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Oxidation Communications 37, No 3, 649–656 (2014) Radical processes in gas-phase oxidation

Kinetics of Slow Oxidation of Methane on the Surfaces of Different Inorganic Salts M. Tsotniashvili*, Z. Dzotsenidze, G. Bezarashvili, M. Katsitadze, M. Kekenadze ‘I. Javakhishvili’ Tbilisi State University, Tbilisi, Georgia E-mail: [email protected] ABSTRACT The paper is dedicated to the study of kinetics of high temperature slow oxidation of methane on the surfaces of different inorganic salts (sulphate, fluoride and sodium carbonate). Experiment was carried out at different temperatures – 783, 798 and 813 K and pressures 340, 360 and 380 mm Hg. In conditions of each test temperature is maintained almost constant. By graphic differentiation of the received kinetic curves maximum (steady-state) values of reaction rates were determined which were achieved in each series of test. The values of effective rate constant of degenerate branching of circuits kIII were calculated in conditions of the given test. The values of kIII received by us greatly differ from the data received by different authors. It is stated that the Arrhenius equations are unfit for adequate description of the given process. Hypothetically the cause is mainly heterogeneous nature of degenerate branchings of reaction circuits. Keywords: kinetics, methane, oxidation, reaction rate, activation energy, empirical model. AIMS AND BACKGROUND For reactions of slow oxidation of hydrocarbons in gas or liquid phase the behaviour in time with reaction chain degenerate branching is characteristic1–3. Such branching is realised with participation of reaction-able intermediate molecular products (aldehydes, hydroperoxides) and causes gradual increase of the number of reaction chains in the system4. Degenerate branching may be realised by homogeneous as well as heterogeneous way which quite complicates its experimental investigation and intercorrelation of kinetic data received by different authors in different conditions. Without dependable experimental and theoretical data on degenerate branching it is *

For correspondence.

649

impossible to build precise kinetic picture of hydrocarbon oxidation. Meanwhile such a task is quite urgent in scientific as well as practical view point; in particular, for studying complicated processes in internal combustion motors3, also with the view of industrial output of valuable intermediate products of hydrocarbon oxidation. In works5–7 the composition of products of slow oxidation of methane on different solid surfaces is investigated. EXPERIMENTAL In the presented work degenerate branching of reaction chains was studied on example of oxidation of simplest hydrocarbon – methane on different surfaces of inorganic salts (sulphate, fluoride and sodium carbonate). Tests were done on static vacuum device; oxidation process was performed in quartz reactor the internal surface of which was covered with the layer of the studied salt. Note that kinetics of the given process going on over the surfaces Na2SO4, NaF and Na2CO3 were almost unstudied earlier. The composition of the initial mixture is 40%CH4 + 60%O2 which corresponds to the basic stoichiometric equation 2CH4 + 3O2 = 2CO + 4H2O.

Tests were carried out at different temperatures (783, 798 and 813 K) and mixture pressures (340, 360 and 380 mm Hg). In conditions of each test temperature was maintained almost constant and the reaction run was tracked by the increase of general pressure of reaction mixture in reactor performed with mercury manometer. Methane consumption rate may be connected with the change of total pressure rise ∆P in time (Fig. 1) W(CH4) =

Fig. 1. ∆P variation in time

650

2

d(∆Ptotal)

RT

dt

.

(1)

RESULTS AND DISCUSSION By graphic differentiation of the received kinetic curves maximum values (steadystate) values of reaction rate which were achieved in each series of experiments were determined. The received results are presented in Tables 1–3. The highest values of rate are stated for the surface of NaF, and lowest for Na2CO3. As is seen from the Tables the increase of pressure or temperature causes the increase of maximum rate; most important is the increase of temperature. The above described process was studied by different authors for different surfaces8,9. The values of reaction rate obtained by us are close to the data obtained by Edgerton on the surface of potassium chloride9. Table 1. Rate values on the surface of sodium sulphate

Т (K) 783 798 813

P0 = 340 mm Hg   9.11 12.57 31.34

Wmax (mmol m–3 s–1) P0 = 360 mm Hg 11.72 16.75 38.71

P0 = 380 mm Hg 14.64 19.16 56.41

Table 2. Rate values on the surface of sodium fluoride

Т (K) 783 798 813

P0 = 340 mm Hg 20.03 31.98 38.04

Wmax (mmol m–3 s–1) P0 = 360 mm Hg 23.37 39.42 47.78

P0 = 380 mm Hg 30.01 47.87 57.14

Table 3. Rate values on the surface of sodium carbonate

Т (K) 783 798 813

Wmax (mmol m–3 s–1) P0 = 340 mm Hg   6.17   7.442 14.492

P0 = 360 mm Hg   7.992   9.052 17.098

P0 = 380 mm Hg 10.66 11.57 22.389

Empirical model reflecting experimental data should qualitatively describe the dependence of reaction rate on pressure P0 and temperature T. At first we have chosen the following equation: Wmax = A P0m e–E/RT.

(2)

In order to determine the value of m the experimental data were presented in coordinates ln P0 – ln W at fixed temperatures (Fig. 2). From the slopes of these straight lines we determined average value of m for each surface: for sodium sulphate m = 4.44, standard deviation S(m) = 0.38, for sodium fluoride m = 3.64, S(m) = 0.012; for sodium carbonate m = 4.26, S(m) = 0.32. With the aim to determine empirical energy of activation, temperature dependence of reaction rate was expressed in the Arrhenius 651

coordinates 1/T versus ln W. It appeared that in these coordinates curved lines are received instead of straight lines (Fig. 3). For mathematical description of function n W(T) we added to equation (2) the correction factor eb(T–Tmin) , where Tmin is minimum temperature of test in our experiment Tmin = 786 K, and b is empirical coefficient. Numerous calculations showed that value n =4 is in good correspondence to experimental data. As to parameter E, for its numerical values we chose experimental values 22.3 kcal/mol for sodium sulphate, 40.22 kcal/mol for sodium fluoride, 10.94 kcal/mol for sodium carbonate received on the basis of analysis of data at 2 temperatures: 783 and 798 K. With consideration of the above said we receive the following expression: for sodium sulphate: Wmax = A P04.4 eb(T–783)4 e–22300/RT;

(3)

Wmax = A P03.64 eb(T–783) e40220/RT;

(4)

Wmax = A P04.26 eb(T–783) e10940/RT.

(5)

for sodium fluoride: 4

for sodium carbonate: 4

Fig. 2. Dependence of rate on pressure on the surface of: Na2SO4 – a; NaF – b, and Na2CO3 – c 1 – T = 783 K; 2 – T = 798 K; 3 – T = 813 K

652

Fig. 3. Dependence of rate on temperature on the surface of: Na2SO4 – a; NaF – b, and Na2CO3 – c 1 – P0 = 340 mm Hg; 2 – P0 = 360 mm Hg; 3 – P0 = 380 mm Hg

By method of the least square10,11 the numerical values of b and lnA were determined: for sodium sulphate – b = 0.9 × 10–6 K–4; S(b) = 4.3 × 10–8 K–4 ln A = 9.109; S(ln A) = 0.020; A = 1.12 × 10–4; for sodium fluorine – b = –0.365 × 10–6 K–4 ; S(b) = 0.018 × 10–6 K–4; ln A = 7.62; S(ln A) = 0.083; A = 2.039 × 10–3; for sodium carbonate – b = –0.67 × 10–6 K–4 ; S(b) = 0.33 × 10–7 K–4; ln A = –15.97; S(ln A) = 0.015; A = 1.16 × 10–7. The finite form of empirical models of methane oxidation is as follows: for sodium sulphate surface: –6

(T–783)4

Wmax = 1.12 × 10–4 P04.4 e0.9×10

e–22300/RT mmol/(m3 s);

(6)

e–40220/RT mmol/(m3 s);

(7)

for sodium fluorine surface: –6

Wmax = 2038.6 P03.64 e–0.365×10

(T–783)4

653

for sodium carbonate surface: –6

Wmax = 1.16 × 10–7 P04.26 e0.67×10

(T–783)4

e–10937.4/RT mmol/(m3 s).

(8)

Curves drawn according to equations (6)–(8) are also presented in Fig. 3. It is seen that these curves are sufficiently close to experimental points. Mean-square deviation between experimental and designed values ln Wmax makes 2.2% which is quite a small value. This means that equations (6)–(8) with satisfactory precision describe experimental results. By data of Enikolopov and Semenov8,12 slow oxidation of methane is done mainly up to carbon monoxide: CH4 + O2 → CH3 + HO2

(0)

CH3 + O2 → H2CO + OH

(I)

OH + CH4 → H2O + CH3

(II)

OH + H2CO → H2O + HCO

(II-a)

H2CO + O2 → HCO + HO2

(III)

HCO + O2 → CO + HO2

(IV)

HO2 + CH4 → H2O2 + CH3

(V)

H2O + H2CO → H2O2 + HCO

(V-a)

CH3 → heterogeneous removal

(VI)

According to the presented diagram the step (III) is a degenerate branching, i.e. interaction of formaldehyde with molecular oxygen. According to Ref. 12, rate constant of degenerate branching can be approximately written in the following form: kIII ≈ Wmax/2ν [CH4]0[O2]0(kII′ kV′/kII kV)–1/2,

(9)

where kI, kV, kII′, kV′ are rate constants of the respective elemental stages, [CH4]0 and [O2]0 – initial molar concentrations of methane and oxygen in reaction mixture and ν – mean length of reaction chains in the system. This value in the given case can be expressed as follows12: ν = kI[O2]/aVI,

(10)

where kI is rate constant of bimolecular interaction of methyl radical with molecular oxygen and aVI – kinetic coefficient (effective rate constant) of heterogeneous removal of methyl radical on the reactor surface. For numerical value of kI the data from Ref. 13 were used while for value aVI the known formula of Frank-Kamenetski was used in diffusion approximation14: aVI ≈ 23.1 D/d12,

(11)

where D is the coefficient of diffusion of methyl radicals in gas mixture, and d – the diameter of reaction vessel. For calculation of numerical values of coefficient D the 654

method of Girshfelder15 was used. The final results of calculations are given in Fig. 4. It is evident that in accordance with temperature and pressure of reaction mixture, numerical values of ν are in the interval 170–270 which is quite acceptable for degenerate branched chain reaction of hydrocarbon oxidation.

Fig. 4. Relation of chain length to temperature: 1 – P0 =340 mm Hg; 2 – P0 = 260 mm Hg; 3 – P0 = 380 mm Hg

Numerical values of rate constants kII, kV, kII′, kV′, were adopted from Ref. 13. It appeared that kII′ kV′/kII kV ≈ 0.78 e15880/RT.

(12)

With consideration of the above said and numerical values of initial concentrations of reagents on the basis of expression (9) the values of effective rate constants of degenerate branching of kIII were calculated in conditions of the given experiment (Table 4). Table 4. Values of kIII on different surfaces (m3/(mol s))

Surface

Na2SO4 NaF Na2CO3

783   5.78 12.02 4.012

Т (K) 798   6.550 16.124   3.775

813 14.010 16.046   6.033

The received values of kIII greatly differ from data received by different authors16,17 (these data also differ from each other). In our opinion the cause of discrepancy of our data with some data of other authors essentially is the heterogeneous nature of degenerate branching of reaction chains at slow oxidation of methane. CONCLUSIONS Slow oxidation reaction of methane which goes on some inorganic salts surfaces are characterised with the non-Arrhenius kinetics. This means that temperature depend655

ence of reaction can not be described with the Arrhenius equation. The reason must be that reaction chains degenerated branching is going on the reaction vessel surface which essentially complicates general kinetic picture. For adequate description of reaction rate temperature dependence it is necessary to use complicated mathematical expressions. The presented work proposes a model which with satisfactory precision describes the dependence of oxidation steady-state rate on initial pressure and temperature of mixture. REFERENCES 1. N. N. SEMENOV: Chain Reactions. Moscow, Nauka, 1986. 535 p. 2. N. M. EMANUEL, G. E. ZAIKOV, V. A. KRITZMAN: Chain Reactions. Historical Aspect. Moscow, Nauka, 1989. 335 p. 3. B. LUIS, G. ELBE: Burning, Flame and Bursts in Gases. Mir, Moscow, 1968. 592 p. 4. V. Ya. SHTERN: Mechanism of Hydrocarbon Oxidation in Gas Phase. Ed. AS USSR, Moscow, 1960. 495 p. 5. A. ERDOHELYI, K. FODOR, F. SOLYMOSI: Partial Oxidation of Methane on Supported Potassium Molibdate. J of Catalysis, 166, 224 (1997). 6. CHAO YAHG, NANPING XU, JUN SHI: Experimental and Modeling Study on a Packed-bed Membrane Reactor for Partial Oxidation of Methane. Ind Eng Chem Res, 37, 2601 (1998). 7. F. B. PASSOS, E. R. OLIVEIRA, L. V. MATTON, F. B. NORONHA: Partial Oxidation of Methane to Synthesis Gas on Pt/CexZr1–xO2 Catalyst. Catalysis Today, 101, 23 (2005). 8. N. S. ENIKOLOPYAN: Kinetics and Mechanism of Methane Oxidation. In: 7th Symposium (Int) on Combustion. Butterworths, London, 1959, 157–164. 9. A. C. EGERTON, G. Y. MINKOFF, K. C. SA LOOJA: The Slow Oxidation of Methane. The Role of the Surface on the Course of the Oxidation of Methane. Combustion and Flame, 1 (1), 25052 (1957). 10. C. L. AKHNAZAROVA, V. V. KAFAROV: Experiment Optimization in Chemistry and Chemical Technology. High School, Moscow, 1978. 319 p. 11. R. L. MASON, K. F. GUNST, I. L. HESS: Statistical Design and Analysis of Experiments. Wiley Publications, New York, 2003. 760 p. 12. N. SEMENOV: On Some Problems of Chemical Kinetics and Reaction Ability. Publ. AS USSR, Moscow, 1958. 686 p. 13. A. A. MANTASHYAN, A. A. SARKISYAN, C. D. AHSENTIEV: Title??? Chem J Armenia, 51 (3, 4), 7 (1998). 14. D. A. FRANK-KAMENETSKI: Diffusion and Heat Transfer in Chemical Kinetics. Nauka, Moscow, 1987. 491 p. 15. J. GIRSHFELDER, Ch. KERTIS, R. BERD: Molecular Theory of Gasses and Liquids. Publ. Foreign Literature, Moscow, 1961. 929 p. 16. R. R. BOLDWIN, R. W. WOLKER: Problems and Progress in Hydrocarbon Oxidation. In: Proc. of the 14th Symposium (Int) on Combustion. Pittsburg, Pennsylvania, 1973, 241–255. 17. I. A. NALBANDYAN, G. A. SACHYAN, A. G. PHILYPOSYAN, A. B. NALBANDYAN: Kinetics and Mechanics of Formaldehyde Oxidation. Combust Flame, 22 (2), 153 (1974). Received 20 July 2011 Revised 11 October 2011

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Oxidation Communications 37, No 3, 657–661 (2014) Radical processes in liquid-phase oxidation

Study of the Influence of Some Additives on the Rate of Cumene Autoxidation in Liquid Phase. 1. Influence of Vitamins Bc and B1 N. M. Beylerian*, P. G. Minasyan, E. R. Saroukhanyan Yerevan State University (YSU), Manoogian Street, 0025 Yerevan, Republic of Armenia E-mail: bnorayr@ mail.ru; parandzemminasyan@ mail.ru ABSTRACT The influence of vitamins Bc (folic acid) and B1 (thiamine) on the rate of initiated with AIBN cumene autoxidation at 353 K by gasometric method in liquid phase is studied. The studied vitamins belong to the same class B. The obtained experimental data show that the influence of both vitamins on the kinetics of the reaction is different, though in both cases in their presence the reaction begins after some induction period (τ). At low concentrations vitamin Bc practically does not affect the reaction stationary rate (Rst). The dependence of τ on vitamin initial concentration may be described by a bell-shape curve. The addition of NaOH to the reaction mixture increases vitamin antioxidant influence. NaOH reacts with COOH groups which are in vitamin Bc (folic acid (FA)) molecule. It results in liberation of amino groups which are very reactive with respect to free radicals formed in the reaction medium. In the case of vitamin B1 (thiamine (TA)) with increase of its initial concentration, τ as well as Rst simultaneously decrease. All observed facts are discussed. Keywords: cumene autoxidation, liquid phase, vitamins Bc and B1, kinetics. AIMS AND BACKGROUND In the surrounding us Nature occur various processes. Among them chemical reactions have their special place. In many reactions free radicals play very important role. Free-radical reactions occur also in the living organisms. If free radicals action remains out of control they can provoke very undesirable consequences, particularly give rise to different diseases. Therefore it is indispensable, at first, to reveal their presence in the reaction system, and then to control their action. One of the ways to solve the cited problem is to use ecologically pure compounds. *

For correspondence.

657

It must be noted that in living organisms controller function is accomplished by vitamins1, enzymes. The conclusion is the following: to use compounds which have natural origin to control free radicals activity. For this purpose, the action of two vitamins – Bc and B1 is studied. As object of investigation the cumene (CU) autoxidation is chosen because it is a model radical-chain reaction. It is a model because its mechanism is very well established. In our experiments only the stationary run of the reaction is considered. After the stationary ryn begins degenerated chain branching. EXPERIMENTAL CU autoxidation rate was determined using a gasometric method. The setting is described in Ref. 2. The oxygen consumed during the reaction proceeding was determined by means of a manometer. The liquid height change (∆h) in the manometer was determined at the reaction run. It is expressed in mm s–1. In all experiments the same manometer was used. The distilled CU was purified as follows: the mixture of CU with fresh portions of ‘chemically pure’ H2SO4 was shaken several times until the added H2SO4 remained colourless. Then the CU was shaken with bidistilled water until the pH of the water became to the pH of the used water. The CU autoxidation was initiated with azobisisobutyronitrile (AIBN). AIBN was twice recrystallised from ethanol solution. Vitamin Bc (folic acid (FA)) was ‘chemically pure’. Vitamin B1 was provided by ‘Aldrich’ firm as thiamine. HCl salt (TA.HCl). Taking into consideration reagents solubility in water the influence of both vitamins was studied in emulsion stabilised by sodium dodecylsulphate (DDS). The DDS purity was > 99%. Bidistilled water was used. Water solutions of vitamins and chlorobenzene solution of CU were used. In all experiments VH2O/VC6H5Cl = 1:1.25 = const, [DDS]0 = 1.5 × 10–2 M. It must be noted that critical concentration of micelles (ccm) of DDS= 8.1×10–3 M (Ref. 3). RESULTS AND DISCUSSION KINETICS OF CU INITIATED AUTOXIDATION IN THE PRESENCE OF VITAMIN Bc (FA)

It must be noted that its mixture with vitamin C exhibits powerful antioxidant activity4,5. Rst dependence on [AIBN]0 and [CU]0 in the presence of FA. In all experiments [FA]0 = 1.0 × 10–4 M = const. At first the Rst = f([AIBN]0) was studied. In this case [CU]0 = 3.0 M = const and [AIBN]0 was changed in the range (2.5÷7.5) × 10–3 M. Then the Rst = f([CU]0) at [AIBN]0 = 5 × 10–3 M = const was studied. [CU]0 changed in the 658

range 2–4 M. The analysis of the obtained experimental data shows that: (i) in all cases the reaction begins after some induction period; (ii) Rst ~ [AIBN]01/2, and (iii) Rst ~ [CU]0. Therefore: Rst = k [AIBN]1/2 [CU]

(1)

which is the classical rate law of radical-chain reactions in the stationary region. This means that FA practically does not affect the considered reaction mechanism. The fact that the reaction begins after induction period (τ) shows that FA reacts with the primary free radicals formed in the reaction initiation step. Influence of [FA]0 on the reaction rate. In this case [AIBN]0 = 5 × 10–3 M = const, [CU]0 = 3 M = const, and [FA]0 changed in the range (0–2.0) × 10–4 M. The function τ = f([FA]0)0 graphically is presented in Fig. 1. 12

IJ PLQ 

10 8 6 4 2 0

0.0

0.5

1.0 1.5 [FA]0î4 0

2.0

Fig. 1. Dependence of τ on [FA]0 VH2O/VC6H5Cl = 1:1.25 = const, [CU]0 = 3 M (in C6H5Cl), [AIBN]0 = 5 × 10–3 M; [FA]0 = (0÷2) × 10–4 M; temperature 353 K

From this figure it follows: (i) FA effectively reacts with primary radicals formed in the chain initiation step, but it may not be considered as classical inhibitor; (ii) when [FA]0 > 1.5 × 10–4 M its inhibitory action decreases. At [FA]0 > 2.0 × 10–4 M, τ ≈ 0 and probably parallel with CU begins FA co-oxidation which favours CU oxidation. For this reason at [FA]0 ≈ 1.5 × 10–4 M τ has a maximum value. Rst dependence on [NaOH]0. The molecule of FA besides amino and OH groups contains two acidic COOH groups. So the intramolecular salt formation between amino and carboxylic groups is inevitable. It is very probable that FA antioxidant property depends to a higher extent on the presence of amino groups. To verify this hypothesis Rst dependence on [NaOH]0 is studied. The neutralisation of COOH group as a result of its reaction with NaOH, is taken into consideration. The experimental data show that with increase of the ratio α = [NaOH]0/[FA]0 > 1 Rst appropriately decreases. The effect is more sensible when α ≥ 2 (Ref. 6). From this fact it follows that the antioxidant property of vitamin Bc very propably depends on the presence of amino groups in its molecule. 659

KINETICS OF CU INITIATED AUTOXIDATION IN PRESENCE OF VITAMIN B1

In the presence of some cations Me(n) H2O2 oxidises luminol. The oxidation is accompanied by chemiluminescence. In the presence of TA the chemiluminescence intensity sensibly decreases7. This fact shows that vitamin B1 effectively reacts with HO• and HOO• free radicals. Now it is well established that free radicals, e. g. O2•, HO•, ROO• which may be formed also in living organisms result in different diseases. Vitamins present in foods favour to overcome their undesirable effects. Among vitamins more positive action shows TA (Ref. 8). TA contributes to the effective removal of oxidative stress9. At the same time TA is a good reductant10. Therefore it is doubtless to assume also that TA may act on the run of radical-chain reactions. To verify this assumption TA action on CU antooxidation kinetics is studied. The reaction conditions are similar as in the case with FA. The reactions are carried out in emulsions: VH2O/VC6H5Cl = 1:1.25 = const, [DDS]0 = 1.5 × 10–2 M, [CU]0 = 3 M, [AIBN]0 = 5 × 10–3 M, T = 353 K.

¨ h (mm)

Study of the effect of [TA.HCl]0 on the reaction kinerics. The kinetic curves are presented in Fig. 2. 18 16 14 12 10 8 6 4 2 0

1 2 3 4 5

0

5 10 15 20 25 30 35 40 45 50 t (min)

Fig. 2. Kinetic curves of CU initiated authoxidation in emulsion in the presence of TA.HCl [TA]0 × 104 (M) = 0 (1); 0.25 (2); 0.75 (3); 1.25 (4); 2.5 (5); T = 353 K

Kinetic curves of CU initiated autoxidation in emulsion in the presence of TA.HCl. Figure 2 shows that in the presence of TA.HCl the reaction begins after a certain induction period τ = f([TA.HCl]0). It is established that when [TA.HCl]0 ≥ 3 × 10–4 M the reaction is completely inhibited. The obtained data confirm the conclusion concerning the vitamin B1 activity as antioxidant. It follows also that TA.HCl decreases the value of Rst. These kinetic data may be explained assuming that: (i) TA reacts with high rate with primary free radicals which are being formed in the reaction initiation step. As a result of those reactions vitamin B1 is transformed into new compounds; (ii) those compounds may compete with CU to react with ROO• free radicals. Namely ROO• free radicals par660

ticipate in the reaction propagation step. If those products are more reactive than CU, this results in Rst decrease. CONCLUSIONS 1. With the help of gasometric method is studied the influence of vitamin Bc and B1 in emulsions and on the kinetics of initiated with AIBN cumene autoxidation at 353 K. 2. Vitamin Bc (folic acid) exhibits inhibitory action at 353 K and at concentrations < 1 × 10–4 M. It does not affect the stationary reaction rate. 3. At [FA]0 > 1 × 10–4 M cumene initiated autoxidation begins without induction period. Probabely occurs vitamins autoxidation parallel with cumene autoxidation. 4. In the presence of NaOH vitamin Bc autoxidation efficiency increases . This effect is more noticeable at [NaOH]0/[FA]0 ≥ 2. This result shows that the antioxidant activity of FA is due mainly to the presence of amino groups in its molecule. 5. Thiamine hydrochloride is a very efficient inhibitor for radical chain reactions. At 353 K and [TA.HCl]0 ≥ 3 × 10–4 M the initiated cumene autoxidation is completely inhibited. It decreases also the stationary rate of the reaction. REFERENCES 1. A. I. KOLOTOVA, E. R. GLOUSHANKOV: Vitamins. Chemistry, Biochemistry and Physiological Role. Leningrad, 1976 (in Russian). 2. N. M. EMANOUEL, E. T. DENISSOV, Z. K. MAIZOUSS: Chain Reactions of Hydrocarbons in Liquid Phase. Moscow, 1965 (in Russian). 3. A. A. ABRAMZON (Ed.): Surfactant Compounds. Handbook. Leningrad, 1979 (in Russian). 4. R. JOSHI, S. ADMIXART, B. S. PATRO, S. CHATTOPADHYAY, T. MUKHERJEE: Free Radical Scavenging Behavior of Folic Acid. Evidence for Possible Antioxidant Activity. Free Rad Biol Med, 30 (12), 1390 (2001). 5. E. BURGUIERES, P. McCUE, YOUNG IN KWON, K. SHETTY: Effect of Vitamin C and Folic Acid on Seed Vigour Response and Phenolic-linked Antioxidant Activity. Biores Technol, 98, 1393 (2007). 6. N. M. BEYLERIAN, E. R. SAROOKHANYAN, P. G. MINASYAN: Influence of Folic Acid on the Rate of Cumene Autoxidation in Emulsion. Sci Trans YSU, (3), 54 (2011). 7. DU JIAN XIU, LI. YINHUAN, LU JIORU: Flow Injection Chemiluminescence. Determination of Thiamine Based on Its Effect on the Luminol–H2O2 System. Talanta, 57, 661 (2002). 8. A. GLISZEZNSKA-SWIGLO: Antioxidant Activity of Watersoluble Vitamins in TEAC (Trolox Equivalent Antioxidant Capacity ) and the FRAP (Ferric Reducing Antioxidant Power) Assays. Food Chem, 96, 131 (2006). 9. D. NANDI, A. PATRA, D. SWARUP: Effect of Cysteine, Methionine, Ascorbic Acid and Thiamine on Arsenic-induced Oxidative Stress and Biochemical Alterations in Rats. Toxicology, 211, 26 (2005). 10. K. N. MOHANA, K. R. RAMYA: Ruthenium (III)-catalyzed Oxidative Cleavage of Thiamine hydrochloride with N-bromosuccinimide in Presence of Hydrochloric Acid Medium. A Kinetic and Mechanistic Approach. J Mol Catal A – Chem, 302, 80 (2009). Received 13 March 2014 Revised 30 April 2014

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Oxidation Communications 37, No 3, 662–668 (2014) Oxidation in the presence of Cr- and Mn-containing oxidants

Kinetics of Oxidation of Glycine by Nicotinium Dichromate in Aqueous Medium in Presence of Perchloric Acid K. Vivekanandan*, R. Lakshmi Narayanan Post-graduate and Research Department of Chemistry, National College, 620 001 Tiruchirappalli, Tamil Nadu, India E-mail: [email protected] ABSTRACT The kinetics of oxidation of amino acid namely glycine by nicotinium dichromate in aqueous medium in the presence of perchloric acid was studied. The reaction was found to be first order with respect to oxidant, fractional order with respect to substrate, and second order with respect to perchloric acid. Amino acid was converted to the corresponding carbonyl compound. The increase in ionic strength had negligible effect on the rate. No polymerisation was observed with acrylonitrile. The addition of Mn2+ ion had a noticeable catalytic effect on the reaction rate. The reaction rates were determined at four different temperatures and the thermodynamic parameters were evaluated. Keywords: kinetics, oxidation, glycine, nicotinium dichromare, mechanism. AIMS AND BACKGROUND Natural amino acids may be defined as the constituents of natural proteins1. They find a number of applications in metabolism, microbiology, nutrition, pharmaceuticals, biological research, fortification of foods and feeds2. Oxidation of amino acids is of great importance both from a chemical point of view and its bearing on the mechanism of amino acids metabolism. Generally the amino and carboxyl functional groups in RCH(NH2)COOH undergo chemical transformations while the hydrocarbon moiety remains inact. This property is attributed to the higher reactivity of the amino group compared to alkyl group. Oxidation of amino acids by different oxidants were reported3–17. Nicotinium dichromate (NDC) is a stable, non-hygroscopic, mild, efficient and selective oxidising reagent in synthetic organic chemistry18. A review of literature *

For correspondence.

662

shows that kinetic studies using this reagent are meager19–26. There is no significant data on the NDC oxidation of glycine. Hence, the kinetics of oxidation of glycine by nicotinium dichromate has been studied in aqueous medium in the presence of perchloric acid. A plausible mechanism has been proposed. EXPERIMENTAL Materials. Nicotinium dichromate (NDC) was prepared by the reported method18 and its purity was checked by an iodometric estimation of the amount of Cr(VI) present in it, and the aqueous stock solution of NDC was prepared, standardized and preserved in dark coloured bottles. Chromatographically pure amino acid – glycine (LOBA) was used. Perchloric acid (AnalaR) was used as a source of hydrogen ions. All other reagents used were of analytical grade. Doubly distilled water was used throughout the study. Kinetic measurements. The reactions were carried out under pseudo-first order conditions by maintaining a large excess of the amino acid over NDC. The reactions were followed at constant temperature (313 K). Kinetic studies have been carried out by colorimetric method using photoelectric colorimeter. The rate constants were evaluated from the linear plot of lg (absorbance) versus time by the least square method and were reproducible within ±3% . Product analysis. The main products of the oxidation of amino acids were the corresponding aldehydes, ammonia and carbon dioxide. The presence of ammonium ions in the reaction mixture was detected by the test with p-nitrobenzene diazonium chloride27. The presence of carbon dioxide in the product was identified by lime water test28. The aldehyde was determined as its 2,4-dinitrophenylhydrazone. In a typical experiment, glycine (0.05 M) and NDC (0.01 M) were made up to 100 ml in 0.5 M perchloric acid. The reaction mixture was allowed to stand for 12 h in the dark to ensure the completion of the reaction. Then it was treated with an excess (250 ml) of a saturated solution of 2,4-dinitrophenylhydrazine in 2 M HCl and kept in a refrigerator for 10 h. The precipitated DNP was filtered off, dried, weighed, recrystallised from ethanol and weighed again. The yield was 85%. The DNP was found to be identical (m.p. and mixed m. p.) with the DNP of formaldehyde. Stoichiometry. The stoichiometry of glycine – NDC oxidation was determined by equilibrating varying ratios of [NDC] to [gly] in the presence of HClO4 at room temperature. After the reaction was complete, the excess of NDC was determined iodometrically, and indicated 1:1 stoichiometry. The observed stoichiometry may be represented as follows: NDC

NH2CH2COOH ––––––→ HCHO + CO2 + NH3 HClO4

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RESULTS AND DISCUSSION Effect of NDC. The concentration of NDC was varied in the range 6.0 × 10–3 to 14.0 × 10–3 mol dm–3 and keeping all other reactant concentrations constant and the rates were measured. The near constancy in the value of k1 irrespective of the concentration of the NDC confirms the first order dependence on NDC (Table 1). Table 1. Rate constants for the oxidation of glycine by NDC at 313 K in aqueous medium in the presence of perchloric acid

[NDC] × 103 (mol dm–3) 6.0 8.0 10.0 12.0 14.0 6.0 6.0 6.0 6.0 6.0 6.0 6.0 6.0 6.0 6.0

[ glycine ]× 102 (mol dm–3) 2.0 2.0 2.0 2.0 2.0 2.0 3.0 4.0 5.0 6.0 2.0 2.0 2.0 2.0 2.0

[ H+] × 102 (mol dm–3) 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 7.5 10.0 12.5 15.0

k1 ×104 (s–1) 2.39 2.37 2.39 2.30 2.38 2.39 2.80 3.34 3.67 4.02 2.39 5.76 12.13 16.30 20.89

Effect of glycine. The concentration of glycine was varied in the range from 2.0 × 10–2 to 6.0 × 10–2 mol dm–3 at 313 K and keeping all other reactant concentrations constant and the rates were measured. The rate constant increased with increase of [glycine] (Table 1). The order of the reaction with respect to [glycine] was found to be fractional as evidenced by the slope (0.48) of the plot of lg k1 versus lg [glycine]. A plot of 1/k1 versus 1/[glycine] (Fig. 1) shows an intercept on the rate axis, indicating the Michaelis–Menten dependence on the concentration of the glycine.

Fig. 1. Plot of 1/k1 versus 1/[S]

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Effect of perchloric acid. The concentration of glycine was varied within the range 5.0 × 10–2 – 15.0 × 10–2 mol dm–3 at 313 K and keeping all other reactant concentrations constant and the rates were measured. The rate constant increased with increase of [perchloric acid] (Table 1). The plot of lg k1 versus lg [H+] (Fig. 2) gave a straight line with a slope (2.018). It has been observed that the order in [perchloric acid] is two.

Fig. 2. Plot of lg k1 versus lg [H+]

Effect of ionic strength. There was no appreciable change in the rate with a change in ionic strength of the medium, affected by sodium perchlorate20 (Table 2). Effect of MnSO4. The addition of Mn2+ ions retared the reaction rate showing the two electron transfer process in the reactions29 (Table 2). Table 2. Effect of NaClO4 and MnSO4 on the oxidation of glycine by NDC at 313 K [glycine ] = 2.0 × 10–2 mol dm–3; [ H+] = 5.0 × 10–2 mol dm–3; [ NDC ] = 6.0 × 10–3mol dm–3

[NaClO4]× 103 (mol dm–3) 0.0 1.0 2.0 3.0 4.0 0.0 0.0 0.0 0.0 0.0

[MnSO4]× 103 (mol dm–3) 0.0 0.0 0.0 0.0 0.0 0.0 1.0 2.0 3.0 4.0

k1 ×104 (s–1) 2.39 2.38 2.39 2.36 2.35 2.39 2.30 2.10 2.00 1.90

Test for free radicals. No polymerisation is observed with acrylonitrile. The reaction mixture when allowed to stand with acrylonitrile does not induce polymerisation suggesting the absence of free radical mechanism. Effect of temperature. The rate of the oxidation of glycine was determined at different temperatures (303–333 K) and the activation parameters were evaluated (Table 3).

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Table 3.Thermodynamic and activation parameters for the oxidation of amino acids by NDC [S] = 2.0 × 10–2 mol dm–3; [H+] = 5.0 × 10–2 mol dm–3; [O] = 6.0 × 10–3 (mol dm–3)

Substrate

k1×104 (s–1)

∆S* ∆G* ∆H* –1 –1 –1 –1 303 K 313 K 323 K 333 K (kJ mol ) (J K mol ) (kJ mol ) Glycine 0.8 2.39 4.83 11.07 69.35 –61.69 88.65

Ea (kJ mol–1) 71.95

Mechanism and rate law. It is clear that the reaction has first order with respect to NDC, fractional order with respect to glycine and second order with respect to H+. The ionic strength of the medium has a negligible effect on the reaction rate, the added Mn2+ had a catalytic effect on the reaction rate and the absence of free radicals during the course of the reaction. The observed data can be explained on the basis of the following mechanism. NDC + H+ S + H+

K1 K2

NDCH+

(1)

SH+

(2)

k3

NDCH+ + SH+ ––––→ [complex]

(3)

slow k4

     [complex] ––––→ HCHO + Cr(IV) fast

(4)

The above mechanism leads to the following rate law: rate = k3 [comlex] =

=

K2k3 [SH+][H+]

(5)

1 + K2 [S] K1K2k3 [NDC][S][H+]2 1 + K2 [S]

.

(6)

Equation (6) accounts for first order dependence on [NDC], second order dependence on [HClO4] and fractional order dependence on [glycine]. Equation (4) indicates the involvement of a two-electrons reduction of Cr(VI) to Cr(IV) which is confirmed by the decrease in the rate constant with the addition of MnSO4 (Table 2).

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CONCLUSIONS The oxidation of glycine by NDC in aqueous medium in the presence of perchloric acid leads to the formation of a complex giving formaldehyde as the final product. The reaction follows pseudo-first order kinetics. Increase in temperature increases the rate of the reaction. The activation parameters are evaluated from the study of oxidation at different temperatures. The mechanism proposed for this oxidation kinetics is in accordance with the observed kinetic facts. The negative sign of the entropy change suggests that the transition state is more orderly when compared with the reactants. ACKNOWLEDGEMENTS The authors thank the Management, National College, Tiruchirappalli, for the facilities provided. REFERENCES 1. KIRK-OTHMER: Encyclopedia of Chemical Technology. Interscience, New York, 1963. 2. PUTTASWAMY, NIRMALA VAZ: Kinetics of Oxidation of Acidic Amino Acids by Sodium Nbromobenzenesulphonamide in Acid Medium. Indian Acad Sci (Chem Sci), 113 (4), 325 (2001). 3. G. IONITA, V. Em. SAHINI, Gh. SEMENESCU, P. IONITA: Kinetics of Oxidation of Amino Acids by Some Free Stable Hydrazyl Radicals. Acta Chim Slov, 47, 111 (2000). 4. S. MEENAKSHISUNDARAM, R. VINOTHINI: Kinetics and Mechanism of Oxidation of Methionine by Chromium(VI): Edta Catalysis. Croat Chem Acta, 76 (1), 75 (2003). 5. H. S. YATHIRAJAN, Ch. R. RAJU, K. N. MOHANA, Sh. SHASHIKANTH, P. NAGARAJA: Kinetics and Mechanism of Oxidation of L-isoleucine and L-ornithine Hydrochloride by Sodium N-bromobenzenesulphonamide in Perchloric Acid Medium. Turk J Chem, 27, 571 (2003). 6. K. VIVEKANANDAN: Oxidative Decarboxylation and Deamination of Proline, Histidine, Arginine, Lysine and Tyrosine by N-chloronicotinamide in Aqueous Acetic Acid Medium. A Kinetic Study. Oxid Commun, 27 (1), 195 (2004). 7. R. SHUKLA, P. K. SHARMA, K. K. BANERJI: Kinetics and Mechanism of the Oxidation of Some Neutral and Acidic α-amino Acids by Tetrabutylammoniumtribromide. J Chem Sci, 116 (2), 101 (2004). 8. N. A. MOHAMED FAROOK, G. A. SEYED DAMEEM, A. MURUGESAN, M. KANAGARAJ: Kinetics of Oxidation of some Essential Amino Acids by N-chlorosaccharin in Aqueous Acetic Acid Medium. E J Chem, 1 (2), 132 ( 2004). 9. D. GARG, S. KOTHARI: Kinetics and Mechanism of the Oxidation of Some α-amino Acids by Benzyltrimethylammoniumtribromide. Indian J Chem, 44B, 1909 (2005). 10. A. J. MOHAMMED, H. HADI: Kinetics and Mechanism Studies of Oxidation of α-amino Acids by N-bromosuccinimide. J Al-Nahrain University, 10 (2), 66 ( 2007). 11. L. PUSHPALATHA, K. VIVEKANANDAN: Oxidation of Acidic Amino Acids by N-bromonicotinimide – A Kinetic Study. Oxid Commun, 31 (3), 598 (2008). 12. L. PUSHPALATHA, K. VIVEKANANDAN: N-Bromonicotinimide Oxidation of Essential Amino Acids – A Kinetic Study. Oxid Commun, 32 (1), 85 (2009). 13. L. PUSHPALATHA, K. VIVEKANANDAN: Kinetics of Oxidative Cleavage of Non-essential Amino Acids by N-bromonicotinimide in Aqueous Acetic Acid Medium. Oxid Commun, 33 (4), 851 (2010). 14. B. L. HIRAN, M. L. MEENA, J. KUNTHWAL: A Kinetics and Mechanistic Study of the Oxidation of Alanine by Chromium (VI) in DMF-Water Medium. Der Pharma Chemica, 2 (5), 470 (2010).

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15. R. SINGH, D. K. TAMTA, S. K. JOSHI, N. CHANDRA, N. D. KANDPAL: Oxidation of Acidic Amino Acids by Manganese (III) in Aqueous Sulphuric Acid. J Chem Pharm Res, 3 (1), 529 (2011). 16. K. M. NAIK, S. T. NANDIBEWOOR: Kinetics and Mechanism of Oxidation of L-leucine by Alkalinediperiodatocuprate (III) – A Free Radical Intervention, Deamination and Decarboxylation. J Chem Sci, 124 (4), 809 (2012). 17. A. SHARMA, A. MEENA, J. KHATRI, P. SWAMI, V. SHARMA: Kinetics and Mechanism of the Oxidation of DL-methionine by Quinolinium Chlorochromate. J Appl Chem, 1 (1), 70 ( 2012). 18. C. LOPEZ, A. GONZALEZ, F. P. COSSIO, C. PALOMO: 3-Carboxypyridinium Dichromate ( NDC) and 4-carboxypyridinium Dichromate (INDC), Two New Mild, Stable, Efficient and Inexpensive Chromium (VI) Oxidation Reagents. Synth Commun, 15 (13), 1197 (1985). 19. C. KARUNAKARAN, V. CHIDAMBARANATHAN: Linear Free Energy Relationships Near Isokinetic Temperature. Oxidation of Organic Sulphides with Nicotinium Dichromate. Croat Chem Acta, 74 (1), 51 (2001). 20. K. G. SEKAR: Oxidation of α,β-unsaturated Alcohols by Nicotinium Dichromate. Asian J Chem, 15 (1), 423 ( 2003). 21. K. G. SEKAR: Kinetic Studies on the Oxidation of Some Para and Meta-substituted Benzaldehydes by Nicotinium Dichromate. J Chem Res (S), 626 (2002). 22. D. S. BHUVANESHWARI, K. P. ELANGO: Correlation Analysis of Reactivity in the Oxidation of Anilines by Nicotinium Dichromate in Non-aqueous media. Int J Chem Kinet, 38, 657 (2006). 23. K. G. SEKAR, K. ANBARASU: Structure and Reactivity of Cyclanols towards Nicotinium Dichromate Oxidation. Oxid Commun, 31 (1), 199 (2008). 24. K. G. SEKAR, K. ANBARASU: Kinetics of Oxidation of Phenoxyacetic Acids by Nicotinium Dichromate. Oxid Commun, 34 (2), 314 (2011). 25. M. VELLAISAMY, K. SURYAKALA, M. RAVISHANKAR: Kinetics and Mechanism of Oxidation of 2-naphthol by Nicotinium Dichromate. J Chem Pharm Res, 3 (5), 678 (2011). 26. A. N. PALANIAPPAN, A. ARULMOZHI VARMAN, S. SRINIVASAN, S. SENTHIL KUMAR: Kinetics and Mechanism of Oxidation of Anthranyl Styryl Ketone by Nicotinium Dichromate in Aqueous Acetic Acid Media. J Chem Pharm Res, 4 (6), 2874 (2012). 27. F. FEIGL: Spot Test. Elsevier, Amsterdam, 1954, p. 20. 28. J. BASSETT, R. C. DENNEY, G. H. JEFFERY, J. MENDHAM: Vogel’s Test Book of Quantitative Inorganic Analysis. Longmans, London, 1978, p. 229. 29. K. B. WIBERG, Th. MILL: The Kinetics of the Chromic Acid Oxidation of Benzaldehyde. J Am Chem Soc, 80 (12), 3022 (1958). Received 5 December 2013 Revised 6 February 2014

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Oxidation Communications 37, No 3, 669–680 (2014) Oxidation in the presence of Cr- and Mn-containing oxidants

Inhibition Effect of Surfactants in Alkaline KMnO4 Oxidation of Xylose and Galactose. A Kinetic Study R. Tripathi, S. K. Upadhyay* Department of Chemistry, H. B. Technological Institute, 208 002 Kanpur, India E-mail: [email protected] ABSTRACT The kinetics of oxidation of reducing sugars, viz. xylose and galactose, by potassium permanganate in the presence of anionic, cationic and non-ionic surfactants has been studied spectrophotometrically by monitoring the concentration of MnO4– at its λmax (520 nm). The reactions exhibit first order dependence of rate with respect to each [sugar], [OH–] and [MnO4–]. The order of reaction in sugar and OH– was found to decrease from unity at higher [sugar] and [OH–], respectively. An inhibition effect of surfactant on the rate, below critical micelle concentration (CMC) of the surfactant in case of anionic and cationic surfactants and above CMC in case of non-ionic surfactant has been observed. A mechanism consistent with kinetic data has been proposed. Keywords: inhibition, surfactants, KMnO4, kinetics, xylose, galactose. AIMS AND BACKGROUND Micelles are known as a colloidal solution which behave as strong electrolyte at low concentration but exhibit colloidal properties at higher concentrations and they affect the rate of reaction by separation the substrate between aqueous and micelles pseudo-phase and also by perturbing the thermodynamic parameters of the reaction1–3. Micellar/premicellar catalysis in various redox reactions is reported in literature4–9. The oxidation of reducing sugars has shown10–12 a wide variety of kinetic results. In few cases, premicellar aggregation has also been observed. It has been observed13 that a small amount of cationic surfactant, viz. cetyltri methyl ammonium bromide (CTAB), retards the rate of oxidation of some sugars by potassium permanganate, which is well known as one-electron oxidant14. The retarding effect of anionic surfactant (sodium lauryl sulphate, NaLS) and non-ionic surfactant (Triton X-100, Tx-100) on the rate of oxidation of reducing sugars by alkaline MnO4– has also been observed13 above the CMC of the surfactant. In order to explore detail mechanistic aspect of the reaction, it *

For correspondence.

669

was thought worthwhile to investigate the kinetics of oxidation of some more reducing sugars, viz. xylose and galactose, by MnO4– in the presence of all three types of surfactants, i.e. anionic, cationic and non-ionic surfactants. In the present communication the kinetic results of the oxidation of xylose and galactose by MnO4– in the presence of NaLS, CTAB and Tx-100 are reported and suitable mechanism is proposed. EXPERIMENTAL Permanganate was prepared by dissolving potassium permanganate (GR grade Loba, Mumbai, India) in doubly distilled water. Freshly prepared solutions of D-xylose and D-galactose (AR Thomas Baker, Mumbai, India) in doubly distilled water were used throughout the experiments. The surfactants NaLS (AR, s.d. fine), CTAB and Tx–100 (AR, Thomas Baker) were used as such. However, their critical micelle concentrations (CMC) were determined by surface tension measurement, obtained as 8.2 × 10–3 mol dm–3, 9.6 × 10–4 mol dm–3 and 1.5 × 10–4 mol dm–3 for NaLS, CTAB and Tx–100, respectively at 25oC. These values of CMC were matching the reported values of CMC as 8.0 × 10–3 mol dm–3, 9.8 × 10–4 mol dm–3 and 1.6 × 10–4 mol dm–3 in case of NaLS (Ref. 15), CTAB (Ref. 16) and Tx–100 (Ref. 17), respectively at 25oC. All other reagents, viz. NaOH, NaNO3 etc. used were of AR grade. All the solutions including that of surfactants were prepared in doubly distilled water. Method. To a reaction mixture containing appropriate quantities of solutions of KMnO4, NaOH, NaNO3, surfactants and required amount of doubly distilled water was added so that the total volume of mixture was 50 ml after adding substrate (xylose/galactose). The reaction mixture was then placed in a water bath maintained at desired temperature ± 0.1oC. The reaction mixture was allowed to attain the bath temperature. The reaction was then initiated by adding requisite amount of xylose/galactose placed separately in the same bath. The rates were measured by monitoring the absorbance due to KMnO4 as a function of time at 520 nm (λmax of KMnO4) on a spectrophotometer (Toshniwal, TVSP-25, India). The absorbance due to other substances, i.e. surfactants, substrate etc. was negligible. The concentration of KMnO4 was kept within the limits of the Beer law. Determination of rate constants. The reactions were studied at different initial concentrations of reactants in the absence as well as in the presence of surfactants. lg (absorbance) versus time plots (Fig. 1) were found to be good straight lines, suggesting a first order dependence of rate with respect to MnO4–. Therefore, the pseudo-first order rate constants in MnO4– (kobs) have been evaluated from the slopes of these straight lines. The reported rate constant data, represented as an average of duplicate runs, were reproducible to within ± 5%.

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Fig. 1. lg (absorbance) versus time, i.e. first order plots in KMnO4 at 35oC [Substrate] = 20.0 ×10–3 mol dm–3; [KMnO4] = 4.0 ×10–4 mol dm–3; [NaOH] = 10.0 × 10–3 mol dm–3; [NaNO3] = 0.2 mol dm–3; a, b, c and d represent the plots in aqueous medium (a), in the presence of [NaLS] = 7.0 × 10–3 mol dm–3 (b), in the presence of Tx–100 (3.0 × 10–3 mol dm–3) (c) and CTAB (0.50 × 10–4 mol dm–3) (d), respectively

Stoichiometry and product analysis. The stoichiometry of the reactions between xylose/ galactose and alkaline KMnO4 in the absence as well as in the presence of surfactants has been studied. The reaction mixtures containing a large excess of MnO4– over xylose/ galactose in alkaline medium in presence and absence of surfactants were kept for 72 h at 40oC until the reaction was complete. Estimation of unreacted MnO4– showed that 1 mol of xylose/galactose consumes 2 mol of MnO4–. The results may be represented by the following stoichiometric equation: RCHOHCHO + 2OH– + 2MnO4– → 2MnO42– + H2O + RCHOHCOOH

The presence of aldonic acid as the oxidation product was identified by the spot test .The formation of aldonic acids as the oxidation products of reducing sugars are also reported in literature19,20. 18

RESULTS Although the oxidation of some monosaccharide’s by alkaline MnO4– is reported, in order to compare the kinetic results of oxidation of xylose and galactose by alkaline MnO4– with those in the presence of surfactants, we have investigated the detailed kinetics of reactions in the absence as well as in the presence of surfactants at constant ionic strength (0.2 mol dm–3) maintained by NaNO3.

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The observed pseudo-first order rate constants in the absence and in the presence of each surfactant at various [MnO4–]0 remained identical (Table 1) confirming a first order dependence of rate with respect to [MnO4–]. Table 1. Effect of [KMnO4] on the observed rate constant (kobs) at 35oC [Xylose] = [galactose] = 20.0 × 10–3 mol dm–3; [NaNO3] = 0.2 mol dm–3; [NaOH] = 10.0 × 10–4 mol dm–3; [CTAB] = 0.50 x 10–4 mol dm–3, [NaLS] = 7.0 × 10–3 mol dm–3 and [ Tx–100] = 3.0 × 10–4 mol dm–3

kobs × 104 (s–1)

[KMnO4] × 104 (mol dm–3) 3.0 4.0 5.0 6.0

kaq. 2.66 2.66 2.65 2.66

xylose kNaLS kCTAB 2.33 1.91 2.32 1.91 2.33 1.91 2.33 1.91

kTx-100 2.08 2.08 2.08 2.08

galactose kNaLS kCTAB 3.25 2.75 3.25 2.75 3.25 2.75 3.20 2.75

kaq. 3.5 3.5 3.5 3.5

kTx-100 3.0 3. 0 3. 0 3. 0

The results of effect of [substrate] and [OH–] on the observed rate constants were identical (Table 2). The plots of kobs versus [substrate] or [OH–] showed derivation from linearly at higher concentration of substrate or OH–. However, the plots of (1/kobs) versus 1/[substrate] (Fig. 2) and plots of (1/kobs) versus 1/ [OH–] (Fig. 3) were linear with positive intercepts. The observed results suggest that the order of reaction in substrate and alkali decreases from unity at higher [substrate] and [OH–], respectively. Table 2. Effect of [substrate] and [OH–] on kobs at 35oC [NaNO3] = 0.2 mol dm–3; [KMnO4] = 4.0 ×10–4 mol dm–3; [NaLS] = 7.0 × 10–3 mol dm–3; [CTAB] = 0.50 × 10–4 mol dm–3, and [ Tx–100] = 3.0 × 10–4 mol dm–3

kobs × 104 (s–1)

[Substrate] × [NaOH] ×104 103 (mol dm–3) (mol dm–3) kaq.   5.0 10.0 15.0 20.0 30.0 40.0 20.0 20.0 20.0 20.0

672

10.0 10.0 10.0 10.0 10.0 10.0   2.5   5.0 15.0 20.0

2.25 2.33 2.66 3.00 3.16 2.08 2.25 2.66 2.91 3.16

xylose kNaLS kCTAB

kTx–100

kaq.

1.75 1.91 2.33 2.83 3.03 1.91 1.83 2.33 2.75 3.00

1.66 1.75 2.08 2.66 2.91 1.60 1.75 2.08 2.36 2.83

3.08 3.16 3.5 3.91 4.08 2.91 3.08 3.5 3.83 3.91

1.38 1.58 1.91 2.58 2.83 1.41 1.58 1.91 2.28 2.58

galactose kNaLS kCTAB

kTx–100

2.91 3.08 3.25 3.66 3.83 2.66 2.91 3.25 3.5 3.66

2.75 2.83 3.00 3.16 3.50 2.50 2.66 3.00 3.25 3.66

2.33 2.50 2.75 2.91 3.00 2.25 2.33 2.75 2.91 3.08

Fig. 2. Plots of 1/kobs versus 1/[substrate] at 35oC [KMnO4] = 5.0 × 10–4 mol dm–3; other conditions are same as in Fig. 1

Fig. 3. Plots of 1/kobs versus 1/[OH–] at 35oC [KMnO4] = 5.0 × 10–4 mol dm–3; other conditions are same as in Fig.1

The effect of each surfactant on the rate of oxidation has been studied at 3 different temperature, viz. 35, 40 and 45oC. The results are represented graphically in form of plots of (kobs) versus [surfactant] in Figs 4–6 in case of NaLS, CTAB and Tx100, respectively. It is observed from the results that kobs decreases on increasing the [surfactant]. The retarding effect of NaLS and CTAB on the rate of oxidation (or on kobs) has been observed below CMC of the surfactants. However, the retarding effect of [Tx-100] on the rate of oxidation (or on kobs) was observed above the CMC (1.6 × 10–4 mol dm–3) of the surfactant. The intercepts of the plots of (kobs) versus [surfactant] were matching with the observed rate constants at [surfactant] = 0. 673

Fig. 4. Plots of kobs versus [NaLS] at 35, 40 and 45oC [Substrate] = 20.0 × 10–3 mol dm–3; [KMnO4] = 4.0 × 10–4 mol dm–3; [NaOH] = 10.0 ×10–4 mol dm–3, and [NaNO3] = 0.2 mol dm–3; a – xylose, b – galactose

Fig. 5. Plots of kobs versus [CTAB] at 35, 40 and 45oC Other conditions are same as in Fig. 4

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Fig. 6. Plots of kobs versus [Tx–100] at 35, 40 and 45oC Other conditions are same as in Fig. 4

The effect of ionic strength on the rate of oxidation was investigated by varying the concentration of NaNO3 and keeping other reactants at a fixed in the absence as well as in the presence of the surfactants. The observed rate constants (kobs) are reported in Table 3. The value of kobs increased with an increase in NaNO3 concentration suggesting an involvement of similar charged ions in the rate-determining step. Table 3. Effect of added [NaNO3] on the observed rate constant at 35oC [KMnO4] = 4.0 × 10–4 mol dm–3; [substrate] = 20.0 × 10–3 mol dm–3; [NaOH] = 10.0 × 10–4 mol dm–3; [NaLS] = 7.0 × 10–3 mol dm–3; [CTAB] = 0.50 × 10–4 mol dm–3, and [ Tx–100] = 3.0 × 10–4 mol dm–3

kobs × 104 (s–1)

[NaNO3] × 102 (mol dm–3)

xylose kaq.

  5.0 10.0 20.0

2.16 2.66 2.83

kNaLS 1.66 2.33 2.58

kCTAB 1.50 1.90 2.00

galactose kTx–100 1.41 2.08 2.41

kaq.

kNaLS

kCTAB

kTx–100

3.25 3.50 3.83

3.08 3.25 3.75

2.16 2.75 2.91

2.66 3.00 3.33

The activation parameters (∆E*act, ∆H* and ∆S*) in the presence and absence of the surfactants have been evaluated using the Arrhenuis and Eyring equations with 675

second order rate constant determined at 35, 40, 45oC and are reported in Table 4. The large values of ∆E*act and ∆H* in the presence of surfactants in comparison to those in aqueous medium are consistent with the accepted view that the slow reaction (in the presence of surfactants) would require a higher ∆E*act, ∆H*. The entropy of activation ∆S* in the presence of surfactants is less negative which suggests that the reactants become relatively more rigid in the presence of surfactants. This is not surprising in view of the binding association of the reactant to the surfactants. Table 4. Activation parameters [KMnO4] = 4.0 × 10–4 mol dm–3; [substrate] = 20.0 × 10–3 mol dm–3; [NaOH] = 10.0 × 10–4 mol dm–3; [NaNO3] = 0.2 mol dm–3; [NaLS] = 7.0 × 10–3 mol dm–3; [CTAB] = 0.50 × 10–4 mol dm–3, and [ Tx–100] = 3.0 × 10–4 mol dm–3

Temperature

Second order rate constants (mol–1 dm3 s–1) xylose           

308 K 313 K 318 K Activation parameters Eact ± 0.25 (kJ mol–1) ∆H* ± 0.25 (kJ mol–1) ∆S* ± 1.00 (J K–1 mol–1) ∆G* ± 0.50 (kJ mol–1)

galactose

aq. 0.66 0.70 0.75

NaLS CTAB Tx–100 aq. 0.58 0.47 0.52 0.87 0.60 0.50 0.54 0.97 0.64 0.52 0.56 0.93

NaLS CTAB Tx–100 0.81 0.68 0.75 0.91 0.75 0.77 0.93 0.77 0.81

9.6 7.0 256.1 85.1

16.3 13.7 235.7 87.5

17.3 14.7 229.9 86.7

11.5 8.9 252.6 88.1

14.6 11.8 242.6 88.1

12.5 9.9 244.2 86.2

13.4 10.8 243.4 87.0

16.3 13.7 233.8 86.9

DISCUSSION A first order dependence of rate with respect to each alkali and substrate at lower [OH–] and [substrate], respectively in the absence and in the presence of the surfactants, indicate the enolisation of sugar i.e. the formation of enediol anion of reducing sugar. In the absence of other reactants these anions undergo epimerisation and isomerisation to form mixture of aldoses and ketoses (the Lobry de-Bruyn alberda Van Ekenstion Transformation). Aldoses and ketoses generally yield a mixture of Z- and E-enediols, the proportion of which differ from sugar to sugar and experimental conditions, viz. strength and nature of alkali and temperature. In the presence of an oxidant or a catalyst, the enediol anion has been considered as the reactive species of the reducing sugar. There are evidences21 for the formation of small complexes between surfactant molecules and substrate/oxidant at the concentration of the surfactant below CMC. In such instances the catalysis or inhibition occurs at the surfactant concentrations lower than that for CMC. The inhibition effect of surfactants on the rate of oxidation of reducing sugars by alkaline hexacyanoferrate(III) has been explained22–24 on the basis of formation of substrate-surfactant aggregate and its inactivity towards Fe(CN)63–. In the present investigations, the reactions have been studied at high ionic 676

strength. The kinetic results are different and premicellar inhibition has been observed in case of both the ionic surfactants, viz. NaLS and CTAB. It, therefore, appears that reactions follow different mechanism. On the basis of the above facts and observed kinetic results, it appears that the surfactant molecules form an inactive aggregate with MnO4– , i.e. {MnO4– – surfactant} aggregate. Thus a common mechanism for oxidation of xylose/galactose by alkaline KMnO4 in the presence of surfactants may be proposed as given in Scheme 1. Scheme 1 K1

sugar + OH– E– (enediol / anion) K2

E– + MnO4– I (intermediate complex) k3

(I) → products K4

surfactant + MnO4– (inactive)

(X) complex

(i)

(ii)

(iii) (iv)

Rate law in the absence of surfactants. Steps (i), (ii) and (iii) of Scheme 1 represent the mechanism in the absence of surfactants. According to Scheme 1, the rate of disappearance of MnO4– in the absence of surfactant may be given as follows: –

d[MnO4–] dt

= k3 [I],

(1)

where [I] = K2 [E–] [MnO4–] , (from step (ii))

(2)

[E–] = K1 [sugar] [OH–], (from step (i))

(3)

[I] = K1K2 [sugar] [OH–] [MnO4–]

(4)

and and, therefore; Further the total concentration of [MnO4–] in the absence of surfactant is given as follows: [MnO4–]T = [MnO4–] + [I]

(5)

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On substituting the value of [I] in equation (5) and solving it, the value of [I] in terms of [MnO4–]T may be obtained as follows: K1K2 [sugar] [OH–] [MnO4–]T

[I] =

1 + K1K2 [sugar] [OH–]

.

(6)

On substituting the value of [I] in equation (1), the rate law becomes: –

d [MnO4–]

=

dt

k3 K1K2 [sugar] [OH–] [MnO4–]T 1 + K1K2 [sugar] [OH–]

.

(7)

At higher concentrations of sugar and OH–, where K1K2 [sugar][OH–]>>1 may be taken as suitable approximation the rate law (7) becomes (8). –

d [MnO4–] dt

= k3 [MnO4–]T.

(8)

Rate laws (7) and (8) explain the experimental results in the absence of surfactants. Rate law in the presence of surfactants. Step (iv) of the proposed mechanism (Scheme 1) is also operative and, therefore, total concentration of [MnO4–]T may be represent as: [MnO4–]T = [MnO4–] + [I] + [X]

(9)

where [X] = K4 [surfactant] [MnO4–] (from step (iv)) Thus, by finding the [I] in terms of [MnO4–]T and substituting its value in equation (1), the rate law become: –

d [MnO4–] dt

=

k3K1K2 [sugar] [OH–] [MnO4–]T 1 + K1K2 [sugar] [OH–] + K4 [surfactant]

.

(10)

Rate law (10) suggests a retarding effect of surfactant on the rate of reaction, which has also been observed experimentally. The formation of aggregate/complex between MnO4– and surfactant depends on electrostatic and hydrophobic forces. In case of NaLS, the electrostatic repulsion between negatively charged MnO4– and anionic surfactant opposes the binding between NaLS and MnO4–. Thus in case of NaLS hydrophobic interaction favours the binding whereas electrostatic repulsion opposes it. Consequently the binding between NaLS and MnO4– should be much less and, therefore, the retarding effect of NaLS should be less. In case of CTAB, the electrostatic attraction between MnO4– and positively charged surfactant favor the binding between MnO4– and CTAB in addition to hydrophobic interactions. Thus the aggregation/complex formation will dominate in case of CTAB. Consequently retarding effect of CTAB on the rate of reaction should be very high.

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There are no electrostatic interaction with polar head groups of non-ionic surfactants, i.e. Triton X-100. The poly(oxyethylene) head groups of non-ionic surfactants play a significant role in favoring the incorporation or solubilation of the MnO4– in the micelles. The binding between the MnO4– and Triton X-100 should be between NaLS and CTAB. The observed experimental rate constants in NaLS, CTAB and Tx–100 are in order of kaq > kNaLS > kTX–100 > kCTAB

suggesting highest retardation of rate by CTAB. The experimental results are in complete agreement with the proposed mechanism. REFERENCES 1. K. U. DIN, K. HARTANI, Z. KHAN: Micellar Catalysis on the Redox Reaction of Glycolic Acid with Chromium (VI). Int J Chem Kinet, 6, 33 (2001). 2. J. PANDA, G. P. PANIGRAHI:Kinetics of Cationic Micelle Catalyzed Oxidation of Cyclohexanone by Vanadium (V). J Ind Chem Soc, 79, 58 (2002). 3. K. U. DIN, A. MORSHED, A. MOHAMMAD, Z. KHAN: Micellar Effect on the Oxidation of Glucose by Chronic Acid in Perchloric Acidic Medium. J Carbohyd Chem, 22, 843 (2003). 4. S. PANDAY, S. K.UPADHYAY: Effect of Cationic Micellar Aggregates on the Kinetics of Oxidation of Amino Alcohols by N-bromosuccinimide in Alkaline Medium. J Colloid Interface Sci, 2, 285 (2005). 5. F. SANCHEZ, M. L. MOYA, A. RODRIQUEZ, R. J. GOMEZ HERENA, C. YANES, P. LOPEZCORNEJO: Micellar Microemulsion and Salt Kinetic Effects upon the Reaction Fe(CN)2(bpy)2+S2O82–. Langmuir, 13, 3084 (1997). 6. J. PANDA, G. P. PANIGRAHI: Studies on Kinetics and Mechanism of Oxidation of D-sorbitol and Dmannitol by Cerium (IV) in a Aqueous Micellar Sulphuric acid Media. Ind J Chem, 42, 11 (2013). 7. A. H. GEMEAY, I. A. MANSOUR, G. R. El-SHARKAWY, B. A. ZAKI: Catalytic Effect of Supported Metal Ion Complexes on the Induced Oxidative Degredation of Pyrocatechol Violet by Hydrogen Peroxide. J Colloid Interface Sci, 263, 228 (2003). 8. E. PANDAY, S. K. UPADHYAY: Effect of Micellar Aggregates on the Kinetics od Oxidation of α-amino Acids by Chloramine-T in Perchloric Acid Medium. Colloid and Surface A: Physicochem Eng Aspect, 269, 7 (2005). 9. N. KAMBO, S. K. UPADHYAY: Inhibition of Tx–100 on the Rate of Hexacyanoferrate (III) Oxidation of Reducing Sugars: A Kinetic Study. J Disp Sci Tech, 27, 6 (2006). 10. S. H. SRIVASTAVA, P. SINGH: Mechanistic Investigation of Pd (II)- Catalysed Oxidation of Maltose by Chloramine-T in Acidic Medium: A Kinetic Study. Oxid Commun, 33 (2), 408(2010). 11. V. K. SHUKLA, K. MITHILESH, R. A. SINGH: Kinetics and Mechanism of Oxidation of Ir(III) Catalysed Oxidation of D-galactose by Potassium Bromate in Perchloric Acid Medium. Oxid Commun, 30 (1), 88 (2007). 12. A. TOMAR, A. KUMAR: Kinetics and Mechanism of Oxidation of D-glucose by Tetraethylammonium Chloromate in Aqueous Acid. Oxid Commun, 30 (2), 368 (2007). 13. R. TRIPATHI, N. KAMBO, S. K. UPADHYAY: Premicellar/Micellar Inhibition in Alkaline KMnO4 Oxidation of Reducing Sugars. J Disp Sci Tech, 33 (10), 1393 (2012). 14. S. DASH, S. PATEL, B. K. MISHRA: Oxidation by Permanaganate: Synthetic and Mechanistic Aspects. Tetrahedron, 65, 707 (2009). 15. F. M. MENER, C. E. PORTNOY: Chemistry of Reactions Proceeding inside Molecular Aggregates. J Am Chem Soc, 89, 18 (1996).

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16. G. S. HARTELY, B. COLLIE, C. S. SAMIS: Transport Numbers of Paraffine-chain Salts in Aqueous Solution Part 1. Measurement of Transport Numbers of Cetylpuridinium and Cetyltrimethylammonium Bromide and Their Interpretation in terms of Micelle Formation with Some Data also for Cetane Sulphonic Acid. Trans Faraday Soc, 32, 795 (1936). 17. I. T. YEON, M. M. GOSH, C. D. COX, K. G. ROBINSON: Micellar Solubilization of Polynuclear. Environ Sci Tech, 29, 3015 (1995). 18. F. FEIGL: Spot Tests in Organic Analysis. Elsevier, New York, 1996. 19. E. O. ODEBUNMI, S. O. OWALUDE: Kinetics and Mechanism of Oxidation of Some Simple Reducing Sugars by Permanganate Ion in Alkaline Medium. J Iranian Chem Soc, 5 (4), 623 (2008). 20. M. C. AGRAWAL, S. P. MUSHRAN: Mechanism of Oxidation of Aldoses by Chloramine-T. J Chem Soc Perkin Trans, 2, 762 (1973). 21. E. H. CORDES. C. GILTER: Progress in Bioorganic Chemistry (Eds E. T. Kaiser, F. J. Kezdy). Willey, New York, 2, 24 (1973). 22. R. SHUKLA, N. KAMBO, S. K. UPADHYAY: Inhibition effect of {Cationic Surfactant-Ascorbicacid} Premicellar Aggregation on the Rate of Hexacyanoferrate (III) Oxidation of Ascorbic Acid: A Kinetic Study. J Disp Sci Tech, 29, 905 (2008) 23. N. KAMBO, S. K. UPADHYAY: Micellar Inhibition in Cysteine–Cystine Transformation by Alkaline Hexacyanoferrate(III): A Kinetic Study. Colloid Surf A: Physicochem Eng Aspects, 296, 117 (2007). 24. R. SHUKLA, S. K. UPADHYAY: Inhibition Effect of {Surfactant–Substrate} Aggregation on the Rate of Oxidation of Reducing Sugars by Alkaline Hexacyanoferrate(III). Int Chem Kinet, 39, 595 (2007). Received 18 April 2012 Revised 26 July 2012

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Oxidation Communications 37, No 3, 681–686 (2014) Electrochemical reduction

Electrochemical Reduction of In(III) with Isoleucine in Non-aqueous Media at Dropping Mercury Electrode J. Singh, O. D. Gupta* Department of Chemistry, University of Rajasthan, 302 004 Jaipur, India E-mail: [email protected] Abstract Studies of In(III) metal with isoleucine have been carried out by polarographic method in non-aqueous (20%, 40%, DMF) medium under varying temperatures at 308 and 318 K. Potassium nitrate was used as a supporting electrolyte. The reduction of In(III) was found to be quasi-reversible in non-aqueous medium. Isoleucine ligand has shown the formation of 1:1, 1:2 and 1:3 complexes. The DeFord and Hume method has been applied for the determination of composition and stability constants of the complex species. The changes in thermodynamic parameters ∆H*, ∆G*, ∆S*, accompanying complexation have been evaluated. The mathematical Mihailov method has also been applied for the comparison of stability constants values. Keywords: In(III), isoleucine, stability constant. AIMS AND BACKGROUND Amino acids are the chemical units or ‘building block’ of the body that makes up proteins, which improve the growth and maintenance of all cells are dependent on them. Isoleucine stimulates the brain in order to produce mental alertness. The coordinated system Cu(II)-neutral–L-isoleucine and Cu(II)–L-isoleucinate ion were studied polarographically, in aqueous medium μ = 1.0 M (NaClO4) and 25 ± 0.1°C and the stability constants of the complexes were determined1. Interaction between Cd(II) and some L-amino acids such as L-lysine, L-serine and other L-amino acids as primary ligands and vitamin–PP (vitamin B3) as a secondary ligand has been studied by direct current (DC) polarography at pH = 7.3 ± 0.01 in 1.0 M KNO3 as a supporting electrolyte at 298 K. The Schaap and McMaster method confirmed the formation of 1:1:1, 1:1:2, 1:2:1 complexes2. *

For correspondence.

681

Polarographic studies carried out on Cd(II) and Pb(II) complexes with DL-serine in aqueous and aqua-DMF and aqua-DMSO media, and the formation of complexes with metal/ligand ratio of 1:1, 1:2, and 1:3, is observed in both cases. The values of overall stability constant of complexes were calculated by the methods of DeFord and Hume as modified by Irving (1960) and Mihailov (1974) (Refs 3 and 4). Studies have been done on mixed-ligand complexes of Pb(II) with some amino acids (isoleucine, valine and other amino acids) at DME by polarography. The stability constants were determined by the method of DeFord and Hume3,4 and the stability constants of mixed-ligand complexes have been evaluated by the method of Schaap and McMasters. Electrochemical study of complexes Cd(II) with antibiotic drug has been carried out at DME in non-aqueous media5–8. Polarographic studies of histidine with p-block elements like Ga(III), In(III), Tl(I) have been carried out at constant ionic strength (μ = 1) by using KCl at 298 and 308 K (Ref. 9). Electrochemical studies have been carried out of penicilin benzyl salt with Pb(II) in nonaqueous media10. Polarographic study was carried out on In(III) complex with 2,2´-oxydiacetic acid in aqueous and aqueous–non-aqueous media (methanol, ethanol)11. Comparative polarographic studies of Cd(II) complex of glycine have been carried out in aqueous and aqueous–nonaqueous media12. Present paper deals with the study of electrochemical reduction of In(III) with isoleucine in non-aqueous media (20% DMF and 40% DMF). Experimental A CL-362 polarographic analyser was used to record polarograms using saturated calomel electrode as the reference electrode and DME as microelectrode. Reagent grade chemicals were used and isoleucine was used as complexing agent. All solutions were prepared in doubly distilled water. Potassium nitrate was used as a supporting electrolyte to maintain constant ionic strength. Triton X-100 was used to suppress the observed maxima. The DME had the following characteristics: m = 4.62, mg/s, t = 2 s and height of the mercury column heff = 43 cm and purified N2 was used for deaeration. Results and Discussion The reduction of In(III) in the presence of isoleucine was found quasi-reversible in non-aqueous (20%, 40%, DMF) medium, under varying temperatures at 308 and 318 K. Direct proportionality of diffusion current to the square root of effective height of mercury column indicates the reduction to be diffusion controlled. The values of half-wave potentials for metal ion and their complexes shifted to more negative values on increasing concentration of ligand. This system has been treated with the Gelling method and Er1/2 values were obtained.

682

The complex ion formed is of much larger size as compared to aqua metal ion. Hence the low values of diffusion currents have been found with the increase of ligand concentration. The values of overall formation constants lg βj were calculated by the graphical extrapolation method. The experimentally determined values calculated for In(III)– isoleucine system in 20% DMF at 308 and 318 K are recorded in Tables 1 and 2, respectively, and the overall formation constant were obtained by extrapolation of Fj[(X)] function to the zero concentration. The formation constants of the three complexes are lg β1 = 3.8994, lg β2 = 3.7240 and lg β3 = 6.2145 at 308 K and the formation constant values at 318 K are lg β1 = 3.4992, lg β2 = 3.6443 and lg β3 = 6.2834. In 40% DMF solvent, the overall formation constants for In(III)-isoleucine system were also calculated and the polarographic parameters are recorded at 308 and 318 K in Tables 3 and 4, respectively. The formation constants of three complex species formed are lg β1 = 3.320, lg β2 = 3.4630 and lg β3 = 7.2630. Table 1. Polarographic measurements and Fj[(X)] function values for In(III)–isoleucine system in 20% DMF at 308 K [In(III)] = 0.1 mM; ionic strength (µ) = 1.0 (KNO3)

CX (mol/l)

id (μA)

0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.007

6.6 6.5 6.4 6.3 6.26 6.20 6.1 6.0

Er1/2 (–V vs SCE) 0.5535 0.5630 0.5865 0.5880 0.5960 0.6040 0.6101 0.6032

F0 [(X)]

F1 [(X)]

–   3.0193   5.2439    7.61001 10.7039 14.0988 18.0370 22.8861

– 2019.30 2121.95 2155.00 2425.97 2619.76 2838.50 3126.15

F2 [(X)] ×103 F3 [(X)] × 106 –   69.325   85.980 102.504 118.980 133.960 148.069 168.032

– 14.2031 14.2520 14.2869 14.3035 14.3635 14.3035 14.3062

lg β1 = 3.8994; lg β2 = 3.7240; lg β3 = 6.2945; CX – concentration of isoleucine (mol l–1). Table 2. Polarographic measurements and Fj[(X)] function values for In(III)–isoleucine system in 20% DMF at 318 K [In(III)] = 0.1 mM; ionic strength (µ) = 1.0 (KNO3)

CX (mol/l) 0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.007

id (μA) 6.8 6.7 6.6 6.5 6.49 6.40 6.3 6.2

Er1/2 (–V vs SCE) 0.5530 0.5625 0.5855 0.5875 0.5955 0.6039 0.6100 0.6220

F0[(X)]

F1[(X)]

–   2.8339   4.5723   7.3860 10.0339 13.2884 17.1631 21.6832

– 1866.8 1968.7 2121.2 2262.1 2456.6 2686.6 2949.5

F2[(X)] ×103 F3[(X)] × 106 –    66.8321   84.330 107.089 115.532 131.299 147.799 164.201

– 13.80 13.88 13.89 13.90 13.88 13.98 13.88

lg β1 = 3.4992; lg β2 = 3.6443; lg β3 = 6.2034.

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Table 3. Polarographic measurements and Fj[(X)] function values for In(III)–isoleucine system in 40% DMF at 308 K [In(III)] = 0.1 mM; ionic strength (µ) = 1.0 (KNO3)

CX (mol/l) 0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.007

id (μA) 6.6 6.5 6.4 6.3 6.26 6.20 6.1 6.0

Er1/2 (–V vs SCE) 0.5568 0.5739 0.5803 0.5854 0.5895 0.5969 0.5983 0.5998

F0[(X)]

F1[(X)]

– 3.58039 6.4054 9.5662 13.2247 17.3836 22.2853 28.009

– 2584.60 2702.48 2858.90 3050.17 3270.72 3547.55 3858.55

F2[(X)] ×103 F3[(X)] × 107 –   84.600 101.210 119.630 139.048 155.394 174.590 194.078

– 15.69 15.65 15.60 14.73 14.70 14.69 14.65

lg β1 = 3.997; lg β2 = 3.5320; lg β3 = 7.287. Table 4. Polarographic measurements and Fj[(X)] function values for In(III)–isoleucine system in 20% DMF at 318 K [In(III)] = 0.1 mM; ionic strength (µ) = 1.0 (KNO3)

CX (mol/l)

id (μA)

0.000 0.001 0.002 0.003 0.004 0.005 0.006 0.007

6.8 6.7 6.6 6.5 6.49 6.40 6.3 6.2

Er1/2 (–V vs SCE) 0.5560 0.5664 0.5738 0.5790 0.5883 0.5870 0.5902 0.5930

F0[(X)]

F1[(X)]

–   3.2730   3.7883   8.6831 11.9790 15.9621 20.4663 25.7890

– 2275.6 2374.5 2553.8 2743.7 2989.5 3242.5 3533.5

F2[(X)] ×103 F3[(X)] × 106 –   75.600   87.259 117.963 135.939 157.501 173.760 190.470

– – – 18.75 18.70 18.69 18.65 18.60

lg β1 = 3.320; lg β2 = 3.460; lg β3 = 7.2630.

It may be concluded from the above results that for the definite composition of the non-aqueous mixture, as the concentration of solvent increases, stability of complexes increases because complexation increases with increased availability of ligand molecules. The overall change in thermodynamic parameters ∆H*, ∆G* and ∆S* on complex formation of In(III)–isoleucine system in 20 and 40% DMF solvent mixtures at 308 and 318 K are recorded in Tables 5 and 6.

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Table 5. Stability constants and thermodynamic parameters of In(III)–isoleucine system in aqueous– DMF (20%) solvent mixtures

Metal complex species MX1 MX2 MX3

lg βj 308 K 3.8994 3.7240 7.2145

∆G* (–) (kcal/mol) 4.4862 5.4380 9.832

318 K 3.4992 3.6443 6.2834

∆H* (–) ∆S* (–) ((kcal/mol) (kcal deg/mol) 41.520 0.124 41.849 0.118 41.953 0.107

M = In(III), X = isoleucine. Table 6. Stability constants and thermodynamic parameters of In(III)–isoleucine system in aqueous– DMF (40%) solvent mixtures

Metal complex species MX1 MX2 MX3

lg βj 308 K 3.997 3.532 7.287

∆G* (–) (kcal/mol) 4.330 6.460 9.832

318 K 3.320 3.460 7.263

∆H* (–) (kcal/mol) 41.098 41.843 41.990

∆S* (–) (kcal deg/mol) 0.122 0.118 0.107

M – In(III), X – isoleucine. Table 7. Stability constants verified by the method of DeFord and Hume, and Mihailov

Solvent

Temperature (K)

lg βj

20% DMF

308

lg β1 lg β2 lg β3 lg β1 lg β2 lg β3 lg β1 lg β2 lg β3 lg β1 lg β2 lg β3

318 40% DMF

308 318

Method of DeFord and Hume 3.8994 3.7240 6.2945 3.4992 3.6443 6.2834 3.9970 3.5320 7.2870 3.3200 3.4630 7.2630

Method of Mihailov 3.4532 3.2412 5.5621 3.0326 4.6344 6.0321 3.8640 3.5330 7.0691 3.3200 3.4630 7.2630

The more negative values of ∆G* for 1:3 complexes show that the driving tendency of the complexation reaction is from left to right and the reaction tends to proceed spontaneously. The negative values of ∆H* suggest that the formation of these complexes is an exothermic process. The values of the stability constants for In(III)–isoleucine system in 20% and 40% DMF solvent mixtures have also been further verified by mathematical method given by Mihailov and data are recorded in Table 7.

685

AcknowledgementS The authors are thankful to the Head, Department of Chemistry, University of Rajasthan, Jaipur, for providing facilities to carry out this research. Jyoti Singh is thankful to UGC for providing fellowship. References 1. V. V. Castro-Aleman, M. T. Sanz-Alaezos, J. C. Rodriguez-Placeres, F. J. Garcia-Montelongo: Polarographic Study of the Complex System Cu(II)–neutral–isoleucine and Cu(II) Isoleucinate Ion. Electrochimica Acta, 35 (6), 999 (1990). 2. A. Khanam, F. Khan: Polarographic Study of Ternary Complexes of [Cd(II)–d-amino acidate vitamin PP] System. J Ind Chem Soc, 85 (1), 89 (2008). 3. D. D. DeFord, D. N. Hume: The Determination of Consecutive Formation Constants of Complex Ions from Polarographic Data. J Am Chem Soc, 73, 5321 (1951). 4. M. H. Mihailov: A Correlation between the Overall Stability Constants of Metal Complexes. 1. Calculation of the Stability Constants Using the Formation Function. J Inorg Nucl Chem, 36, 107 (1974). 5. K. D. Gupta, S. C. Baghel, J. N. Gaur: Studies on Mixed Complexes of Cd(II) with Propylenediamine and Trimethyl Amine at the Dropping Mercury Electrode. J Electrochem Soc, 35 (3), 26 (1977). 6. P. S. FernaNedes: Polarogrpahic Study of Few Metals Complex with L-hydroxy Prolline. Res J of Chem Environ (Research Laboratory, St. Xaviers, Mumbai), 59 (2001). 7. G. S. Kalawati, R. S. Panday: Electrochemical Study of Complex Cd(II) with Antibiotic Drug at DME in 20% Methanol-Water and Ethanol–Water Mixture. J Chem Soc, 83, 495 (2006). 8. P. ZUMAN: Role of Mercury Electrodes in Contemporary Analytical Chemistry. Elect An, 912, 1187 (2000). 9. C. Karadia, O. D. Gupta: Determination of Stability Constants of As(III) Complexes with Glycine in DMF and DMSO at Droping Mercury Electrode., Rasayan J Chem, 2 (1), 18 (2009). 10. Meena, S. Sharma, O. D. Gupta: Electrochemical Study of Penicillin Benzyl Salt with Pb(II) in Non-aqueous Media. Asian J Chem, 21, 4346 (2009). 11. S. Sharma, A. Sharma, Meena: Electrochemical Studies of In(III) Complexes with 2,2oxydiacetic Acid in Aqueous and Aqueous–Non-aqueous Media (Methanol/Ethanol). Ultra Chemistry, 4 (2), 165 (2008). 12. S. Kumar, Meena, A. K. Barjatya, O. D. Gupta: Comparison of Stability Constant Values of Cu(II) Complexes with Amino Acids by DeFord & Hume’s and Mihailov’s Methods. Rasayan J Chem, 2, 371 (2009). Received 3 August 2011 Revised 15 October 2011

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Oxidation Communications 37, No 3, 687–695 (2014) Thermooxidative degradation of inorganic complexes

Thermal Decomposition of Indium(III) bis-Oxalatoindate(III) T. Kebede, B. B. V. Sailaja*, M. S. Prasada Rao Bio-inorganic Chemistry Laboratories, School of Chemistry, Andhra University, Visakhapatnam, India E-mail: [email protected] ABSTRACT Indium(III) is precipitated with oxalic acid in the presence of indium nitrate in slightly acidic solution of nitric acid. Chemical analysis of the complex salt obtained indicates the formula In3+[In(III)(C2O4)2(H2O)2]3.2H2O. The thermal decomposition behaviour of the complex was studied using TG, DTA and DTG techniques. These studies indicated the formation of anhydrous indium oxalate and then to a mixture of oxides through the formation of indium(III) oxide as intermediates. Isothermal study, X-ray diffraction pattern and IR spectral data support the proposed thermal decomposition mechanism. Keywords: thermal decomposition, indium oxalate, bis-oxalates, X-ray diffraction, IR data. AIMS AND BACKGROUND The thermal decomposition of some solid oxalates was explained in a review1. The crystalline phases and surface properties of zinc oxalate and ferric oxalate were reported by Dollimore and Nicholson2–4. In decomposition of solid oxalates, it was reported that the surface area will vary with temperature5. Synthesis and characterisation of indium complexes with salen ligands and indium complexes containing two different chelate ligands was also explained6,7. Decomposition mechanisms of oxalato complexes of various metals were reported by Dollimore et al.8–10 Oxalate ion is used in the automatic thermogravimetric analysis of various metal oxalates11–13. The synthesis and characterisation of thallium oxalates was also reported14. Drouane et al. 15 explained the thermal decomposition of oxalato complexes of Ca, Sr and Ba. Decomposition mechanisms of oxalate complexes of cobalt and nickel were described in Refs 16 and 17. The crystal structure of indium complexes *

For correspondence.

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was also reported18. It was established that indium complexes acts as water-tolerant Lewis acids19. The indium bis-oxalatodiaquaindate(III) complex is prepared by adding oxalic acid directly with continuous stirring to a mixture containing an indium salt and an excess of the precipitating metal ion maintaining slightly acidic pH (2–4) conditions. The thermal decomposition of the complex has been studied and a mechanism has been proposed on the basis of the thermal data, infrared spectroscopic and X-ray diffraction studies. EXPERIMENTAL INSTRUMENTATION

Thermal analysis unit. SEIKO combined thermal analysis system (TG/DTA-32), temperature programmable thermal balance, made in Japan and platinum crucible as container were used for taking thermograms in air. The rate of heating was fixed to10oC/min, and sensitivity of the instrument was 0.1 mg. Infrared spectroscopy. The infrared spectra of the complexes were recorded on a Shimadzu FTIR-8201 PC infrared spectrophotometer in KBr pellets. X-ray diffraction data. X-ray diffractometer of Rich Seifert & Co. (made in Germany) attached to a microprocessor was used for taking X-ray diffraction patterns at wavelength of CuKα1 = 1.540598 Å. PREPARATION AND ANALYSIS

The indium bis-oxalatodiaquaindate(III) complex was prepared by adopting the following procedure: About 80 cm3 of 0.05 M indium(III) nitrate in 0.625M HNO3 were taken in a 500-cm3 beaker to which about 60 cm3 1 M ammonium nitrate and 180 cm3 triply distilled water were added slowly while stirring the contents. Then about 80 cm3, 0.1 M oxalic acid were added very slowly (drop-wise from a burette) while stirring the contents vigorously. The complex salt formed was allowed to settle and filtered through a G4 sintered glass crucible. Then the precipitate was washed several times (with a solution obtained by mixing 0.3 M NH4NO3 and 0.25 M HNO3 in 1:1 ratio) to free the excess of oxalic acid. Finally the precipitate was washed with 0.1 M HNO3 to remove NH4NO3. The precipitate was dried in a vacuum desiccator over silica gel. The compound thus obtained was tested to confirm the absence of nitrate. The bis-oxalato complex was isolated from the sulphate medium (containing 0.0625 M H2SO4 and 0.05 M (NH4)2SO4) and purified following the procedure similar to that adopted in nitrate medium.

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RESULTS AND DISCUSSION Chemical analysis. The indium(III)bis-oxalatodiaquaindate(III) complex is prepared according to established procedure. The analyses of the oxalate and indium contents are carried out. The analytical results suggest the composition as In3+[In(III) (C2O4)2(H2O)2]3.2H2O. Table 1 summarises the analytical data. Table 1. Chemical analysis data of indium(III) bis-oxalatoindate(III)

Methods of preparation Moeller1 Deichman2 Present work Calculated *

In3+ 34.24 – 34.05 34.07

Composition (%) Ratio C2O42– total water C2O42–/In3+ 39.20 – 39.08 39.19

26.56* – 26.88* 26.74

1.49 1.50 1.50

Possible formula In2 (C2O4)3.10H2O In2 (C2O4)3.nH2O** In2 (C2O4)3.10H2O In2 (C2O4)3.10H2O

Calculated from the formula; ** n varied from 2 to 8 depending on the time of standing.

Thermogravimetric analysis (TGA). The thermogram of indium(III) bis-oxala­to­ di­aquaindate(III) tetrahydrade and the data obtained from it are given in Fig. 1 and Table 2, respectively. The curve shows four temperature ranges where the complex loses some weight. The first weight loss from 130 to 180oC is due to loss of crystal water for which the observed loss is 4.41% (ca. 4.69%) of the original weight. This is followed by loss of coordinated water from 180 to 230oC. The loss corresponding to this is 14.11% on the thermogram against the calculated 14.09%. The anhydrous indium(III) oxalate so formed is very unstable and subsequently decomposes possibly in a step-wise manner, from 230 to 331oC until the final decomposition product, In2O3, is obtained. The overall weight loss at this stage is 51.71% as is observed on the TG curve (ca. 52.40%).

Fig. 1. TG/DTA of indium(III) bis-oxalatoindate(III)

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Table 2. Summary of the thermal decomposition of indium(III) bis-oxalatoindate(III)

Weight of Step Temperature Loss in weight Possible decomposition product comp. No start (oC) end (oC) obs. (%) calc. (%) (intermediate) 3+ 11.00 1 130 180 4.41 4.69 In [In(C2O4)2(H2O)2]3 2 180 208 14.11 14.09 In3+[In(C2O4)2]3 3 208 297 19.88 24.66 In2O3 + 3In2(C2O4)3 4 297 331.7 51.71 52.40 4In2O3

Differential thermogravimetric analysis (DTG). The DTG data on indium(III) bisoxalatodiaquaindate(III) are shown in Fig. 1. From the figure it is evident that there are 2 significant weight losses. The first being due to loss of water, which occurs in 2 steps as indicated by the two adjacent small peaks on the DTG at about 181.9 and 203.6oC corresponding to the loss of crystal and coordinated water, respectively. The second sharp and very strong peak (68.9%/min) at 327.4oC indicates a continuous decomposition involving the anhydrous indium(III) oxalate to ultimately indium(III) oxide. Differential thermal analysis (DTA). The DTA results of the complex are also shown in Fig. 1 in which the broad endothermic peak with ∆Tmin. at 184.1oC and the shoulder with ∆Tmin. at 214.5oC indicate the 2-step dehydration of the complex to give the anhydrous indium(III) oxalate. The latter being unstable may undergo a step-wise decomposition as indicated by the prominent peaks in the DTA curve. The first endothermic peak with ∆Tmin. at 316.5oC represents decomposition of the complex to form indium oxide corresponding to one-fourth of the indium in the anhydrous salt. The exothermic peak with ∆Tmax. at 329.5oC and the shoulder (also an exothermic peak) with ∆Tmax. at 342.5oC mark the subsequent decomposition of the possible intermediate leading to the final product, In2O3. Isothermal decomposition of indium(III) bis-oxalatoindate(III). The pyrolysis curve of indium(III)bis-oxalatoindate(III) shown in Fig. 1 indicates that complete dehydration of the complex may take place up to about 210oC. In studying the decomposition behaviour of this complex, ca. 210 mg of the sample was taken and heated up to a selected temperature of 200oC. The heating continued isothermally for half an hour. The process may be described by the following equation: 30–200oC

In3+[In(III)(C2O4)2]3.nH2O –––––→ In3+[In(III)(C2O4)2]3 + nH2O

The residue obtained in this manner was analysed for its composition and the ratio of C2O42– to In(III) becomes 2:1 suggesting the probable formula of the intermediate to be In3+[In(III)(C2O4)2]3. Similarly the water content of the original complex salt was determined based on the experimental data. The result of this calculation shows that there are 10

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water molecules in the complex indicating that the possible formula is In3+[In(III) (C2O4)2]3.10H2O or In3+[In(III)(C2O4)2(H2O)2]3.4H2O. Infrared spectra of indium(III) bis-oxalatoindate(III). The infrared spectra of indium(III) bis-oxalatoindate(III) complex salt and the product obtained by heating it to 200oC and cooling to room temperature are given in Figs 2 and 3, respectively. A very broad absorption band at 3456 cm–1 in the spectrum of the original complex indicates the presence of water. However, the absence of such a band in the neighbourhood of 3500 cm–1 in the spectrum of the heated product suggests that the intermediate is anhydrous. A very sharp and medium peak at 809 cm–1 as well as the band at 718 cm–1 in the spectrum of the complex confirm the presence of water in the coordinated form. The corresponding band at 800 cm–1 in that of the heated product appears to be much less intense.

Fig. 2. Infrared spectrum of indium(III) bis-oxalatoindate(III)

Fig. 3. Infrared spectrum of indium(III) bis-oxalatoindate(III) after heating to 200oC

Moreover, the thermal decomposition study of the complex salt under discussion clearly shows that an intermediate corresponding to the empirical formula, In2(C2O4)3 is obtained in the range of 180 to 230oC.

691

The infrared spectral data of the original complex and the product by heating the complex to 200oC are presented in Table 3 for comparison. Table 3. Infrared absorption data of indium(III) bis-oxalatoindate(III) and the heated product at 200oC

original 3600 sp 3450 vb 1720 sh,w 1610 b,s

Complex

1445 w 1366 sp,s 1350 sp,m 809 sp,m 718 sp,m 520 sp,m 440 sp,s

350 sp,s

heated at 200oC

Band assignment

νas,s (H–O–H) 1630 b,s 1520 w 1450 w

1300 vw 800 w 500 w 440 w 400 sh 390 w 370 sp,w 350sp,s

νa (C=O) + δ (H–O–H) νs (C–O) + ν (C–C) νs (C=O) + δ (O–C=O) coordinated water and δ(O–C=O) + ν(M–O)

ν(M–O) + δ(O–C=O) δ (O–C=O) + ν (C–C)

Note: b – broad; m – medium; s – strong; sp – sharp; sh – shoulder; w – weak.

X-ray diffraction data. The X-ray diffraction data of indium(III) bis-oxalatoindate(III) and that of the product obtained after heating the original complex to 200oC and cooling are given in Table 4. The data in Table 4 clearly show that the product obtained by heating the original complex to 200oC has more or less different d-spacings suggesting that it is anhydrous compound. Table 4. X-ray diffraction data of indium(III) bis-oxalatoindate(III) and the product after heating it to 200oC

In3+[In(III) complex] 1 5.6511 5.507x 5.0470 4.6541 4.2380 4.0111 3.9221

692

Complex after heating to 200oC 2 5.6713 5.515x 5.3571 5.0922 5.0352 4.7941 4.6481

to be continued

Continuation of Table 4

1 3.8650 3.8080 3.6751 3.5570 3.0482 3.0010 2.7691 2.7370 2.5510 2.5371 2.4660 2.3349 2.2470 2.2231 2.0091 1.9791 1.9131 1.8491 1.8070 1.7801 1.6830 1.6591 1.6320

2 4.0262 3.9212 3.8073 3.6772 3.5552 3.3061 3.2312 3.1431 3.0504 3.0012 2.7731 2.7402 2.5302 2.4312 2.3342 2.2181 2.1681 2.1341 2.0052 1.9521 1.9372 1.8401 1.7802 1.6931 1.5301

On the basis of the results obtained from these investigations the following thermal decomposition mechanism is proposed:   130–180oC    180–230oC

2In3+[In(III)(C2O4)2(H2O)2]3.4H2O ––––––→ 2In3+[In(III)(C2O4)2(H2O)2]3 ––––––––→



230–316oC   316–329.5oC

2In [In(III) (C2O4)2]3 –––––→ In2O3 + 3In2(C2O4)3 –––––––––→ 3+

   329.5–342.5oC

In2O3 + 3In2O(C2O4)2 ––––––––––––→ 4In2O3

Overall: 2In3+[In(III)(C2O4)2(H2O)2]3 . 4H2O → 4In2O3 + 12CO2 + 12CO + 20H2O

The above mechanism suggests that the compound should be represented by a different formula instead of In2(C2O4)3.10H2O. In the proposed formula three-fourths 693

of the indium should have different chemical environment when compared with the remaining one-fourth of the indium. Hence the structure similar to the other salts is proposed which is also a bis-oxalatoindate(III) complex. The proposed mechanism suggests the most probable structural formula of the complex with indium(III) at the centre of the octahedron as follows: In3+[In(III)(C2O4)2(H2O)2]3.2H2O

REFERENCES   1. K. V. KRISHNA MURTHY, G. M. HARRIS: The Chemistry of the Metal Oxalato Complexes. Chem Rev, 61, 213 (1961).   2. D. DOLLIMORE, J. DOLLIMORE, D. NICHOLSON: The Thermal Decomposition of Oxalates. Part IV. Interrelation of Crystalline Phases in the Thermal Decomposition and Dehydration of Zinc Oxalate. J Chem Soc, 380, 2132 (1965).   3. D. DOLLIMORE, D. NICHOLSON: The Thermal Decomposition of Oxalates. Part III. The Decomposition and Surface Properties of Zinc Oxalate. J Chem Soc, 178, 908 (1964).   4. D. DOLLIMORE, D. NICHOLSON: The Thermal Decomposition of Oxalates. Part VI. The Decomposition and Surface Properties of Ferric Oxalate. J Chem Soc A, 281 (1966).   5. D. DOLLIMORE, D. NICHOLSON: The Thermal Decomposition of Oxalates. Part I. The Variation of Surface Area with the Temperature of Treatment in Air. J Chem Soc, 179, 960 (1962).   6. XING GAO, CHENG JAIN, FANG YUAN, YU HUA ZHU, YI PANI: Salen Ligands. Chinese Chem Lett, 14, 138 (2003).   7. D. L. REGER, S. S. MASON, B. L. REGER, L. A. RHEINGOLD, L. R. OSTRANDER: Synthesis and Characterization of Indium(III) Complexes Containing Two Different Chelate Ligands of the Type [Hydrotris(3,5-dimethylpyrazolyl)borate]In[ligand]X. Inorg Chem, 33, 1811 (1994).   8. D. DOLLIMORE, J. DOLLIMORE, P. D. PERRY: The Thermal Decomposition of Oxalates. Part VIII. Thermogravimetric and X-ray Analysis Study of the Decomposition of Aluminium Oxalate. J Chem Soc A, 448 (1967).   9. D. BROADBENT, D. DOLLIMORE, J. DOLLIMORE: The thermal Decomposition of Oxalates. Part IX. The Thermal Decomposition of the Oxalate Complexes of Iron. J Chem Soc A – Inorg, Phys Theoretical, 451 (1967). 10. D. DOLLIMORE, J. DOLLIMORE, J. LITTLE: The Thermal Decomposition of Oxalates. Part X. Nitrogen Adsorption Data on Solid Residues from the Isothermal Heat Treatment of Manganese(II) Oxalate Dihydrate. J Chem Soc A – Inorg, Phys, Theoretical, 2946 (1969). 11. C. DUVAL: Inorganic Thermogravimetric Analysis. 2nd ed. Elsevier Publishing Company, New York, London, 1963, 148–162. 12. ZHIYONG YANG, S. E. GOULD: Reactions of Tris(oxalato)cobaltate(III) with Two-electron Reductants. Dalton Trans, 3601 (2004). 13. SIHAIYANG, GUOBAO LI, SHUJIANTIAN, FUHUILIAO, JIANHUALIN: An Open-framework Three-dimensional Indium Oxalate: [In(OH)(C2O4)(H2O)]3·H2O. J Solid State Chem, 178, 3703 (2005). 14. T. KEBEDE, K. V. RAMANA, M. S. P. RAO: Some Studies on Thallium Oxalates – XIV. Indium(III) Bis-oxalatodiaquathallate(III) Hexahydrate. Thermochimica Acta, 381, 31 (2002). 15. E. D. DROUANE, Z. GABELICA, R. HUBIN, M. J. HUBIN FRANSKIN: Etude des mécanismes de décomposition thermique des oxalates de baryum, strontium, et magnesium. Thermochimica Acta, 11, 287(1975).

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16. D. BROADBENT, D. DOLLIMORE, J. DOLLIMORE: The Thermal Decomposition of Oxalates. Part VII. The Effect of Prior Dehydration Conditions Upon the Subsequent Decomposition of Cobalt Oxalate. J Chem Soc A, 1491(1966). 17. D. BROADBENT, D. DOLLIMORE, J. DOLLIMORE: The Thermal Decomposition of Oxalates. Part V. The Thermal Decomposition of Nickel Oxalate Dehydrate. J Chem Soc A, 278 (1966). 18. M. A. A. F. de C. T. CARRONDO, A. R. DIAS, M. H. GARCIA, P. MATIAS, M. FATIMA, M. PIENDADE, M. J. VILLA de BRITO: Indium(III) Thiolate-bridged Molybdenocene Complexes: Crystal Structure of [InCl2MoCp2(η-SEt)22][BPh4]· (CH3)2CO. J Organomet Chem, 466, 159 (1994). 19. TECH-PENG LOH, GUAN-LEONG CHUA: An Open-framework Three-dimensional Indium Oxalate: [In(OH)(C2O4)(H2O)]3.H2O. Chem Com, 14, 2739 (2006). Received 15 December 2011 Revised 20 February 2012

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Oxidation Communications 37, No 3, 696–712 (2014) Oxidation in biological systems

Oxidation of Kinetin by Cu(III) Complex in Alkali Media: A Kinetic and Mechanistic Approach A. M. Tatagara, S. D. Lamanib, S. T. Nandibewoora* S.D.M. College of Engineering and Technology, Dharwad, India E-mail: [email protected] b Post-graduated Department of Studies in Chemistry, Karnatak University, 580 003 Dharwad, India a

ABSTRACT The kinetics of oxidation of kinetin (KNT) by diperiodatocuprate(III) (DPC) in alkaline medium at a constant ionic strength of 0.02 mol dm–3 was studied spectrophtometrically. The reaction between DPC and kinetin in alkaline medium exhibits 1:3 stoichiometry (KNT:DPC). Intervention of free radicals was observed in the reaction. Based on the observed orders and experimental evidences, a mechanism involving the monoperiodatocuprate (III) (MPC) as the reactive oxidant species has been proposed. The products were identified by spot test and characterised by spectral studies. The reaction constants involved in the different steps of mechanism were calculated. The activation parameters with respect to slow step of the mechanism were computed and discussed. The thermodynamic quantities were also determined for different equilibrium steps. Keywords: kinetics, mechanism, oxidation, kinetin, diperiodatocuprate(III), thermodynamic parameters. AIMS AND BACKGROUND Kinetin, 6-(furfuryl amino)-purine, is known as plant growth regulator and has profound effect on cell division, cell enlargement, shoot formation, root development and seed germination. It has showed an increase in alkaloid synthesis and nitrogen metabolism (Ambrose and Scikchetti, 1962; Miller et al. 1955). KNT is often used in plant tissue culture for inducing formation of callus (in conjunction with auxin) and to regenerate shoot tissues from callus (with lower auxin concentration). Kinetin exists naturally in the DNA of almost all organisms tested so far, including human cells, and various plants. Kinetin has been thoroughly tested for its powerful anti*

For correspondence.

696

ageing effects in human skin cells and other systems. At present, kinetin is one of the widely used components in numerous skin care cosmetics and cosmeceuticals, such as Valeant products kinerase1. There are some reports published on other biological effects of kinetin in human beings, for example its effects as anti-platelet aggregation factor reducing thrombosis formation. In recent years, the study of highest oxidation state of transition metal has intrigued many researchers. Transition metals in a higher oxidation state can be stabilised by chelation with suitable polydentate ligands. Metal chelates, such as diperiodatocuprate(III) (Refs 2 and 3), diperiodatoargentate(III) (Refs 4–6) and diperiodatonickelate(IV) (Ref. 7) are good oxidants in a medium with an appropriate pH value. Periodate and tellurate complexes of copper in its trivalent state have been extensively used in analysis of several organic compounds8. The kinetics of self-decomposition of these complexes was studied in some details9. Copper(III) is shown to be an intermediate in the copper(II)-catalysed oxidation of amino acids by peroxydisulphate10. The oxidation reaction usually involves the copper(II)-copper(I) couple and such aspects are detailed in different reviews11,12. The use of diperiodatocuprate(III) (DPC) as an oxidant in alkaline medium is new and restricted to a few cases due to its limited solubility and low stability in aqueous medium. DPC is versatile one-electron oxidant for various organic compounds in alkaline medium and it is used as an analytical reagent. Copper complexes have occupied a major place in oxidation chemistry due to their abundance and relevance in biological chemistry13–15. Copper(III) is involved in many biological electron transfer reactions16. They have also been used17 in the differential titration of organic mixtures, in the estimation of chromium, calcium, and magnesium from their ores, and antimony, arsenic and tin – from their alloys. Since multiple equilibria between different copper(III) species are involved, it would be interesting to know which of the species is the active oxidant. In the earlier reports18 on DPC oxidation, periodate had retarding effect and the order in [OH–] was found to be less than unity in most of the reactions. However in the present study, we have observed entirely different kinetic behaviour. A literature survey reveals that there are no reports on the oxidation of KNT by diperiodatocuprate(III). The present study deals with the title reaction to investigate the redox chemistry of DPC in alkaline media, to compute the thermodynamic quantities of various steps of Scheme 1 and to arrive at a plausible mechanism on the basis of kinetics and spectral studies. EXPERIMENTAL Materials and reagents. All chemicals used were of reagent grade and Millipore water was used throughout the experiment. The copper(III) periodate complex was prepared19,20 and standardised by standard procedure21. The UV-vis. spectrum with maximum absorption at 415 nm was characteristic of the copper(III) complex. Solutions of KNT (HI media) and copper sulphate (BDH) were prepared by dissolving 697

known amounts of the samples in distilled water. Periodate solution was prepared and standardised iodometrically22. Required alkalinity and ionic strength were maintained by KOH (BDH) and KNO3 (Analar), respectively, in the reaction solution. Kinetic studies. The kinetics was followed under pseudo-first order condition where [KNT] >[DPC] at 298±0.1 K, unless specified using a Varian CARY 50 Bio UV-vis. spectrophotometer. The reaction was initiated by mixing the DPC to KNT solution which also contained required concentration of KNO3, KOH and KIO4. The progress of the reaction was followed spectrophotometrically at 415 nm by monitoring the decrease in absorbance due to DPC with the molar absorbency index, ε to be 6.231 ± 100 dm3 mol–1 cm–1 (literature ε = 6.230 (Ref. 23)). It was verified that there is a negligible interference from other species present in the reaction mixture at this wavelength. The pseudo-first order rate constants, kobs, were determined from the lg(absorbance) versus time plots. The plots were linear up to 80% completion of reaction under the range of [OH–] used. The order for various species were determined from the slopes of plots of lg kobs versus respective concentration of species except for [DPC] in which non variation of kobs was observed as expected for the reaction condition. During the kinetics a constant concentration, viz. 5.0 × 10–5 mol dm–3 of KIO4 was used throughout the study unless otherwise stated. Since periodate is present in excess in DPC, the possibility of oxidation of KNT by periodate in alkaline medium at 298 K was tested. The progress of the reaction was followed iodometrically. However, it was found that there was no significant reaction under the experimental conditions employed compared to the DPC oxidation of KNT. The total concentration of periodate and OH– were calculated by considering the amount present in the DPC solution and that additionally added. Kinetics runs were also carried out in N2 atmosphere in order to understand the effect of dissolved oxygen on the rate of reaction. No significant difference in the results was obtained under N2 atmosphere and in the presence of air. In view of the ubiquitous contamination of carbonate in the basic medium, the effect of carbonate was also studied. Added carbonate had no effect on the reaction rates. The spectral changes during the reaction are shown in Fig. 1. It is evident from the figure that the concentration of DPC decreases at 415 nm.

698

Fig. 1. Spectroscopic changes occurring in the oxidation of KNT by DPC at 298 K [DPC] = 5.0 × 10–5 mol dm–3; [KNT] = 5.0 × 10–4 mol dm–3; [OH–] = 0.50 mol dm–3; and I = 0.5 mol dm–3 with scanning time interval of 1 min

RESULTS Stoichiometry and product analysis. Different sets of reaction mixtures containing excess of DPC to kinetin in the presence of constants amounts of OH– and KNO3 were kept for 6 h in closed vessel under inert atmosphere. The remaining DPC concentration was estimated spectrophotometrically at 415 nm. The results, 1:3 stoichiometry, are as given in equation (1). H N

H2C N

N N

+ 3 [Cu(H2IO6)(H2O)2] + 2OH–

O

N

NO2 N

N N

N

CH2OH + O

+ 3Cu(I) + 3H2IO63– + 5H2O + 4H+

   

(1)

The main oxidation products were identified as para-nitro-pteridine and furon-2methanol, and were characterised by their melting point, IR and GC-MS. The nature of furon-2-methanol was confirmed by its IR spectrum, which showed a –OH stretch at 3386 cm–1 indicating the presence of alcohol and para-nitro-pteridine was also confirmed by the presence of NO2 group, stretching frequency at 1332 cm–1. Further the products were subjected to GC-mass spectral analysis. GC-MS data were obtained on a QP-2010S Shimadzu gas chromatograph mass spectrometer. The mass spectral data showed a base peak at 98 m/z, thus confirming the presence of furon-2-methanol (Fig. 2) and another product was also confirmed, mass spectral data showed a peak at 699

182 m/z, hence confirming the product para-nitro-pteridine (Fig. 3). All other peaks observed in GC-MS can be interpreted in accordance with the observed structure of para-nitro-pteridine and furon-2-methanol. The presence of Cu(I) was confirmed by UV-vis. spectra.

Fig. 2. Mass spectrum of reaction product, furon-2-methanol at 98 amu

Fig. 3. Mass spectrum of reaction product, para-nitro-pteridine at 182 amu

Regression analysis of experimental data to obtain the regression coefficient r and standard deviation s from the regression line was performed using Microsoft Excel-2003. Reaction order. The reaction order was determined from the slope of lg kobs versus lg [concentration] plots by varying the concentration of KNT, alkali in turn while keeping all other concentrations and conditions constant. 700

Effect of [diperiodatocuprate(III)]. The oxidation of DPC concentration was varied in the range of 1.0 × 10–5 to 1.0 × 10–4 mol dm–3 and the fairly constant value of kobs indicates that the order with respect to [DPC] was unity (Table 1). This was also confirmed by linearity of the plots of log [absorbance] versus time to 80% completion of reaction. Table 1. Effect of [DPC], [KNT], [OH–] and [IO4–] on the oxidation of kinetin by DPC in alkaline medium at 298 K I = 0.1 mol dm–3

[DPC]× 105 (mol dm–3) 1.0 3.0 5.0 8.0 10.0

[KNT] ×104 (mol dm–3) 5.0 5.0 5.0 5.0 5.0

[OH–] (mol dm–3) 0.05 0.05 0.05 0.05 0.05

[IO4–] ×105 (mol dm–3) 1.0 1.0 1.0 1.0 1.0

kobs×103 (s–1) 4.0 4.0 3.9 4.0 4.1

kcal×103 (s–1) 4.0 4.0 4.0 4.0 4.0

5.0 5.0 5.0 5.0 5.0

1.0 3.0 5.0 8.0 10.0

0.05 0.05 0.05 0.05 0.05

1.0 1.0 1.0 1.0 1.0

1.2 3.0 3.9 5.4 6.0

1.1 2.9 4.0 5.3 5.9

5.0 5.0 5.0 5.0 5.0

5.0 5.0 5.0 5.0 5.0

0.01 0.03 0.05 0.08 0.10

1.0 1.0 1.0 1.0 1.0

2.8 3.6 4.0 4.2 4.5

2.6 3.7 4.0 4.3 4.4

5.0 5.0 5.0 5.0 5.0

5.0 5.0 5.0 5.0 5.0

0.05 0.05 0.05 0.05 0.05

0.5 0.8 1.0 3.0 5.0

4.3 4.1 4.0 3.0 2.4

4.4 4.2 4.0 3.0 2.4

Effect of [kinetin]. The effect of KNT on the rate of reaction was studied at constant concentration of alkali, DPC and periodate at constant ionic strength of 0.02 mol dm–3. The substrate KNT was varied in the range of 1.0 × 10–4 to 1.0 × 10–3 mol dm–3. The kobs values increased with increase in concentration of KNT. The apparent order with respect to [kinetin] was found to be less than unity (Table 1) (r ≥ 0.998, S ≤ 0.02). This was also confirmed by the plots of kobs versus [KNT]0.69 which is linear rather than the direct plot of kobs versus [KNT] (Fig. 4).

701

[KNT]0.69 × 10 (mol dm–3) 7

0

1

2

3

4

5

7.4

6

6.4

5

5.4

a

4

4.4

3

b

3.4

2

2.4

1

1.4

0

0

0.2

0.4

0.6

0.8

1

kobs × 103 (s–1)

kobs × 103 (s–1)

6

0.4 1.2

[KNT] × 103 (mol dm–3)

Fig. 4. Plots of kobs versus [KNT]

0.69

(a) and kobs versus [KNT] (b) (conditions as in Table1)

Effect of [alkali]. The effect of increase in concentration of alkali on the reaction was studied at constant concentration of KNT, DPC and periodate at constant ionic strength of 0.02 mol dm–3 at 298 K. The rate of reaction decreases with increase in alkali concentrations (Table 1), indicating positive fractional order dependence of rate on alkali concentration (r ≥ 0.996, S ≤ 0.03). This was also confirmed by the plots of kobs versus [OH–]0.23 which is linear rather than the direct plot of kobs versus [OH–] (Fig. 5) [OH–]0.23 × 10 (mol dm–3) 0

2

4

6

8

10

12 5.0

5.4

4.0 a

b

3.0

3.4 2.4

2.0

1.4

1.0

0.4

0

1

2

3

4

5

6

kobs × 103 (s–1)

kobs × 103 (s–1)

4.4

0.0

[OH–] × 10–2 (mol dm–3)

Fig. 5. Plots of kobs versus [OH ]

– 0.23

(a) and kobs versus [OH–] (b) (conditions as in Table 1)

Effect of [periodate]. The effect of increasing concentration of periodate was studied by varying the periodate concentration from 1.0 × 10–5 to 1.0 × 10–4 mol dm–3 keeping all other reactants concentration constant. It was found that added periodate had retarding effect on the rate of reaction.

702

Effect of ionic strength (I) and dielectric constant of medium (D). The addition of KNO3 at constant [DPC], [KNT], [OH–] and [IO4–] was found that increasing ionic strength of the reaction medium did not effect the rate of reaction. Varying the t-butyl alcohol and water percentage varied dielectric constant of the medium D. The D values were calculated from the equation D = DwVw + DBVB, where Dw and DB are dielectric constants of pure water and t-butyl alcohol, respectively, and Vw and VB – the volume fractions of components of water and t-butyl alcohol, respectively, in the total mixture. The decrease in dielectric constant of the reaction medium decreased the rate of reaction. The plot of lg kobs versus 1/D was linear with negative slope. Effect of initially added products. The externally added products, Cu(I), para-nitropteridine and furon-2-methanol, did not have any significant effect on the rate of reaction. Polymerisation study. The intervention of free radicals in the reaction was examined as follows. The reaction mixture, to which a known quantity of acrylonitrile monomer initially added, was kept for 2 h in an inert atmosphere. On diluting the reaction mixture with methanol, a white precipitate was formed, which indicated the intervention of free radicals in the reaction23. The blank experiments either of DPC or kinetin alone with acrylonitrile did not induce any polymerisation under the same conditions as those induced for the reaction mixture. Initially, added acrylonitrile decreased the rate of reaction indicating free radical intervention, which is the case in earlier work24,25. Effect of temperature. The kinetics was studied at 6 different temperatures (15, 20, 25, 30, 35 and 40oC) under varying concentrations of KNT, alkali and periodate keeping all other conditions constant. The rate constants (k) of the slow step of the reaction mechanism were obtained from the slopes and intercepts of the plots of 1/kobs versus 1/[KNT] at 4 different temperatures and were used to calculate the activation parameters. The energy of activation corresponding to these constants was evaluated from the Arrhenius plot of lg k versus 1/T (r≥ 0.9904, S ≤0.025) and other activation parameters obtained are tabulated in Table 2. Table 2. Thermodynamic activation parameters for the oxidation of kinetin by DPC in aqueous alkaline medium with respect to the slow step of Scheme 1 (a) Effect of temperature

Temperature (K) 288 298 308 318

k ×102 (s–1) 0.4 0.6 1.0 1.9

703

(b) Activation parameters (Scheme 1)

Parameters Ea (kJ mol–1) ΔH* (kJ mol–1) ΔS* (J K–1 mol–1) ΔG* (kJ mol–1) lg A

Value   73.71   71.24 –43.91   84.31 10.9

(c) Effect of temperature to calculate K1, K2 and K3 for the oxidation of kinetin by DPC in alkaline medium

Temperature (K)

K1 (dm3 mol–1)

K2 × 104 (mol dm–3)

K3 × 10–3 (dm3 mol–1)

288 298 308 318

1.4 2.3 3.5 4.6

1.2 1.6 2.3 3.3

2.83 2.24 1.49 1.05

(d) Thermodynamic parameters using K1, K2 and K3

Thermodynamic parameters ∆H (kJ mol ) ∆S (J K–1 mol–1) ∆G (kJ mol–1) –1

Values based on K1 Values based on K2 Values based on K3   57.98 48.83   –49.14 204.5 91.60 –104.36   –0.453 17.83   –17.23

Thus, from the observed experimental results the rate law for the reaction is given as follows: rate = kobs [DPC]1.0 [KNT]0.69 [OH–]0.23 [IO4–]–0.26.

DISCUSSION The water-soluble copper(III) periodate complex is reported26 to be [Cu(HIO6)2 (OH)2]7–. However, in an aqueous alkaline medium and at a high pH range employed in the study, periodate is unlikely to exist as HIO64– (as present in the complex) as is evident from its involvement in the multiple equilibria27 depending on the pH of the solution. Periodic acid exists as H5IO6 in acid medium and as H3IO62– near pH 7. Hence, under alkaline conditions as employed in this study, the main species are expected to be H3IO62– and H2IO63–. Thus, at the pH employed in this study, the soluble copper(III) periodate complex might be [Cu(OH)2(H3IO6)2]3–, a conclusion also supported by earlier work 2,3,18. The reaction between the diperiodatocuprate (III) complex and kinetin in alkaline medium has the stoichiometry 1:3 (KNT:DPC) with a first order dependence on [DPC] and an apparent order of less than unit order in [substrate], [alkali] and a negative fractional order dependence on [periodate]. No effect of added product was 704

observed. Based on the experimental results, a mechanism is proposed for which all the observed orders in each constituent such as [oxidant], [reductant], [OH–] and [IO4–] may be well accommodated. In most reports18 on DPC oxidation, periodate had a retarding effect and OH– had an increasing effect on the rate of reaction. However, in the present kinetic study, different kinetic results have been obtained. In this study OH– had less than unit order and periodate retarded the rate of reaction with increase in alkalinity (Table 1) can be explained in terms of prevailing equilibrium of formation of [Cu(OH)2(H3IO6)]3– from [Cu(OH)2(H3IO6)(H2IO6)]4– hydrolysis as given in the following equation: [Cu(H3IO6)2]– + [OH–]

K1

[Cu(H2IO6)(H3IO6)]2– + H2O

(2)

Also, decrease in the rate of reaction with increase in [H3IO62–] (Table 1) suggests that equilibrium of copper(III) periodate complex to form monoperiodatocuptrate(III) (MPC) species as given in equation (3) is established. [Cu(H2IO6)(H3IO6)]2– + 2H2O

K2

[Cu(H2IO6)(H2O)2] + [H3IO6]2–

(3)

Such equilibria (2) and (3) have been well documented in literature2,3. Periodate complex such as monoperiodatocuptrate(III) (MPC) is more important in the reaction than the DPC. The inverse fractional order in [H3IO62–] might also be due to this reason. Therefore, MPC might be the main reactive form of the oxidant. The less than unit order in [KNT] presumably results from formation of a complex (C) decomposes slowly in a slow step to form a free radical derived from kinetin. This free radical species further reacts with another molecule of MPC in a fast step to form the products as given in Scheme 1. Scheme 1 Detailed scheme for the oxidation of kinetin by alkaline DPC (III)

  [Cu(H3IO6)2]– + [OH–] [Cu(H2IO6)(H3IO6)]2– + 2H2O H

N

N

K2

[Cu(H2IO6)(H3IO6)]2– + H2O [Cu(H2IO6)(H2O)2] + [H3IO6]2–

H2C N

N

K1

N

O

+ [Cu(H2IO6)(H2O)2]

K3

complex (C) (I)

705

. HN

complex (C)

H N

N

k slow

H2C

+

+

N

N

O

(II) NHOH

N

N

(III) NH+

H N

N

+ Cu 2+ + H2O + H2IO63– + H+

fast

+ [Cu(H2IO6)(H2O)2]

N

N

N

N

(II) +

N (IV)

fast

+ H2O

N

N

N

N (V)

NO fast

+ 2[Cu(H2IO6)(H2O)2]

H N

N

N

N NO2

H

N

fast

+ 2[Cu(H2IO6)(H2O)2]

+ 2Cu 2+ + 2H2IO63– + 4H2O + 2H+

H N

N

N

N

2+ + 2Cu + 2H2IO63– + 2H+ + 3H2O

(VII)

(VI) +

CH 2 O

+ OH



fast

CH2OH O

(II)

(VIII) fast

6H+ + 6OH– ––––→ 6H2O

706

H+

(VI)

N

N

+ N

(V) NO

H N

N

H N

N

NHOH

H N

N

N

+ Cu2+ + H2IO63– + 2H2O

(VI) NH

NHOH

H

Since Scheme 1 is in accordance with the generally well-accepted principle of non-complementary oxidations taking place in sequence of one-electron steps, the reaction between the substrate and the oxidant would afford a radical intermediate. A free radical scavenging experiment revealed such a possibility. This type of radical intermediate has also been observed in earlier work28. A direct plot of kobs versus [KNT] was drawn to characterise the parallel reaction if any along with interaction of oxidant and reductant. However, the plot of kobs versus [KNT] was not linear. Thus, in Scheme 1, the parallel reaction and involvement of 2 mol of kinetin in the complex are excluded. The fractional order with respect to KNT presumably results from the complex formation between MPC and KNT prior to the slow step. Indeed it is to be noted that a plot of 1/kobs versus 1/[KNT] was linear and shows an intercept in agreement with the complex formation which slowly decomposes to form the product. In the rate-determining stage, this monoperiodatocuprate(III) (MPC) combines with molecule of kinetin to give a complex (I) which decomposes in a slow step to give 1H-imidazole pyridine-4-yl amine radical (II), furon-3-yl-methane carbocation (III) and Cu(II) species. Imidazole pyridine 4-yl-amine radical (II) combines with Cu(II) species in a fast step to give 1H-imidazole pyridine 4-yl-amine cation (IV) and Cu(I) species, further it reacts with water molecule in fast step to give 1H-imidazole Npyrimidine 4-yl-hydroxyl amine (V), which further reacts with another mol of MPC in a fast step to give 1H-imidazole 4-nitroso-pyrimidine (VI) and Cu(I) species. In further fast step the 1H-imidazole-4-nitroso pyrimidine (VI) reacts with another mol of MPC to give the product 1H-imidazole-nitro pyrimidine (VII) and Cu(I). In further fast step the furon-3-yl-methane carbocation(II) reacts with MPC molecule to the give the another product furon-3-yl-methanol (VIII). All these results may be interpreted in the form of Scheme 1. Spectroscopic evidence for the complex formation between oxidant and substrate was obtained from UV-vis. spectra of kinetin (5 × 10–4), DPC (5 × 10–5), [OH–] = 0.002 mol dm–3 and a mixture of both. A hypsochromic shift of about 8 nm from 291 to 283 nm in the spectra of DPC was observed. The Michaelis–Menten plot also proved the complex formation between DPC and kinetin. Such a complex between an oxidant has been observed in other studies29. Scheme 1 leads to the following rate law: rate = –

kobs =

d[DPC] dt rate [DPC]

=

=

kK1K2K3[DPC][KNT][OH–] [HIO62–] + K1[OH–][H3IO62–] + K1K2[OH–] + K1K2K3[OH–][KNT] kK1K2K3[KNT][OH–] [HIO62–] + K1[OH–][H3IO62–] + K1K2[OH–] + K1K2K3[OH–][KNT]

(4)

(5)

This explains all the observed kinetic orders of different species. In equation (6) the appearance of [KNT] term both in numerator and denominator explains the observed less than unit order in [kinetin]. Similarly the appearance of [H3IO62–] and [OH–] in the denominator agrees with the observed negative less than unit order 707

[H3IO62–] and [OH–], respectively. This explains all the observed kinetic orders of different species. Rate law (5) can be rearranged into the following form which is suitable for verification: 1

=

kobs

[H3IO62–]

+

kK1K2K3[OH ] –

[H3IO62–] kK2K3[KNT]

+

1 kK3[KNT]

+

1 k

.

(6)

According to equation (6), other conditions being constant, plot of 1/kobs versus 1/[KNT] (r ≥ 0.9998, S ≤ 0.016), 1/kobs versus 1/[OH–] (r ≥ 0.996, S ≤ 0.019) and 1/ kobs versus [H3IO62–] (r ≥ 0.9312, S ≤ 0.017) should be linear and are fond to be so (Fig. 6a, b and c). The slopes and intercepts of these plots lead to the value of K1, K2, K3 and k as 2.3 dm3 mol–1, 1.6 × 10–4 mol dm–3, 2.2 × 103 dm3 mol–1 and 0.60 × 10–2 s–1, respectively. The values of K1 and K2 are in good agreement with those reported earlier in literature23. These constants were used to calculate the rate constant and compared with the experimental kobs values and found to be in reasonable agreement with each other which fortifies Scheme 1. 18

a

16

288 K

(1/kobs) × 10–2 (s)

14 12

293 K

10 298 K

8 6

303 K

4 2 0

0

20

40

60

80

100

120

(1/[KNT]) × 10–2 (dm3 mol–1) 10.0

b

9.0

288 K

(1/kobs) × 10–2 (s)

8.0 7.0 6.0 293 K

5.0 4.0

298 K

3.0 2.0

303 K

1.0 0.0 0

20

40

60

80

(1/[OH–]) × 10–2 (dm3 mol–1)

708

100

120

12.0

c

(1/kobs) × 10–2 (s)

10.0

288 K

8.0 6.0

293 K

4.0

298 K

2.0 0.0

303 K 0

1

2

3

4

5

6

[H3IO62–] × 10–5 (mol dm–3)

Fig. 6. Verification of rate law (5) for the oxidation of KNT by DPC 1/kobs versus 1/[KNT] at 4 different temperatures (a); 1/kobs versus 1/[OH–] at 4 different temperatures (b); 1/kobs versus [H3IO6]2– at 4 different temperatures (c) (conditions as in Table 1)

The negligible effect of ionic strength on the rate explains qualitatively the reaction between one negatively charged ion and neutral molecule, as seen in Scheme 1. The effect of solvent on the rate has been described in details in literature. Increasing the content of t-butyl alcohol in the reaction medium leads to an increased effect on the rate of reaction, which seems to be contrary to the expected reaction between neutral and anionic species in media of lower relative permittivity. However, an increase in the rate of reaction with decreasing relative permittivity may be due to stabilisation of the complex (C) at relative permittivity, which is less solvated than DPC at higher relative permittivity because of its larger size. The thermodynamic parameters for the first, second and third equilibrium steps of Scheme 1 can be evaluated as follows: [KNT], [OH–] and [IO4–] (Table 1) were varied at 4 different temperatures. The plot of 1/kobs versus 1/[KNT] (r ≥ 0.9998, S ≤ 0.016) 1/kobs versus 1/[OH–] (r ≥ 0.996, S ≤ 0.019), and 1/kobs versus [H3IO62–] (r ≥ 0.9312, S ≤ 0.017) should be linear. From the slopes and intercepts, the values of K1, K2 and K3 were calculated at 4 different temperatures and these values are given in Table 2. The vant Hoff plots were made for variation of K1, K2 and K3 with temperatures (lg K1 versus 1/T) (r ≥ 0.957, S ≤ 0.008) (lg K2 versus 1/T) (r ≥ 0.9908, S ≤ 0.006) and (lg K3 versus 1/T) (r ≥ 0.9732, S ≤ 0.006) and the values of enthalpy of reaction ΔH, entropy of reaction ΔS and free energy of reaction ΔG, were calculated for the first, second and third equilibrium steps. These values are given in Table 2. A comparison of the thermodynamic parameters of Scheme 1 with those obtained for the slow step of the reaction shows that these values mainly refer to the rate-limiting step, supporting the fact that the before rate-determining step is fairly fast and involves low activation energy30. The values ΔH* and ΔS* were both favourable for electron transfer processes. The negative value of ΔS* suggests that the intermediate complex is more ordered than the reactant31. The observed modest enthalpy of activation and a higher rate constant for the slow step indicates that the oxidation presumably occurs via an inner-sphere mechanism. This conclusion is supported by earlier observations32,33. 709

The activation parameters for the oxidation of some amino acids by DPC are summarised in Table 3. CONCLUSIONS Among the various species of DPC in alkaline medium, MPC, i.e. [Cu(H2IO6)(H2O)2] is considered as active species for the title reaction. The results indicated that the role of pH in the reaction medium is crucial. Rate constant of slow step and other equilibrium constants involved in the mechanism were evaluated and activation parameters with respect to slow step of reaction were computed. The overall mechanistic sequence described is consistent with product studies, mechanistic and kinetic studies. Appendix According to Scheme 1, rate = –

d[DPC] dt

= k[C] =

k1K1K2K3[DPC][KNT][OH–] [H3IO6]2–

.

(A1)

The total concentration of [DPC]T is given by [DPC]T = [DPC]f + [Cu(H2IO6)(H3IO6)]2– + [Cu(H2IO6)(H2O)2] + [C],

(A2)

where T and f refer to total and free concentrations. [DPC]f =

[DPC]T[H3IO6]2– [H3IO6]2– + K1[H3IO6]2–[OH–] + K1K2[OH–] + K1K2K3[OH–][KNT]

.

Similarly,

[KNT]T = [KNT]f + C

= [KNT]f + K1K2K3[DPC]f[KNT]f[OH–]f/[H3IO6]2–



= [KNT]f[1 + K1K2K3[DPC]f[OH–]f/[H3IO6]2–].

In view of low concentration of [DPC] and [H3IO6]2– second term can be neglected [KNT]T = [KNT]f.

(A3)

Similarly, [OH–]T = [OH–]f + [Cu(H2IO6)(H3IO6)]2– + [Cu(H2IO6)(H2O)2]      = [OH–]f + K1[OH–][DPC] + K1K2[DPC][OH–]/[H3IO6]2–. In view of low concentration of [DPC] and [H3IO6]2– used [OH–]T = [OH–]f.

(A4)

Substituting the values of [DPC]f, [KNT]f and [OH–]f in equation (A1) and omitting subscripts, we get kobs =

710

rate [DPC]

=

kK1K2K3[KNT][OH–] [H3IO62–] + K1[OH–][H3IO62–] + K1K2[OH–] + K1K2K3[OH–][KNT]

.

ACKNOWLEDGEMENTs One of the authors (M. T. Asma) thanks the Director/Principal Dr. M. N. S. Rao, and Dr. A. A. Kittur and other colleagues of Department of Chemistry, SDMCET, Dharwad, for the co-operation during research work. REFERENCES   1. D. W. S. MOK, M. MOK: Cytokinins: Chemistry, Activity and Function. C RC Press Inc., Boca Raton, 1994.   2. B. REDDY, B. SETHURAM, T. N. RAO: Kinetics of Oxidation of Benzaldehydes by Copper(III) in t-butanol Water Medium. Indian J Chem, 23A, 593 (1984).   3. R. B. CHOUGALE, G. A. HIREMATH, S. T. NANDIBEWOOR: Kinetics and Mechanism of Oxidation of L-alanine by Alkaline Permanganate. Pol J Chem, 71, 1471 (1997).   4. A. KUMAR, P. KUMAR, P. RAMAMURTHY: Kinetics of Oxidation of Glycine and Related Substrates by Diperiodatoargentate(III). Polyhedron, 18 (6), 773 (1999).   5. A. KUMAR, P. KUMAR: Kinetics and Mechanism of Oxidation of Nitrilotriacetic Acid by Diperiodatoargentate(III). J Phys Org Chem, 12 (2), 79 (1999).   6. A. KUMAR, A. VAISHALI, P. RAMAMURTHY: Kinetics and Mechanism of Oxidation of Ethylenediamine and Related Compounds by Diperiodatoargentate(III) Ion. Int J Chem Kinet, 32 (5), 286 (2000).   7. H. SHAN, J. QIAN, M. Z. GAO, S. G. SHEN, H. W. SUN: Kinetics and Mechanism of Oxidation of n-propanolamine by Dihydroxydiperiodatonickelate(IV) in Alkaline Medium. Turk J Chem, 28, 9 (2004).   8. W. NIU, Y. ZHU Y, K. HU, C. TONG, H. YANG: Kinetics of Oxidation of SCN − by Diperiodatocuprate(III) (DPC) in Alkaline Medium. Int J Chem Kinet, 28 (12), 899 (1996).   9. V. RAGAINI, L. FORNI: Kinetic Study and Catalyst Structural Analysis for the Dehydrogenationhydrogenolysis of n-pentane over Ru-Al2O3catalysts. Kinet Catal, 37, 339 (1975). 10. M. G. R. REDDY, B. SETHURAM, T. N. RAO: Effect of Copper(II) on Kinetics and Mechanism of Silver(I) Catalysed Oxidation of Some Amino acids by Peroxydisulfate ion in Aqueous Medium. Indian J Chem A, 16, 313 (1978). 11. K. D. KARLIN, Y. GULTNEH: In: Progress in Inorganic Chemistry (Ed. S. J. Lipard). Wiley, New York, 1997, p. 35. 12. W. B. TOLMAN: Making and Breaking the Dioxygen O–O bond: New Insights from Studies of Synthetic Copper Complexes. Acc Chem Res, 30 (6), 227 (1997). 13. K. N. KITAJIMA, Y. MOROOKA: Copper-dioxygen Complexes. Inorganic and Bioinorganic Perspectives. Chem Rev, 94 (3), 737 (1994). 14. K. D. KARLIN, S. KADERLI, A. D. ZUBERBUHLER: Kinetics and Thermodynamics of Copper(I)/ Dioxygen Interaction. Acc Chem Res, 30 (3), 139 (1997). 15. J. L. PIERE: One Electron at Time Oxidations and Enzymatic Paradigms: From Metallic to Nonmetallic Redox Centers. Chem Soc Rev, 29, 251 (2000). 16. J. PEISACH, P. ALSEN, W. E. BLUMBERG: The Biochemistry of Copper. Academic Press, New York, 1996, p. 49. 17. B. SETHURAM: Some Aspects of Electron Transfer Reactions Involving Organic Molecules. Allied Publishers (P) Ltd., New Delhi, 2003, p. 73. 18. S. NADIMPALLI, J. PADMAVATHY, K. K. M. YUSUFF: Determination of the Nature of the Diperiodatocuprate(III) Species in Aqueous Alkaline Medium through a Kinetic and Mechanistic Study on the Oxidation of Iodide Ion. Trans Met Chem, 26 (3), 315 (2001). 19. P. K. JAISWAL, K. L. YADAV: Determination of Sugars and Organic Acids with Periodato Complex of Cu(III). Indian J Chem, 11, 837 (1973).

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20. C. P. MURTHY, B. SETHURAM, N. RAO: Kinetics of Oxidation of Some Alcohols by Diperiodatocuprate(III) in Alkaline Medium. Z Phys Chem, 262, 336 (1981). 21. G. H. JEFFERY, J. BASSETT, J. MENDHAM, R. C. DENNY: Vogel’s Text Book of Quantitative Chemical Analysis. 5th ed. ELBS Longman, Essex UK, 1996 p. 455. 22. G. P. PANIGRAHI, P. K. MISRO: Kinetics and Mechanism of Osmium (VIII)-catalyzed Oxidation of Aromatic Aldehydes by Sodium Periodate. Ind J Chem A, 15, 1066 (1977). 23. G. C. HIREMATH, R. M. MULLA, S. T. NANDIBEWOOR: Mechanistic Study of the Oxidation of Isonicotinate Ion by Diperiodatocuprate(III) in Aqueous Alkaline Medium. J Chem Res, 3 (5), 197 (2005). 24. I. M. KOLTHOFF, E. J. MEEHAN, E. M. CARR: Mechanism of Initiation of Emulsion Polymerization by Persulfate. J Am Chem Soc, 75 (6), 1439 (1953). 25. S. BHATTACHARYA, P. BANERJEE: Kinetic Studies on the Electron Transfer between Azide and Nickel(IV) Oxime Imine Complexes in Aqueous Solution. Bull Chem Soc Japan, 69 (12), 3475 (1996). 26. K. B. REDDY, B. SETURAM, T. N. RAO: Photon Cross Sections in Copper, Platinum and Gold at 81 keV. Z Phys Chem, 268, 706 (1987). 27. J. C. BAILAR, H. J. EMELEUS, S. R. NYHOLM, A. F. TROTMAN-DIKENSON: Comprehensive Inorganic Chemistry. Pergamon Press, Oxford, Vol. 2, 1975, p. 1456. 28. S. A. FAROKHI, S. T. NANDIBEWOOR: Kinetic, Mechanistic and Spectral Studies for the Oxidation of Sulfanilic Acid by Alkaline Hexacyanoferrate(III). Tetrahedron, 59 (38), 7595 (2003). 29. C. ORVIG, M. J. ABRAMS (Eds): Medicinal Inorganic Chemistry. Special Issue of Chem Rev, 99 (9), 2201 (1999). 30. K. S. RANGAPPA, M. P. RAGHVENDRA, D. S. MAHADEVAPPA, D. CHANNEGOUDA: Sodium N-Chlorobenzenesulfonamide as a Selective Oxidant for Hexosamines in Alkaline Medium: A Kinetic and Mechanistic Study. J Org Chem, 63 (3), 531 (1998). 31. A. WEISSBERGER, E. S. LEWIS (Eds): Investigations of Rates and Mechanism of Reaction in Techniques of Chemistry. Wiley, New York, 1974, p. 421. 32. F. M. MOORE, K. W. HICKS: An Investigation of Lewis Base-monoiodoboranes and Their Reactions with Group V Donors in Solution. J Inorg Nucl Chem, 38, 379 (1976). 33. D. C. HIREMATH, K. T. SIRSALMATH, S. T. NANDIBEWOOR: Osmium(VIII)/Ruthenium(III) Catalysed Oxidation of L-lysine by Diperiodatocuprate(III) in Aqueous Alkaline Medium: A Comparative Mechanistic Approach by Stopped Flow Technique. Catal Lett, 122 (2), 144 (2008). Received 19 September 2011 Revised 26 October 2011

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Oxidation Communications 37, No 3, 713–721 (2014) Antioxidant and antimicrobial activity in biological systems

Comparative Studies on Polyphenols Profile and Antioxidative Activity of Some Berry Fruits N. Gougoulias Department of Plant Production, Technological Educational Institute of Larissa, Larissa, Greece E-mail: [email protected] ABSTRACT The fruits of raspberry, blackcurrant, gooseberry, redcurrant, and blueberry were evaluated for their content of total phenols (TP) and phenol fractions – nonflavonoid and flavonoid phenols (NFP, FP), total flavanols (F-3-ols), non-tannin and tannin phenols (NTanP, TanP), and total anthocyanins (АС ). The antiradical (DPPH●) activity and ferric reducing antioxidant power (FRAP) of the ethanol fruit extracts were assayed as well. It has been found that the fruits studied are characterised by high content of phenol compounds and manifest pronounced antioxidant activity. The content of TP and phenol fractions varies depending on the fruit genotype and the antioxidant activity correlates with TP and phenol fractions. The highest values of TP content were obtained from blueberry fruits – 3500 mg/kg GAE, and the lowest 1350 mg/kg GAE – from the red currant. The antiradical activity of these fruits is 6.18 and 5.13 mM TEAC, and the ferric reducing antioxidant power – 4.80 and 2.11 mM AAE, respectively. FRAP value was highly correlated with anthocyanins content (r2 = 0.9978), whereas TP content was relatively less correlated with DPPH● (r2 = 0.5524). Keywords: berry fruits, total phenols, phenol groups, DPPH● activity, FRAP. aims and background Many studies show that plant phenolics, due to their high diversity, are still not enough studied and assayed in terms of their nutritive and health benefits1,2. The intake of fruits, vegetables and cereals is highly recommended by the dietologits against current diseases, which is closely related to their higher contents of many natural bioactive nutrients including phenol3,4. It is known that the consumption of phenol compounds with human food exerts strong antioxidant properties in the human organism and neutralise the active free oxygen- and nitrogen-containing radicals (ROS, RNS). Being strong oxidisers they damage the biological molecules and induce the occurrence of many diseases5–8. The tendency for growing more berry fruits shows permanent 713

increase in the last decade. These fruits are characterised by high content of phenol compounds which are strong natural antioxidants9–11. According to some researchers, the nutrition effects of these fruits, used in traditional medicine, are due not only to their phenol content, but to the complex character of the functional biologically active substances, the high content of vitamin C, vitamin E, carotenoids, etc., which in combination with phenols form ‘highly efficient antioxidant system’ for cell protection against the action of free radicals. The content and the qualitative composition of phenol compounds in small berry fruits depend on the genotype, natural climatic and agrotechnical conditions which imposes the necessity to evaluate their potential use as food or raw material for the purposes of pharmaceutical and food industry12,13. The aim of the present study is to determine the content of phenol compounds, their major fractions and the antioxidant activity of some small berry fruits, wildgrown and cultivated in one and the same region. experimental Material. The this study 5 small berry fruit cultivars – raspberry (Rubus idaeus), blackcurrant (Ribus nigrum), redcurrant (Ribus rubrum), gooseberry (Ribus uva-crispa) and blueberry (Vaccinium mirtyllus) were studied. The raspberry, blackcurrant, redcurrant and gooseberry were cultivated in garden conditions in the village of Boikovo – Plovdiv district, and the blueberry were harvested from natural habitat at 1200–1300 m a.s.l., near the same village. The fruits of the selected cultivars were harvested at complete physiological maturation stage in the summer months of 2007–2008. Extract preparation. The extracts of the studied berry fruits were obtained after single treatment of 10 g sample with 80% ethanol after 1-h storage at dark and room temperature. After centrifugation the extracts were brought to 20 ml with aqueous ethanol and used for further chemical analysis14. Determination of total polyphenols content. Total polyphenols (TP) contents were determined using the Folin–Ciocalteu reagent according to the method of Singleton and Rossi15 and were expressed as gallic acid equivalent (GAE). Determination of phenol fractions. The non-flavonoid phenols (NFP) were determined with the F.–C. reagent after the removal of flavonoid phenols (FP) with formaldehyde according to the method of Kramling16. FP content was determined as the difference between the content of TP and NFP. The separation of tannin (ТanP) from non-tannin phenols (NTanP) was carried out by removing the tannin fraction with quinine sulphate solution and determination of the non-tannin phenols in the liquid phase. This allows the representation of all fractions as GAE for comparison purposes17,18 The total flavanols (F-3-ols, catechins and procyanidins) were assayed using p-DMACA reagent after the method of Li et al.19 and were presented as catechin equivalent (CE). 714

Anthocyanins were determined using the method of Ribereau-Gayon and Stonestred, modified by Burns et al.20 Determination of antioxidant activity. The antiradical activity of the methanol extracts was determined according to the method of Brand-Williams et al.21 using the stable free radical 2,2′-diphenyl-1-picrylhydrazyl (DPPH). The activity was evaluated in µmol DPPH/g sample and in Trolox equivalent (synthetic vitamin E) as well. The ferric reducing antioxidant power (FRAP) was evaluated according to the method of Benzie et al.22 and was expressed in µmol FRAP reagent/g grape. The activity was also presented as ascorbic acid equivalent (AAE) in mg% per fresh matter. Statistical analyses were performed using a statistical program MINITAB (Ref. 23). The results are means of three parallel samples deviation (±sd). Student test at a confidence level p 0.05). 717

Table 7. Anthocyanins content in the fruits under study (cyanidin-3-glucoside)

No 1 2 3 4 5

Fruit type raspberry blackberry redcurrant gooseberry blueberry

μmol/kg 1114.9 1359.5 1039.9 1089.0 1544.1

mg/kg 500.6 610.4 466.9 489.0 693.3

The determination of the antiradical activity of ethanol fruit extracts was carried out by using the stable free radical DPPH•, through its decolourisation in the presence of polyphenols. The antiradical activity is expressed as μmols DPPH/g fresh matter and as Trolox equivalent, which is a synthetic vitamin E (Table 8). The antiradical activity of the fruits varies from 1.42 to 1.71 μmol/DPPH/g fresh matter. The blueberry fruits are distinguished by the highest antiradical capacity and that of the redcurrant – by the lowest one. The antiradical activity was correlated to the amount of ТР, FР and anthocyanins and the following significance levels were obtained (r2) – 0.5524, 0.8072, and 0.6606, respectively. Table 8. Antiradical (DPPH•) activity of the fruits under study

No 1 2 3 4 5

Fruit type raspberry blackberry redcurrant gooseberry blueberry

μmol DPPH/g 1.50± 0.09 1.60±0.11 1.42±0.08 1.55±0.10 1.71±0.12

mM TEAC/kg 5.38 5.81 5.13 5.62 6.18

The results of the FRAP antioxidative assay of the fruits reveal that the blueberry fruits exert the highest activity – 27.4 μmol FRAP/g fresh matter and those of the redcurrant – the lowest – 4. 8 μmol FRAP/g (Table 9). The antioxidative activity of the fruits, presented as ascorbic acid equivalent (mM ААЕ ), ranges from 5.13 to 8.18 mM AAE, the highest manifested by the blueberry fruits and the weakest – by the redcurrant fruits. Determining the correlations between the antioxidant FRAP activity of the fruits and the contents of ТР, FР and anthocyanins, the following correlation coefficients (r2) were obtained: 0.7030, 0.8772 and 0.9989, respectively. Table 9. Ferric reducing antioxidant power (FRAP) of the fruits under study

No 1 2 3 4 5

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Fruit type raspberry blackberry redcurrant gooseberry blueberry

μmol FRAP / g 5.01±0.12 7.30±0.14 4.28±0.10 4.63±0.10 9.65±0.17

mM ААЕ 2.49 3.62 2.27 2.11 4.80

The blueberry and blackberry fruits belong to the group of small berry fruits – wild-grown and cultivated fruit species, which as described in literature are abundant in phenol compounds11–13,26. The contents of ТР, as well as of the individual phenol compounds in the fruits of the separate cultivars and genotypes are strongly affected by the genotype, preharvest conditions, maturity, and postharvest handling, soil-climatic and other conditions10,12,27. The variety differences are sometimes so strongly expressed that the fruits of one cultivar are comparable to the fruits of another cultivar. Here should be mentioned also the effect of climatic conditions, the altitude, soil characteristics and agrotechnical handling in growing of some cultivated species. This was one of the reasons to study the fruits of plant species grown under similar garden or natural habitat conditions in the mountainous region of the village of Boikovo-the Plovdiv district. Benvenuti et al.13 established different TP content in the fruits of the genera Rubus, Ribis and Аroniа. Analysing 26 types of fruits from different forest and fruit plants they have found the dominance of one and the same group of compounds in some of them, although they have different family and genotype, and in other, different in quality and quantity compounds of one family, but different genera. They have established that in the raspberry of the Rosaceae family, the genus Rubus еlagotannins are dominant while they are second in the strawberry of the same family and the genus Fragaria. Our studies confirm the results of many authors which have established that the wild and cultivated berry fruits are an excellent source of polyphenol compounds, which together with other biologically active compounds as vitamin C, vitamin E and carotenoids cane be regarded as a source of natural antioxidants9,28–30. The increased interest towards the use of these fruits in the last years is provoked by the large number of researches demonstrating their role in reducing or eliminating the development of many diseases such as cardiovascular, arteriosclerosis, cancer ands others induced by the oxidative stress and the action of oxygen- and nitrogen-containing active free radicals (ROS, RNS) in the cells of the different organs and tissues in the human body7. In addition to the pharmaceutical industry which uses in many countries berry fruits for the preparation of various food supplements with healthy effects, the food technology also is directed to the preparation and standartisation of beverages on the basis of these fruits, rich in polyphenols and natural antioxidants, of great significance for the human organism10,31,32. CONCLUSIONS The study on total phenols content, group composition and antioxidant activity of the fruits of raspberry, blackberry, gooseberry, redcurrant and blueberry, grown under comparative climatic and soil conditions show that they are rich in phenol compounds and exert strong antioxidative activity as compared with other fruits and vegetables. Among the studies fruits, the wild-grown blueberry and the cultivated blackberry accumulate the highest amount of TP, NFP, FP, AC, and TanP in their fruitss. The in 719

vitro tests reveal their antiradical (DPPH), and antioxidant (FRAP) capacity, as 6.18 and 5.13 mM ТЕАС, and 4.80 and 2.11 mM AAE. The small berry crops, such as raspberry, blackberry, gooseberry, redcurrant and blueberry are grown in the area of Bulgaria and Greece and their fruits are source of bioactive nutritive chemical compounds which can be included in the composition of functional foods and beverages for human being. REFERENCES   1. L. BRAVO: Polyphenols: Chemistry, Dietary Sources, Metabolism and Nutritional Significance. Nutr Rev, 56 (11), 317 (1998).   2. J. J. Machiex, A. FLEURIET, J. Billot: Fruit Phenolics. CRC Press, Boca Raton, FL, 1990.   3. C. A. RICE-EVANS, N. J. MILLER: Antioxidants: The Case for Fruit and Vegetables in the Diet. British Food J, 97 (9), 35 (1995).   4. F. Shahidi: Natural Antioxidants: Chemistry, Health Effects and Applications (Ed. F. Shahidi). AOCS Press, Champaign Illinois, 1997, 1–11.   5. B. J. F. Hudson: Food Antioxidants. Elsevier Applied Science, London, 1990.   6. M. J. Morello, F. Shahidi, Chi-Tang Ho (Eds): Free Radical in Food. Chemistry, Nutrition, and Health Effects. Am. Chem. Society, Washington, DC, 2002.   7. B. Halliwell, J. M. C. Gutteridge: Free Radicals in Biology and Medicine. 2nd ed. Clarendon Press, Oxford, 1989, 1–21.   8. J. Pokorny, N. Yanishlieva, M. Gordon: Antioxidants in Food. CSC Press, England, 2001.   9. V. Kondakova, I. Tsvetkov, R, Batchvarova, I. Badjakov, T. Dzhambazova, S. Slavov: Phenol Compounds – Qualitative Index in Small Fruits. Biotechnol Biotechn EQ, 23, 1444 (2009). 10. I. Badjakov, M. Nikolova, R. Gevrenova, V. Kondakova, E. Todorovska, A. Atanassov: Bioactive Compounds in Small Fruits and Their Influence on Human Health. Biotechnol Biotecnol EQ, 22 (1), 581 (2008). 11. H. Wang, G. Cao, R. L. Prior: Oxygen Radical Absorbing Capacity of Anthocyanins. J Agric Food Chem, 45, 304 (1997). 12. S. Y. Wang: Antioxidant Capacity and Phenolic Content of Berry Fruits as Affected by Genotype, Preharvest Conditions, Maturity, and Postharvest Handling. In: Berry Fruit (Ed. Y. Zhao). CRC Press, 2007, 148–178. 13. S. Benvenuti, F. Pellati, M. Melegari, D. Bertelli: Polyphenols, Anthocyanins, Ascorbic Acid and Radical Scavenging Activity of Rubis, Ribes and Aronia. J Food Sci, 69 (3), 164 (2004). 14. J. Kanner, E. Frankel, R. Granit, B. German, J. Kinsella: Natural Antioxidants in Grapes and Wines. J Agric Food Chem, 42, 64 (1994). 15. V. L. Singleton, S. A. Rossi: Colorimetry of Total Phenolics with Phosphomolibdic-phosphotungstic Acid Reagents. J Enol Viticult, 16, 144 (1965). 16. T. E. Kramling, V. L. Singleton: An Estimate of the Nonflavonoids Phenolics in Wines. Am J Enol Viticult, 20, 86 (1969). 17. A. Pirie, M. G. Mullins: Changes in Anthocyanin and Phenolic Content Grapewine Leaf and Fruit Tissues Treated with Sucrose, Nitrate and Abcisic acid. Plant Physiol, 58, 462 (1976). 18. A. Brugirard, J. Tavernier: Les matieres tanoides dans les cidres et les poires. Annal Technol, ΙΙΙ, 311 (1952). 19. Y.-G. Li, G. Tanner, P. Larkin: The DMACA–HCl Protocol and the Threshold Proantocyanidin Content for Bloat Safety in Forage Legumes. J Sci Food Agric, 70, 89 (1996). 20. J. BURNS, P. GARDNER, S. O’NEIL, S. Crawford, I. MORECROFT, D. B. McPHAIL, C. LISTER, D. MATTHEWS, M. R. McLEAN, M. E. S. LEAN, G. G. DUTHIE, A. CROZIER:

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Relationship among Antioxidans Activity Vasodilation Capacitty and Phenolic Content of Red Wines. J Agric Food Chem, 48, 220 (2000). 21. W. Brand-Williams, M. E. Cuvellier, C. Berset: Use of a Free Radical Method to Evaluate Antioxidant Activity. Lebensm Wiss Technol, 28, 25 (1995). 22. I. F. F. Benzie, J. J. Strain: Ferric Reducing Antioxidant Power Assay. Method Enzymol, 299, 15 (1999). 23. B. F. RYAN, B. L. JOINERAND, J. D. CRYER: MINITAB Handbook. Update for Release 14, 15th ed, 2005. 24. S. Hakkinen: Flavonols and Phenolic Acids in Berries und Berry Products. Kuopio Univ. Publ. D. Med. Sci., Kuopio, 2000, p. 221. 25. V. Dragovic-Uzelac, b. Levaj, D. Bursac, S. Pedisic, I. Radojcic, A. Bisko: Total Phenolics and Antioxidant Capacity Assays of Selected Fruits. Agr Conspectus Scientificus, 72 (4), 279 (2007). 26. J. B. Harborne: Plant Phenolycs. In: Encyclopedia of Plant Physiology. Vol. 8. Secondery Plant Products. Springer, Berlin , 1980, 329–402. 27. K. Manning: Soft Fruit. In: Biochemistry of Fruit Ripening (Eds G. B. Seymour, J. E. Taylor, G. A. Tucker). Chapman and Hall, London, 1993. 28. J. Beattie, A. Crozier, G. G. Duthie: Potential Health Benefits of Berries. Curr Nutr Food Sci, 1, 71 (2005). 29. E. Sicora, E. Cieslik, K. Topolska: The Sources of Natural Antioxidants. Acta Sci Pol Technol Aliment,7 (1), 5 (2008). 30. O. Yildiz, S. P. Eyduran: Functional Components of Berry Fruits and Their Usage in Food Technologies. Afr J Agric Res, 4 (5), 422 (2009). 31. M. Sojka, S. Guyot, K. Kolodziejczyk, B. Krol, A. Baron: Composition and Properties of Purified Phenolics Preparations Obtained from an Extract of Industrial Blackcurrant (Ribes nigrum L. ) pomace. J Horticult Sci Biotechnol, ISAFRUIT Special Issue, 100 (2009). 32. N. Seeram, M. Aviram, Y. Zhang, S. Henning, L. Feng, M. Dreher, D. Heber: Comparison of Antioxidant Potency of Commonly Consumed Polyphenol-rich Beverages in the United States. J Agric Food Chem, 56, 1415 (2008). Received 23 July 2010 Revised 18 September 2010

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Oxidation Communications 37, No 3, 722–732 (2014) Antioxidant and antimicrobial activity in biological systems

Polyphenol Contents and Antioxidant Activity of Taraxacum officinale from Some Regions in Thessaly, Greece N. Gougoulias Department of Plant Production, Technological Educational Institute of Larissa, Larissa, Greece E-mail: [email protected] ABSTRACT Methanol extracts of leaf mass of Dandelion (Taraxacum officinale) grown in 7 regions in Thessaly, Greece: Farkadona – Trikala; Trikala; Tyrnavos – Larissa; Terpsithea – Larissa; Kalogriana – Karditsa; Karditsa and Nea Anchialos – Magnesia were screened for total phenols (TP), flavonoid phenols (FP), nonflavonoid phenols (NFP) and flavanols-3-ols (F-3-ols) and antiradical activity. The antiradical activity was evaluated using the stable free radicals DPPH and ABTS, and the FRAP reagent. The correlations between ТР and FP from the one side and the antiradical and ferric reducing power, from the other were, studied. ТР, NTP and FP in the methanol extracts vary from 14.70 tо 17.93, from 3.33 tо 6.60, and from 9.68 tо 12.05 mg DAE/ g dw, respectively. The antiradical activity determined with DPPH, ABTS and FRAP ranges from 8.12 tо 12.24, from 16.28 to 18.99, and from от 3.32 to 5.46 µg TEAC/g dw, respectively. The comparison of the values of IC50 of the methanol extracts assayed by the three methods revealed that the dandelion grown in the region of Tyrnavos – Larissa is distinguished by the highest content of TP and antioxidant activity. The correlation between the results for polyphenols fractions, antiradical activity and ferric reducing power and some soil-climatic characteristics of the regions is discussed. Keywords: Taraxacum officinale, polyphenols, DPPH, ABTS, FRAP. AIMS AND BACKGROUND The plant Dandеlion (Taraxacum officinale W e b) is known for a long time in the traditional medicine in China, India, Russia and many other countries as a remedy against various liver and stomach diseases, as diuretic, laxative and tonic agent. Its generic name Taraxacum originates from the Greek word ‘taraxos’ meaning ‘disorder’ and ‘akos’ meaning ‘remedy’. Nowadays many species of this genera and particularly Dandelion are studied intensively due to their popularity in the Complementary and 722

Alternative System of medicine (CAM), Ayurveda, Sddha, the traditional Chinese medicine and sources of natural antioxidants for food technology. A number of bioactive chemical compounds have been identified in many species of the Taraxacum genera and probably due to their combination they exert positive effects at treating of many diseases. Today the extracts from the whole plants, their organs and parts isolated and pharmaceutical preparations from dandelion are used for treatment at stomach and liver problems, against high temperature, kidney diseases, various forms of diabetes, eye problems, breast, lung and skin cancer, etc.1–4 It has been found that extracts from many fruits, vegetables, seeds, kernels, plant spices, herbs and medicinal plants possess antimicrobial, anti-inflammatory, antivirus, anticancer as well as strong antioxidant properties. The strong antioxidant effects of plant foods and spices are related to the high content of natural antioxidants such as vitamins C, E, carotenoids, polyphenols and many other compounds5–7. Dandelion is characterised by a great diversity and high content of various phenol structures2,8,9. Many of their health effects are associated with their ability to scavenge and remove the harmful free oxygen- and nitrogen-containing radicals (ROS, RNS), to inhibit the oxidation of major biological molecules – proteins, lipids, nucleic acids. The rapid production of free radicals can lead to oxidative damage of the cell membranes, to reducing the activity of the cell enzyme and genetic apparatus which cause the occurrence of many diseases10–12. The biosynthesis of polyphenols depends not only on genetic, but on many environmental factors and the microclimatic conditions of plants growth13,14. In view of this the purpose of the present study was to: to determine polyphenols content and evaluated their antioxidant activity of the surface part of Taraxacum officinale growing in 7 regions of Thessaly – Greece and due to its wide use in the Greek culinary and medicine to establish and use new and promising sources for the preparation of natural antioxidants. EXPERIMENTAL Plant material. The plant material consisted of the leaf mass of dandelion grown in 7 natural areas not fertilised with mineral fertilisers in Thessaly – Greece: Farkadona – Trikala; Trikala; Tyrnavos – Larissa; Terpsithea – Larissa; Kalogriana – Karditsa; Karditsa and Nea Anchialos – Magnesia. The regions differ in some soil-climatic characteristics and have same altitude, except Nea Anchialos – Magnesia (Table 1). In Table 1 are shown the mean results of 10-year observation (1995–2005), for April, May and June, which are characteristic for the spring vegetation of dandelion. The leaf mass was gathered in April–May 2010 in the phase of active growth. The samples were dark-dried, at room temperature, finely ground and kept at 4oC in dark until tested. Some of the elements of the soil-climatic characteristics of the regions are listed in Table 1.

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Table 1. Soil and climatic characteristics of the studied regions

Region

Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos Magnesia

Аltitude (m) 100 120 100 150   80 110   15

Soil organic nitrogenmatter (%) available (mg/kg) 9.14 290 1.36   79 1.27 144 0.74 232 11.2 106 1.10   85 1.40 205

рН 6.75 6.59 7.70 8.13 7.51 6.60 7.00

Climate rainfall T (oC) max. mean (mm/m2) 23.15 23.15 24.95 23.18 31.20 31.20 21.53

45.10 45.10 27.65 36.15 56.65 56.65 45.10

Preparation of the methanol extracts. 500 mg of the finely ground sample were 2-fold extracted with 20 ml 80% aqueous methanol. At the first treatment the sample was incubated for 24 h in the extragent at stirring and the second one – continued for 2 h at stirring at ambient temperature. The extract was collected after centrifugation or filtration and the volume was made up to 50 ml with 80% aqueous methanol. DETERMINATION OF POLYPEHNOLS

Total polyphenols (TP). The amount of total polyphenols (TP) was determined with the Folin–Ciocalteu (F.–C.) reagent according to the method of Singleton and Rossi15 using the microvariant proposed by Baderschneider et al.16, and were expressed as gallic acid equivalent (GAE) in mg/g dry weight. Nonflavonoid phenols (NFP). The content of NFP was determined with the F.–C. reagent after removing the flavonoid phenols (FP) with formaldehyde according to the method of Kramling17 and was expressed as gallic acid equivalent (GAE) in mg/g dry weight. Flavonoid phenols (FP). Flavanoid phenols were determined as a difference between the content of total phenols (TP) and nonflavonoid phenols (NFP). Their amount was evaluated as gallic acid equivalent in mg/g dry weight. Total flavanols (F-3-ols). The amount of total flavanols (catechins and procyanidins) was determined using р-dimethylaminocinnamaldehyde (p-DMACA) reagent after the method of Li et al.18 and was presented as catechin equivalent (CE), in μg/g dry weight. DETERMINATION OF ANTIOXIDANT ACTIVITY

DPPH• assay. The radical scavenging activities by antioxidants in the dandelion extracts were evaluated using the stable free radical 2,2′-diphenyl-1-pycrylhydrazyl radical (DPPH•), as a reagent, according to the method of Brand–Williams19. The

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activity was expressed in μmol DPPH•/g dry weight, as well as in mg/g dry matter and μmol Trolox (synthetic vitamin Е )/g dry weight. ABTS•+ assay. The ABTS•+ radical scavenging activity of the extracts was determined by bleaching the colour of the stable free cation ABTS•+ (2,2-аzinobis-(3ethylbenzothiazolin-6-sulphonic acid ) using the method of Re et al.20, and the results were expressed as μmol Trolox (TAEC)/g dry weight. Ferric reducing antioxidant power assay (with FRAP reagent). The ferric reducing antioxidant power (FRAP) of the dandelion extracts was evaluated according to the method of Benzie et al.21 and the results were expressed as µmol FRAP reagent/g dry weight. The activity was also presented as a Trolox equivalent (TEAC) and ascorbic acid equivalent (AAE) in μmol/ g dry weight. Determination of the inhibition coefficient (IC50). The inhibition coefficient (IC50), represents 50% reduction in the colour intensity of the DPPH and ABTS radicals by the total phenols (mg/g) in the studied extracts after plotting the dependence of the TP content on the bleaching of DPPH• and ABTS•+ solutions. The inhibition coefficient (IC50) was calculated using the following equation: % inhibition = [(E0 – Ex)/E0 ] × 100,

where Е0 is the extinction of the radical solution before the reaction and Ех – after polyphenols addition22. Statistical analyses were performed by the use of statistical program MINTAB (Ref. 23). Data were reported as mean arithmetic for at least three replications and (±sd) deviation. The correlation coefficients (R2) were determined using dispersion analysis. RESULTS AND DISCUSSION Depending on the areas conditions of dandelion growth, the content of total phenols (ТР) in the extracts varies from 14.57 tо 17.93 mg GAE/g dw (Table 2). The samples from the Tyrnavos region are characterised by the highest value, and those from Nea Anchialos – Magnesia – by the lowest TP concentration, and the difference between them amounts to more than 20% (р Tyrnavo > Kalogriana > Farkadona > Terpsithea > Karditsa > Nea Anchialos – Magnesia. Table 3. Сontent of nonflavonoid phenols (NFР) in methanol extracts of dandelion leaf mass (gallic acid equivalent – GAE)

No 1 2 3 4 5 6 7

Region Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos – Magnesia

mg/g dw 5.20 ± 0.07 6.60 ± 0.09 5.88 ± 0.05 5.14 ± 0.06 5.55 ± 0.07 4.58 ± 0.06 3.33 ± 0.05

µmol/ml extract 0.277 0.351 0.313 0.273 0.295 0.244 0.177

The content of flavonoid phenols (FP) in the extracts under study is within the range 9.68 tо 12.05 mg GAE/mg dw (Table 4), and the difference between them is 15.90%. The FP content represents 68.5% of the TP amount. The majority of the analysed samples contain similar values of FP. The differences in the FР in the extracts from the Dandelion leaves in the regions of Trikala, Tyrnavos, Kalogriana, Karditsa and Karditsa do not differ substantially (р>0.05).

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Table 4. Content of flavonoid phenols (FР) in methanol extracts of dandelion leaf mass (gallic acid equivalent – GAE)

No 1 2 3 4 5 6 7

Region Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos – Magnesia

mg/g dw 10.36 ± 0.14 11.46 ± 0.16 12.05 ± 0.15   9.68 ± 0.14 11.36 ± 0.16 11.15 ± 0.18 11.24 ± 0.15

µmol/ml extract 0.551 0.610 0.641 0.515 0.604 0.593 0.598

The fraction of flavanols (F-3-ols) in the methanol extracts of the leaf mass varies from 93 tо 144 μg CE/g dw (Table 5). Relative to TP its amount is the lowest – on average smaller than 0.1%. Although their small amounts the total flavanols in the leaf mass of the analysed species from the 7 regions differ and the differences varies from 13 tо 35%. The lowest content of flavanols (on average 103 μg/g) was observed in the leaves of dandelion grown in the regions of Nea Anchialos – Magnesia and Terpsithea whereas the extracts of dandelion grown in the regions of Tyrnavos, Kalogriana and Trikala showed the highest flavanol content (on average 141 µg/g). Table 5. Сontent of total flavanols (F-3-ols) in methanol extracts of dandelion leaf mass (catechin equivalent – СЕ )

No 1 2 3 4 5 6 7

Region Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos – Magnesia

µg/g dw 124 ± 1.6 133 ± 1.8 145 ± 2.0 112 ± 1.4 144 ± 2.3 123 ± 1.8 93 ± 0.12

µmol/ml extract 4.27 4.59 5.00 3.86 4.97 4.24 3.21

Dandelion and its parts – leaves, flowers and roots contain relatively high content of TP, which as many authors underline could be regarded as major contributor to their antioxidant activities. Our results on TP content agree well with the data of some authors2,24, and differ from those of other researchers25,26 . Most probably, these variations in TP content in the Dandelion leaves, as well as in many other medicinal plants could be explained with the differences in the genetic characteristics of the various species and varieties, the characteristics of the growth region, the season for gathering the leaf mass, the extraction method, as well as the nature of the extragent24,27,28. Some authors point that the flavonoid fraction is the major constituent of TP in the Dandelion leaves which is in accordance with our data. The antiradical activity, assayed with the stable free radical DPPH, varies from 26.43 tо 30.49 μmol DPPH/g dw, and the difference between them is 13.28% (р < 727

0.05). The leaves of Dandelion grown in the regions of Тyrnavos display the highest antiradical activity whereas those from the Nea Anchialos – Magnesia region – the lowest one. The results expressed as TEAC and IC50 are demonstrated in Table 6. The antiradical activity of the methanol extracts of the investigated plants from the various regions determined as the inhibition coefficient (IC50), correlates with ТР (Table 6). With regard to the value of IC50 they could be arranged as follows: Tyrnavos > Trikala > Kalogriana > Karditsa > Farkadona > Terpsithea > Nea Anchialos – Magnesia. Table 6. Antiradical activity (DPPH) of methanol extracts of dandelion leaf mass

No

Region

1 2 3 4 5 6 7

Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos – Magnesia

DPPH (µl mol/g dw) 26.63 ± 0.31 29.43 ± 0.39 30.49 ± 0.43 26.37 ± 0.37 27.76 ± 0.43 26.83 ± 0.34 26.03 ± 0.35

TEAC (µmol/g dw) 10.84 ± 0.17 11.50 ± 0.14 12.41 ± 0.13   9.50 ± 0.12 11.45 ± 0.15 10.11 ± 0.13   8.12 ± 0.10

IC50 (µg/ml extract) 301 180 150 340 220 280 390

The antiradical activity determined with the free stable radical in anion form (ABTS•─) varies from 16.28 tо 18.99 μg Trolox/g dw (Table 7). The antioxidant capacity of the extracts depends strongly on the TP content (R2 = 0.9859), whereas this dependence determined as the inhibition coefficient (IC50) is relatively weaker (R2 = 0.8825). The highest antiradical activity is manifested by the leaf mass of the dandelion grown in the region of Tyrnavos, and the lowest – from the region of Nea Anchialos – Magnesia, which expressed by IC50 are 103 and 155 μl extract/ml (p < 0.05), respectively. Table 7. Аntiradical activity (ABTS) of methanol extracts of the dandelion leaf mass

No

Region

1 2 3 4 5 6 7

Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos – Magnesia

TEAC (µmol/g dw) 16.85 ± 0.22 18.45 ± 0.24 18.99 ± 0.27 16.55 ± 0.25 18.01 ± 0.27 17.34 ± 0.26 16.28 ± 0.19

IC50 (µg/ml extract) 127 ± 2.1 115 ± 2.4 103 ± 1.5 137 ± 1.9 120 ±2.2 125 ±1.1 155 ± 3.4

1/IC50 (µg/ml extract) 7.84 × 10–3 8.70 × 10–3 9.71 × 10–3 7.30 × 10–3 8.83 × 10–3 8.00 × 10–3 6.45 × 10–3

The ferric reducing power of the methanol extracts of the analysed samples is given in Table 8. It varies in the range from 9.37 tо 15.03 μl FRAP/g dw, from 3.32 tо 5.46 µmpl Trolox/g dw and from 4.26 tо7.37 μmol AAE/g dw. The lowest reducing power is exhibited by the extract from the dandelion leaves from the region of Nea 728

Anchialos – Magnesia, and the highest – from the region of Тyrnavos. The antioxidant capacity of dandelion growing in the regions of Farkadona and Terpsithea, as well as from Trikala and Kalogriana expressed as FRAP, ТЕАС and ААС, although differing in values, is insignificant (р > 0.05). Our results for the high polyphenol content and antioxidant activity of the leaf mass of dandelion agree with the results of other authors who have established higher content in comparison with other herbs and medicinal plants. Stef et al.29 studying the total activity of 11 medicinal plants with DPPH found that dandelion, with respect to the antioxidant capacity, belongs to the group possessing the highest activity, such as Artemisia absanthinum, Phoeniculus, Epilobium montanum, etc., and with respect to the FRAP assay, is characterised by medium ferric reducing power. Studying 17 common herbs Yoo et al.25 established that dandelion is characterised by high content of ТР, FP, antiradical capacity expressed as ААЕ, high superoxide dismutase activity (SOD), and catalase activity (CAT), as well as high protective effect against Н2О2 – induced oxidative stress of gap-junction intrаcellular communication (GJIC) in V79-4 cells. Some of these activities are higher than that of the common thyme, rooibos, fennel, rosmarine, black tea, etc. In addition they found that the antioxidant activity (DPPH, ABTS) is in a strong correlation with the content of ТР and FP. The results obtained show that the antioxidant activity of the methanol extracts from the leaf mass of dandelion from the tested areas and determined by the three assay methods (DPPH, ABTS, FRAP), is in strong correlation dependence on the content ТР, and in weaker on the content of FP (Table 8). The data also reveal that the antioxidant activity of the dandelion leaves from the Tyrnavos area, determined by the three assay tests, is the highest, while that of the dandelion extracts from Nea Anchialos – Magnesia – the lowest. Most probably the similar antioxidant activity and polyphenols content in the extracts from the dandelion leaves from the other areas is due to the similar conditions of their growth. According to may authors, the environmental conditions have a strong effect on the polyphenols biosynthesis, and respectively on their antioxidant capacity12,13. In view of this more research is required to clarify the influence of the analysed microregions on the polyphenols biosynthesis and their antioxidant potential. Table 8. Ferric reducing antioxidant power (FRAP) of methanol extracts of the dandelion leaf mass

No 1 2 3 4 5 6 7

Region Farkadona – Trikala Trikala Tyrnavos – Larissa Terpsithea – Larissa Kalogriana – Karditsa Karditsa Nea Anchialos – Magnesia

FRAP (µmol/g dw) 10.73 ± 0.12 13.95 ± 0.14 15.03 ± 0.18 10.73 ± 0.13 12.70 ± 0.15 11.26 ± 0.13   9.60 ± 0.12

TEAC (µmol/g dw) 4.07 ± 0.06 4.92 ± 0.07 5.46 ± 0.08 3.99 ± 0.05 4.79 ± 0.05 4.31 ± 0.06 3.32 ± 0.07

AAE (µmol/g dw) 5.35 ± 0.09 6.57 ± 0.12 7.35 ± 0.15 5.31 ± 0.13 6.23 ± 0.14 5.69 ± 0.16 4.26 ± 0.10

729

A strong correlation is established between the radical scavenging capacity of the chemical compounds present in dandelion, as demonstrated by DPPH and ABTS and ferric reducing power (with FRAP reagent), expressed in same units – ТЕАС mg/g gw, where the correlation coefficients are as follows: DPPH/FRAP – R2 = 0.9067, and ABTS/FRAP – R2 = 0.9480 (Table 9). This shows that the methanol extracts of Taraxacum officinale contain compounds possessing radiсal scavenging capacity and compounds with ferric reducing power. Table 9. Correlation coefficients (R2) between ТР, FP and antioxidant activity assayed with DPPH•, ABTS•+ and FRAP, expressed as ТЕАС, DPPH and FRAP

TP/ТЕАС FP/TEAC TP/DPPH TP/TEAC FP/TEAC TP/FRAP TP/ TEAC FP/TEAC

DPPH• – R2

ABTS•+ – R2 FRAP – R2

0.8465 0.2336 0.8890 0.9839 0.4258 0.9577 0.9152 0.4393

CONCLUSIONS The study on the dandelion leaves from 7 areas in middle Greece shows that they are characterised by relatively high content of total polyphenols and high antioxidant activity. Depending on the microclimate and soil peculiarities, some differences in the content of total polyphenols, flavonoid phenols, nonflavonoid phenols and flavan3-ols are observed. This affects the antiradical activity determined by the stable free radicals DPPH and ABTS, and the ferric reducing power assayed by the FRAP reagent. A strong positive correlation is observed (R2) between ТР and the antioxidant activity and a weak dependence on FP. The leaf mass of the dandelion grown in the Tyrnavos area is distinguished by the highest content of polyphenols and antioxidant activity whereas that from the Nea Anchialos – Magnesia Aghialos – Magnesia shows the lowest antioxidant activity. REFERENCES   1. G. K. SCHUTZ, R. CARLE, A. SCHIEBER: Taraxacum: A Review on Its Phytochemical and Pharmacolodical Profile. J Ethnopharmacol, 107 (3), 313 ( 2006 ).   2. A. SINGH, S. MALHOTRA, R. SUBBAN: Dandelion (Taraxacum officinale) – Hepatoprotective Herb with Therapeutic Potential. Pharmacognosy Rev, 2 (3), 163 (2008).

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  3. J. BARNES, L. A. ANDERSON: Dandelion Herbal Medicines. 2 ed., Pharmaceutical Press, London, UK, 2003, 171–173.   4. J. BURNETON: Pharmacognosie, Phytochemie Plants Medecinales. 2 ed. Lavoisier TecDoc, Paris, 1993, 78–79.   5. F. Shahidi: Natural Antioxidants: Chemistry, Health Effects and Applications (Ed. F. Shahidi). AOCS Press, Champaign Illinois, 1997.   6. C. A. RICE-EVANS, N. J. MILLER: Antioxidants: The Case for Fruit and Vegetables in the Diet. British Food J, 97 (9), 35 (1995).   7. J. Pokorny, N. YanIshlieva, M. Gordon: Antioxidants in Food. CSC Press, England, 2001.   8. С. WILLIAMS, F. GOLDSTONE, J. GREENHAM: Flavonoids, Cinnamic Acids and Coumarins from the Different Tissues and Medicinal Preparations of Taraxacum officinale. Phytochem, 42 (1), 121 (1996).   9. L. BRAVO: Polyphenols: Chemistry, Dietary Sources, Metabolism and Nutritional Significance. Nutr Rev, 56 (11), 317 (1998). 10. B. Halliwell, J. M. C. Gutteridge: Free Radicals in Biology and Medicine. 2nd ed. Clarendon Press, Oxford, 1989, 1–21. 11. L. PACKER, M. Hiramatsu, T. Yoshikawa: Antioxidant. Food Supplements in Human Health (Eds L. Packer, M. Hiramatsu, T. Yoshikawa). Academic Press, 1999. 12. J. B. HarborNE: Plant Phenolycs. In: Encyclopedia of Plant Physiology. Vol. 8. Secondary Plant Products. Springer, Berlin, 1980, 329–402. 13. J. J. MacheIx, A. FlEuriet, J. BillIot: Fruit Phenolics. CRC Press, Boca Raton, FL, 1990. 14. G. P. P. KAMATOU: Indegenus Salvia Species and Investigation of Their Pharmacological Activity and Phytochemistry. Ph.D. Tessis, Johannesburg, 2006. 15. V. L. Singleton, S. A. Rossi: Colorimetry of Total Phenolics with Phosphomolibdic-phosphotungstic Acid Reagents. J Enol Viticult, 16, 144 (1965). 16. B. BADERSCHNEIDER, D. LUTHRIA, A. L. WATERHOUSE, P. WINTERHALTER: Antioxidants in White Wine (cv. Riesling): 1. Comparison of Different Testing Methods for Antioxidant Activity. Vitis, 38 (3), 127 (1999). 17. T. E. Kramling, V. L. Singleton: An Estimate of the Nonflavonoids Phenolics in Wines. Am J Enol Vitic, 20, 86 (1969). 18. Y.-G. Li, G. Tanner, P. LaRKin: The DMACA–HCL Protocol and the Threshold Proantocyanidin Content for Bloat Safety in Forage Legumes. J Sci Food Agric, 70, 89 (1996) 19. W. Brand-Williams, M. E. Cuvellier, C. Berset: Use of a Free Radical Method to Evaluate Antioxidant Activity. Lebensm Wiss Technol, 28, 25 (1995). 20. R. RE, N. Pellegrini, A. Proteggente, A. Pannala, C. Min Yang, C. Rice-Evans: Antioxidant Activity Applying an Improved ABTS Radical Cation Decolorization Assay. Free Radical Bio Med, 26 (9/10), 1231 (1999). 21. F. F. Benzie, J. J. Strain: Ferric Reducing (Antioxidant Power Assay). Methods in Enzymology, 299, 15 (1999); F. F. Benzie, J. J. Strain: Red Wines. J Agric Food Chem, 48, 220 (2000). 22. G. C. YEN, P. D. DUH: Scavenging Effect of Methanolic Extracts of Peanut Hulls on Free Radical and Active-oxygen Species. J Agric Food Chem, 42, 629 (1994). 23. B. F. RYAN, B. L. JOINERAND, J. D. CRYER: MINITAB Handbook, Updated for Release 14. 15th ed., 2005. 24. M. KRATCHANOVA, P. DENEV, M. CIZ, A. LOJEK, A. MIHAILOV: Evaluation of the Antioxidant Activity of Medicinal Plants, Containing Polyphenol Compounds, Comparison of Two Extraction Systems. Acta Biohim Pol, 57 (2), 229 ( 2010 ). 25. K. M. YOO, C. H. LEE, H. LEE, B. K. MOON, C. Y. LEE: Relative Antioxidant and Cytoprotective Activities of Common Herbs. Food Chem, 106, 929 (2008).

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26. W. ZHENG, Z. Y. WANG: Antioxidant Activity and Phenolic Compounds in Selected Herbs. J Agric Food Chem, 49, 5165 (2001). 27. W. WANGCHAROEN, W. MORASUK : Antioxidant Capacity and Phenolic Content of Holy Basil. Songklanakarin J Sci Technol, 29 (5), 1407 (2007). 28. S. ERCISLI, E. ORHAN, O. OZDEMIR, M. SENGUL, N. GUNGOR: Seasonal Variation of Total Phenolic, Antioxidant Activity, Plant Nutritional Elements, and Fatty Acids in Tea Leaves (Camomillia sinensis var. Sinensis clone Derepazari 7) Grown in Turkey. Pharm Biol, 46 (10–11), 683 (2008). 29. D. S. STEF, I. GERGEN, T. I. TRASCA, M. HARMANESCU, L. STEF, R. BIRON, M. G. HEGHEDUS: Total Antioxidant and Radical Scavenging Capacities for Different Medicinal Herbs. Romanian Biotechnol Lett, 14 (5), 4704 (2009). Received 15 November 2011 Revised 25 March 2012

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Oxidation Communications 37, No 3, 733–740 (2014) Antioxidant and antimicrobial activity in biological systems

Antioxidant Activity, Total Phenolic and Flavonoid Content of Some Local and Cultivated Almonds (Prunus dulcis L.) H. Yildiza, H. Seyfettin Atlib, M. Tosunc, S. Ercislid*, H. Pinare Department of Food Engineering, Nevsehir Haci Bektas Veli University, Nevsehir, Turkey b Department of Horticulture, Faculty of Agriculture, Siirt University, Siirt, Turkey c Oltu Vocational School, Ataturk University, Oltu-Erzurum, Turkey d Department of Horticulture, Faculty of Agriculture, Ataturk University, Erzurum, Turkey E-mail: [email protected] e Alata Horticultural Research Station Directorate, Erdemli-Mersin, Turkey a

ABSTRACT In this study, bioactive content (total phenolic and total flavonoid content, antioxidant activity) of 11 local and 13 foreign almond accessions sampled from a national almond repository collection in Southern Anatolia region in Turkey were investigated. The total phenolic content of methanol extract of almond kernels was determined according to the Folin–Ciocalteu method. Antioxidant activities of the extracts were evaluated by a 2,2′-diphenylpicrylhydrazyl (DPPH), ferric reducing antioxidant power (FRAP) and β-carotene bleaching assays. Total phenolic content was observed in kernel of almond accessions between 45.58 (cv. Primorski) and 93.64 mg GAE/g (cv. Garrigues), respectively. Total flavonoid content was the highest in cv. Garrigues (51.15 mg CE/g) while the lowest value was observed in cv. Ferraduel (11.71 mg CE/g). A wide variation of antioxidant activity of almond kernels was found in all antioxidant determining method and cv. Garrigues in general showed the highest antioxidant activity values in all three methods. Keywords: almonds, total phenolics, FRAP, DPPH, antioxidant activity. AIMS AND BACKGROUND More recently, much attention has been paid on the roles of horticultural crops, including fruits, vegetables and grapes in the promotion of human health1,2. A lot of *

For correspondence.

733

epidemiological studies reported that consumption of horticultural crops has been associated with reduced risk of chronic diseases3–5. All these studies point out that phytochemicals, a class of plant-derived molecules, endowed of strong antioxidant properties6. The studies also revealed that phytochemicals are primarily responsible food components of protective effect on chronic diseases and thus they have received great attention7. In order to properly investigate the role of dietary antioxidants in disease prevention, a complete database of antioxidant-rich foods is required in each horticultural growing region throughout world. Almonds are native to the Mediterranean climate region of the Middle East. Humans spread cultivation of the tree in ancient times along the shores of the Mediterranean into northern Africa and then into southern Europe8. All around the Mediterranean, hillside almond culture became well established with some areas developing important industries based on the nut, including Spain, Turkey, Greece, Portugal, Morocco and Tunisia8. Among these countries, Turkey has unique position from the viewpoint of almond culture and almond genetic resources. Richness in almond diversity in Turkey is mainly due to rich variety of topography and climate over short distances9. Almonds have been cultivated from seeds for centuries in Turkey. The existence of a large number of trees grown from seeds under various ecological conditions provides an invaluable source for varietal selections. The number of mono-crop almond orchards is limited and most almond trees are planted at field borders9. It is well known that most commercial almond cultivars throughout the world were selected by chance from seedling sources8. Previously almond germplasm in Turkey was poorly characterised and only morphological parameters were considered for identification10–12. However, bioactive content of almond germplasm including local cultivars and genotypes have not yet been investigated. Moreover, to our knowledge, there have been no comparative studies on bioactive content of local and foreign almond accessions grown in same locations in Turkey. Therefore, the aim of this study is to compare local and foreign almond accessions in terms of bioactive contents. EXPERIMENTAL Plant material. Almond fruits were harvested from Prunus dulcis accessions from Gaziantep province (latitude 37o03′ N, longitude 37o22′ E, and altitude 842 m) in Turkey. All almond fruits were picked in ripe stage. The fruits were selected according to uniformity of shape and colour. The fruits were then transported to laboratory for analysis. Samples were air-dried and were ground to a fine powder with a mortar and pestle and kept at room temperature prior to extraction. The dried samples were packed into new plastic bags and stored in a dessicator for a maximum of 3 days until antioxidant activity, total phenolics and total flavonoid analyses.

734

Preparation of the methanol extracts. For bioactive content extraction, a fine dried powder (20 mesh) of sample (3 g) was extracted using 50 ml of methanol at 25°C for 60 min. The extracts were filtered through Whatman no 4 paper and evaporated at 40°C to dryness. All the samples were re-dissolved in water at a concentration of 20 mg/ml and analysed for their contents in phenols, flavonoids and antioxidant activity (DPPH, FRAP and β-carotene bleaching assays). Determination of total phenolic content. Total phenolic constituents of almond kernels were performed employing literature methods involving the Folin–Ciocalteu reagent and gallic acid as standard13. Basically, 1 ml of sample was mixed with 1 ml of the Folin and Ciocalteu phenol reagent. After 3 min, 1 ml of saturated sodium carbonate solution was added to the mixture and adjusted to 10 ml with distilled water. The reaction was kept in the dark for 90 min, after which the absorbance was read at 725 nm. The same procedure was repeated to all standard gallic acid solutions and the results were expressed as mg of gallic acid equivalents per g of extract. Determination of total flavonoids. Flavonoid contents in the extracts were determined by a colorimetric method described by Jia et al.14 with some modifications. The almond extract (250 µl) was mixed with 1.25 ml of distilled water and 75 µl of a 5% NaNO2 solution. After 5 min, 150 µl of a 10% AlCl3.H2O solution were added. After 6 min, 500 µl of 1 M NaOH and 275 µl of distilled water were added to the mixture. The solution was mixed well and the intensity of pink colour was measured at 510 nm. (+)-catechin was used to calculate the standard curve and the results were expressed as mg of (+)-catechin equivalents (CE) per g of extract. ANTIOXIDANT ACTIVITY

DPPH radical-scavenging assay. Various concentrations of almond extracts (0.3 ml) were mixed with 2.7 ml of methanol solution containing DPPH radicals (6×10–5 mol/l). The mixture was shaken vigorously and left to stand for 60 min in the dark (until stable absorption values were obtained). The reduction of the DPPH radical was determined by measuring the absorption at 517 nm. The radical scavenging activity (RSA) was calculated as a percentage of DPPH discolouration using the following equation: %RSA = [(ADPPH–AS)/ADPPH] × 100, where AS is the absorbance of the solution when the sample extract has been added at a particular level, and ADPPH – the absorbance of the DPPH solution15. The extract concentration providing 50% of radicals scavenging activity (EC50) was calculated from the graph of RSA percentage against extract concentration. FRAP assay. The stock solutions included 300 mM acetate buffer (3.1 g CH3COONa and 16 ml CH3OOH), pH 3.6, 10 mM TPTZ (2,4,6-tripyridyl-s-triazine) solution in 40 mM HCl, and 20 mM FeCl3·6H2O solution. The fresh working solution was prepared by mixing 30 ml acetate buffer, 3 ml TPTZ, and 3 ml FeCl3·6H2O. Plant extracts (100 μl) were allowed to react with 2750 μl of the FRAP solution and add 150 μl pure water 735

for 30 min in the dark condition. Readings of the coloured product (ferrous tripyridyltriazine complex) were taken at 593 nm. The standard curve was linear between 100 and 1000 μM FeSO4 (r2 = 0.999). Results are expressed in μmol FeSO4 per g. β-carotene bleaching assay. The antioxidant activity of almond extracts was evaluated by the β-carotene linoleate model system. A solution of β-carotene was prepared by dissolving 2 mg of β-carotene in 10 ml of chloroform. 2 ml of this solution were pipetted into a 100-ml round-bottom flask. After the chloroform was removed at 40°C under vacuum, 40 mg of linoleic acid, 400 mg of Tween 80 emulsifier, and 100 ml of distilled water were added to the flask with vigorous shaking. Aliquots (4.8 ml) of this emulsion were transferred into different test tubes containing 0.2 ml of different concentrations of the almond extracts. The tubes were shaken and incubated at 50°C in a water bath. As soon as the emulsion was added to each tube, the zero time absorbance was measured at 470 nm using a spectrophotometer. Absorbance readings were then recorded at 20-min intervals until the control sample had changed colour. A blank, devoid of β-carotene, was prepared for background subtraction. Lipid peroxidation (LPO) inhibition was calculated using the following equation: LPO inhibition = (β-carotene content after 2 h of assay/initial β-carotene)/100 (Ref. 15). Statistical analysis. The experiment was a completely randomised design with 4 replications. Data were subjected to analysis of variance (ANOVA) and means were separated by the Duncan multiple range test at p < 0.01 significant level. RESULTS AND DISCUSSION Total phenolic and total flavonoid content. The differences in total phenolics and total flavonoid content among different almond accessions were statistically significant (p < 0.01, Table 1). The content of total phenolic ranged from 45.58 mg GAE/g (cv. Primorski) and 93.64 mg GAE/g (cv. Garrigues), respectively (Table 1). The accessions seemed to influence the extent of total phenolic content in the almond fruits. The results for total phenolics clearly showed that fruits of almonds are rich in terms of total phenolic content. Amarowicz et al.16, Siriwardhana and Shahidi17, Wijeratne et al.18, and Siriwardhana et al.19 reported total phenolic content in almond kernel extract without the brown skin to be 16.1 mg CE/g, 8.1 mg CE/g, 8 mg quercetin equivalents/g and 8 mg quercetin equivalents/g, respectively. Esfahlan and Jamei20 reported total phenolic content of the 10 wild almond species kernel extracts between 184 mg GAE/g (A. urumiensis) and 482 mg GAE per g extract for (A. pabotti). Barreira et al.21 evaluated 10 almond cultivars (both commercial and regional) and showed that the phenolic content of extracts from almond kernels retaining the brown skin can range from 9.22 to 163.71 mg GAE/g. The total phenolic content in the different accessions investigated here was higher than those of Amarowicz et al.16, Siriwardhana and Shahidi17, Wijeratne et al.18 and Siriwardhana et al.19. However, our values were lower than those for wild almonds20. 736

Table 1. Bioactive content ant antioxidant activity of the tested almond fruits

Accessions

T1 T2 T3 T4 17–4 48 Akbadem Cristomorto Dokuzoguz-2 Drake D. Largueta Farraduel Femagne Garrignes Gulcan-1 Gulcan-2 H.V. Albay Nikitski Nonpareil Picantili Primorski Texas Tuono BHA

Total phenolics (mg GAE/g) 54.75c 60.31bc 48.36cd 53.36cd 50.86cd 46.81cd 48.92cd 47.80cd 54.75c 50.30cd 50.02cd 66.97bc 58.92bc 93.64a 57.26bc 67.53b 53.64cd 64.47bc 46.97cd 50.03cd 45.58d 50.86cd 50.03cd

Total flaAntioxidant activity vanoids FRAP β-carotene (%) DPPH (EC50) (mg CE/g) (mg/ml) (μmol FeSO /g) 33.40ab 30.81ab 27.10ab 18.60bcd 24.22bc 18.11bcd 22.44bc 20.15bcd 33.81ab 30.40ab 27.10abc 11.71cd 18.40bc 51.15a 29.22ab 38.71ab 31.40ab 38.60ab 19.40bcd 15.11c 23.31bc 30.47ab 27.33b

82.90cd 90.77b 70.47de 79.83d 66.20fg 81.37cd 80.50cd 68.80ef 67.60f 61.37fgh 68.40ef 71.63de 84.00cd 91.60b 85.73c 85.00c 71.15de 77.43de 50.10h 62.07g 72.17e 81.23cd 62.71g 96.74a

2.41ab 1.14bcd 2.08ab 1.78b 1.85ab 2.72a 2.09ab 2.00ab 2.24ab 2.39ab 1.99ab 0.66c 1.63bc 0.53d 1.13bcd 1.06bcd 2.55ab 1.02bcd 2.06ab 1.86ab 2.48ab 1.66bc 1.63bc

4

26.25bcd 55.41b 31.61bcd 33.22bcd 32.84bcd 23.12bcd 31.10bcd 32.94bcd 38.62c 29.16bcd 27.59bcd 76.32a 41.88cd 61.46ab 20.88d 38.57c 28.33bcd 61.69ab 30.86bcd 43.55bcd 30.61bcd 45.70bcd 42.14bcd

Means within a column followed by the same letter are not significantly different at p