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The theory behind this experiment is that of rate laws and rate expressions. Arsenious acid, the reducing agent present in this experiment, reacts with iodine at the rate it is formed.
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The theory behind this experiment is that of rate laws and rate expressions. Arsenious acid, the reducing agent present in this experiment, reacts with iodine at the rate it is formed. Therefore, by measuring the time it takes for the consumption of arsenious acid to occur, one can determine the rate of iodine formation. This would be considered the rate-controlling step. The rate-controlling step of the reaction is the step that proceeds so much more slowly than all the others that it effectively controls the overall reaction rate. The order of a reaction is determined by the reaction mechanism, which is also related and equal to the number of reactant molecules in the rate-controlling step of the reaction. In the initial-rate method, the reaction is run only long enough to determine an initial rate.
The reaction being considered is :
IO3- + 8I- + 6H+ ? 3I3- + 3H2O ( Equation 1)
The rate law associated with this reaction can be written as:
d [IO3-] =([IO3-],[I-],[H+],[I3-],[H2O]......)
dt (Equation 2)
however the most frequently type of rate law for this reaction is of the form;
d [IO3-] =k([IO3-]
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