The redox reaction between permanganate ion MnO4- and oxalic acid H2C2O4 can be represented by the following balanced equation:2 MnO4- + 5 H2C2O4 + 6 H+ 2 Mn2+ + 10 CO2 + 8 H2OThis reaction occurs under acidic conditions, as indicated by the presence of H+ ions in the balanced equation. The permanganate ion acts as an oxidizing agent, while oxalic acid acts as a reducing agent.Now, let's consider the effect of changing the pH of the reaction mixture from acidic to neutral. In a neutral solution, the concentration of H+ ions is significantly lower than in an acidic solution. Since the reaction requires H+ ions to proceed, a decrease in H+ ion concentration will affect the reaction rate and equilibrium position.According to Le Chatelier's principle, if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system will adjust its equilibrium position to counteract the change. In this case, the decrease in H+ ion concentration will cause the system to shift its equilibrium position to the left, in an attempt to produce more H+ ions. This means that the reaction between permanganate ion and oxalic acid will be less favorable under neutral conditions compared to acidic conditions.In summary, changing the pH of the reaction mixture from acidic to neutral will result in a shift of the equilibrium position to the left, making the redox reaction between permanganate ion and oxalic acid less favorable. This is due to the decreased concentration of H+ ions in a neutral solution, which are required for the reaction to proceed.