Mechanism of catalytic ascorbic acid oxidation system Cu2+–ascorbic acid–O2

Mechanism of Catalytic Ascorbic Acid Oxidation

System Cu2+-Ascorbic Acid-0%

E. V. SHTAMM, A. P. PURMAL, and YU. I. SKURLATOV Institute of Chemical Physics, Academy of Sciences of the USSR, Moscow 11 7334, USSR

Abstract

Analysis is made of reported results on the kinetics and mechanism of ascorbic acid oxidation with oxygen in the presence of cupric ions. The diversities due to methodological reasons are cleared up. A kinetic study of the mechanism of Cu2+ anaerobic reaction with ascorbic acid (DH2) is carried out. The true kinetic regularities of catalytic ascorbic acid oxidation with oxygen are established at 2.7 6 pH < 4,5 X 6 [Cu2+] d 10-3M, and

d [DHz] 6 10-2M, Q [Oz] Q 10-3M.

Wo = K~[CU~+][DH~][O~]~.~/[H+]

where ~1 (25OC) = 0.13 f 0.01 M-o.5.sec-1. The activation energy for this reaction is E1 = 22 f 1 kcal/mol. It is found by means of adding Cu+ acceptors (acetonitrile and ally1 alcohol) that the catalytic process is of a chain nature. The Cu+ ion generation at the interaction of the Cu2+ ion with ascorbic acid is the initiation step. The rate of the chain initiation at [Cu2+] 3 10-4M, [DH2] Q 1 0 3 4 , 2.5 < pH < 4, is

Wi = K ~ , ~ [ C U ~ + ] ~ [ D H ~ ] / [ H + ]

where K ~ J (25OC) = (1.8 f 0.3)M-1-sec-1, Ei,l = 31 f 2 kcal/mol. The reaction of the Cu+ ion with 0 2 is involved in a chain propagation, so that the rate of catalytic ascorbic acid oxidation for the system Cu2+-DHz-02 is

Wo = k1[Cu+l[OzI

where k l (25°C) = (5 f 0.5) X lo4 M-l-sec-'. The Cu+ ion and a species interacting with ascorbate are involved to quadratic chain termination. By means of photochemical and flow electron spin resonance methods we obtained data characteristic of the reactivities of ascorbic acid radicals and ruled out their importance for the catalytic chain process. A new type of chain mechanism of catalytic ascorbic acid oxidation with oxygen is proposed

(1) initiation 2Cu2+ + DH- - 2Cu+ + D + H+ (2) propagation (3) propagation CuOi + DH- - Cu+ + D + HOT

(4) termination CuOJ + Cu+ -+ 2Cu2+ + HzOz

Cu+ + 0 2 - CuO;

2H+

International Journal of Chemical Kinetics, Vol. XI, 461-494 (1979) 0 1979 John Wiley & Sons, Inc. 0538-8066/9/0011- 0461$01.00

462 SHTAMM, PURMAL, AND SKURLATOV

Introduction

The important part played by vitamin C in the processes of vital activity [l] accounts for the large number of papers on ascorbic acid (DH2) oxidation. The classical studies of Barron and co-workers [2] and Weissberger and co-workers [3] have shown convincingly that there is no direct interaction between DH2(DH-) and 0 2 . DH2 oxidation with molecular oxygen is a typically catalytic process which is very sensitive to microamounts of heavy-metal ions [4,5]. According to Barron and co-workers [2] the Cu2+ ions are the most effective DH2 oxidation catalysts. A t pH < 5 dehydroascorbic acid (D) appears to be the only end product [2]. Hydrogen peroxide is an intermediate [6,7].

The kinetic data on DH2 oxidation catalyzed by copper ions are very different. The orders with respect to the catalyst vary from 0.5 [8] to 1 [6, 7,9, lo], for ascorbic acid from 0.5 [7] to 1 [6,8,9,10], for H+ from -2 [6] to -0.5 [8], and for 0 2 from 0.4 [7] to 1 [3,9]. In certain cases the kinetic data are different under identical conditions. A dependence of the rate constant of DH2 oxidation on the initial ascorbic acid concentration, or a “drift” of the constant in the course of the reaction has been reported [7,

There is no widely accepted mechanism of ascorbic acid oxidation. A change in the copper-ion valence in the course of the reaction has been reported by Barron and co-workers [2], and then by the authors of [6,11-131. A relevant mechanism of alternating reduction-oxidation of copper ions was proposed. Neglecting secondary radical reactions, this mechanism can be written as

101.

-e- Cu2+ + DH-- Cu+ + D- + H+ D---+D

H+ +e-,H+ CU+ + 0 2 --+ Cu2+ + H02 HO2 ----+ H202

As it follows from the scheme, the oxidation of ascorbic acid occurs with the participation of only Cu2+ ions, and the role of oxygen consists in the reoxidation of emerging Cui- ions.

On the other hand, on the basis of the linear dependence between the catalytic reaction rate and the 0 2 pressure, Khan and Martell [9] and Ogata and co-workers [lo] propose a mechanism implying that DH2 oxidation occurs via a ternary metal-substrate-oxygen complex:

Cu2+ + DH- + CuDH+

0 2 + CuDH+ F= OzCuDH+ - C U ~ + + D- + HO2

It is assumed in this case that the metal ion performs a function of an electron transfer agent between the substrate and the oxidant and does not undergo the resulting valency change.

CATALYTIC ASCORBIC ACID OXIDATION 463

Neither of the mechanisms proposed is consistent with the ensemble of data available. Thus the data on the kinetics of catalytic DH2 oxidation must be reconsidered and the process mechanism subjected to analysis.

The purpose of this paper was to carry out a detailed kinetic analysis of various aspects of the system Cu2+-DH2-02. This analysis involves the study of the interaction of Cu2+ with ascorbic acid under anaerobic conditions; the study of the kinetic features of the catalytic reaction of DH2 oxidation with oxygen; the determination of the general mechanism of the catalytic process (in the present case the chain mechanism is in action); the characterization of the chain transfer agents and of the initiation and termination steps; quantitative kinetic studies of alternative mechanisms satisfying formal kinetic dependences; finally, the utilization of auxiliary techniques (in the present case the use of photochemical and flow electron spin resonance methods) for the simulation and quantitative description of separate steps of the process. As a result we propose a new type of chain ion-molecular mechanism embracing the general feature of all those proposed earlier. Tentative communications have been published before [14-161.

Experimental

Methodological Features

Analysis has shown the following sources of errors responsible for in- correct kinetic results.

(a) Most researchers erroneously assume a prion‘ that the reaction occurs in a “kinetic regime,” that is, that the oxygen concentration in solution in the course of DH2 oxidation is the same as 0 2 solubility. As a rule, the kinetic-regime condition is not fulfilled for the system Cu2+-DH2-02 [8,17]. Obviously, correct allowance for this factor would require control of the oxygen concentration in the course of the reaction.

(b) Various buffer additions containing complexing anions have been used for maintaining constant pH, and changes in the latter induced relevant changes in the catalyst composition [18]. Certain researchers studied the catalytic activity of copper ions in the presence of chloride ions, also neglecting the effects of the C1- complexing properties. To rule out these errors, we conducted the reaction in an acid medium in the absence of buffer or other complexing impurities, under conditions with indifferent changes in pH (<0.1 unit). The pH changes in the course of the reaction were investigated in parallel experiments.

(c) Most authors used a first-order rate constant as the kinetic characteristic, and oxidized DH2 to a great extent. The effects connected with a possible H202 contribution to ascorbic acid oxidation were disregarded [6,7,19]. In this connection we measured only initial rates of ascorbic acid oxidation, Wo.


(d) In certain cases, for instance [9,20], the spontaneous oxidation rate of DH2 in the absence of a catalyst was relatively high. The difference in rates in the presence and absence of a catalyst was taken accordingly as the rate characteristic. The high spontaneous oxidation rate was probably due to ions of heavy metal impurities. The formation of a nonadditive mixed catalytic system is possible in the presence of a catalyst. For instance, Fe3+ impurity ions have a marked effect on the catalysis of DH2 oxidation by Cu2+ ions [6]. A t times the reaction was conducted at high ionic strength, usually 0.1 [9], and this provided an additional source of impurities since the most careful purification of the initial reactants did not ensure complete absence of iron ions [2,3]. We have also observed DH2 autooxidation on impurity ions in the absence of a catalyst. Addition of 10-5M EDTA was found to completely eliminate the background reaction in ascorbic acid oxidation. In order to reduce errors in determining the “active” catdyst concentration ([Cu2+] = [Cu2+]0 - 1 0 - 5 ~ ) , most experiments were conducted at [Cu2+]0 > lOV4M at an intrinsic ionic strength usually not ex- ceeding 0.001. The “standard” solution contained 10-4M Cu2+, 10-5M EDTA, and 10-3M DH2 at pH = 3.4.l

(e) No paper gives a parallel kinetic investigation of the Cu2+-DHy02 system for 0 2 absorption and DH2 oxidation. We have used both the gasometric method and the sampling analysis for DH2. The DH2 concentration is, as a rule, determined by “slow” titration of ascorbic acid with 2,6-dichlorophenolindophenol, or by titration of dehydroascorbic acid [20]. We have used the sampling technique together with a spectrophotometric “rapid” determination of DH2 concentration. For this purpose, a “sample” of the solution was introduced into a 0.1M solution of buffer acetate (pH = 5.1) with excess EDTA. In this buffer ascorbic acid shows a maximum at 265 nm with an extinction coefficient of 12,500 M-l-cm-l. Dehydro- ascorbic acid does not absorb light at 265 nm.

The accuracy of spectrophotometric WO determination was f5%. (f) In discussing the mechanism of DH2 oxidation in the Cu2+-

DH2-02 system, assumptions were often made about the mechanism of interaction between Cu2+ and DH2, or about reactions of ascorbic acid radicals. In this connection, along with a kinetic study of the catalytic reactions, we have conducted a kinetic analysis of the Cu2+ reaction with DH2 under anaerobic conditions, and we have also studied the reactivity of ascorbic acid radicals in a simulated photochemical system, in which the radicals were generated under the action of HO2 radicals formed by pho- tolytic decomposition of H202. The contribution of free radicals to DH2 oxidation was estimated by electron spin resonance combined with the flow technique.

In the absence of copper ions the concentration of impurity iron ions decreases kFeEDTA

[EDTAJo times; however, in the presence of Cu2+ at [Cu2+]0 >> [EDTAIo it decreases ~F~EDTA[EDTA]~/~c,EDTA[CU~+]~ times; under typical conditions about lo5 times.


Gasometric Technique

A gasometric constant-pressure apparatus [21] was used for studying the kinetics of 0 2 consumption. The amount of absorbed oxygen AV, in cubic centimeters, was recalculated in terms of moles per liter of the solution volume V,, in milliliters, using the expression

where Po stands for the 0 2 pressure in torr in the measuring burette. A procedure for the control of the deviation of diluted oxygen concen-

tration from equilibrium in a gas-liquid chemical reaction has been de- veloped [22]. The following relation between the reaction rate W and the gas concentration C was found to hold for turbulent intermixing (that is, for intermixing frequencies of = 103 min-l):

(11) where W , stands for the initial rate of physical gas absorbance under the same conditions and C* for equilibrium gas concentration. Relationship (11) is applicable at WIW, d 0.7 [22]. Fulfillment of the inequality WIW, << 1 is a criterium for the so-called “kinetic regime.” The initial rate of physical 0 2 absorption and the oxygen solubility were measured before each kinetic experiment (see Fig. 1). The experimental C* value always coincided with the calculated C* = mPb,, where m(25”C) = 1.26 X 10-3M/atm and Pb, is the partial 0 2 pressure over the solution (PO = P, + Pb2, P, being the pressure of saturated water vapor over the solution). The accuracy of measuring the initial DH2 oxidation rate WO by oxygen consumption was f3%. The error in the W , measurement did not exceed 10%.

Sampling was made under identical conditions at V, = 20 ml. The W , value of 2 X 10-6Mlsec was measured gasometrically. As sampling induced changes in W,, the latter method appeared to be inadequate at WIW, > 0.1.

(1) AC(M) = 5.4 X IO-’POAVIV,

c = c* (1 - W/W,)

Anaerobic Technique

0 2 was removed from the solution either by bubbling with nitrogen or argon, or by degassing in vacuum. The residual 0 2 concentration was estimated from Cu+dipy, decoloration (A,,, = 435 nm, E = 4.5 X 103M-l. cm-l [23]) and was no higher than 10-6M. The deoxygenated solutions were poured out by means of hermetic syringes and vessels covered by a rubber membrane.

Cu+ in the Cu2+-DH2 system was detected by means of a,a’-dipyridyl. The kinetics of Cu+ generation was studied by quenching the reaction with excess EDTA and a,a’-dipyridyl in a cell with an optical length of 1 cm (see Fig. 1). The accuracy in determining the initial rate of Cu+ generation, Wi was 5%-10%. The value of Wi in optical density units (AD,,dAt)O was recalculated in molar units using the expression


timc, min

Figure 1. Curve I-Kinetics of oxygen consumption A V cm3 in the system Cu2+-DH2-02. [DHz] = 5 X 10-3M, [Cu2+]0 = 10-4M, pH = 3.4, [EDTA] = 10-5M, PO = 600 torr, V, = 8 ml, T = 25’C. Curve 2-Kinetics of physical absorption of oxygen during saturation of degassed reaction solution before the introduction of the catalyst. Curve 3-Kinetics of the Cu2+ ion reduction by DHz (as the optical density of the reaction solution at the absorption maximum of Cu+dipyZ, X = 430 nm). [DH& = 2 X 10-3M, [Cu2+]~ = IO-*M, pH = 3.45, T = 22OC (anaerobic conditions). Dashed lines measured initial rates of DH2 oxidation Wo = 5.8 X loe2 cm3/min = 3.9 X 10-6M/sec, physical 0 2 absorption W, = 0.35 cm3/min = 2.5 X 10-5M/sec, and Cu+ formation Wi = 2.3 X min-’ = 8.5 X 10-8M/sec.

Photochemical Technique

€ 4 2 0 2 was photolyzed with filtered X = 313 nm light a t [H202] = 10F2M or X = 365 nm at [H202] = 1 td 5 M . The light beam was calibrated by a ferric-oxalate actinometer [24]. The OH or HO2 photoinitiation rate Ri,ph was measured spectrophotometrically by means of paranitrosodimethylaniline or tetranitromethane [24], respectively. The DH2 oxidation kinetics was followed spectrophotometrically. The runs were multiply re- peated for various irradiation times or by the method of consecutive sampling from the same solution. In analyzing the “sample” for DH2, the H202 was removed by means of a catalase. The error in measuring Ri,ph and the initial rate of photochemical DH2 oxidation WO,ph does not exceed 10%. The experimental temperature was 25°C.

Electron Spin Resonance Combined with Flow Technique

Electron spin resonance spectra and the decay kinetics of ascorbic acid radicals were recorded by a radiospectrometer with a response of -2 X 1011 ~ m - ~ , combined with a flow system [25]. An automatic recording of


changes in the signal amplitude at the derivative maximum, starting from -10 msec after intermixing, was made by a two-coordinate recording device on a smooth change of the distance from the mixer to the cavity. The experiments using the flow technique were conducted with Shuvalov and Moravskii.

Reagents, Equipment

High-grade ascorbic acid and cupric nitrate Cu(N0&3Hz0 were purified by recrystallization from twice distilled water. The initial Cu2+ concentrations were determined by iodometric titration. Ally1 alcohol and acetonitrile were thoroughly purified. Nitrogen and argon of high purity contained 0 2 < 0.003 and < 0.001%, respectively. Doubly recrystallized sodium nitrate, purified tetranitromethane, and paranitrosodimethylaniline were also used. The a,a’-dipyridyl and Na2-ethylanediaminetetraacetate of the trademark “Reanal” as well as standard solutions of HN03, NaOH, crystalline catalase, sodium acetate, and acetic acid were used without additional purification. Hydrogen peroxide was purified by distillation under vacuum at 60°C. The concentration of H202 was determined by KMn04 titration.

The “pH-340” instrument was for measuring pH (fO.O1) . The solutions were stirred by a magnetic mixer MM-3 (intermixing frequency -lo3 min-I). The solutions and gasometric burettes were thermostatted by U-8 and U-10 thermostats (*O.l”C). The optical density of the solutions was measured by the spectrophotometer SP-4A. The absorption spectra of DH2 and Cu+dipy2 were recorded by the spectrophotometer “Specord.” 500-W and 1000-W lamps provided ultraviolet light for the photolysis of H202.

Results

Interaction of Cu2+ with DHz

As it will be clear from the subsequent discussion, the oxidation of ascorbic acid in the system Cu2+-DH2-02 follows a chain mechanism, an essential role being played by Cu+ ions formed at the initiation step as a result of the interaction of Cu2+ with DH2. Due to this fact, the presen- tation of the data on kinetic studies of the catalytic system per se is pre- ceded by the description of the results of studies of interaction of cupric ions with DH2 in the absence of oxygen [ E l . Under anaerobic conditions an equilibrium is attained in the system Cu2+DH2

K , (5) 2Cu2+ + DH2 F+ ~ C U + + D + 2H+

and the experimental value


[15] is consistent with the value log K , = -8.2 calculated from the redox potentials cpo(Cu2+lCu+) = 0.15 V [26] and cpo(DDH2) = 0.39 V [27, 281. The Cu2+-DH2 solution becomes turbid and with time a dark-brown de- posit appears on the wall, probably Cu20 [13] generated by the disproportionation of Cu+ ions

kd H2O (6) CU+ + CU+ + (CU’ + Cu2+) CUZO + 2H+

The rate constant k d FS 3 M-l-sec-l [15] was estimated from changes in “equilibrium” values of [Cu+Ie7 [DHz], a t high copper concentrations.

The kinetics of Cu+ formation was studied by the method of the initial rate under conditions for which the disproportionation of Cu+ ions is in- essential ( [Cu2+]0 < 10-3M) at excess DH2. The plots given below show the primary experimental results for Cu+ generation at 22°C in terms of the optical density of Cu+dipyz (see Experimental). Wi values can be recalculated into units of Mlsec in accordance with relationship (111).

Figure 2 shows the initial Cu+ generation rate in the system Cu2+-DH~ as a function of [Cu2+]0. The quadratic nature of this function indicates a simultaneous contribution of two Cu2+ ions to an elementary DH2 oxidation step. The experimental dependence of Wi on [Cu2+] can be lin- earized in the coordinates Wil[Cu2+] versus [Cu2+]. Within error limits,

Figure 2. Curve 1-Dependence of the initial rates of Cu+ formation Wi = (m.l3O/At)O min-’ during Cu2+ interaction with DH2 on the total concentration of cupric ions [Cu2+]0. Curve 2-Same dependence expressed as Wil[Cu2+] versus [Cu2+] ([Cu2+] = [Cu2+]o - 10-5M). Curve 3-Probable shape of the graph (see text). [DH& = 2 X lOP3M, pH = 3.45, T = 22°C (anaerobic conditions).


either a straight line starting at the origin or a straight line cutting a segment on the Y axis (solid line 2 and dashed line 3, respectively) can be drawn across the experimental points. The slope of the line characterizes the contribution of the second-order reaction with respect to [Cu2+] and the intersection point represents the contribution of the first-order reaction with respect to [Cu2+]. It is evident that at [Cu2+] 3 10-4M the relative share of the latter is small. A similar conclusion has been drawn by the authors of [29] which deals with the interaction of Cu2+ with ascorbate in an acetate buffer.

It follows from Figure 3 that the rate of Cu+ formation at 2.5 < pH < 4 is proportional to IH'1-I. Since in this pH range the hydrolysis of copper ions can be neglected and ascorbic acid exists mainly in nonionized form ( K , = (5.0 f 0.5) X lO+M [30]), it follows from the dependence of Wi on pH that a t pH < 4 an ascorbate ion takes part in the Cu2+ reduction.2

At pH remaining constant and at [DH2]o < 10-2M, the rate of Cu+ formation is proportional to [DH& [Fig. 4(a)]. Thus under the conditions of the main amount of catalysis experiments, that is, at [Cu2+]0 3 10-4M, [DH& d 10-2M, 2.5 < pH < 4 (see following section), the expression for the initiation rate for Cu+ can be roughly represented as

d [Cu'] [Cu2+] [DHz] wi=---- [H+l

- K i , l dt

2 4 6 8 10

fo-' [H y-', M -' Figure 3. (anaerobic conditions).

Wi as a function of [H+]-'. [DH& = 2 X 10-3M, [Cu2+]~ = 10-4M

When analyzing the pH dependence, we limited ourselves (at pH < 3.5) to the approximate equality [DHz] = [DH& instead of the equation [DH& = [DHz] + [DH-] + [CuDH+].


(b)

Figure 4. (a) Dependence of Wi on [DH& at [DHzlo d 2 X 1OW2M, [Cu2+]o = IO-hM, pH = 3.55, T = 22OC (anaerobic conditions). (b) Same dependence at [DH& > 2 X 10-2M. Dashed squares-Wi values extrapolated to zero ionic strength in accordance with the empirical relationship (V).

where ~ i , l (25'C) = 1.8 f 0.3 M-l~sec-l.~ Under this condition, the effective activation energy of Cu+ initiation is E ~ , J = 31 f 2 kcal/mol (see [15]).

In accordance with eq. (IV), it can be concluded that the reduction of Cu2+ occurs in the course of a pseudo-termolecular reaction of two Cu2+ ions and a DH- anion. It is natural to assume that this reaction involves an intermediate cupric-ascorbic complex with no intermediate formation of free radicals of ascorbic acid, that is, the cooperative action of two Cu2+ ions leads to the two-electron oxidation of DH2:

According to [29] a similar constant for 0.1M acetate buffer at 3OoC equals = 1.2 M-' sec-I.


(la) Cu2+ + CuDH+ k’.t 2Cu+ + D + H+

It follows from the experimental data in Figure 4(b) that the Cu+ initiation rate passes through a maximum at [DHZ], = 5 X 10-2M. However, at [DH2] > 0.1 M the growth of the Cu+ formation rate is again observed as [DHZ] increases. We have explained this growth by the increase in the intrinsic ionic strength of the reaction solution that is determined by the amount of NaOH introduced to the soIution for maintaining pH constant while the ascorbic acid concentration increases (pin = [NaOHIo = [DH-] = K,[DH2]$[H+]). Indeed Figure 5, illustrating the effect of NaN03 addition at [DH& = 0.3M, shows that the rate of Cu+ initiation growth with the increase in ionic strength p = pin + [NaN03]. In this case the graph log Wi versus 4 is represented by a straight line with the slope ap- proaching unity. Accordingly, we have extrapolated experimental values of Wi at high concentrations of ascorbic acid to zero ionic strength on the basis of the empirical equation

Extrapolated Wi values are given in Figure 4(b) as dashed squares. I t

* 0.6

* 0.4 ’ 0.2

0 3- ’ -0.2 $’ . -0.4

’ -0.6

’ -0.8

I 1

as LO 1.5 hfn NO,S, M

Figure 5. Curve 1- Wi as a function of the NaN03 amount. Curve 2-Anamor- phosis of this function as log Wi versus 6. [DH& = 0.3 M, [Cu2+]0 = 10-4M, pH = 3.55. The ionic strength was taken as jt = [NaOHIo + [NaN03], where [NaOHIo is the amount of alkali added to the DH2 solution.


can be seen from the dashed line in Figure 4(b) that the introduction of a correction for ionic strength at high [DHz] leads to the independence of the rate of Cu+ formation on [DH&

The data in Figures 4(b) and 5 prove the participation of both free Cu2+ ions and the cupric-ascorbic CuDH+ complex in the two-electron reaction. Indeed, the presence of a maximum in the Wi versus [DH2] graph testifies to the participation of a free Cu2+ ion and of a complex with ascorbic acid ( [Cu2+]~ - [Cu2+] + [CuDH+]) in the reaction, and the growth of Wi with the increase in ionic strength indicating the involvement of particles with the same charge sign, that is, Cu2+ and CuDH+, in the act of Cu+ forma- t i ~ n . ~

From the dashed line of Figure 2 and from the fact that the Wi values extrapolated to p - 0 are limited at high [DH-1, when the free ion Cu2+ concentration is small, [Cu2+]0 = [CuDH+] [Fig. 4(b), dashed line], it is concluded that the reaction of the first order on [Cu2+] and [DH-] can become important at low [Cu2+]:

(1b) Cu2+ + DH- CU+ + D- + H+

Direct measurements of the concentration of radicals in the system* Cu2+-DH2 have been conducted in our experiments by means of the flow electron spin resonance method [31]. A typical electron spin resonance spectrum of radicals of dehydroascorbic acid D- is recorded practically immediately after mixing (=15 msec) deaerated solutions of Cu2+ and DH2 at pH > 4 or in more acidic media in the presence of acetonitrile. The studies of the dependence of the “initial” amplitude of the signal on [Cu2+], [DHZ], and pH made it possible to characterize in detail the mechanism of formation of D- radicals and their further transformations [31]. It has been established, in particular, that reaction (lb) also involves the formation of intermediate cupric-ascorbic complexes, CuDH+ and CUD (at pH > 4), as copper ions participate not only in the generation but also in the termination of D-:

k 3

H+ (9) k3 = 3 X 104M-’.sec-l

(4a) D- + Cu+ + Cu2+ + DH-, k t , l = 108M-l-sec-’

D- + Cu2+ +Cu+ + D,

If, together with reaction (la), the reaction of decomposition of the cupric-ascorbic complex

( 1 C ) CuDH+ Cu+ + D- + H+

At pH > 4, a complex of Cu2+ with a dianion of ascorbic acid, CUD, can also interact with Cu2+ (see [29]).


is taken into account, then the maximum Wi value will be reached at

where

and

[CuDH+] [Cu2+] [DH-]

K, =

is the stability constant of the cupric-ascorbic complex. On the basis of the dependence of Wi on [Cu2+] in the presence of acetonitrile, ki,l/hi,3 was evaluated to be equal to = 2.5 X 104M-l (see [30] and Fig. 2).5 Substituting [DHz], = 5 X 10-2M and this value into the expression for [DH21m, we find that K, = 200M-1.6

Since the factor K ~ , J in expression (IV) is determined by the combination of the constants 2KaK,ki,l, substituting numerical values of Ka, K,, and K ~ , J we obtain ki,l = 1O2M-l.sec-l. Accordingly, we find with respect to ki,Jki,3 that ki,3 = 4 X sec-l. This value agrees with the rate of the single-electron reaction at high [DH2] on the assumption that [CuDH+] = [Cu2+]o [dashed line in Fig. 4(b)].

Thus the interaction between Cu2+ ions and DH2 follows both the mechanism of single-electron transfers and the cooperative pathway with the resulting two-electron oxidation of ascorbic acid. Relatively low values of the rate constants of single-electron transfers (ki,z = lM-’.sec-l, k 3 = 3 X 104M-l.sec-l) testify to a low efficiency of ascorbic acid and its radical as single-electron donors7 At the same time, a relatively high efficiency of the cooperative reaction testifies to the readiness of the two-electron reduction of ascorbate.

The quadratic character of the dependence of Wi on [Cu2+] in the presence of acetonitrile is retained; owing to the parallel increase of the rates of single- and two-electron reactions, their relative contributions can be separated with greater accuracy.

On the basis of electron spin resonance we found that the stability constant of the CuDH+ complex at an “intrinsic” ionic strength < 0.01 equals 210 f 40 M-’ (see (311). In [9] log K , (0.4OC) = 1.57 at an ionic strength of 0.1. In [29] the authors obtained log K , (25°C) = 2.32 in 0.1M KN03.

According to the evaluation by Shtamm [30] the redox potentials of (D- + H+)/DH- and D/D- are 0.7 V and -0.18 V, respectively.


P,, t= (a)

Figure 6. (a) Curve 1-Dependence of the initial rate of 0 2 consumption (AV/ At)o cm3/min on the pressure of oxygen PO in the measuring gas burette (atmosphere of pure oxygen). Curve 2-Dependence of the initial rate of oxidation of ascorbic acid WO = (-A[DHz]/At)o Mlsec on PO under the same conditions. Curve 3-Dependence of PO X (AV/At)o (this product is proportional to the molar rate of oxygen consumption) on (Pb,)0.5. Curve 4-Dependence of Wo on (Pb,)0.5, where P&= PO - P, is the 0 2 pressure over the solution and P, = 23 torr is the pressure of saturated water vapors at 25OC. (- A[Oz]/At)o = 1.36 X 10-6 (Pb, atm)0.5 Mlsec, (- A[DHz]/At)o = 1.4 X (Pb, atm)0.5 Mlsec. (b) Curve 1-Dependence of (AV/At )o cm3/min on the oxygen partial fraction in the measuring gas burette, p = Po,/Po. (The burette contains a mixture of nitrogen and oxygen, PO = PO, + Curve 2-Dependence of Wo on p. Curves 3,4-graphs of dependence of (AV/At)o and WO on q, where p’ is the 0 2 partial pressure over the solution, p’ = p(1 - P,/Po), PO = 745 torr. (- A[Oz]/At)o = 1.35 X (p’)0.5 Mlsec; (- A[DHz]/At)o = 1.40 X (p’)0.5 Mlsec. a,b-[DHz]o = 10-3M, [ C U ~ + ] ~ = 10-4M, pH = 3.5, T = 25OC, V , = 8 ml, W / W , < 0.1.

Kinetic Features of a Catalytic Reaction

Under anaerobic conditions the rate of DH2 oxidation with Cu2+ ions is low and equals 0.5 W; (see [15]), while in the presence of oxygen the rate of the reaction grows to a considerable extent and the oxidation of DH2 is accompanied by 0 2 consumption. It can be seen from reported data that the dependence of a catalytic reaction on oxygen concentration has not been sufficiently studied. We investigated this dependence over a Po2 range from 0 to 1 atm in pure oxygen [Fig. 6(a)], and in N2 + 0 2 mixtures [Fig. 6(b)], both with respect to 0 2 consumption and to DH2 oxidation. Ex- periments were conducted for C = C*, that is, W/W, d 0.1 (see Experi-


e

P (b)

Figure 6. (Continued from previous page)

mental). The four dependences obtained were consistent with each other.

It will be seen from Figure 6(b) that the reaction rate is proportional to the square root of the partial 0 2 amount over the solution p’ = p(1- P,/Po), where (3 = Po2/(Pon + PNJ = Po2/Po is the partial 0 2 amount in the gasometric burette and P, is the pressure of saturated water vapors. Similar conclusions can be drawn from Figure 6(a). The rates of both DHz oxidation and 0 2 consumption expressed in molar units [according to expression (I) Wo is proportional to the product Po X (AVlAt)] are seen to increase in proportion with [OZ]O.~.

Thus the reaction order with respect to oxygen is 0.5. The initial 0 2

consumption rate coincides with that of DH2 oxidation (see caption to Fig. 6). In other words, the initial stage of the catalytic reaction is the 0 2 reduction to HzO2, and hydrogen peroxide takes no part in DHz oxidation:

CU2f DH2 + 02-D + HZ02

This conclusion is consistent with the results of [7], showing that H202 accumulation in the system Cu2+--DH2-O2 at t - 0 occurs in the stoi- chiometric ratio D:H202 = 1:l.

A WO dependence on Po2 correlating with our results has been reported in [’i, 111 and also, quite recently, in [32]. Ogata and co-workers [ lo , 331 extrapolated the Po2 dependence of DH2 oxidation by a straight line that


1.0 2.0

2 4 6 8 i o 10'fWe. M

Figure 7. Curve 1-Dependence of the initial rate of 0 2 consumption (AVlAt)o cm3/min on [DHzlo (dashed line-effect of the formation of the cupric-ascorbate complex a t high [DHz]). Curves 2, 3-Dependence of the initial oxidation rate of ascorbic acid WO Mlsec on [DHzlo. Curves l', 2'-Without allowance for (VII). Curve 3-Under inapplicability conditions (11). [Cu2+]o = 10-4M, pH = 3.4, T = 25OC. Curve l-V, = 8 ml, PO = 600 torr, W , = 0.35 cm3/min. Curves 2,3-V, = 20 ml, air.

does not pass through the coordinate origin, and the intersection length comprising a noticeable fraction of the rate of aerobic reaction is ascribed [lo] to anaerobic oxidation of DH2. It follows from our data in Figure 6 and the data of other authors [7,11,18] who have been studying the dependence of the rate of oxidation of ascorbic acid on oxygen concentration in a broader range than [lo, 331, and also from the comparison of expressions (IV) and (VIII) (see below), that under typical conditions of the catalytic reaction, at [O,] = 10-3M, [Cu2+] = lO-*M, the rate of DH2 oxidation in the anaeriobic reaction (at Poz - 0) is tens of times lower than that of DH2 oxidation in 0 2 atmosphere.

Apparently Ogata and co-workers erroneously extrapolated by the linear section the branch of the root dependence of WO on Po2 at 0.2 d p d 1. Khan and Martell suggest [9] that the catalytic reaction rate is proportional to the oxygen concentration; however, this does not correlate with the


primary results given in the same paper since the straight line drawn through experimental points does not pass through the coordinate origin.

As it has been pointed out under Experimental, relationship (11) makes it possible to follow the decrease in 0 2 concentration in the reaction solution in the course of a chemical reaction involving oxygen consumption. The deviation of the oxygen concentration in the solution C from the oxygen solubility C* during the reaction will, evidently, be the stronger the higher the reaction rate. Consequently, in the studies of the rate of catalytic reaction as a function of [Cu2+], [DHz], pH, or temperature the correction for the derivation of C from C* will vary in each case. Since the reaction rate is proportional to C0.5, we arrive, by expressing C via C* in accordance with eq. (11), a t the following relationship between the measured (experimental) rate of 0 2 consumption Wo,, and the rate with respect to C*, that is,

where W , is the initial rate of the physical absorption of 0 2 measured independently.

With the help of relationships (11) and (VII) we have avoided errors re- lated to the dependence of oxygen concentration in the solution on the rate of reaction. As it can be seen from Figures 7-9, the linear character of true dependences of Wo on [DHz], [Cu2+], and [H+] becomes clear only when these relationships are taken into account. These figures also reveal that the lower the rate of physical absorption of 0 2 (in experiments with sampling in the course of reaction with large solution volume and relatively small gas-liquid interface area) the more important the corrections introduced into the measured values of the rate of catalytic reaction WO,, (dashed symbols). Figure 7 illustrates the case when the dependence of WO on DH2 was measured under the conditions of a certain inapplicability of relationship (11), that is, a t WIW, > 0.7 (curve 3). Here the reaction takes place mainly near the gasliquid interface [22]; however, a practically proportional dependence between W O , ~ and [DHz] is observed in the experiment. In this case the effective rate constant of the reaction proves to be the function of noncontrollable diffusion parameters.

In this way the analysis of the dependence of WO on Poz, [Cu2+], [DHz], and pH within the ranges 0.1 d Poz d 1 atm, 5 X 6 [DHz] d 10-2M,

d [Cu2+]0 d lO-3M, and 2.7 d pH < 4 gives the following kinetic expression for the initial rate of ascorbic acid oxidation:

(VIII)


Figure 8. Curve 1-(AV/At)n cm"/min. Curve 2-Wo Mlsec as a function of [CU~+]O. Curves l', 2'-Without allowance for (VII). [DH& = 10-3M, T = 25'C. 1-V, = 6 ml, pH = 3.35, PO = 600 torr, W , = 0.24 cm3/min. Curve 2-V, = 20 ml, pH = 3.5, air.

where ~1 (25OC) = (4.6 f 0.3) X atm-O%ec-l = 0.13 f 0.01M-0.5- sec-l.8

The activation energy for DH2 oxidation was found to be E = 22 f 1 kcal/mol from the temperature dependence of the rate of 0 2 consumption AVlAt (Fig. lo), taking into account the temperature dependence of the coefficient of solubility of oxygen m, pH, and P,, and correction (11) for the deviation of C from C*. The reported activation energies range over 15-25 kcal/mol [2, 7, 8, 101.

* In a recent paper [32] Jameson and Blackburn found a similar dependence by monitoring the variation of 0 2 concentration in the course of the reaction with the help of the oxygen ana- lyzer

where k' = 1.29 X 105M-3//2-min-1 (25OC, 0.1M KN03). According to our data, a t an inherent ionic strength of the solution smaller than 0.01 the constant k' = K~K; ' equals (1.56 f 0.25) X 105M-3/2.min-1.


A

10 '(H ?-', M -' Figure 9. Curve l-(AV/At)o cm"/min. 2-Wo Mlsec as a function of [H+]-l. Curves l', 2'-Without allowance for (VII). [Cu2+]~ = 10-4M, T = 25°C. 1-V, = 8 ml, [DH& = 10-3M, PO = 600 torr, W , = 0.35 cm3/min. Curve 2-V, = 20 ml, [DHzlo = 7.5 X 10-4M, air.

General Mechanism

Barron and co-workers [2] found for CO that Cu+ acceptors can have an essential effect on the DH2 oxidation rate in the system Cu2+-DHz-02. In order to establish the Cu+ role of the catalytic reaction mechanism, we thoroughly studied the effects of other Cu+ acceptors as well: acetonitrile (AN) and ally1 alcohol (AA).

Acetonitrile forms weak complexes with Cu2+ and complexes with Cu+ in a ratio of 2:l:

K,tl CU+ + 2 AN --+ CU+(AN)~, log K,, = 4.35 [34]

The redox potential of the Cu2+/Cu+ pair in acetonitrile is -1 V [35] , so that the Cu2+AN complex is a much stronger oxidizer than the aquoion. At the same time the large value of this redox potential shows that the constant of Cu2+ complex formation with acetonitrile is very small. Ap- parently, in the whole range of AN concentrations the fraction of copper ions bound into the cupric-acetonitrile complex is negligibly small com- pared to the concentration of free Cu2+ ions. Nevertheless, it has been revealed in experiments with AN added that even a t low concentrations of acetonitrile the rate of Cu+ formation in the interaction of Cu2+ with DH2 grows substantially (Fig. 11). In these experiments the orders of the anaerobic reaction with respect to [Cu2+], [DH2], and [H+] are retained [16,


3.2 3.3 64 3.5 3.6 10’TyK 4

Figure 10. Temperature dependence (AV/At)o cm3/min in Arrhenius coordinates. [DH& = 10-3M, [Cu2+]0 = 10-4M, pH (25OC) = 3.45, PO = 600 torr, V, = 8 ml; W , = 0.35 cm3/min is virtually temperature independent. Dashed line re- fers to measured values of (AV/At)o solid line plotted with allowance for (VII) and with correction for temperature dependence of C* = mPb, and pH as log(AV/At)O,, + 0.5 log m + 0.5 log (PO - P,) + pH.

Figure 11. Dependence of Cu+ generation rate Wi = (AD4so/At)o min-’ on the acetonitrile concentration. [DHzlo = 10-3M, [Cu2+]o = 10-4M, pH = 3.5, T = 22“C, anaerobic conditions.

30, 311, and the expression for the initial rate of Cu+ formation in the presence of acetonitrile can be written as


1 x- B

1.0

I >*

0.2

t - . -“a,. - .=. I

1 2 3

Figure 12. Curve 1-Dependence of the initial rate of oxygen consumption WO = (AV/At)o cm3/min in the system Cu2+-DHz-02 on the amount of acetonitrile added. Curve 2-Initial portion of this dependence. Curve 3-graph of ( WO,,/WO)~ versus [AN] on the initial portion of the dependence (broken line indicates the deviation from linearity because of the formation of the Cu+(AN)2 complex). Curve ~ - W O , ~ ~ [ A N ] ~ values on the portion of the decrease of WO with [AN]; the decrease of the values (dashed lines) a t low [AN] is due to incomplete transformation of Cu2+ into Cu+(AN)2 while the decrease at high [AN] results from the possible formation of Cu+(AN)3 [36]. ~ - W O , ~ ~ / W O (1 + b[AN])0.5M-0.5 versus Wo,,,[ANI2 M3/sec where b = 1.5 X lo2 M-l (dashed lines indicate the deviation of the graph from linearity as a result of binding of all copper into the form of the Cu+(AN)2 complex). [Cu2+]~ = 10-4M, [DH& = 10-3M, pH = 3.5, Pb, = 0.76 atm, V , = 8 rnl, T = 25OC, W/W, << 1.

where a = (1.8 f 0.2) X 102M-1.9 The mechanism of the promoting action of acetonitrile is insufficiently clear; however, the simultaneous enhance- ment of single-electron and two-electron pathways of the reaction of Cu2+ with DH2 (see [16,31]) makes it possible to put forward a suggestion that the accelerating effect of acetonitrile is due not to the complexing with the Cu2+ aquoion but to its interaction with the cupric-ascorbic complex which can be regarded as a “partial charge transfer” complex [30] favoring the transfer of electrons from DH- to Cu2+ in this complex.

In 116,301 we established that under the same conditions as in Figure 2, the presence of 5 X 10-2M AN the quadratic dependence of Wi against [Cu2+] is just the same as shown in Figure 2. Furthermore in [31] we showed that for the one-electron reaction the equation of type (IX) is fulfilled with the same value of the coefficient a.


Acetonitrile turned out to strongly change the surface properties of water so that even at low concentrations of this compound the rate of physical absorbance of 0 2 grows considerably. On the other hand, the solubility of oxygen in water-acetonitrile solution does not practically change with the increase of the AN concentration up to 5M. Due to this fact, in studies of the system Cu2+-DH2-02-AN the concentration of 0 2 in solution practically coincided with the equilibrium concentration (the condition WoIW, << 1 was always fulfilled).

With the increase in acetonitrile concentration, the rate of DH2 oxidation first grows, reaches its maximum at [AN] = 0.1M, and then falls practically to zero (Fig. 12). This dependence can be qualitatively interpreted as follows. At low concentrations acetonitrile affects only the rate of Cu+ initiation Wi, essentially without variation in catalyst composition (the concentration of cuprous-acetonitrile complexes is small). Hence the growth of the rate of oxygen consumption in DH2 oxidation testifies t o the relationship between the rate of catalytic reaction Wo and the rate of Cu+ initiation. With the increase in [AN] at the expense of Cu+ binding with AN, a progressively larger amount of copper ions transform into a non- reactive complex Cu+(AN):! ([Cu2+]0 = [Cu2+] + [Cu+(AN)2]). Since the concentration of this complex is proportional to [ANI2 while W ; is proportional to the first power of [AN], the combined action of these two factors leads to a maximum on the graph of Wo versus [AN]. Finally, at high [AN] the decrease in the rate of DH2 oxidation occurs as a result of the rapid transformation of practically all the catalyst into Cu+(AN)2. A t this stage the catalytic function is being performed by Cu+ ions which are in equilibrium with Cu+(AN)2.

The formation of the Cu+(AN)2 complex in the system Cu2+--DH2-O2 in the presence of acetonitrile has been proved in experiments with the addition of 2,9-dimethyl-l,lO-phenantroline, a specific complexing agent forming a very strong intensively colored complex with Cu+ (€455 = 7890 M-l-cm-l, see [23]) which, due to its high redox potential [26], is relatively stable against oxidation in air. Therefore the measurement of Cu+ concentration can be conducted directly at the moment of stoppage of the catalytic reaction by simultaneous introduction of EDTA and 2,9-dimethyl-1,lO-phenantroline into the reaction solution. In the absence of AN the steady-state concentration of Cu+ ions proved to be too small to be measured by this method, while with the increase in [AN], the transition of copper into the form of the cuprous-acetonitrile complex is observed.

A t the initial section of the dependence of WO on [AN] under the conditions when [Cu+(AN)2] << [Cu2+]o ([AN] < 2 X 1OV2M), the relation between the quadratic rate and the acetonitrile concentration is linear (Fig. 12):

(X) W$,, = W i (1 + b[AN]) where b = (1.5 f 0.1) X 102M-l.


Comparing coefficients a and b in expressions (IX) and (X), we can see that a = b, that is, the following relationship holds true:

--- - constant W,,, - Wi w b a n w:

Thus at constant concentrations of the catalyst, substrate, and oxidant, the rate of oxidation of ascorbic acid in the system Cu2+-DHz-02 is proportional to the square root of the rate of Cu+ initiation Wi:

(XII) WO = c (Wp.5

where c is, generally speaking, a function of the reagent concentrations. (According to kinetic expressions (IV) and (VIII), c is proportional to ([O,] [DHz]/[H+])0.5-see below.)

A formal analogy can be drawn between the activating effect of small AN additives (when the composition of the catalytic system does not practically change, at the same time, while the rate of Cu+ generation varies) and the supplementary initiation of chain transfer agents in usual radical chain reactions. In the latter case the dependence of the type (XII) relationship between the reaction rate and the rate of initiation of radicals serves as a proof of the chain mechanism of the reaction with quadratic chain termination. Likewise, in our case dependence (XII) testifies to the chain mechanism of DH2 oxidation with quadratic termination, the generation of Cu+ as a result of the Cu2+ interaction with DH2 being the initiation step.

As it is demonstrated by experiments with the addition of &,a'-dipyridyl [16,30] and 2,9-dimethyl-l,lQ-phenantroline, at [AN] 3 0.3M a practically instantaneous reduction of Cu2+ to Cu+ occurs so that in the course of the

-i mb.

1 2 3 4 5 6 i04[Gul,, M

Figure 13. Dependence of the initial rate of 02 consumption (AVlAt ) , cm3/min on the total concentration of copper ions [Cu2+]~ in the presence of 0.5M acetonitrile. [DH& = 10-3M, pH = 3.5, Pb, = 0.76 atm, T = 25'C, V, = 8 ml, W / W , << 1.


catalytic reaction copper exists in the form of cuprous-acetonitrile complexes inert toward 0 2 . Under these conditions the rate of the reaction is independent on pH (2.5 < pH < 4) and[DHz] (zero order with respect to the substrate concentration), directly proportional to the initial total concentration of copper ions (Fig. 13) and to the oxygen concentration (Fig. 14), and inversely proportional to the square of the acetonitrile concentration (Fig. 12):

(XIII)

where d = 2.2 f 0.2M-sec-l. The character of the dependence of WO on [Cu2+]0 and [AN] is evidence that at 0.3 Q [AN] Q 2.5M the copper ions are mostly in the form of a bis-acetonitrile complex Cu+ ([Cu2+]0 = [Cu+(AN)z]), and the rate of DH2 oxidation in this case proves to be proportional to the concentration of Cu+ ions:

[CU'] = [ CU+ (AN) 21 /Kan [AN] [CU'+] o/Kan [AN]

Hence the rate of oxidation of ascorbic acid is determined by the rate of Cu+ oxidation with oxygen:

(XIV) wo = k 1 [CU+l [Ozl where k l (25OC) = d K A N = (5 f 0.5) X 1O4M-l.sec-l.

Figure 14. Dependence of the initial rate of oxidation WO Mlsec of DH2 on the pressure of oxygen Pb, over the reaction solution in the presence of 0.5M acetonitrile. [DH& = 10-3M, Point 1-in air; Point 2-from results in Fig. 12. [cU*+lO = 1 0 - 4 ~ , PH = 3.5, T = 25oc, w/w, << I.


The k l value virtually coincides with that of Cu+ interaction with 0 2 ,

measured independently by Zuberbuhler [36] in the absence of DH2: k1 (2OOC) = 3.5 X lO*M-l-sec-l.

Since acetonitrile affects only the rate of Cu+ initiation, it can be assumed that the rate of DH2 oxidation is determined by the interaction between Cu+ and 0 2 , even in the absence of acetonitrile, that is, that relationship (XIV) for the rate of DH2 oxidation holds true for all other conditions. This assumption is confirmed by the results of the analysis of the effect of allyl alcohol on the rate of oxidation of ascorbic acid in the system Cu2+- DH2-02.

Ally1 alcohol does not interact with Cu2+ and forms stable complexes with Cu+ in the ratio 1:1, log Kaa = 4.7 f 0.2 [37]. Accordingly it does not affect the initial rate of Cu+ formation during Cu2+ interaction with DH2. A t the same time allyl alcohol affects the catalytic system Cu2+--DH2-O2 as an inhibitor of DH2 oxidation (Fig. 15). In this case, as it has been demonstrated with the help of a,cr'-dipyridyl[16], the degree of inhibition proves to be proportional to the share of the catalyst bound into the Cu+AA complex inactive toward 0 2 . This result testifies to the formation of Cu+ in the system Cu2+-DHz-02 also in the absence of acetonitrile.

The plot of ( Wg,aa)-l versus [AA] proves to be rectilinear (Fig. 15). On the basis of the equation [Cu2+]0 = [Cu2+Iaa + [Cu+AA], assuming that

Kaa[Cu+] [AA], we find that [ C ~ ~ + ] a a = Wo,aa[Cu2+]o/Wo, [C~+]aa = Wo,aa/k1[02], and [Cu+AA] =

(XV) (Wo,aa)-' = (W01-l + K ~ ~ [ A A ] / ~ ~ [ C U ~ + ] ~ [ ~ Z ]

From the slope of the graph in Figure 15 we found that KaJk1[C~~+]o[02] = (1.0 f 0.1) X 107M-2-sec-1. Thus on the basis of studies of the inhibition

0.2 OA 0.6 0.8

1% M Figure 15. Left-Dependence of WO = (AV/At)o cm?/min on the allyl alcohol addition. Right-This dependence in coordinates W;' versus [AA]. [DHzlo = 10-3M, [cu2+]o = 10-4M, pH = 3.5, Poz = 0.76 atm, V, = 8 ml, W , = 0.35 cm3/ min.


of the rate of oxidation of ascorbic acid and the value k l = 5 X 104M-l-sec-l, we determined the value of the stability constant of the Cu+AA complex, namely, Kaa = (3.6 f 0.8) X 104M-l. This value is in a good agreement with the figure (5 f 2) X 104M-l obtained in polarographic experiments [38]. This coincidence is a criterion of the applicability of relation (XIV) for the catalytic reaction in the absence of acetonitrile.

In a similar way a linear graph can be obtained in the coordinates Wo,an/Wo(l + b[AN])0.5 at the maximum for [AN] < 0.2 (Fig. 12), that is, under the condition when some amount of copper in the course of the catalytic reaction exists in the form of the Cu+(AN)2 complex, but the concentration of Cu+ is not determined by the dissociation equilibrium of the complex. In this case [Cu2+Ian = Wo,a,[Cu2+]~/W~(l + b[AN])0.5 = [Cu2+]o - Kan[Cu+]an[AN12 or, if expression (XIV) is used

The experimental value of the slope equal to (8 f 2) X 106M-3.sec is in fair agreement with the calculated KaJI21[Cu2+]o[02] = (6.3 f 0.7) X 106M-3-sec. This coincidence proves once again the validity of relationship (XIV) under the conditions when the catalyst exists mainly in the form of the Cu2+ aquoion.

Thus the rate of catalytic DH2 oxidation is controlled by that of Cu+ interaction with 0 2 , regardless of the fact that the rate of Cu+ generation Wi is considerably lower than that of the catalytic reaction, and the steady-state Cu+ concentration represents only a small fraction of the total catalyst concentration. These data serve as an unequivocal evidence of the chain mechanism of oxidation of ascorbic acid in the presence of copper ions, where Cu+ acts as a chain transfer agent and the Cu+ reaction with 0 2 is one of the stages of chain propagation.'O

On the basis of the data obtained, the kinetic nature of the chain transfer agents participating in the quadratic termination can be established. In- deed, in a steady-state chain reaction the rate of the chain termination Wt should definitely coincide with the rate of Cu+ initiation Wi. Therefore the kinetic expression for Wt can be written as follows, with relationships (IV), (VIII), and (XII) in mind:

lo In [32] a chain mechanism of the ascorbic acid oxidation for the system Cu2+-DH2-02 is also suggested. However, the role of the Cu+ ions is excluded from the reactions of chain propagation. It is suggested that the initiation rate is determined by the interaction of the dimer complex (CuDH);' with 0 2 , that is Wi = ki[C~~+]~[DH-]~[02] . In the opinion of the authors of [32] the chain carrier is the complex DCuHO2 that dissociates in the reaction of the chain propagation to Cu2+, HOz-, and D-. Radical D- interacts in the chaii propagation reaction with the oxygen dimer [(CuDH)2O2l2+. The chain termination step is in that case the recombination of DCuHOz. This mechanism corresponds to the formal kinetic data, but it contradicts our data, indicating the formation of cuprous ions during the catalytic reaction.


For a chain reaction with a sufficiently long chain the rates of all propagation reactions are the same and equal to the measured rate of catalytic reaction. Accordingly, the steady-state concentration of the chain transfer agents is proportional to Wo and inversely proportional to the concentration of the reagent which brings about the disappearance of the chain transfer agent of the propagation stage.

It follows from relationships (XIV) and (XVII) that the step of the chain termination involves the Cu+ ions and the chain carrier X interacting in the chain propagation step with the substrate anionic form (the concentration of X is proportional to the ratio Wol[DH-]):

(XVIII) Wt = 2kt[C~+][X]

where 2kt = Ka~i,1klk2/~?, k2 is the rate constant of the chain propagation involving X (Wo = kz[X] [DH-I). Substituting ~1 = 0.13 M-0.5.~e~-1, K ~ J

= 1.8M-hec-l, k l = 5 X 104M-’.sec-l, K, = 5 X 10-5M, we obtain the ratio of the rate constants for the chain propagation to the chain termination k21kt = 7.5 X Obviously, the chain carrier X is generated by the interaction of Cu+ with 0 2 . Thus the problem of the detailed catalytic reaction mechanism comes to that of the X nature and reactivity.

Simulated Photochemical System

Zuberbuhler studied the kinetics of Cu+ autooxidation [36] and suggested that the radicals HO2, 0; were intermediate in the reaction of Cu+ with 02. In this case DH2 oxidation in system Cu2+--DH2-O2 at pH < 4 would occur under the action of HO2 radicals with the generation of ascorbic acid radicals as intermediates:

k”z (3”) HO2 + DH2 + H202 + D- + H+

The H02 radicals were generated by photochemical decomposition of H202 in concentrations, such that OH radicals took no part in DH2 oxidation [38]. The rate of photochemical DH2 oxidation WO,ph was found to be practically independent of [H202], pH, and [DH2] both under anaerobic and aerobic conditions (Table I) and is determined by the relation

(XIX) WO,ph = 0*5Ri,ph

Thus the ascorbic acid radicals, probably D- [38], take no part in the


TABLE I. Photochemical oxidation of ascorbic acid with HOz radicals.

Number [DHz] (M) pH [HzOzI ( M ) WO,ph (M/SeC) WO,ph/Ri,ph

l a 5 x 10-4 2.65 5.25 0.9 0.45 f 0.05 2b 5 x 10-4 2.65 5.25 0.17 0.55 3 5 x 10-4 2.93 1.05 0.22 0.55 4b 5 x 10-4 2.93 1.05 0.03 0.5 5a 5 x 10-4 2.8 2.6 0.52 0.5 f 0.05 6 5 x 10-4 2.55 2.6 0.55 0.55 7 3 x 10-3 1.25 5.25 0.95 0.45 8 2 x 10-3 0.3 5.25 0.85 0.4

[HzOz] = 5.25 M ; Ri,ph (HOz) = (2 f 0.2) X a Experiments conducted both in anaerobic conditions and with solutions saturated with

Mlsec.

oxygen at 1 atm. Irradiation intensity 15% (in other cases 100%).

reactions involving 0 2 and H202, as suggested earlier [6,9,39], but decay be recombination (k, = 108M-l-sec-l [40-431). Marked stabilization of unpaired electrons by the tricarbonyl system [44-46] seems to account for the inertness of D- toward 0 2 , Hz02, and Cu2+ [31].

The conclusion about the low reactivity of D- radicals toward Cu2+ made by using the electron spin resonance flow method (as discussed earlier) was confirmed by experiments with photocatalytic oxidation of ascorbic acid in the system Cu2+-DH2--H202, yielding the estimate k3(Cu2+ + D-) << 106M-l.sec-l [38].

0.81 1 F

0.4 *a

0.6 3 c\I

t h e , mLL

Figure 16. Kinetics of photochemical DH2 oxidation at the photolysis of Hz02. [HzOz] = 5.25M, Ri,ph = 2 X 10-6M/sec. 1-[DH2]o = 5 X 10-4M, pH = 2.65; 2-[DHz]o = 2 X 10-3M, pH = 0.3. Optical density of the DHz solution at 265 nm was measured after 10- and 40-fold dilution of the “sample” in acetate buffer, respectively. Dashed line-graphical determination of [DHZ),, (see text).


In strong acid media at low [DHz] the HO2 radicals may recombine ( k i = 10GM-l-sec-l [47]), and this can result in an increasing reaction order on [DH2] in the course of ascorbic acid oxidation (Fig. 16). Then, instead of relationship (XIX), we have the relation (XX) Ri,ph = 2W0,ph -k 2k,”H02I2

From known k l , Ri,ph, and the [DHz], value at which the DH2 oxidation rate becomes two times lower, that is, Wj,ph = 0.25Ri,ph = 0.5k2”[HOzl- n[DH2]n, [HOz], = (Ri,~h/qki)’.~, using the data in Figure 16, we obtain the rate constant of HO2 interaction with DH2: 122’’ = (1.8 f 0.3) X 103M-l-sec-l.

It follows from [43,48] that the reactivities of oxidants (except OH [43]) toward DH2 and DH- differ by 4.5pKa, as consistent with Marcus’s theory [49]. Consequently, the rate constant of HO2 reaction with DH- can be estimated approximately as Iz2’ = 3 X 1O5M-l-sec-’.

Discussion The most important result of the kinetic study is, in our opinion, the

establishing of an interconnection between the catalytic reaction rate, the Cu+ generation rate, and the steady-state concentration of Cu+. Quali- tatively, the same follows from reported results.

Dekker and Dickinson 161 have found that the rate of DH2 oxidation with Cu2+ ions in the absence of 0 2 is 50-70 times lower than that of the catalytic reaction, and Mapson [ l l ] reported that the rate of Cu+ generation in- creased in the presence of C1-. In [5,6,50] it was found that small amounts of C1- ions considerably accelreated the catalytic reaction. Since the chloride ion is a weak complexing agent relative to Cu2+, but forms a strong bis-complex with Cu+ 1511, cpo = 0.487 V [52], an analogy can be drawn between the effect of the “initiating” additive AN and that of C1-. As with AN, there must be an optimum for [Cl-1, as at high [Cl-] the copper ions will be bound in the form of CuC1; complexes of low activity toward 0 2 [5, 521. Such an optimum was actually observed by Mapson [5] and Ogata and co-workers [lo].

It has been found [ll-131 that the higher the catalytic reaction rate, the sooner a Cu2O residue is formed. Since the rate of Cu20 generation as a result of Cu+ disproportionation increases with the steady-state Cu+ concentration, these results point to an interconnection between the catalytic reaction rate and the Cu+ concentration.

The reaction of Cu+ with 0 2 was found to yield a species contributing to DH- oxidation. Assuming, according to Zuberbuhler [36], that HOa radicals are formed in this reaction, the chain mechanism of DH2 oxidation will be

(1) chain initiation 2Cu2+ + DH- - 2Cu+ + D + H+


H+ CU+ + 0 2 --+ Cu2+ + HO2

(3’) propagation HO2 + DH- - D- + H202

(2’)

(9)

(4’) termination Cu+ + HO2 - Cu2+ + HO, This mechanism satisfies the kinetic data characteristic of the initiation

and of the overall catalytic process (see the first three sections of Results). A quantitative criterion of its consistence with experimental data is the comparison of the rate constants of reactions (3’) and (9) with values obtained independently (see the first and last section of Results).

The rate constant of the Cu+ interaction with HO2, k( = 2.9 X 10gM-l-sec-l, was measured by Kozlov and Berdnikov [53]. Consequently, knowing the ratio k2’/k( = 7.5 X (see the third section of Results), we obtain the rate constant of the HO2 interaction with DH-, a t which the reaction could contribute to the chain propagation k2’ = 2.2 X 1O7M-l-sec-l. Its difference from the estimate made from photochemical results (k2’ = 3 X 1O5M-l.sec-l) is beyond possible error. Moreover, even at such a value of k2’ it would be impossible to neglect the competing HOz radical reaction

HOz + Cu2+ - CU+ + 0 2 + H+

{ D- + Cu2+ -+ Cu+ + D

the rate constant of which is 3.4 X 1O7M-l.sec-l [53] . In terms of the proposed mechanism the recombination of ascorbic acid

radicals plays no part in the chain termination, that is, the inequality h,[D-I2 < k,”Cu+][HO2] must always be fulfilled. Since [D-3 = WO/ k3[Cu2+], we get

Substituting kl’ = 5 X 1O4M-l.sec-l, k2’/k( = 7.5 X k, = 108M-l-sec-l [40-431, and “standard” conditions (see Experimental), we obtain the estimate k3 >> 106M-l.sec-l, at which value it is possible, in principle, to accept the assumption that D- takes no part in the chain termination. This estimate explicitly contradicts those obtained by photocatalytic 1381 and electron spin resonance flow [31] experiments (123 N 3 X 104M-l-sec-l).

Thus neither HO2 nor the radical anion D- contribute to the DH:! oxidation chain mechanism. Consequently, accumulation of H202 in the system Cu2+-DH2-02 occurs without intermediate HO2 formation by two-electron reduction of 0 2 . Since the rate of 0 2 reduction coincides with the DH2 oxidation rate, and is controlled by the rate of Cu+ reaction with 02, it will be concluded that the interaction of Cu+ with 0 2 yields some species suffering a two-electron reaction with DH-, without the formation of a D- radical.

This statement was confirmed by experiments using an electron spin resonance flow apparatus. The reaction of Cu+ ( W ~ O - ~ M ) with 0 2


(-10-3M) was conducted in the presence of DH2 in 0.1M AN, simulating in this way the catalytic reaction in the presence of AN. The rate of DH2 oxidation coinciding under these conditions with that of the Cu+ + 0 2 reaction was -10-5M-sec-1. With the recombination decay of radicals, if they had been formed in the process of oxidation of ascorbic acid, the generation of -3 X 10-7M radicals would have been expected. However, it appeared that the D- signal was not recorded, though the apparatus response ensured the detection of -10-8M& radicals [31]. Under the same conditions mixtures of Cu2+ or (Cu2+ + Cu+) with DH2 in the presence of 0.1M AN exhibited a marked electron spin resonance signal from D- radicals, testifying to the predominantly recombinational termination of these radicals [31].

As shown in [54], autooxidation of the cup’-dipyridyl Cu+ complex in a weak acid or neutral involves an intermediate “partial charge transfer” complex. Similarity between kinetics of the Cu+dipyz and the aquoion Oxidation suggests that at pH > 2.5 the CuO: complex, rather than 0; or HO2 radicals, is formed as an intermediate in the case of the aquoion as well.

Assuming that X = CuO: (see the third section of Results) the DH2 chain oxidation in the system Cu2+-DH2-02 will be written as

(11) wi - c u t

(12)

(13)

c u + + 0 2 - cuo:

CuO;+ DH- - Cu+ + D + HOT 2H+

(14) CUO: + CU+ - 2Cu2+ + H202

The rate constant of the CUO; interaction with DH-, k2 FZ 107M-’.sec-l, can be estimated assuming that k t = 10gM-l-sec-l [54]. An intermediate ternary copper-ascorbate-oxygen complex that decays without changing the ion metal valence seems to be formed in reaction (13). Unlike the suggestion made in [9, lo], a reduced copper ion Cu+ enters this complex composition. Probably, simultaneously with the electron transfer by the metal ion-oxygen, substrate-metal ion sequence in the ternary complex, there also occure a transfer of the H atom from the substrate to the 0 2 molecules (via H20?):


The proposed new type of the chain ion-molecular mechanism of oxidation of ascorbic acid encompasses all the most essential features of the mechanisms discussed earlier. An inherent step of the reaction is the generation of a reduced metal ion; ascorbic acid oxidation occurs in the ternary metal-substrate-oxygen complex. Such multielectron oxidation processes probably simulate the action of Cu+-containing oxidases.

It will be interest to note that with a natural ascorbic acid oxidation catalyst, ascorbate oxidase, the reaction yields intermediate ascorbic acid radicals [40,41] that decay by recombination. Obviously, as in the case of cup’-dipyridyl copper complexes [55], steric hindrance interferes with two-electron (cooperative) oxidation of the ascorbate, so that the Cu2+ ions in the active center of the ascorbate oxidase independently interact with DH-. At the same time reoxidation of reduced ascorbate oxidase and Cu+dipyz complexes [54] by 0 2 can be cooperative [56].

The contribution from the reducing metal ion to the oxidation process seems to be essential for the 0 2 molecule activation by the generation of the “charge transfer’’ complex. In strong acid media (in the case of the Cu2+ aquoion, probably, at pH < 2 [36]) protons of the media [54] take part in a direct reduction of 0 2 to the HO2 radical. For this reason the chain DH2 oxidation involving no free radicals is, in principle, impossible in such media.

The part played by HzOz on later stages of catalytic DH2 oxidation is not yet clear. Our investigations [30,57-591 have shown that the mechanism of combined DH2 oxidation with 0 2 and H202 is made very complicated by the simultaneous occurrence of several chain processes.

Acknowledgment

The authors wish to thank Dr. Yu. N. Kozlov, Institute of Chemical Physics, Academy of Sciences of the USSR, Moskow, for valuable discussion and help in the assay of the system Cu2+-DH2-O~-H20~ by the photochemical method, Dr. V. F. Shuvalov and Dr. A. P. Moravskii, Institute of Chemical Physics, Academy of Science of the USSR, Chernogolovka, for performing the experiments of the flow electron spin resonance. They also wish to thank Dr. S. S. Yufit, Institute of Organic Chemistry, Academy of Science of the USSR, Moskow, for contributing the high-purity acetonitrile and ally1 alcohol.

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Received April 1,1976 Revised January 4,1977

Documents

Mechanism of catalytic ascorbic acid oxidation system Cu2+–ascorbic acid–O2