**5. Discussion**

The characteristic sigmoid shape of oxygen consumption in the autoxidation of tartaric acid is observed at all pH conditions. The initiation step is due to oxygen activation to hydrogen peroxide by a Fe(II)-tartrate complex. It is rapid with a time scale of minutes to hours. This reaction step is pH sensitive, slower at lower pH (2.5) than higher (4.5). The extent of peroxide formation is limited to the pool of the Fe(II) state, free and complexed. Initiation is the critical feature of autoxidation, both in these solutions and in wine.

The propagation step is due to the free Fe(II) and hydrogen peroxide oxidation of tartaric acid to produce a tartaric radical which then goes further to form dihydroxymaleic acid, regenerating Fe(II) and consuming additional oxygen as hydrogen peroxide is reformed.

The termination step cannot be explained by the direct decomposition of residual hydrogen peroxide by Fe(III) since oxygen formation is not observed. This termination step may also involve intermediates and end products such as the tartaric acid radical and dihydroxymaleic acid, which adds to the complexity of tartaric oxidation and the accuracy of this model to explain the termination as observed at low pH.

The mechanism proposed here, (**Figure 6**), is slightly different from that previously presented [10]. The difference lies in the propagation steps leading to the exponential growth of one hydrogen peroxide leading to two. The current version does not as accurately fit the loss of Fe(III) observed at the end of oxygen consumption, in the 2.5 pH cases.

The reactivity of dihydroxymaleic acid and the initiation of subsequent radical chain reactions makes describing and interpreting of the termination stage in this system complicated; however, it can be linked to termination of the chain propagation conditions (a negative eigen-value of the kinetic matrix) of our pseudo-first order reaction analysis of Section 3.

#### **5.1 The role of the acid**

The unique properties of tartaric acid in these autoxidation reactions can be seen if other major organic acids are substituted for it in the same reaction medium. **Figure 8** shows the individual oxidation of malic, citric and succinic acids at pH 2.5 and 4.5.

#### **Figure 8.**

*Wine acids and oxygen consumption. Autocatalytic reaction initiated with 265 μM Fe(II)n air-saturated 26.7 mM tartaric acid ( ), malic acid ( ), succinic acid ( ), citric ( ) and hydrochloric acid ( ) at (a) pH 2.5 and (b) pH 4.5. Reproduced from [10], with the permission of AIP Publishing.*

At low pH none of these acid show any significant propagation stage while at high pH malic and citric acid show a slow propagation stage but still much slower than tartaric.

The propagation kinetic phase is associated with a chain reaction that requires regeneration of radicals with multiplicity factor greater than one. In our scheme, *two* hydrogen peroxide molecules are regenerated in the chain propagation reaction with tartaric acid (at low pH) for each hydrogen peroxide entering the cycle. Thus multiplicity factor is two. **Figure 9** illustrates the importance of multiplicity factor for a chain reaction. As seen in the figure, upon addition of hydrogen peroxide to citric and succinic acids, the same amount of oxygen is consumed, with ratio 1:1. Here one acid radical is produced per hydrogen peroxide, which reacts further with oxygen to produce peroxy species, but no addition radical or hydrogen peroxide is regenerated. Thus multiplicity factor is zero.

Indeed, the reaction overall oxygen consumption is given by the sum of probabilities of the chain, where the multiplicity factor q (probability of radical generation):

$$\text{Oxygen consumption} = \mathbf{1} + \mathbf{q} + \mathbf{q}^2 + \mathbf{q}^3 + \dots = \mathbf{1}/(\mathbf{1} - \mathbf{q})\tag{9}$$

For malic acid, interestingly, *four* oxygen molecules are consumed for each hydrogen peroxide added, with ratio 1:4. It can be interpreted as a reaction with multiplicity factor q = 3/4, but still less than one, needed to ignite the chain reaction. The overall number of radicals and oxygen consumed is summed to four, as observed. The factor 3/4 is clearly related to the chemical nature of malic acid that has one OH group out of 4 possibilities on C2 and C3 carbons, and 3/4 occupied by hydrogens.

For tartaric acid the multiplicity factor is obviously greater than one and the overall oxygen consumption is as much as 10 in this trial, and the chain could run without stopping until all oxygen or Fe(II) is consumed. That is what we observe indeed for tartaric acid at low pH.

#### **Figure 9.**

*Wine acids and oxygen consumption initiated with hydrogen peroxide. Autocatalytic reaction initiated with 265 μM Fe(II) and addition of 26.5 μM hydrogen peroxide in air-saturated 26.7 mM tartaric acid ( ) offset to 1 hour to align with: malic acid ( ), succinic acid ( ), and citric acid ( ) at pH 2.5. Reproduced from [10], with the permission of AIP Publishing.*

*The Kinetics of Autoxidation in Wine DOI: http://dx.doi.org/10.5772/intechopen.103828*

By comparison in case of succinic and citric acids, the overall probability of oxygen consumption equals to 1.

#### **5.2 The role of ethanol**

The addition of ethanol to this system (**Figure 10**), shows that even at concentrations of 26.5 mM and both pH of 2.5 and 4.5, the propagation reaction is not established, even when oxygen and Fe(II) are available. This effect, due to ethanol at 100 times the Fe(II) concentration is similar across pH. This can be interpreted as a competition of ferryl ion or hydroxyl radical for ethanol over tartaric acid or the depletion of tartaric radicals in the proposed cycle due to hydroxyethyl radical formation rather than a solvent dielectric effect.

This indicates that the hydrogen peroxide is limiting the propagation stage but the external addition overcomes this limitation in rate (**Figure 11**). The extent of reaction when hydrogen peroxide is available, is independent of pH. In the absence of ethanol,

#### **Figure 10.**

*Ethanol and tartaric acid. Autocatalytic reaction with 0 mM ( ), 0.265 mM ( ), 2.65 mM ( ), 26.5 mM ( ), 265 mM ( ), and 2.56 M or 15% (v/v) ( ) ethanol with 265 μM Fe(II) in air-saturated 26.7 mM tartaric acid at (a) pH 2.5 and (b) pH 4.5.*

#### **Figure 11.**

*Ethanol and tartaric acid initiated with hydrogen peroxide. Autocatalytic reaction with 0 mM ethanol and 0 μM H2O2 ( ), 26.5 mM ethanol and 0 μM H2O2 ( ),0 mM ethanol and 26.5 μM H2O2 ( ), 26.5 mM ethanol and 26.5 μM H2O2 ( ) with 265 μM Fe(II) in air-saturated 26.7 mM tartaric acid at (a) pH 2.5 and (b) pH 4.5.*

100% of the oxygen would be consumed within 15 minutes and this result is also pH independent.
