3.2.3.1  Zero-Order Reactions

In zero-order reactions, the rate is independent of the concentration. This may occur in two different situations: (a) when intrinsically the reaction rate is independent of the concentration of reactants and (b) when the concentration of the reacting compound is so large that the overall reaction rate appears to be independent of its concentration. Many catalyzed reactions fall in the category of zero-order reactions with respect to the reactants. On the other hand, the reaction rate may depend upon the catalyst concentration or other factors unrelated to the concentration of the compound under investigation.

Thus, for a zero-order reaction at constant density, the overall expression would be as follows:

where C is the concentration; t is time; k ₀ is the zero-order reaction rate constant, which by integration would result in:

where C is the concentration at time (t) and C ₀ is the initial concentration.

According to this mathematical expression, a distinguishing feature for this type of reaction is a linear decrease in concentration as a function of time as illustrated in Figure 3.2.

Figure 3.2   Graphical representation for the determination of a zero-order rate constant (k 0).

Typical reactions that have been represented by zero-order reactions include some of the autooxidation and non-enzymatic browning reactions. It is clear that zero-order reactions do not appear to occur as frequently in food systems as other reaction orders. In most cases, it is evident that the most common situation for this type of reaction is when the concentration of the reactants is so large that the system appears to be independent of concentration.

3.2.3.2  First-Order Reactions

A large number of reactions occurring in food systems appear to follow a first-order reaction. A mathematical expression for this behavior would be as follows:

where k ₁ is the first-order reaction rate constant. By integration, this equation becomes:

Thus, according to this mathematical expression ln C vs. time will be a linear function where the slope corresponds to –k ₁ as shown in Figure 3.3. The half-life (t _1/2) is given by:

Figure 3.3   Graphical representation for the determination of a first-order rate constant (k 1).

The mathematical expressions above clearly indicate that the half-life and the reaction rate for a true first-order reaction are independent of the initial concentration. However, although in a number of systems this may be the case, formulated products will not necessarily follow true first-order reaction kinetics, but rather a pseudo-first-order reaction. In fact, in formulated systems the presence of breakdown products may strongly influence the order of the reaction; however, the reaction may follow apparent first-order kinetics for only a given value of initial concentration. To determine if a given reaction does indeed follow a pseudo-first-order kinetics, conditions for the kinetic study can be chosen to follow the technique of flooding. Through this approach, all but one of the concentrations are set sufficiently high that, compared to the one reagent present at lower concentration, the others are effectively constant during the time of the experiment. Since only one of the concentrations changes appreciably during the run, the effective kinetic order is reduced to the reaction order with respect to that one substance. If the order of the reaction is determined to be one, the reaction is said to follow a pseudo-first-order reaction. The degradation of ascorbic acid, for instance, has been primarily found to follow first-order kinetics in food systems. On the contrary, degradation of ascorbic acid in model systems has frequently been found to follow pseudo-first-order kinetics. It appears that the presence of breakdown products modifies the kinetics of deterioration of ascorbic acid and thus its initial concentration will influence its rate of degradation.

3.2.3.3  Second-Order Reactions

Two types of second-order reaction kinetics are of importance:

Type I: A + A → P

and Type II: A + B → P

where C _A is the concentration of reactant species (A) at time (t), C _B is the concentration of reactant species (B) at time (t), and k ₂ is the second-order reaction rate constant. For Type I, the integrated kinetic expression yields:

which in terms of the half-life becomes:

For Type II, the integrated form yields:

where C _A0 and C _B0 are the respective initial concentrations and C _A and C _B are the respective concentrations at time (t). It should be stressed, however, that Type II reactions do not necessarily have to follow a second-order reaction. For instance, for the particular case where component A is present in large amounts as compared with component B, the reaction may follow first-order kinetics with respect to B. A typical plot of second order kinetics is presented in Figure 3.4.

Figure 3.4   Graphical representation for the determination of a second-order rate constant (k 2) for a type I reaction.

3.2.4.1  Types of Intermediates

3.2.4.1.4  Transition Complexes

Chain-Reactions:

Catalyzed reactions may fall in the category of chain reactions. In these reactions an intermediate is formed in the initiation step. Such an intermediate then interacts with the reactant to obtain a product and more intermediate to participate in the reaction. A key consideration is that the intermediate may catalyze a series of reactions before being destroyed.

A classic example for the case of chain reactions is the degradation of β-carotene, an autooxidation reaction involving three main periods: (1) induction (formation of free radicals), (2) propagation (free radical-chain reactions), and (3) termination (formation of non-radical products). According to this pathway, the reaction may be expressed as follows:

$$\begin{array}{c}\begin{array}{c}2\text{RH}\stackrel{k_{\text{i}}}{\to }2{\text{R}}^{•}\\ {\text{R}}^{•}+{\text{O}}_2\stackrel{k}{\rightleftarrows }{\text{RO}}_2^{•}\end{array}\}\left(1\right)\text{induction}\\ \begin{array}{c}{\text{RO}}_2^{•}+\text{RH}\stackrel{k_{\text{p}}}{\to }{\text{R}}^{•}+\text{ROOH}\\ {\text{R}}^{•}+R^{\prime }H\stackrel{{k^{\prime }}_{\text{p}}}{\to }{\text{R}}^{•}+\text{RH}\end{array}\}\left(2\right)\text{propagation}\\ \begin{array}{c}{\text{R}}^{•}+{\text{RO}}_2^{•}\stackrel{k_{\text{t}}}{\to }\text{products}\\ {\text{R}}^{•}+{\text{R}}^{•}\stackrel{{k^{\prime }}_{\text{t}}}{\to }\text{products}\end{array}\}\left(3\right)\text{termination}\end{array}$$

Although different approaches have been taken to mathematically describe the kinetic behavior of β-carotene, a simplified free-radical recombination has been suggested by a number of authors (Alekseev et al., 1968; Gagarina et al.; 1970; Finkel’shtein et al., 1973, 1974).

According to this approach, the rate of consumption of a hydrocarbon, i.e. carotenes, in a chain process with a second-order chain termination can be described by:

3.17 $$-\frac{dC}{dt}=aC\sqrt{w_i}$$

By replacing the value corresponding to w _i , the rate of formation of free radicals, the rate of consumption of the hydrocarbon can be expressed as being:

3.18 $$-\frac{dC}{dt}=aC\sqrt{b_{0}C+b\left(C_0-C\right)}$$

where C is the carotenoid concentration at time (t); C ₀ is the initial carotenoid concentration; a is the constant derived from Equation 3.19; b is the initiation rate constant of the products; b ₀ is the initiation rate constant of unoxidized carotenoids; and t is time.

The initiation rate is considered to be the sum of the initiation rates of radicals formed by the unreacted carotene and by the intermediate products. The value of the constant “a” can be represented by:

3.19 $$a=\frac{k_p}{\sqrt{k_t}}\cdot \sqrt{k{K}_S{P}_{O_2}}$$

where k _p , k _t , and k are rate constants as previously shown; K _s is the solubility coefficient of oxygen in carotenoids; and P_O2 is the partial pressure of oxygen. Integrating Equation (3.19) using the dimensionless variables suggested by Gagarina et al. (1970) yields:

3.20 $$\ln \left(\frac{1+\sqrt{1-C/C_0}}{1-\sqrt{1-C/C_0}}\right)=a\sqrt{\left(b-b_0\right)}\cdot \sqrt{C_{0}t}$$

If all the constants on the right side of the previous equation are lumped together to denote the effective rate constant, namely (σ),

3.21 $$\sigma =a\sqrt{\left(b-b_0\right)}\cdot \sqrt{C_0}$$

Equation (3.20) can be simplified to yield:

3.22 $$\ln \left(\frac{1+\sqrt{1-C/C_0}}{1-\sqrt{1-C/C_0}}\right)=\sigma t$$

which corresponds to the equation for a straight line. It is evident that the slope of a graph of $\ln false[false(1+\sqrt{false(1-C/C_{0}false)}false)/false(1-\sqrt{false(1-C/C_{0}false)}false)false]$ vs. t will give the value corresponding to the effective rate constant.

The aforementioned simplified models have been successfully used by various investigators to describe the reaction kinetics of carotenoids assuming a scheme of an unbranched chain as previously described. It is obvious that in systems where oxygen accessibility is limited due to the density of the material, higher rates of oxidation will take place close to the surface. Hence, different rates of degradation of the carotenoids will take place simultaneously, thus, highly complicating the analysis of the kinetics for the overall system.

Non-Chain Reactions:

Non-chain catalyzed reactions may involve the interaction of the substrate with the catalyst to form a complex, followed by its decomposition to form the product. Upon decomposition, the catalyst is then regenerated and is capable of taking part in the reaction once again. A typical example of this behavior is the acid-catalyzed hydrolysis of pyranosides according to the pathways presented in Figure 3.5. According to this diagram, the mechanism involves rapid reversible protonation of the glycosidic oxygen atom to produce a protonated oligosaccharide (2), which undergoes a slow unimolecular decomposition to a stable monosaccharide and an acyclic carbonium ion (3). It is considered that the carbonium ion is stabilized by resonance with the oxonium ion (4). Nucleophilic addition of water would yield a protonated reducing sugar (5), which through the loss of a proton would result in the hydrolytic products (6) and reappearance of the catalyst.

Figure 3.5   Illustration of a non-chain catalyzed reaction: the acid-catalyzed hydrolysis of pyranosides.

Enzyme Kinetics:

Another example of transition complex reactions is that of enzyme catalyzed reactions. It should be mentioned that most of the reactions occurring in biological systems are catalytic in nature. The basic principles for enzyme-catalyzed reactions have been presented by Michaelis-Menten, who proposed the theory of complex formation according to the following equation:

$$\text{E}+\text{S}\underset{k_{-1}}{\overset{k_1}{\rightleftarrows }}\text{ES}\underset{k_{-2}}{\overset{k_2}{\rightleftarrows }}\text{E}+\text{P}$$

where E is the enzyme; S is the substrate; ES is the enzyme-substrate complex; and P is the product. Both reactions are considered to be reversible. In this equation, k ₁, k _–1, k ₂, and k _–2 are the specific constants for the designated reactions.

Although the general principle of chemical kinetics may apply to enzymatic reactions, the phenomenon of saturation with substrate is unique to enzymatic reactions. In fact, at low substrate concentrations the reaction velocity is proportional to the substrate concentration, and thus the reaction is first-order with respect to the substrate. As the substrate conc entration increases, the reaction progressively decreases, being no longer proportional to the concentration of the substrate and deviating from any first-order kinetics. The reaction follows zero-order reaction kinetics, due to saturation with the substrate (Figure 3.6).

Figure 3.6   Illustration of reaction rate dependence on substrate concentration for the initial phase of an enzyme-catalyzed reaction, indicating maximum velocity (V max ), ½V max and Michaelis constant (K m ).

For the particular case of enzyme kinetics, the cases of competitive and noncompetitive inhibition need to be considered. In the first case, the competitive inhibitor (I) is able to interact with the enzyme to generate a complex (EI), according to the reaction:

$$\text{E}+\text{I}\underset{k_{-1}}{\overset{k_1}{\rightleftarrows }}\text{EI}$$

In this type of reaction, the complex EI does not break down to create products. However, the reaction can be reversed by increasing the substrate concentration.

On the other hand, in the case of noncompetitive inhibition, the inhibitor may bind to the enzyme on a locus different from the active site of the enzyme, and thus it may bind to the free enzyme or to the complex according to the equations presented below:

$$\begin{array}{c}\text{E}+\text{I}\stackrel{}{\rightleftarrows }\text{EI}\\ \text{ES}+\text{I}\stackrel{}{\rightleftarrows }\text{ESI}\end{array}$$

where the forms EI and ESI are inactive.

The most common type of competitive inhibition is that created by compounds that bind reversibly with the sulfhydryl groups of cysteine residues that are essential for the catalytic activity of some enzymes. Such groups may be located at or near the active site. In this second case, the catalytic activity may be related to stearic hindrance or the ability to maintain the three-dimensional conformation of the enzyme.

It is of critical importance to be able to identify competitive and irreversible inhibition. For the particular case of irreversible inhibition, the inhibitor binds to the enzyme irreversibly, and some may modify its molecular structure. It is obvious that for this particular case the use of the Michaelis-Menten equation is not possible since this approach assumes that the interaction between the enzyme and the inhibitor is reversible. A typical case of irreversible reactions would be the case of the trypsin inhibitors found in soybeans, namely, the Kunitz and the Bowman-Birk inhibitors. Chymotrypsin has been found to be strongly inhibited by the Bowman-Birk inhibitor, while only weakly inhibited by the Kunitz inhibitor. Both inhibitors have also been shown to be active against bovine trypsin. The activity of human trypsin has been observed to be inhibited to a significant extent by the Kunitz inhibitor.

3.2.4.2  Consecutive Reactions

Consecutive reactions form another category of reactions of importance in food products. Intermediates are formed in such reactions, which may decompose or react to create other compounds. In many cases, the intermediate may have a short life, and thus simplifications may be taken to describe their kinetics. Since decomposition of the intermediates may proceed under different reaction kinetics, complex situations may arise.

For the particular case of reactions in sequence, following first-order or pseudo-first-order kinetics:

and for the particular case where the aforementioned reactions are not reversible, the rate of disappearance of A can be expressed as follows:

which becomes:

The concentration of the intermediate can be determined by:

By substituting Equation 3.24 into Equation 3.25 and multiplying each term by exp (k ₂ t), the following expression is derived:

Integration of this equation with [B] = 0 at t = 0 yields:

It can be easily illustrated that the final product (C) does not form immediately as A is decomposed due to the formation of B. This period is normally termed an induction period.

It is evident that the aforementioned case is a simple case for these particular types of reactions. More complicated situations correspond to the cases for consecutive reactions with a reversible step or when the individual steps do not follow the same order reaction kinetics. If the rate of disappearance of B is fairly rapid, it is evident that if one monitors the appearance of C only based on the concentration of A, inaccuracies may be introduced in the kinetic analysis.

3.2.4.3  Reversible First-Order Reactions

For the most part, we have considered reactions whose rate constant consists of a single value with an integral reaction order. However, now we will consider the reaction:

in which the rate of disappearance of A is given by:

To solve this equation, two considerations are to be taken into account: (1) from the stoichiometry:

and (2) from the condition −d[A]/dt = 0 at equilibrium:

Through substitution and rearrangement:

Integration between the corresponding limits will give:

According to this equation, a plot of ln ([A]−[A_∞]) vs. time will be a straight line, whose slope corresponds to −(k ₁ + k ₋₁). Figure 3.7 describes the concentration of reactant A as a function of time for two different situations. In the first case, the reaction will reach certain equilibrium with retention values leveling off. On the other hand, if a reactant is added to rapidly consume B, this will prevent its return to A. In this situation, only the forward reaction controls –d[A]/dt with a continuous depletion of reactant A. For both cases, the initial rate of the reaction is the same. Although this type of behavior may be possible in food systems, as in the interconversion of pyridoxamine and pyridoxal, where these compounds may undergo further degradation, most information reported in the literature will consider the rate for the overall reaction with no simultaneous information for the forward and the reverse reactions. Huang and von Elbe (1985), however, developed a kinetic reaction model accounting for the forward and reverse reactions for the degradation and regeneration of betanine in solution and its degradation products, betalamic acid, and cyclodopa-5-O-glycoside. This model could predict the amount of betanine remaining before and after regeneration of the pigment under different experimental conditions. Since in most practical situations dealing with food products the majority of investigators have not explored the mechanisms of degradation of the compound under question to minimize the amount of work involved, it becomes simple to visualize the limitations of available kinetic information if one tries to extrapolate to other systems.

Figure 3.7   Illustration of time dependence of a hypothetical reversible first-order reaction.

3.2.4.4  Simultaneous Competitive Reactions

In reference to complex reactions, another case of significance corresponds to the degradation of a single compound to different products, following different pathways. The degradation of ascorbic acid is a typical example of a complex reaction where several pathways may operate simultaneously. Seldom have investigators taken the time to identify the contribution of each pathway to the degradation of this vitamin, but rather have reported an overall value. In the case of chlorophylls, Heaton et al. (1996a,b) developed a general mechanistic model for rates of chlorophyll degradation to pheophorbides via either pheophytins or chlorophyllides. Depending upon the contribution of each pathway when dealing with competitive reactions, variable levels of inaccuracy may result since each reaction will proceed at a different rate. Thus, if the reaction is expressed by the following mechanism:

where P₁, P₂,…, P _n   = products; and k ₁, k _2,…k _n   = rate constants. The rate of disappearance of A is given by:

which by integration results in:

where

Since the products are generated with different yields, it is obvious that information on product formation is needed to evaluate the rates of the reaction. Moreover, it is evident that solely monitoring the rate of disappearance of A and assigning an overall reaction rate may be highly inaccurate, since each pathway may have a different reaction order and a different associated rate constant. This approach has been commonly taken by many investigators in determining reaction rates for the degradation of vitamins, pigments, etc. This discussion should emphasize the need for a better understanding of the reaction mechanisms to properly report kinetics. Since this is a more time-consuming approach, it is understandable why many investigators circumvented the complexity of the reaction kinetics.

3.3.1.1  Vitamin C (Ascorbic Acid)

Vitamin C (ascorbic acid) is chemically known as L-3-keto-threo-hexuronic acid lactone, is naturally found as the L-isomer in various citrus fruits, hip berries, and fresh tea leaves. It also exists in a stereoisomeric form referred to as D-isoascorbic acid (also called D-araboascorbic or erythorbic acid) and has only a twentieth of the bioavailability of L-ascorbic acid. L-ascorbic acid can reversibly convert to dehydroascorbic acid in the presence of mild oxidants and may subsequently and irreversibly convert to 2,3-diketogulonic acid, which has no bioavailability. This makes it important for proper differentiation during its analysis for nutritional purposes. Originally discovered for its ability to cure scurvy, ascorbic acid is currently well known as a biological cofactor that plays an essential role in a broad array of physiological pathways, fundamental to cellular functions. Its biochemical actions as an antioxidant and universal reducing agent in vivo have been well documented and reviewed (Eitenmiller et al., 2008; Johnston et al., 2014). Considerable amounts of ascorbic acid may be lost during processing and storage of food products. In fact, ascorbic acid is readily destroyed by heating and oxidation, and its protection is particularly difficult to achieve. Other factors influencing the degradation of this vitamin include water activity or moisture content, pH, and metal traces, especially copper and iron. In general, it has been observed for a wide number of products containing ascorbic acid that the reaction appears to follow first-order kinetics. It should be mentioned, however, that different pathways exist for the degradation of ascorbic acid. Such pathways give origin to different breakdown products, and, therefore, affect the overall rates of vitamin degradation. In fact, the reaction may proceed under aerobic or anaerobic conditions, or through catalyzed or uncatalyzed aerobic pathways (Figure 3.9). Understanding of the mechanisms involved facilitates the handling of kinetic data. However, because of the fact that many parameters will influence the kinetics of ascorbic acid decomposition, it is difficult to establish a precursor–product relationship, except for the earlier part of the reactions. For instance, in stored canned products, the reaction may occur at the beginning through catalyzed or uncatalyzed aerobic mechanisms. Upon storage after the disappearance of the free oxygen, subsequent losses may be due to anaerobic decomposition of the compound. It is also possible that various mechanisms of deterioration can operate simultaneously, thus highly complicating the treatment of the kinetic data. For instance, in the presence of oxygen, it is possible that the anaerobic pathway will take place, although its contribution appears to be less significant than even the pathway for uncatalyzed degradation. Eison-Perchonok and Downes (1982) used second-order kinetics to describe the degradation of ascorbic acid under limiting oxygen concentrations. Finholt et al. (1963) indicated a maximum rate for anaerobic degradation of ascorbic acid at pH 4 and attributed this behavior to a 1:1 complex of ascorbic acid molecules and hydrogen ascorbate ions which would be present at the highest concentration at a pH near 4.0. Since at normal temperatures the anaerobic degradation proceeds at low rates, normally its contribution can be considered insignificant in the presence of excess oxygen.

Figure 3.9   Degradation pathways of ascorbic acid (adapted from Bauernfeind and Pinkert, 1970 and Tannenbaum et al., 1985).

It is a well-known fact that the autooxidation of ascorbic acid to dehydroascorbic acid is a reversible process, which takes place in two steps with the formation of a free radical as an intermediate. Early studies of ascorbic acid radicals, reported by Yamazaki et al. (1959 and 1960) and Yamazaki and Piette (1961), indicated its decay according to second-order reaction kinetics. Bielski et al. (1971) indicated that complex reactions occur for the decay of ascorbic acid radicals, despite the fact that the decay follows strictly second-order reaction kinetics. Huelin (1953) investigated the stability of ascorbic acid in a variety of food products. The author observed that decomposition of the compound proceeded faster in the pH range of 3–4. Under anaerobic conditions, it has been observed that a maximum degradation occurs at approximately pH 4, followed by a rate decrease upon a decrease in pH to 2.0 (Huelin et al., 1971). Van Bree et al. (2012) looked at the effect of oxygen headspace on oxidation rates of ascorbic acid in fruit juice model systems and the subsequent formation and breakdown of dehydroascorbic acid during storage at 22°C up to 60 days. The authors found ascorbic acid degradation to be successfully described by a first-order kinetic model across the entire oxygen range tested (0.03–20.9%), but only at low oxygen levels below 0.63% O₂ when using a zero-order reaction; similar results were found for both fruit juice model systems and commercial orange juice. More recently, the affinity of ascorbic acid for free radicals was shown by Giroux et al. (2001) to improve the color stability of beef during storage, following gamma irradiation treatment. The vitamin C-scavenging of free radicals (such as ^•OH and ^•SH) showed a reduction of pigment oxidation in meats. Similarly, addition of ascorbic acid was found to improve stability of color of dry chili powder (carotenoids) during storage at different water activities (Bera et al., 2001).

With regard to metal-catalyzed reactions of ascorbic acid, it has been found that the reaction can be catalyzed by transition metals such as copper, iron, and vanadium, both free and bound in complex compounds (Khan and Martell, 1967a,b, 1968). The reaction may also be catalyzed by the copper containing metal enzyme, ascorbate oxidase. Two different mechanisms have been proposed for the non-enzyme catalyzed degradation of ascorbic acid by free and complex ions. Differences in their reaction order, rate limiting step, and products of oxygen reduction were observed. Schwertnerová et al. (1976) reported that the autooxidation of ascorbic acid, catalyzed by copper ions, followed the Michaelis-Menten law in the presence of an inhibitor. Pekkarinen (1974) observed that the autooxidation of ascorbic acid catalyzed by iron salts in citric acid solutions had a clear induction period, particularly at lower concentrations. Although the addition of copper salt decreased the induction period, it was also observed that the rate of the reaction slowed down at a later stage. It is possible, according to the author, that the copper salt may destroy radicals produced by the reaction catalyzed by the iron salts. Spanyár and Kevei (1963) had previously reported that copper was very destructive to ascorbic acid in air but had an insignificant effect in nitrogen. The authors also indicated that although iron had a prooxidant activity, iron and copper combined accelerated the degradation of ascorbic acid at a lower rate than either of the two acting alone.

In most situations it has been observed that an increase in temperature also increases the destruction rates of ascorbic acid. On the other hand, contradicting results have been suggested for conditions at sub-freezing temperatures. Grant and Alburn ( 1965a,b) studied the degradation of ascorbic acid in frozen and unfrozen solutions containing 10⁻⁴ M ascorbic acid in a 0.02 M acetate buffer at pH values of 5.0 and 5.5. At both pH values, the rates of ascorbic acid oxidation were observed to be significantly higher at −11°C as compared to the system at 1°C, and more pronounced for the system at pH 5.5. A number of factors have been suggested to be responsible for the observed trends. An increase in concentration of the reactants occurs upon freezing of the system (Pincock and Kiovsky, 1966). Other factors, such as a catalytic effect exerted by the ice crystals, favorable orientation of the reactants in the partially frozen system, and a decrease in dielectric constant or an increase in proton mobility, have also been given as possible reasons for the enhanced rates of ascorbic acid degradation at sub-freezing temperatures. It is also possible that oxygen availability is another factor to be taken into consideration. Results reported by Thompson and Fennema (1971) indicated that, in fact, an increase in reactant concentration, pH, and oxygen availability could explain higher rates of degradation or a smalle r than expected decrease in rate in partially frozen systems. Changes of pH as a result of freezing should also be taken into consideration. In fact, Van den Berg and Rose (1959) reported significant changes in pH during the freezing of well-buffered phosphate solutions upon freezing from 0 to −10°C.

In general, taking into consideration the possible reaction pathways, as summarized by Bauernfeind and Pinkert (1970), the specific kinetic parameters may be dependent on more factors than normally accounted for. In fact, it is a common assumption to measure the rate of degradation of ascorbic acid in food products, natural or formulated, for only a given initial concentration. Work done by various investigators, indicates that for the case of this particular vitamin the degradation may follow pseudo-first-order reaction kinetics. In fact, the presence of breakdown products will alter the kinetic rates and, possibly, the mechanisms of degradation (Villota, 1979). In fact, depending on the composition of the system and considering that some of the steps involved in the degradation of ascorbic acid are reversible, it is clear that although the reaction may still follow first-order reaction kinetics the rates of degradation will vary depending on the equilibrium created in the system and the levels of breakdown products. Lavelli and Giovanelli (2003) further exemplified the effect of initial concentration in tomato products. They found higher rates of degradation of ascorbic acid at lower initial concentration levels (Table 3.3).