Kinetics and Mechanism of Oxidation of Leucine and Alanine by Ag(III) Complex in Alkaline Medium

Abstract

Kinetics and mechanism of oxidation of leucine and alanine by Ag(III) complex were studied spectrophotometrically in alkaline medium at constant ion strength. The reaction was in first order with respect to Ag(III) complex and amino acids (leucine, alanine). The second-order rate constant, 𝑘⁻, decreased with the increasing in [OH⁻] and [IO₄₋]. A plausible mechanism was proposed from the kinetics study, and the rate equations derived from mechanism can explain all experimental phenomena. The activation parameters were calculated at 298.2 K.

1. Introduction

Transition metals in a higher oxidation state can be stabilized by chelation with suitable polydentate ligands, such as diperiodatargentate(III) [1–4], [ethylenebis(biguanide)silver(III)] [5, 6]. These complexes were stable in a suitable aqueous solution. Ag(III) complexes were stabilized in alkaline medium by periodate or tellurate ions [7, 8]. These complexes such as Ag(III), Cu(III), and Ni(IV) complex can be used as an oxidation reagents in organic chemistry and analytical chemistry [9–11]. As a kind of oxidation reagents, those complexes have been used widely in kinetic study [12].

The kinetics and mechanism of oxidation of some amino acids have been studied by some researchers. However, the mechanism is different in the different reaction systems [13–23]. In order to perfect the mechanism of oxidation of amino acids, the kinetics and mechanism of oxidation of alanine and leucine by [Ag(HIO₆)₂]⁵⁻ have been studied in this paper.

2. Experimental

All the reagents used were of A.R. grade. All solutions were prepared with doubly distilled water. Solution of [Ag(HIO₆)₂]⁵⁻ (DPA) was prepared and standardized by the method reported earlier [7, 8]. Solution of DPA was always freshly prepared before used.

2.1. Apparatus and Kinetics Measurements and Product Analysis

All kinetics measurements were carried out as previously [12]. The way for product analysis was similar to previous study [24]. After completion of the reaction, the products of oxidation were identified as ammonia and the corresponding ketone by their characteristic spot test [25].

3. Result and Discussion

3.1. Kinetics

The pseudo-first-order rate constants 𝑘_obs were calculated by the method as previously [12].

𝑘_obs values increased with the increase in [AA], the plots of 𝑘_obs versus [AA] are straight line which pass the grid origin at different temperatures (Table 1), the order with respect to [AA] was unity. Thus, an overall second-order rate law described by (1) is established here, where 𝑘^ represents observed second-order rate constant: 𝑟=𝑘_obs[Ag(III)]=𝑘^[Ag(III)][AA].(1)𝑘^ values decreased with the increase in [IO₄₋]. The plot of 1/𝑘^ versus [IO₄₋] was a straight line (Figures 1 and 2). The values of 𝑘^ decreased with the increase in [OH⁻] and the plots of 1/𝑘^ versus 𝑓(OH)[OH⁻] were line (Figures 3 and 4).

Table 1 

𝑘_obs varied with different [AA] at different temperatures.

Figure 1 

The second-order rate constant, 𝑘, as a function of the concentration of IO4− at different temperatures for oxidation of leucine.

Figure 2 

The second-order rate constant, 𝑘, as a function of the concentration of IO4− at different temperatures for oxidation of alanine.

Figure 3 

The second-order rate constant, 𝑘, as a function of the concentration of OH− at different temperatures for oxidation of leucine.

Figure 4 

The second-order rate constant, 𝑘, as a function of the concentration of OH− at different temperatures for oxidation of alanine.

3.2. Reaction Mechanism

In aqueous periodate solution, (2)–(4) were detected and the corresponding equilibrium constants at 273.2 K were determined by Aveston [26] as follows: IO2IO₄₋₄₋+2OH⁻⇌H₂I₂O₁₀₄₋,log𝛽₁=15.05,(2)IO+₄₋OH⁻+H₂O⇌H₃IO₆₂₋,log𝛽₂=6.21,(3)+2OH⁻⇌H₂IO₆₃₋,log𝛽₃=8.67.(4)

The distribution of all species of periodate in alkaline solution can be calculated from (2)–(4). In the [OH⁻] range used in this work, the dimer and IO₄₋ species of periodate can be neglected. The main species of periodate are H₂IO₆₃₋ and H₃IO₆₂₋. Equation (5) can be obtained from (2), (3), (4) as follows: H²IO₆₃₋=𝛽³OH⁻²1+𝛽₂OH⁻+𝛽³OH⁻²IO⁴⁻_ex=𝑓OH⁻IO⁴⁻_ex,(5) where [IO₄₋]_ex represents the concentration of original overall periodate ion.

From the discussion above, the reaction mechanism was proposed as below: AgHIO⁶₂⁵⁻+3H₂O_𝐾1Ag⇌H₂AgH₂O₂HIO⁶⁻+H₂IO₆₃₋+OH⁻,(6)O₂HIO⁶⁻+RCHNH₂COO^−𝐾→AgH₂O₂HIO⁶RCHNH²COO⁻²⁻,(7) R means (CH₃)₂CHCH₂—for leucine and H—for alanine AgH²O₂HIO⁶RCHNH²COO⁻²⁻_fast[Ag(I→II)]A_totg(I)+RCOCOO⁻+H₂IO₆₃₋+H₂O+NH₃,=[Ag(OH⁻)₂H²IO₆)₂]_e⁵⁻+[Ag(OH⁻)₂HIO⁶)]_e³⁻.(8) Here “e” and “tot” mean the equilibrium concentration of the all kinds of Ag(III) complexes and the total concentration of all kinds of Ag(III) complexes, respectively. The reaction (7) is the rate-determining step, k means the rate-determining step constant, the rate equation can be expressed as below: −dAg(III)_T=d𝑘t_obs⋅Ag(III)_tot==𝑘^𝑘AKg₁H²O₂HIO⁶⁻⋅RCHNH²COO⁻[R]Ag(III)_tot𝑘K₁+H₂IO₆₃₋OH⁻,(9)_obs=𝑘K¹[R]K₁+H₂IO₆₃₋OH⁻.(10) Then substitute (5) for [H₂IO₆₃₋ in (10) is as follows: 𝑘_obs=𝑘K¹[R]𝑘K₁+𝑓OH⁻[IO⁴⁻_exOH⁻,(11)_=𝑘K¹K₁+𝑓OH⁻[IO⁴⁻_exOH⁻.(12) From (12), the plot of 1/𝑘^ versus 𝑓(OH)⋅[OH⁻] is a straight line, and 1/𝑘^ versus [IO₄₋]_ex is straight line too, which can explain all the experimental phenomena, and the activation parameters according to k are calculated [27] (see Table 2).

Table 2 

Thermodynamic activation parameters for the oxidation of leucine and alanine by Ag(III) complex in alkaline medium.

4. Conclusion

Through the kinetics study, we supposed a simple reaction mechanism, which can explain the reaction very well. Activation parameters and rate determining step constant were calculated from the rate law derived from the proposed reaction mechanism. The negative value of △𝑆^# suggested that intermediate complex is more in order than the reactants. The higher rate constant for the rate determining step indicates that the oxidation presumably occurs via innersphere mechanism. The pre-equilibrium constants, K₁, are calculated at 298.2 K, which are smaller values. This also shows that rate law from equilibrium hypothesis is suitable for the reaction mechanism.