Final answer:
Calculating the rate of reaction using H₂SO₃ instead of IO₃⁻ could affect the determination of reaction orders, as it would require considering the stoichiometries of individual steps and the reaction mechanism.
Step-by-step explanation:
If the rate of reaction had been calculated using the concentrations of H₂SO₃ instead of IO₃⁻, while taking stoichiometry into account, it could potentially affect the determination of the reaction orders. In kinetics, the rate law expresses the rate of a reaction in terms of the concentration of reactants, and each reactant can contribute differently to the rate depending on its order in the reaction. Using a different reactant for rate determination means understanding the stoichiometric relationship between that reactant and the other species involved in the rate-determining step.
The reaction orders are deduced by varying the initial concentration of one reactant while holding the others constant and measuring the effect on the reaction rate. If the stoichiometry between H₂SO₃ and IO₃⁻ is not 1:1, then changing the reactant used to calculate the rate would involve adjusting the rate law to account for the stoichiometric coefficients. Therefore, it is essential to consider the balanced equation of the reaction and the stoichiometry of the individual steps in a reaction mechanism, which could reveal that the reaction does not proceed in a single step and that intermediates such as IO⁻ play a role.
Moreover, if the reaction were to be second-order with respect to a particular reactant based on its coefficient in the balanced equation, and yet the experimental data suggest another reaction order, it implies that the reaction mechanism is more complex. In assessing the reaction order, one must consider the reaction mechanism, which often involves elementary steps with their own stoichiometric relationships that do not necessarily mirror the overall balanced equation.