The oxidation of nitrite by dissolved oxygen to form nitrate is known to be accelerated ca. 105 times by the freezing of the aqueous solution. Here we report a detailed study on the acceleration mechanism of the above-mentioned oxidation. The reaction was studied at pH values between 3.0 and 5.6 at various freezing rates, by different freezing methods, and with and without additional salts. The effect of freezing which induced concentration (freeze concentration) of reactants into the unfrozen bulk solution was too small to explain the acceleration factor of ca. 105. Nitrate formations were completely prevented by addition of salts, such as NaCl and KCl, which make the freezing potential of ice negative, while the reaction was not affected by addition of salts, such as Na2SO4 and NH4Cl, which make the freezing potential of ice positive. When a sample solution was frozen in such a way as to form a single crystal of ice, most nitrite was exclusively liberated from the ice to the gas phase. This observation suggests the importance of ice in the polycrystalline form to retain nitrite during freezing. When freezing begins, grains of crystalline ice begin to grow. The solutes are rejected from the ice and concentrated in the interfacial water layer by assistance of the electrostatic force generated by the freezing potential. At a certain stage of freezing, the water layer is completely confined by the walls of some ice grains. Protons move from the ice phase to the unfrozen solution surrounded by the ice walls to neutralize the electric potential generated, and thus the pH of the unfrozen solution decreases. As a result, the reactant species, HNO2, increased more in the unfrozen solution. After this stage, the concentrations of the reactants in the unfrozen solution abruptly increase resulting in the acceleration of the rate of formation of nitrate. On the basis of the above mechanism, the concentration factor for nitrite was calculated as 2.4 × 103. The validity of this mechanism is further discussed.

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Some reactions are accelerated in ice compared to aqueous solution at higher temperatures. Accelerated reactions in ice take place mainly due to the freeze-concentration effect of solutes in an unfrozen solution at temperatures higher than the eutectic point of the solution. Pincock was the first to report an acceleration model for reactions in ice,1 which successfully simulated experimental results. We propose here a modified version of the model for reactions in ice. The new model includes the total molar change involved in reactions in ice. Furthermore, we explain why many reactions are not accelerated in ice. The acceleration of reactions can be observed in the cases of (i) second- or higher-order reactions, (ii) low concentrations, and (iii) reactions with a small activation energy. Reactions with a buffer solution or additives in order to adjust ion strength, zero- or first-order reactions, or reactions containing high reactant concentrations are not accelerated by freezing. We conclude that the acceleration of reactions in the unfrozen solution of ice is not an abnormal phenomenon.

Oxidative decomposition of gallic acid occurs in alkaline solutions but hardly arises in acidic solutions. We have found that the addition of sodium chloride promotes the decomposition of gallic acid caused by freezing even under neutral and acidic conditions. Even at pH 4.5, gallic acid was decomposed by freezing in the presence of NaCl; however, in the absence of NaCl, it was hardly decomposed by freezing at pH lower than 7. Chloride ions are more easily incorporated in ice than sodium ions when the NaCl solution is frozen. The unfrozen solution in ice becomes positively charged, and as a result, protons transfer from the unfrozen solution to the ice. We measured the pH in the unfrozen solution which coexists with single-crystal ice formed from a 5 mmol dm(-3) NaCl solution and determined the pH to be 8.6 at equilibrium with CO(2) of 380 ppm or 11.3 in the absence of CO(2) compared to pH 5.6 in the original solution. From the model calculation performed for gallic acid solution in the presence of 5 mmol dm(-3) NaCl, it can be estimated that the amount of OH(-) transferred from the ice to the solution corresponds to 1.26 x 10(-5) mol dm(-3). The amount of OH(-) transferred is concentrated into the unfrozen solution and affects the pH of the unfrozen solution. Therefore, the pH in an unfrozen gallic acid solution in ice becomes alkaline, and the decomposition of gallic acid proceeds. It is expected that other base-catalyzed reactions in weakly acidic solutions also proceed by freezing in the presence of NaCl without the need for any alkaline reagents.

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