Iodine Clock

Wizzbang Chemistry・2 minutes read

The iodine clock lab explores reaction kinetics through experiments with iodine and sodium thiosulfate to determine rate law constants, illustrating the impact of varying reactant concentrations and temperature on reaction speed and rate calculations. By conducting reactions with different concentrations of reactants and analyzing the impact of temperature changes on the rate constant using the Arrhenius equation, the lab allows for the determination of activation energy and assessment of experimental data fitting within the calculations.

Insights

  • Varying reactant concentrations impact the speed of the iodine clock reaction, with molarity calculations and color changes indicating the reaction's progress.
  • The determination of rate law constants m, n, and k involves conducting multiple reactions with different reactant concentrations, utilizing logarithmic formulas and the Arrhenius equation to calculate activation energy and reaction rates accurately.

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Recent questions

  • What is the purpose of the iodine clock lab?

    The iodine clock lab aims to study reaction kinetics and the impact of varying reactant concentrations on reaction speed.

  • How are reactant concentrations determined in the lab?

    Reactant concentrations are determined using the mixing formula m1v1 = m2v2, with the molarity of iodine set at 0.200 molar.

  • What factors affect the color change indicating the end of the reaction?

    The color change indicating the end of the reaction is delayed by adding sodium thiosulfate to prevent immediate reaction with starch, affecting the final time of the reaction.

  • How are the rate law constants m and n determined in the lab?

    The m and n values in the rate law equation are determined by comparing rates of reactions with different reactant concentrations using logarithmic calculations.

  • How is activation energy calculated in the lab?

    Activation energy is calculated by finding the slope of the line graphed with 1 over temperature and natural log of k, using the slope form of the Arrhenius equation.

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Summary

00:00

Iodine Clock Lab: Reaction Kinetics Analysis

  • Professor Amanda Hudson introduces the iodine clock lab, focusing on reaction kinetics and the impact of varying reactant concentrations on reaction speed.
  • The lab involves a reaction between iodine and sodium thiosulfate, aiming to determine the rate law constants k, m, and n.
  • Three reactions are conducted to determine the order of iodine and sodium thiosulfate by keeping one reactant constant in each pair of reactions.
  • The temperature is maintained at 19 degrees Celsius throughout the experiments to ensure consistency.
  • Concentrations of iodine and sodium thiosulfate are determined using the mixing formula m1v1 = m2v2, with molarity of iodine set at 0.200 molar.
  • The color change indicating the end of the reaction is delayed by adding sodium thiosulfate to prevent immediate reaction with starch.
  • The final time of the reaction varies based on the rate of iodine production, affecting the overall reaction rate.
  • The rate of the reaction is calculated by dividing the change in iodine concentration by the final time obtained from the stopwatch.
  • The m and n values in the rate law equation are determined by comparing rates of reactions with different reactant concentrations using logarithmic calculations.
  • The m and n values are calculated using the logarithmic formula log(rate2/rate1) = log(iodine2/iodine1)^m and log(rate2/rate1) = log(sodium thiosulfate2/sodium thiosulfate1)^n, respectively.

19:33

Determining Iodine Concentration and Activation Energy

  • To determine concentrations of iodine, solve for m and n in the rate law equation rate equals k concentration of i to the m times s2o8 2 minus to the n. Once m and n are found, input them based on the order of the reaction (first order - ones, second order - twos) to calculate rates. Choose an experiment, input rate, iodine concentration, and s2o8 concentration to find m and n values, then solve for k using these values for experiments one, two, and three.
  • In the second part of the lab, analyze how changing temperature impacts the rate constant. Calculate rates using given times, noting different temperatures for each experiment. Use the Arrhenius equation in linear form, lnk equals negative ea over r times 1 over t plus ln of a, to graph 1 over t and natural log of k for a straight line. Determine activation energy by finding the slope of the line, which equals negative ea over r, using two points on the graph.
  • Utilize the slope form of the Arrhenius equation to calculate activation energy by finding the slope of the line graphed with 1 over temperature and natural log of k. Compare experimental data points to the line to assess how they fit within the activation energy calculation.
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