In summary, our computational study reveals that reactions Q2 and Q7 describe the most likely hydroquinone's oxidation reactions for the given experimental conditions in solution and with the presence of Al 3+ . They differ -0.07 and -0.01 eV with the experimental reference, respectively, and, energetically, both results are equiprobable. However, given the acidic experimental conditions (pH=1.5), [Al(EtOH) 6 ] 3+ is a more likely populated species than [Al(EtOH) 5 EtO -] 2+ , and therefore, Q2 should be considered the preferred oxidation process. Both reactions entail a PCET process; which it is the same oxidation process than the one observed without the presence of Al 3+ ions in solution (b in Table ). If we compare oxidation potentials reported in Table with E ox (b) of Table , the oxidation potential shift due to the presence of aluminum (∆E ∅→Al ox ) can be studied. We can observe that when Al 3+ is present, more negative results are obtained. This trend is also observed experimentally (Table ) and our results replicate it quantitatively. Then, we can conclude that QH 2 oxidation reactions are more endergonic and unfavored when Al 3+ ion is present in solution. The reason for this increase in endergonicity is a balance between the promotion of proton release and the disfavor of electron lose in the presence of aluminum. Thus, on one hand, aluminum has a tendency to favor deprotonation of the ligands that are attached to it, so that it can interact with a negatively charged species, however for the same reason, it tends to disfavor the lose of negative charge of the ligands, since this implies a decrease of the electrostatic interaction between the ligand and aluminun. In PCET processes, where both reactions coexist, the lose of an electron has a higher energetic cost than the lose of a proton in these aluminum complexes, and the overall effect is a more negative oxidation potential in the presence of aluminum. Notice for example that in the case of Q4 where one electron is released but two protons are unbound instead Assuming that the oxidation process will evolve following the same trend as for QH 2 , only the key oxidation reactions determined before (i.e., Q2, Q7, and Q9) have been studied here. However, as discussed before, the existence of the reagent [Al(EtOH) 5 (Me 4 QH -)] 2+ of reaction MeQ9 is less likely than the existence of the other Al 3+ reagents involved in MeQ2 and MeQ7 reactions. The [Al(EtOH) 5 (Me 4 QH -)] 2+ reagent complex of MeQ9 is more unstable (∆G of 10.8 kcal/mol) than the equivalent nonmethylated form [Al(EtOH) 5 (QH -)] 2+ (∆G of 8.7 kcal/mol) compared with the most stable form at pH=1.5, [Al(EtOH) 6 ] 3+ , (Table ). Moreover, it can neither be formed directly by [Al(EtOH) 6 ] 3+ + Me 4 QH -because the deprotonated form of Me 4 QH 2 is not present in solutions at pH=1.5 as its pKa value is > 11. Overall, we consider that MeQ2 and MeQ7 are the oxidation reactions of Me 4 QH 2 that better represent experimental results. Then, as for the QH 2 case, the most likely oxidation reactions of Me 4 QH 2 comparing to the experimental oxidation potential are dehydrogenation reactions (PCET reactions) that also imply the coordi-nation of Me 4 QH 2 with the Al 3+ complex.