pH 4-7) as an intermediate under either aerobic or anaerobic conditions.

t (mm) A) (=UT (1.0 X lo-’ i%.+zatiyzed 0) and anaerobic (----, Figure2. Aerobic (-, hydrolygis of cw&-phosphate (8.92 X lO_’ iK) at 4-06 (27.3q *IL= 0.2), where OD, ie directlyproportional to orthophosphate concentration.

S. J. BENKOWC

118

AND E. M. MILLEB

Q#’

-

,’ : : :

Oo

10 ,

20 1

30 t

40I

50 I

-4

__b__-*------

A__-----

I\ir Itltroduced

601

70 I

00 t

90I

100 I

110 8

120 1 .

t (min)

0) and anaerobic (- - - -, n) CUT (1.0x lo-'a%caWyzed Figure 3. Aerobic (-, hy&oiysisof cysu&phosphate (8.92 x 1O-4M) at pH 7.10 (26-O”,P = O-21,whereODr is directly proportional to orthophosphate concentration.

The rate of the Cu(I1) -catalyzed hydrolysis of cysteamine-S-phosphate WLS shown to be inhibited by increasing concentrations of the product, cysteamine (pH 4.10, Fig. 4), and necessitated the initial rate method- Kinetic runs were begun with known added amounts of cysteamine, and measurements recorded

before production of cysteamine via hydrolysis of the ester could contribute significantly to the total concentration of the inhibitor_ The inhibition exhibited by cysteamine is in accord with its complexation of Cu(H~O)~+, resulting in a decrease in the elective concentration of this species in solution4 The decrease in kOb3(PH 4.1) may be expressed quantitatively by setting Cur = Wu(HoOM+]

+ [:(h=z-R-S-)

.Cu(H20)4=+],

(5)

where [(N&-RAF) -Cu (H20) 2+] represents the assumed monocysteaminea(n) chelate. Partitioning to CU(H,O)~OH+ is negligible at this PH. Utilizing

the apparent acid dissociation constants and chelate formation constant, KS, defined in Table 2, and solving as before leads to

The ~II term is insignifkant at this pH_ Inspection of Fig. 4 reveals that v~W.5 of hbs generated from Eq. (6) and Table 2 are in satisfactory agreement with

the experimental observations_ At PH 7, where the rate of release of orthophosphate is sensitive to anaerobic or aerobic conditions, the rate apparently is unaltered by the addition of cqS_ t=mine, although an initial lag phase in the kinetics was observed. This 6nding

MEZ’AL ION CATALYSIS

0

IO

OF THIOLPHOSPHA~

2.0

3.0

HYDROLYSIS

119

40

CCysteamrne I x 10v5 M

Figure4. The effect of inapasing concentrations of cystamine on k.,b for the &I(U)Cat&~ hydrolysis Of cysteamine&phosphate; conditions: CUT (1.0 X lo-' &f), ET (8.92 X lo-( M), pH4.10, 25“, p = 0.2. The solid curve is theoretical, calculated from Eq. 6 and the values listed in Table 2.

t (min)

Figure 5. Plots of the release of orthophospbate (ODd veraus time for varying incubation timea before initiating the kinetic run; the solid curvy are experimental; conditions: CUZ (1.0 X 1W4 M), cyateamine (1.25 X l(ra M), pH ?.10,24.5q p = 0.2. For kin&c run A (0), Preincubationtime24 min; for B (A), 10 min; for C (o), 1 min.

S. J. BENKOWC

120

AND E- IS%-MIULER

TABLE

-

3

Per at of Methyl Phosphate Fon.ned on %l~~lysis of Ethy1-Sph~phat.o (la), cy&.~&-phosphate (lb), and 2-Carboxyeth&Sphosphat.e (1~)

& Methauol-Water” (35”, fi = 0.20)

SU-bstrste la la + Cu(u)lb lb + Cu(II)’

lb + Cu(II)= :: -I- Cu(Iw Ic f Cu(II)s

PH 3.60 3.76

Mole percent inorganic phosphate in productb

for methanol

55.4 55.2

2.1

selectivity

l-8

3.65 3.70 7.00

53.8 45.6 37-9

2-7 3.7

3-64 3.75 5.73

56-7 47.7 44-2

l-7 2-5 2.8

l-9

= 50y0 V/V CH,OH-H@ = 69.1 mole Y0 H&-349 moie % CHaOHbNonselective solvolvsis would yield a mole per cent product ratio of 69.1 :30.9.

'Cur = 3.4 x lo-=-k_ be rationahsed by preoxidation of added cysteamine to cystsmine by CU(II), a mechanism which apparently involves the thiol anion (Rs) [32]. Since increasing concentrations (up to 6 X 1O-3 Al) of the disdfide product, cystamine, do not alter the rate of hydrolysis (pE 4-7), the inhibition by added cysteamine diminishes as oxidation to disulfide proceeds5 [33]- This hypothesis is supported by the fkliug that the length of the lag period is directly proportional to the preincubation time of Cu(I1) and cysteamine before the addition of ester to initiate the kinetic run (Fig_ 5). Solvoly-tic product distributions in methanol-water for la, lb, and lc in the presence and the absence of metal ion are given in Table 3. In each case the mole per cent inorganic phosphate is less than the mole per cent E&O in the SOL vent. The selectivity factor for phosphorylation of methanol is co. 2-3.7, and mithin experimental error is the same in the absence and the presence of metal ion for la but increases slightly in the presence of Cu(I1) for lb and lc. The results of total hydrolysis of cystean&e-S-phosphate in the absence and the presence of Cu(I1) and Fe (III) in isotopically enriched water are given in Table 4. The data shorn no exchange of ‘80 into the phosphate group of lb during hydroIysk Control experiments employing potsssium dihydrogcn phosphase in the absence and the presence of metal ions similarly showed no exchange under the hydrolysis conditions. The facile hydrolysis of cysteamine-S-phosphate catalyzed by Cu(II) was inves%ted in terms of its synthetic capabfi&s_ Ad&&n of alcohols to the can

TARLE5 A.&4 PhosphateFormationin the Cu(II)-Catalyzed Hydrol~ of Cyskamin&phosphste @icyclohexylammoniumSalt) in Acetonitzil~+H@~ (350, IL = 0.20) Moles Ha/ mole of ROE in solvent

hloles orthophosphate/Mole of fUky1 Phosphate*

Mole Percent A&y1 Phosphate in Product

1.21 0.34

0.59

0.17

62.9 85.3

1.15 0.81

1.00 0.84

54.4

0.30

0.28

78.1

1-Butmlol

2.54 1.15

3.92 1.78

20-O 35.9

z-Propanol

2 -45 1.14

5.16 1.51

16.2 39.9

3.82 1.91 1.27

8.42 3.74 2.21

10 3 20.8

2.51 1.14

0.47

12.07 1.70 0.57

7.4 37.0 63.3

2.16 l-08

4.92 2.08

32.5

0.94 0.39

0.0 51.5 71-S

Alcohol Methanol

%&&Uld

cycIoperLtallol

Cyclohexmol Phenol

3.93 1.54 0.78

b Nonselective phosphorylation would yield numbers identical to those in

50.0

31.2

16.6

column 1.

Cu(I1) -cysteamin&-phosphate system in aqueous acetonitrile produced alkyl phosphates in yields ranging from 10 to 90%. These results are shown in Table 5. Discussion The spontaneous hydrolysis of the neutral and monoanionic thiolphosphate monoester species has been discussed in detail elsewhere [4, 29, 343. It is suBicient to mention here that the usual mechanistic mode for hydrolysis of the neutral species involves nucleopbilic attack by a mater molecule assisted by partial proton transfer to the departing thiolate moiety, whereas that of the

KmAL

ION CATALYSIS OF THIOLPHOSPHATE

HYDROLYSIS

123

monoanion features partial or pm-equilibrium proton transfer to sulfur, followed by metaphosphate expulsion. The juxtaposition of the amino and carboxyl functions in the neutral species derived from lb and the monoanion from lc presents the possibility of intramolecular cychzation through the loss of water to yield the hydrolytically IabiIe intermediates (2) and (3)) thus ahering the mechanism of hydrolysis.

(2)

(3)

However, in the csse of lb, the lack of 180incorporation greater than I atom of the isotope per molecule in the orthophosphate product, combined with the inability to detect an intermediate phosphoramidate, mitigates aga.&t formation of species (2). Furthermore, a structure-reactivity correlation established for the hydrolysis of a series of thiolphosphate monoanions [35] accommodates the rates of hydrolysis of lb and lc (35”) with no significant positive deviation. Consequently, although entry into five-membered cyclic phosphate chemistry appears feasible for lb, the predictably low concentration of the nonzwitterionic neutral species @&-R+P03HLJ6 may result in this cyclization being kinetically unimportant. Cyclization of lc is made Iess likely by the requirement of a sixmembered transition state. In short, there is no evidence for intramolecular catalysis of hydrolysis in these systems. The effects of metal ions on the rates of hydrolysis of tbiolphosphate monoesters may be divided into two categories; the first featuring metal ions &ich inhibit; the second, metal ions which accelerate the rate of hydrolysis. In class I are Fe(II1) and Al(III) . In regard to the inhibitory properties of Fe(III) and Al(III), (a) the presence of an additional hgand is not required, since the rates of hydrolysis of all three thiolphosphates, including the ethyl derivative, are retarded, and (b) rate inhibition occurs with both the monoanion mPO3H-] and neutral mPO&] species, although in the case of Al(II1) the degree of inhibition is greater for the monoanion. In view of the metal ions employed, these two possess the greatest afhnity for ligands which coordinate via oxygen [36]. As a result this rate retardation may be attributed to metal ion chelation with the phosphoryl oxygens, precluding the normal hydrolytic mechanisms. The lack of rate inhibition by Ca(IIj , Mg(II), and Mn(II), which aho preferentially bind at, oxygen, may be ascribed to the decreased stability of such follows the order FeWI) > a(HI) 2 complexes, which qualitatively MGU > Mg(I.U = Ca(H) ~37, 38, 32-j. The difference in stabiity between the tervalent and divalent metd ion-phosphate complexes, unfortunately, k

124

S. .J. BENKOVIC

AND

E. M.

b%JJLER

not known quantitatively. In terms of chelste stability, Fe(II1) should be more effective as an inhibitor than Ai(III), assuming that the decreasein free energy re&cted in increased chelate stability is likewisemanifestedin decreasedgroundstate free energy of the thiolphosphate monoester-metal ion complex. It is striking that apparent strong chelation to the phosphoryl oxygens leads to an inhibitionrather than a catalysis. In the searchfor a rationaIe,let us assume that in the limiting case coordination with Fe(II1) approachesesterificationof the phosphoryl oxygens, that is, total charge neutrahzation. Because of the lack of structure-reactivity correlations for O-substituted thiolphosphate triesters, we will employ similar correlations obtained for O-phosphate esters, partidy justifying their utilization on the basis of a close similarity betKeen the hydrolytic mechanismsfor the neutral and monoanion speciesof the respectivemonoesters. From the data obtained by Khan and Kirby [40], O-phosphate triesternhyhdroIyaeat rates at least W-fold more slowly than the correspondingmonoester monoanions when the leaving group corresponds to a pK, = &the approximate p& of the thiolphosphates (la-c) _ Apparently O-phosphate triesters only approach O-phosphate monoester monoanions in terms of hydrolytic reactivity for leaving groups with p& < 2, owing to the operation of the Iatter’s unique metaphosphate mechanism. Consequently, conversion of the hydrolysis of a monoester monoanion from the metaphosphate mechanism to one involving nucleophilic attack by water as a consequence of metal ion chelation is kinetically less efficient. Such reasoning, obviously, should not be applied to nucleophilic attack by anionic species. The absence of more than 1 atom of I*0 per mole of orthophosphate product suggests that Fe(II1) is not catalyzing hydrolytically unproductive exchange into the phosphoryl moiety (P-O but not P-S bond cleavage) through possible pentacovdent intermediates. Hydrolysis probably occurs

through the normal metaphosphate mechanism involving uncomplexed ester. In essence, there appears to_be no catalytic advantage in complexation of this type for bioIogica1 transfer processes involving uncharged nucleopbiles. In the second category are those metal ions which catalyze the hydrolysis,

including Zn(II), l&(11), and Cu(I1). Of these, the catalysis exhibited by Cu(II) is more dramatic and hencereceivedmost of our attention, in particular, the reaction with cysteamine-S-phosphate. Catalysis by Cu(I1) requires the presence of additional chelating groups other than the phosphoryl oxygens. Since Cu(II) does not acceI&ate the hydrolysis of ethyls-phosphate, it is inferred that the amino and carboxy moieties of lb and lc flmction in forming a hydroIyticahy labile chelate. Precedent for the latter has been established by study of the Cu(II) complex of LLmercaptopropionic acid? [41-J_ Attempts to isolate the CWI) complex of 8-mercaptoethylamine have been less successful because of the rapid oxidation of thiol to disulfide. However, methyl 2+3min0_ ethyl sulfide ([email protected]$H&), which would appear to serve as a good model for cysteami--S-phosphate, forms stable mono and his chelates Mth Cu(II) CW.

METAL ION CATALYSIS OF THIOLPHOSPHATE HYDROLYSIS

125

Before postulating a mechanism for the Cu(II)-catalyzed hydrolysis, it is necessary to review six experimental facts. (1) At pH 4-6 the rate (k,,bs) of the observed catalysis is not affected by aerobic or anaerobic conditions, the major products (anaerobically) being cysteamine (95r0), orthophosphate, and Cu (II) _ (2) At pH 7 under anaerobic conditions the rate of Cu(II)-catalyzed hydrolysis of cysteamine-S-phosphate is markedly reduced; Cu(1) is detected; and the yield of cysteamine (70$&) is significantly less than the theoretical amount. (3) The observed rate constant for hydrolysis (pH 4 and 7) is directly proportional to Cu(I1) concentration up to an ester Cu(I1) ratio of 9:1-no saturation is observed. (4) The pH-rate profiie at a constant Cu(II)/ester ratio (1: 9) reveals that J&b,is nearly invariant with pH over a pH range of S-8. (5) The rate of the reaction is inhibited by the product cysteamine, but not by the disulfide cystamine. (6) At pH 7 under aerobic conditions, rate inhibition caused by addition of known quantities of cysteamine can be removed by preincubation with Cu(II), the required preincubation time being directly proportional to the added cysteamine. Let us 6rst consider statements 1 and 2. It is apparent that continuous release of orthophosphate at pH < 7 does not depend on the reoxidation of an intermediate Cu(1) species. Only at pH 27 does the liberation of orthophosphate require reoxidation of the metal ion to the catalytically active Cu(I1) state. These observations are in accord with the known reaction of Cu(I1) m-ith mercaptide anions, which pertains in the absence of air and in the presence of excess RS- [33]: Cu’* + 2RS- s $ (RS)

2

+ CuSR.

(7)

Since the pK, of the thiol group of cysteamine is cu. S-3 [43], this equilibrium is displaced to disulfide formation with increasing pH under anaerobic conditions, apparently becoming kinetically important in the present system at pH The presence of oxygen, of course, favors Cu(I1) formation [32]: 2Cu+ + 402 f

Hz0 ti 2Cu*+ + 2OH-.

27.

(8)

To approximate the stability of the presumed Cu(II)-cysteamine-S-phosphate complexes, recourse to the coordination chemistry of methyl Zaminoethyl sulfide was necessary. The latter forms stable mono and bis complexes with

Cu(I1) with formation constants of 3.8 X 105 (KS) and 1.3 X 105 (KS’) at 30” [42], respectively. Because of the similarity of these constants, the stoichiometry of the labile Cu(II)-ester complexes cannot be ascertained by simple calculation. Attempts to monitor reaction rates at varying Cu(I1) concentrations in order to determine kinetically the stoichiometry of the trarkient complexes were unsuccessful_ Since the kinetic rate law, however, is f&t order in Cu(I1) over the pH range considered (statement 3), me will depict the mechanism, for simplicity, in terms of mono complexes, although the involvement of bis complexes cannot be discounted.

126

S. J. BENKOVIC

AND

E. M. MII;L;ER

in &c&g the kinetic expression (Es. 4 and Staten?& 4) that predicts the observed behavior of k,,& as a function of PH, we have i=mmed Steady-state wncenbfiom for the two labile intermediate complexes formed from Cu (H,O) s2f md ~~R+pO&L-and NH~R$POsz-. On chemical grounds, complexation & anticipated with the free rather than the protonated amino group. The invariance in kob. at pH > 5 can readily be assigned to the compensating ionizations : Cu(EI~O)lr XI&+-R+PO$-

;

Cu(E~O)~OII+

2

Nl&-RSPO~2-

+ H30+

(9)

+ HaO*

(W The lack of catalytic activity of hydroxo pentaaquo &(II) has ample precedent, rationalized mai&- on the basis of unfavorable interactions [44]. Further &scussion of the mechanism for hydrolysis of the above complexes and the identity of the rate-determinin g step Gil be deferred until Scheme 1 is presented. Finally, let us consider the inhibition by cysteamine described in statements 5 and 6. The data of Fig. 4 suggest that the rate inhibition is due to complex formation by the product, cysteamine with Cu(I1) thus lowering the available catalyst concentration_ The calculated formation constant (Kp) for the mono cysteamine-Cu(II) complex from kinetic data8 at pH 4 is 1.6 X 10’5. The validity of the high formation constant found for this compIex is supported to some extent by a reported value9 of CQ. lW” E453. Unfortunately there is little

,,r?-!-Li A .F.s-.2 , ‘-tX~H*O'~

*

+ Cu(T1~*(aerobacally,

+ *ii-

or cuv.1

(anaemb~cally~

P03H

NH-2

>=3-

K2

+

2+ cucfi2016

Scheme 1

4

A,

solvolys+s products

METAL

ION

CATALYSIS

OF THIOLPHOSPHAT2

HYDROLYSIS

I27

other precedent for the stability of amin_*+“‘yd+Cu( II) complexes. At pH 7 inhibition by cysteamine can be removed by its conversion to disulfide via Eq. 7, a~ shown in Fig. 5. Under aerobic conditions Cu(1) is oxidized to Cu(I1) (Eq. 8)) allols-hz the hydrolysis to proceed normally, in contrast to anaerobic experiments in which the rate of orthophosphate release decreases with increasing Cu (I) concentration. All of the above facts can be rationalized according to Scheme 1. The pertinent details of this mechanism may be summarized as follows: (1) formation of two hydrolytically labile complexes (4) and (5) from monoanion (NHRSPOJH-) and/or dianion (NH,RSPOs*-) and CU(H~O)~~+; (2) chelation by the product, cysteamine, to cause rate inhibition by decreasing the effective Cu(I1) concentration; and (3) relief of product inhibition at high pH owing to oxidation of cysteamine to cystamine. Let us now inquire into the possibility that formation of either complex may become rate determining over the pH range investigated. The measured rate constant for the substitution of ammonia for a water molecule in the inner coordination sphere of Cu(II) is 1.2 X lo9 M-l min-’ [46]; this corresponds to the rate steps 121 and kl’ in Scheme 1. Comparing this value to the experimentally determtied values of kM (2.6 X lOlo) and kD (5.6 X 107), we find that the value of kx, within tolerance limits of KT, is in close agreement but that the value of kD is somewhat less than would be expected for rate-determining formation of complex (5). A reasonable explanation assumes that interaction of Cu(HzO)Z+ and the dianion (NH,RSPOZ-) d oes not lead to complex (5) alone. The formation of an unproductive complex species-binding of Cu(I1) with the ionized phosphoryl oxygens would tend to lower the actual concentration of (5) so that the calculated value of kD based on kObSis decreased proportionally. Unfortunately the concentration of unproductive complex is presently unknown, and no correction can be applied. Therefore, we conclude that formation of both complexes is rate deter mining over the pH range investigated.10 A minimal estimate for the rates of hydrolysis of (4) and (5) finds kj 2 10’ and k3’ 2 l(r.” Error in ka’ is anticipated since its calculation depends on the assumed values of kl’ and k’. Since hydrolysis of complex (4) via a protonated

metaphospbate mechanism is unlikely [47], this complex is presumed to hydrolyze via a bimolecular mechanism involving attack of water on phosphorus (Scheme 2)) a mechanism previously postulated for O-phosphate diester mono-

=

il,OH

NH -----P -OH2 .\ 2 ,” \ ‘\ d 2+ 0, Cu (H20) 4

-->

Scheme2

H3P04

-----po3\

I ‘C’U (H20)

c-------$

I’

ns

N\t2 “--p02\ / ‘CU (H20) ;’

Scheme3

eons_ Complex (5) is thought to decompose through a transition state (Scheme 3) wtich aen-ise mvolves strong polarization of the P-S bond as a result of a sulfur-cu(~I) interaction that lowers the free-energy requirements for P-S bond cleavage. HoTever, the latter features metaphosphate expulsion based on analogy to ~QI&-Q mec&&ms for monoester dianion hydrolyses with leaving groups of p~(, < 7 [48]. Insofar as structure-reactivity correlations for O-phosphate esters are applicable, the rate of hydrolysis of O-phosphate monoester dianions via metaphosphate expulsion is ahvays expected (for leaving groups of p& < 15) to be greater than the rate of hydrolysis of O-phosphate diester monoanions, xbich hydrolyze via a bimolecular mechanism [40]. Therefore, values of h should be viewed only 8s minimal estimates and should not be contrasted. The similarity in selectivity for methanol of species (4) and (5) is attributed to increased reactivity of the phosphoryl moiety as a result of strong Cu(I1) interaction with the leaving group in both complexes. The activated phosphoryl moiety may be expected to exhibit little solvent selectivity regardless of the hydrolysis mechanism, a situation which has been previously encountered in the Hg(II)-catalyzed hydrolysis of phosphoenolpyruvate [49]. In addition to the unique kinetic and mechanistic aspects of the Cu(II)catalyzed hydrolysis, it was apparent that this reaction might have synthetic utility. Synthetic runs xere typically carried out homogeneously in alcoholaqueous acetonitrile (89% v/v) in the presence of catalytic amounts of Cu(I1) under aerobic conditions. This phosphorylating agent did not appear to be very sensitive to steric influences, phosphorylating 1” and 2O isomeric alcohols tith equal facility. However, the yield of alkyl phosphate is sensitive to the chemical nature of the alcohol and the solvent composition (Table 5). Comparing the product ratio (column 2) to the solvent ratio (column 1) yields a crude index of the selectivity of the phosphorylating agent. For methanol, the product ratio < the solvent ratio, that is, the phosphorylating agent is selective for methanol, whereas for 1-butanol the product ratio > the solvent ratio (i.e., phosphorylation of water is preferred). In the case of ethanol, product ratio = solvent ratio; there is no selectivity. This is typical of a metaphosphate species. In most cases the unfavorable selectivity may be overcome by increasing the alcohol concentration (column 3) _ High yields of lo~-_molec~~-~eight phosPbt% therefore, were readily obtainable with a minimum of worhup. The

MJ3TA-L ION CATALYSIS

OF THIOLPHOSPHATR

HYDROLYSIS

122

application of this synthetic procedure to more complex materials is hindered at present by the low solubility of the dicyclohexylammonium salt of cysteamine S-phosphate in most organic media. Acknowledgment

We should like to acknowledge the generous support of the National Institutes of Health (GM 37812). Footnotes 1Since Cu(I.I) forms stable insoluble complexes with ,&mercaptopropionic acid [41], precipitation has been attributed to chelate formation. No attempt was made to further characterize the precipitate. * Tank nitrogen (Wolf) wss passed through a scrubbing and drying train consisting of (1) two separate solutions (250 ml) of acidic vanadyl sulfate over Zn-Hg amalgam; (2) 250 ml 0.1 N NaOH; (3) 250 ml reduced lumiflavin (co. lO+ M, reduced by sodium hydrosuhite) in 0.1X NaOH; (4) a heated tube of copper turnings; (5) 250 ml concentrated HSO,; and (6) a U-tube containing Drier& and Ascarite. a It will be assumed throughout this paper that aqueous Cu(lI) is hexacoordinatecl. See Ref. 32. 4 At pH 4.1 cysteamine exists predominantly ss the monocationic species in acp?ou~ solution. The p& values for the amino and sulfhydryl moieties are 10.6 and 8.3, respectively. See Ref. 43, p_ 383. 6 The -S-Slinkage does not strongly interact with Cu(I1). Evidently the concentration of disuliide is insufllcient to lower k,,, via complexation through the amino moiety. 5 Assuming the dissociation constants for the phosphoryl and amino functions are mutually independent, the concentration of nonxwitterionic species is CLL. 10-r3 at pH 3, utilizing the data from Table 2. f The authors report a stable compound containing 1 mole of reagent per mole of metal ion and formulate the structure as follows.

2-

CU*+

CH2 \ co-0

S -

CH2

* By utilizing the kinetic data in Table 2 it can be shown that KF = %/KA

KB, where

Rp = KNHrRs-)~Cu(HzO)r*l [Cu(H:O)2+1 DE&-R-S-1 and KA and KB are macroscopic dissociation constants (~KA = 8.3, pKn = 10.6) for the sulfhydryl and amino groups, respectively, of 2-aminoethanethiol (Ref. 43, p. 383).

130

S. J. BENKOVIC

AND E. M. MIL?XlX

* The v&e reported is an apparent stability constant obtained under obvious oxporimenta~ difsxdty. mInorderfor& = k,endkD = kl’, the steadydaie derivation assumes that ka > > & a& k: >> kz’. = kl’&‘, idere k, = kx and kxt = ko, the Vah33 Of 4 ad kt’can rea&ly u since rr, = k& be calculated. An estimate for & (the formation constant for the mono compkx) is 3.8 X 105, taken from the data of Ref_ 42 for methyl 2-aminoethyl sulfide.

References 1. 2. 3. 4.

T. Wfe&md and R. Lambert, Clkz Bm-, 3% 2476 0956). D. C. Dittmer and V. 0. Siiverstein, J. Org. CM-, 2% 4766 (1961). S. Akerfcldt, Actn Chem. Shad, 17,319 (1963). T. C. Bece and S. J. Benkovic, Bioorgunic i%fechan+&?ns, Vol. n, chap. 5. w. A. Benjamin: New York (1966). 5. H-Neumann and L. Boross, Isrue J- Chpm., 3,92p (1965). 6. S. Almrfeldt, Soensk.Kent. Tidskt., 75,231 (1963). 7. S. Akerfeldt, Ada Ckern.Stand., 14, 1019 (1969). 8. F. Binkley, J_ Btil. Ckn., 181,317 (1949). 9. E. F. Korman, J. H. Shaper, 0. Cernichieri, and R. A. Smith, &CL B&x&n. Biophgs., IO!?, 284 (1965). 10. H. Neumann and S. Yaroslaaby, IsraeZJ- Chem., 4, 86~ (X%6).11. J. B. Martin and D. I&f.Doty, Anal Chem., 21, 965 (1949). 12. W. P. Jencks sod 11. Gilchrist, J. Am. Chem. Sot., 86,141O (1964). 13. C_ S. Hanes and F. A. Isherwood, _&&we, X%1107 (1949). 14, H. A_ Fla~chlra, EDTA Tiltatbns: An Introdudan to Thcorg and Pradce- Pergamon: New York (1959). 1% G. Schwarzenbach,&mplexomet+icTifraLions. Interscience: New York (1957). 16_ E_ J_ 31Flls,Jr., esd M_ T. Bogert, J. Am. Chen. Sot., 62, 1173 (1940). 27. S_ Akerfeldt, dd3 Ch. Stand, 16, 1897 (1962). 18. S. Akerfeldt, Ada Chem. Scan& 15,575 (1961). 19. S, Akrfeldt, Adu Chem. Scad, 13, 1479 (1959). 20. S. AkerfeIdt, A& Chem. Stand., 14, 1980 (1960). 21- S. AkerfeIdt, Ada Chtm. &and., 13,627 (1959). 22. S. J. Benkovic and P. A. Benkovic, J. Am. Chem. Sot., 88,55W (1966). 23. J_ Ho&e, And Chim. A&L, 4,23 (1950). 24. H. Erknmeyer et aL, Eelv. Chim. A&, 47, 792 (1964). 25. S. J. Benkotic and P- A. Benkovic, J- Am. Chem. Sot., 89,4714 (1967). 26. P- C. Haake and F. H. Westheimer, .T. Am. Chem. Sot., 83,1102 (1961). 27. P- D. Bayer, D. J. Graves, C. H. Suelter, and M. E. Dempsey, A&. Chrpm.,33, 1906 (1961). 28. D- C. Dittmer, 0. B. Ramsay, and R. E. SpaMing, J. Org. Chem., 28,1273 (x%3). 29. E. B. Herr, Jr-, and D. E. Koshland, Jr., Biochim. Bwphys. A&, ~5,219 (1957). 30. D- C. Dittmer and 0. B. Rammy, J. Org. Chem.. 28,126s (xx%)_ 31. FF. TV. Butcher ad F- H. yestheimer, J. Am. Chum. f&c., ?7,2420 (1955). 32. C. F. Cull& and D. L. Trimm, z& cus&ns Far&y Sot., No. 46, 144 (1968). 33. J. Peisach, P. Risen, and W. E. Blumberg, Tire B&&mof cqp~~. Academic: New York (19661. 34. S. Miktien~and T. H. Fife, .T. Am. Chem. SOL, 89,5320 (1967).

METAL

ION

CATALYSIS

OF THIOLPHOSPHATE

HYDROLYSIS

131

35. T. H. Fife and S. ~Milstien,J. Org. Ch., 34,4007 (1969). 36. F. A_ Cotton and G. Wiion, A&uncd Inorganic Chemistry, 2nd ed. Interscience: London (1966). 37. A. Hoh-oyd and J. E. Salmon, J. Chem. Sot., 269 (1956). 38. J. A. R. Genge and J. E. Salmon, J. C’hem.SOL, 256 (1957). 39. R. M. Smith and R. A. Alberty, J. Am. Chem. Sot., 78,2.376 (1956). 40. S. A. Khan and A. J. Kirby, J. Chem. Sot., B1172 (1970). 41. Q. Feman do and H. Freiir, J. Am. Chem. So-z., 80,4928 (1958). 42. E. Gonick, IV. C. Femelius, and B. E. Douglas, J. Am. Chem. Sot., 76,467l (1954). 43, “Stabiity Constants,” Spec. P&Z_ 17. The Chemical Society: London (1964). 44. M_ Tetas and J. bf. Lowenstein, Biochemistry, 2,350 (1963). 45. E. C. Knoblock and W. C. Purdy, Radialion Rex, 15,94 (1961). 46. M_ Eigen and R. G. Wilkins in “Mechanisms of Inorganic Reactions,” A&an. Chem. Ser. 49, Chap. 3. American Chemical Society: Wa&ington, DC. (1965). 47. S. J. Benkovic, “Hydrolysis of Inorganic Esters,” in Comprehen&ue Ciwtnical Kinetics (ed. by C. H. Banforci and C. F. Tipper), to be published. 48. A. J. Kirby and A. G. Varvoglis, J. Am. Chem. Sot., 69,415 (1967). 49. S. J. Benkovic and Kc.J. S&ray, Btichemtifzy, 7,4697 (1968). (ReceivedDecember 1370)