Calculate the equilibrium constant of the reaction \(\mathrm{I}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{g})\) at 1000 K from the following data for \(\mathrm{I}_{2}\), \(\tilde{v}=214.36 \mathrm{~cm}^{-1}\), \(\tilde{B}=0.0373 \mathrm{~cm}^{-1}\), \(h c \tilde{D}_{e}=1.5422 \mathrm{eV}\). The ground state of the I atoms is \({ }^{2} \mathrm{P}_{3 / 2}\), implying fourfold degeneracy.

Text Transcription:

I_2(g) rightleftharpoons 2 I(g)

I_2

tilde v=214.36 cm^-1

tilde B=0.0373 cm^-1

hc tilde D_e=1.5422 eV

^2P_3/2

1 Kinetics: Iodination of Acetone name ______________________ names Lab Section number TA’s name Date it’s due 2 Introduction: Kinetics in chemistry deals with the rate at which a chemical reaction occurs. This rate, which is referred to as the reaction rate, is defined as the change in concentration of a reactant or product with time, and is measured in M/s. The rate of a reaction is proportional to the concentration of reactants. An equation called the rate law expresses the relationship of the reaction rate to the rate constant, k, and the concentrations of the reactants raised to some powers, x and y, found experimentally. The rate law is expressed as, rate = k [A] [B] . The constant k is equal to the rate divided by the concentration of a certain substance. The purpose in this lab was to experimentally determine the rate constant k, as well as the exponential values of x and y in the rate law. Experimental: The procedure of this lab was obtained from the Chemistry 188 Labratory course website, which can be found directly by the following URL . Results: Table 1: Reagent Volumes + Trial I2 Acetone H dH 2 Total Volume 1 0.5 mL 0.8 mL 0.8 mL 1.9 mL