The reaction between hydrogen peroxide H2O2 and iodide ions I- is a classic example of a reaction that proceeds through a reaction intermediate. The overall reaction can be represented as:H2O2 + 2I- + 2H+ 2H2O + I2This reaction occurs in two steps:1. H2O2 + I- HOI + OH-2. HOI + I- + H+ I2 + H2OIn this reaction, HOI hypoiodous acid is the reaction intermediate. The reaction rate is determined by the slowest step, which is usually the first step in this case. The rate law for this reaction can be written as:Rate = k[H2O2][I-]where k is the rate constant.Now, let's consider the effect of the concentration of the reaction intermediate HOI on the reaction rate. Since HOI is formed in the first step and consumed in the second step, its concentration remains relatively constant during the reaction. This is known as the steady-state approximation.The steady-state approximation allows us to write the rate law for the second step as:Rate2 = k2[HOI][I-]where k2 is the rate constant for the second step.By combining the rate laws for both steps and using the steady-state approximation, we can derive an overall rate law that includes the concentration of the reaction intermediate:Rate = k1k2[H2O2][I-]^2 / k2[I-] + k1[H2O2] In this expression, the concentration of the reaction intermediate HOI is implicitly included in the rate law. The effect of the concentration of HOI on the reaction rate depends on the relative magnitudes of the rate constants k1 and k2 and the concentrations of H2O2 and I-.In general, if the concentration of the reaction intermediate HOI increases, the reaction rate will increase as well, since more HOI is available to react with I- in the second step. However, the steady-state approximation assumes that the concentration of HOI remains relatively constant, so any changes in its concentration will be small and may not have a significant impact on the overall reaction rate.