Activation Energy Lab Report

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57 THE IODINATION OF ACETONE Part One: Determining the Rate for a Chemical Reaction The rate of a chemical reaction depends on several factors: the nature of the reaction, the concentrations of the reactants, the temperature, and the presence of a possible catalyst. In this experiment we will first determine the rate law for a reaction by changing some of the above variables and measuring the rate of the reaction. During Part Two, we will explore the relation between the rate constant and temperature to discover the activation energy for this reaction. In this experiment we will study the kinetics of the reaction between iodine and acetone: O C H3C CH3 + I2(aq) H3C H+ O C CH2I + HI(aq) The rate of this reaction is found to depend on the concentration of the hydrogen ion (acid, HCl) as well as the concentrations of the reactants (acetone and iodine). The rate law for this reaction is rate = k[acetone]m[H+]n[I2]p where k is the rate constant for the reaction and m, n, and p are the orders of the reaction with respect to acetone, hydrogen ions (acid), and iodine, respectively. Although orders of reaction can be any value, for this lab we will be looking only for integer values for the orders of reaction (0, 1, 2 are acceptable but not 0.5, 1.3, etc.) The rate of the reaction can also be expressed as the change in the concentration of a reactant divided by the time interval: rate = - Δ[ I 2 ] Δt The iodination of acetone is easily investigated because iodine (I2) has a deep yellow/brown color. As the acetone is iodinated and the iodine converted to the iodide anion, this color will disappear, allowing the rate of the reaction to be easily monitored. We can study the rate of this reaction by simply making I2 the limiting reactant in a large excess of acetone and H+ ion. By measuring the time required for the initial concentration of iodine (I2) to be

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