Final answer:
To determine the number of moles of persulphate ions needed to oxidize one mole of manganese ions, we would need to know the specific oxidation states involved. In an example scenario, if manganese ions were oxidized from Mn²+ to MnO₄⁻, 2.5 moles of persulphate ions would be required to provide the necessary 5 electrons per manganese ion.
Step-by-step explanation:
The question relates to the stoichiometry of a redox reaction involving manganese ions and persulphate ions. Specifically, it asks how many moles of persulphate ions are needed to oxidize one mole of manganese ions. Using redox chemistry, we can deduce that each persulphate ion can donate two electrons in a redox reaction, as shown in the half-reaction S₂O₈²- (aq) + 2e⁻ → 2SO₄²- (aq).
To fully answer this question, we would need the specific oxidation states involved in the reaction, which are not provided. However, as an example, if we consider the oxidation of Mn²+ (starting oxidation state of +2) to MnO₄⁻ (ending oxidation state of +7), we can determine the number of electrons transferred and thus the stoichiometry of the reaction.
If manganese is taken from an oxidation state of +2 to +7, it loses 5 electrons per manganese ion. The persulphate ion donates two electrons, so 2.5 moles of persulphate ions would be required to provide 5 electrons for each mole of manganese ions. Hence, in this hypothetical scenario, you would need 2.5 moles of persulphate to oxidize 1 mole of manganese ions from Mn²+ to MnO₄⁻.