## 4 Discussion
\[\mathrm{H_{2}A_{ox}}\rightleftharpoons\mathrm{HA_{ox}}^{-}+\mathrm{H^{+}}\left( \mathrm{ pKa=1.27}\right) \tag{1}\]  
\[\mathrm{HA_{ox}}^{-}\rightleftharpoons\mathrm{A_{ox}}^{2-}+\mathrm{H^{+}}\left( \mathrm{ pKa=4.67}\right) \tag{2}\]  
\[\mathrm{H_{3}A_{cit}}\rightleftharpoons\mathrm{H_{2}A_{cit}}^{-}+\mathrm{H^{+}} \left(\mathrm{ pKa=3.13}\right) \tag{3}\]  
\[\mathrm{H_{2}A_{cit}}^{-}\rightleftharpoons\mathrm{HA_{cit}}^{2-}+\mathrm{H^{+} }\left(\mathrm{ pKa=4.76}\right) \tag{4}\]  
\[\mathrm{HA_{cit}}^{2-}\rightleftharpoons\mathrm{A_{cit}}^{3-}+\mathrm{H^{+} }\left(\mathrm{ pKa=4.67}\right) \tag{5}\]  
\[\mathrm{H_{2}A_{mal}}\rightleftharpoons\mathrm{HA_{mal}}^{-}+\mathrm{H^{+}} \left(\mathrm{ pKa=1.92}\right) \tag{6}\]  
\[\mathrm{HA_{mal}}^{-}\rightleftharpoons\mathrm{A_{mal}}^{2-}+\mathrm{H^{+}} \left(\mathrm{ pKa=6.27}\right) \tag{7}\]  
The sharp increase in pH during the first 3 days for the oxalic acid treatment indicates rapid consumption of H\({}^{+}\) generated from dissociation of oxalic acid (Equations (1) and (2)). This could involve protonation of variably charged organic and inorganic soil colloids (including oxides of Fe, Mn, and Al) and reactions with carbonate and silicate minerals [30; 31]. The slight decrease in solution-borne Fe, Mn, and Al prior to the 5 th day of experiment suggests that release of soil-borne Fe, Mn, and Al to the solutions was inhibited. The reaction between the soil-borne Fe/Mn/Al oxides/hydroxides and individual organic acids could lead to the release of these metals into the solution, as illustrated in the following example for the reaction of aluminium hydroxide with citric acid:  
\[\mathrm{nAl(OH)_{3(g)}+3H^{+}_{(aq)}+C_{6}H_{5}O_{7}}^{3-}{}_{(aq)}\to(n-1) \mathrm{Al(OH)_{3(g)}+AlC_{6}H_{5}O_{7(g)}+3H_{2}O} \tag{8}\]  
\[\mathrm{AlC_{6}H_{5}O_{7(g)}+C_{6}H_{5}O_{7}}^{3-}{}_{(aq)}\to\mathrm{Al(C_{6} H_{5}O_{7})_{2}}^{3-}{}_{(aq)} \tag{9}\]  
In Equation (8), the attack of aluminium hydroxide by citric acid results in the formation of insoluble aluminium citrate, which further reacts with another citrate ion to form soluble aluminium citrate complex (Equation (9)). Therefore, Equation (9) represents the rate-limiting step for the organic acid-driven dissolution of aluminium hydroxide. At least before the 7 th day of the experiment, the reaction equilibrium was not reached because the concentration of dissolved Al kept increasing from the 7 th day to the 15 th of the experiment. The malic acid treatment showed a similar trend to the citric acid, suggesting that both citric acid and malic acid behaved similarly in terms of solubilizing the Al in the investigated soil. The delayed release of Al in the oxalic acid treatment could be attributed to the smaller acid dissociation constant (Ka), as compared to the other two organic acids [27]. The close correlation among the solution-borne Fe, Mn, and Al suggests that these metals were solubilized consistently from the oxides of these metals upon contacts with the organic acids. The reduced correlation coefficient during the later stage of the experiment may reflect differential rate of precipitation for these metals. By comparison, the amount of three metals mobilized during LMWOA extraction was in the following decreasing order: Al \(>\) Fe \(>\) Mn. This was consistent with the total concentration of these three metals contained in the investigated soil (Table 1), suggesting the strong control of the soil-borne metals on the solution-borne metals. However, when considering the ratio of solution-borne metal to soil-borne metal (M\({}_{solution}\):M\({}_{soil}\)), it is clear that Mn had much stronger affinity to the LMWOAs, as compared to either Al or Fe; M\({}_{solution}\):M\({}_{soil}\) was more than 0.23 for Mn while M\({}_{solution}\):M\({}_{soil}\) for either Al or Fe was only 0.07. The fact that solution-borne Pb was not related to any of the solution-borne Fe, Mn, and Fe suggests that Pb was not predominantly bound to the oxides of these metals.