Materials Included In Kit

Additional Materials Required

Prelab Preparation

1. In a fume hood, make a 4 M solution of acetone. Add approximately 148 mL of acetone, using a graduated cylinder, into a 600-mL beaker and fill with approximately 352 mL of distilled or deionized water. Add a magnetic stir bar and mix on a stirrer.
2. Make a 0.005 M solution of iodine. Add 20 mL of 0.05 M iodine solution, using a graduated cylinder, into a 400-mL beaker. Fill with 180 mL of distilled or deionized water. Add a magnetic stir bar and mix on a stirrer.
3. Provide 10-mL each of 4 M acetone solution, 0.005 M iodine solution, 1 M hydrochloric acid solution, and distilled or deioized water to the student groups for a total of 40-mL of chemicals for the iodination of acetone.

Safety Precautions

The iodine solution causes skin and eye irritation. Acetone is a highly flammable liquid and vapor, causes serious eye irritation and may cause drowsiness or dizziness. Keep acetone away from heat, sparks, open flames and hot surfaces. Hydrochloric acid solution causes severe skin burns and eye damage and respiratory irritation. Hydrochloric acid may be harmful if swallowed. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Collect leftover student solutions in a beaker and save the halogenated hydrocarbon mixture for a licensed hazardous waste disposal facility according to Flinn Suggested Disposal Method #4b. If desired, all unused solutions may be saved for future experiments. The leftover iodine solution may be disposed of by reducing with 4% sodium thiosulfate according to Flinn Suggested Disposal Method #12a. The leftover acetone solution may be flushed down the drain with excess water according to Flinn Suggested Disposal Method #18a. The leftover hydrochloric acid solution may be neutralized with base and flushed down the drain with excess water according to Flinn Suggested Disposal Method #24b.

Lab Hints

• The major structural form of crystal violet is the monovalent cation, abbreviated CV⁺. CV⁺ is the predominant form of crystal violet in the solid state and in aqueous solution across a broad range of pH values from pH 1 to 13. The positive charge shown on the central carbon atom in Figure 1a is delocalized via resonance to the three nitrogen atoms. See Figure 1b for one of the three additional resonance forms with the positive charge on a nitrogen atom. Delocalization of the charge across the system of double bonds in the benzene rings stabilizes the carbocation and is responsible for the vibrant purple color of the dye.
  {12311_Hints_Figure_1}
• The purpose of running the cold reaction is for students to observe the effects of temperature in a kinetics reaction. Students should conclude that the reaction rate is much slower.

Further Extensions

Alignment to AP^® Chemistry Curriculum Framework

Enduring Understandings and Essential Knowledge
Reaction rates that depend on temperature and other environmental factors are determined by measuring changes in concentrations of reactants or products over time. (4A)
4A1: The rate of a reaction is influenced by the concentration or pressure of reactants, the phase of the reactants and products, and environmental factors such as temperature and solvent.
4A2: The rate law shows how the rate depends on reactant concentrations. Elementary reactions are mediated by collisions between molecules. Only collisions having sufficient energy and proper relative orientation of reactants lead to products. (4B)
4B1: Elementary reactions can be unimolecular or involve collisions between two or more molecules.

Many reactions proceed via a series of elementary reactions. (4C)

Reaction rates may be increased by the presence of a catalyst. (4D)
4D1: Catalysts function by lowering the activation energy of an elementary step in a reaction mechanism, and by providing a new and faster reaction mechanism.

Learning Objectives
4.1 The student is able to design and/or interpret the results of an experiment regarding the factors (i.e., temperature, concentration, surface area) that may influence the rate of a reaction.
4.2 The student is able to analyze concentration versus time data to determine the rate law for a zeroth-, first-, or second-order reaction.
4.4 The student is able to connect the rate law for an elementary reaction to the frequency and success of molecular collisions, including connecting the frequency and success to the order and rate constant, respectively.
4.7 The student is able to evaluate alternative explanations, as expressed by reaction mechanisms, to determine which are consistent with data regarding the overall rate of a reaction, and data that can be used to infer the presence of a reaction mechanism.
4.8 The student can translate among reaction energy profile representations, particulate representations, and symbolic representations (chemical equations) of a chemical reaction occurring in the presence and absence of a catalyst.

Science Practices
1.4 The student can use representations and models to analyze situations or solve problems qualitatively and quantitatively.
2.1 The student can justify the selection of a mathematical routine to solve problems.
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
5.1 The student can analyze data to identify patterns or relationships.
6.4 The student can make claims and predictions about natural phenomena based on scientific theories and models.

Answers to Prelab Questions

1. A student designed an iodine clock experiment by mixing two colorless solutions and immediately measured the time required for a blue color to suddenly appear. Refer to the following data table to answer a–d.

   2I^–(aq) + S₂O₈^2–(aq) → I₂(aq) + 2SO₄^2–(aq)

   {12311_PreLab_Table_1}
   1. In each trial, the blue color appeared after 0.0020 M iodine (I₂) had been produced. Calculate the reaction rate for each trial.

      Trial 1: Rate = 0.0020 M/270 sec = 7.4 x 10^–6 M/sec
      Trial 2: Rate = 0.0020 M/138 sec = 1.5 x 10^–5 M/sec
      Trial 3: Rate = 0.0020 M/142 sec = 1.4 x 10^–5 M/sec

   2. Compare trials 1 and 2 and determine the order of reaction with respect to iodide ions. How did the concentration of iodide ions change in these two trials, and how did the rate change accordingly? What is the reaction order with respect to iodide?

      In Trials 1 and 2, the concentration of persulfate ions was held constant while the concentration of iodide ions was doubled. The rate increased by a factor of two when [I^–] was doubled. The reaction is first order in iodide.

   3. Which two trials should be compared to determine the order of reaction with respect to persulfate ions? What is the reaction order for persulfate?

      Comparing the rates of Trials 1 and 3 will show how the rate of the reaction depends on the concentration of persulfate ions. In Trials 1 and 3, the concentration of iodide ions was held constant while the concentration of persulfate ions was doubled. The rate increased by a factor of two when [S₂O₈^2–] was doubled. The reaction is first order in persulfate.

   4. Write the rate law for the iodine clock reaction. Could the rate law have been predicted using the coefficients in the balanced chemical equation? Explain.

      Rate = k[I^–][S₂O₈^2–]
      The rate law cannot be predicted simply by looking at the balanced chemical equation—the exponents are not the same as the coefficients in the balanced equation.

2. In a separate kinetics experiment, the student analyzed a color fading reaction spectroscopically. In sodium hydroxide solution, purple crystal violet (CV⁺), cation slowly combines with hydroxide ions to form the neutral product, CVOH, which is colorless. The rate of this reaction is slower than typical acid–base proton transfer reactions and depends on the initial concentration of both crystal violet and hydroxide ions. Answer a–d.
   {12311_PreLab_Equation_5}
   1. Assume that the reaction of CV⁺ with OH⁻ ions (Equation 5) proceeds to completion, that is, the solution turns colorless. What percentage of OH⁻ ions will remain at the end of reaction if the initial crystal violet to sodium hydroxide mole ratio is 1:1? What if the initial ratio is 1:1000?
      {12311_Answers_Figure_2}
   2. In Equation 5, the rate law has the general form of rate = k[CV⁺]ⁿ[OH^–]^m. If the reaction is carried out under certain conditions, then the rate will reduce to the form of rate = kʹ[CV⁺]ⁿ, where kʹ = k[OH^–]^m. The constant kʹ is a new “pseudo” rate constant incorporating both the “true” rate constant k and the [OH^–]^m term. The pseudo-rate law is valid when the concentration of OH^– ions is much greater than the concentration of CV⁺ ions. From 2a, which mole ratio can be used to ensure that the reaction between CV⁺ and OH^– ions can be treated using a pseudo-rate law to determine the reaction order n with respect to [CV⁺]?

      In reaction 1, there will be 0% of both reactants remaining. The mole ratio of CV⁺:OH^– is 1:1; therefore, all of the reactants would be consumed in the reaction. In reaction 2, there will be 0% of CV⁺ remaining and 99.9% OH– remaining. The mole ratio remains the same but with the large excess of OH^– in reaction 2, the amount changes very little. Reaction 2 fits the criterion to be used to find the pseudo-rate law because the concentration of hydroxide ions is effectively constant in the reaction.

   3. Based on observations from 2a and 2b, the student mixed crystal violet and sodium hydroxide solutions in a beaker, transferred to a sample cell, placed into the spectrophotometer, and collected absorbance readings every 20 seconds for 15 minutes. Graph the student’s data below and generate four plots: A vs. time, [CV] vs. time, ln[CV] vs. time, and 1/ [CV] vs. time.
      {12311_Answers_Figure_3}
   4. Select the appropriate linear graph from 2c that provides the best fit straight line through the data points. Consult your textbook and determine the order of the reaction (Equation 5). What is the value of the pseudo-rate constant kʹ? The pseudo-rate constant kʹ can be calculated from Graph In[CV] vs. Time by adding a line of best-fit to the data points and taking the negative of the slope. The value of kʹ = slope of the graph: kʹ = 0.0013.
3. On lab day, you will determine the rate of the iodination of acetone reaction (Equation 5). The iodination of acetone is a simple and measurable reaction, it proceeds from yellow in color to colorless. The observable color change is slow enough to measure in minutes, but not too slow that it will exceed a lab class period.
   {12311_PreLab_Equation_6}
   We can use the following equation to calculate the rate of this reaction.
   {12311_PreLab_Equation_7}
   Because iodine is the limiting reagent and the solutions of acetone and hydrochloric acid are in excess, we can use the initial and final concentrations of iodine to calculate the rate of the reaction. Look at the chemical formula (Equation 6). What is the initial concentration of iodine? What is the final? The initial concentration of iodine is 0.005 M and its final concentration is 0 M, therefore we can determine the rate of the reaction (Equation 6) using the concentration of iodine. Students should show their work from Equation 7. Students should also readily recognize that the other two reactants, acetone and hydrochloric acid, are in excess, thus making iodine the limiting reagent. In addition, the rate of the reaction does not depend on the concentration of iodine because the reaction is zero order with respect to iodine. You will be provided with 10-mL, each of three solutions: 4 M acetone, 0.005 M iodine and 1 M hydrochloric acid, in addition to distilled (or deionized) water. Run five reactions, four at room temperature and one at 0 to 10 °C. Write your own procedure to be preapproved by your instructor prior to lab day. Student procedures will vary, helpful tips were provided. See student sample data tables in Sample Data section.

Sample procedure:

1. Obtain three reactants from the instructor: 4 M acetone solution, 0.005 M iodine solution, 1 M hydrochloric acid solution and deionized water. Make sure the outside container of each solution is properly labeled. Insert a 1-mL beral-type pipet into each.
2. Obtain a test tube rack and insert five small test tubes into it.
3. Setup an ice bath in a small beaker.
4. Setup the cold run first because this reaction will take the longest. For the cold reaction run, measure with the pipet 1-mL each of 4 M acetone solution, 0.005 M iodine, 1 M hydrochloric acid, and water, into to a small test tube submerged in the ice bath. Write down the start time from the classroom clock as the stopwatch will be used for the next runs.
5. In a small test tube, measure with the pipet 1-mL each of 4 M acetone solution, 0.005 M iodine solution, 1 M hydrochloric acid solution, and water. Immediately start the stopwatch. Observe the solution’s color change from yellow to colorless.
6. Setup three more reactions where one of each reactant’s concentration is changed. See data table in Sample Data section for examples.

Sample Data

Example Calculations

Run 1: [acetone], [HCl] and [I₂] determination: M₁V₁ = M₂V₂

Acetone: (4 M)(1 mL) = (M₂)(4 mL)
M₂ = 1 M

HCl: (1 M)(1 mL) = (M₂)(4 mL)
M₂ = 0.25 M

I₂: (0.005 M)(1 mL) = (M₂)(4 mL)
M₂ = 0.00125 M

Run 1 Reaction rate (M/min) determination: (0.00125 M/3.0 min) = 4.17 x 10^–4

Answers to Questions

1. Determine the rate law of the iodination of acetone reaction and the order of the reaction with respect to acetone, iodine and hydrochloric acid.

   Rate = k[acetone][HCl]
   Compare the standard run 1 to run 2, where acetone concentration was doubled, the reaction rate doubled as well. When comparing run 1 to run 3, where the acid concentration was doubled, the reaction rate, again, doubled. Thus, the iodination of acetone is first order with respect to acetone and hydrochloric acid. When iodine was doubled, compare run 1 to run 4, the rate of the reaction did not significantly change. Thus, the reaction is zero order with respect to iodine.

2. Calculate the rate constant, k.

   4.17 x 10^–4 = k(1 M)(0.25 M)
   4.17 x 10^–4 = k(0.25)
   k = 0.0017

3. The iodination of acetone occurs in a series of elementary reactions, or reaction mechanism. Seek education resources and draw the reaction mechanism of the iodination of acetone.
   {12311_Answers_Figure_4}
4. What is the purpose of the hydrochloric acid in the reaction?

   The acid catalyzes the reaction. It is not consumed in the reaction; it speeds it up.

References

AP^® Chemistry Guided-Inquiry Experiments: Applying the Science Practices; The College Board: New York, NY, 2013.

Introduction

Experience and learn the concepts you need to help you succeed on the AP^® Chemistry exam with this guided-inquiry activity! Measuring and controlling reaction rates makes it possible for chemists and engineers to create a variety of products, everything from antibiotics to fertilizers, in a safe and economical manner. How fast will a chemical reaction occur? If a reaction is too slow, it may not be useful. If the reaction is too fast, it may be harmful or explosive. The purpose of this experiment is to investigate how the rate of a reaction can be measured and how varying conditions can affect the reaction rate of the iodination of acetone.

Concepts

• Kinetics
• Reaction rate
• Collision theory
• Reaction mechanism
• Catalyst

Background

Kinetics is the study of the rates of chemical reactions. As reactants are transformed into products in a chemical reaction, the amount of reactants will decrease and the amount of products will increase. The rate of the reaction describes how fast the reaction occurs. The greater the rate of the reaction, the less time is needed for a specific amount of reactants to be converted to products. Some of the factors that may affect the rate of a chemical reaction include temperature, the nature of the reactants, their concentrations, and the presence of a catalyst. Reaction rates are therefore inversely proportional to time (rate ∝ 1/time).

In general, the rate of a reaction increases as the concentration of reactants increases. The relationship between the rate of a reaction and the concentration of reactants is expressed in a mathematical equation called a rate law. For a general reaction of the form

A + B → C

the rate law can be written as

Rate = k[A]ⁿ[B]^m

where k is the rate constant, [A] and [B] are the molar concentrations of the reactants, and n and m are exponents that define how the rate depends on the individual reactant concentrations.

The exponents n and m are also referred to as the order of reaction with respect to each reactant. In the above example, the reaction is said to be nth order in A and mth order in B. In general, n and m will be positive whole numbers—typical values of n and m are 0, 1 and 2. Note that when n = 0, the rate does not depend on the concentration of the reactant. When n = 1, the reaction will occur twice as fast when the reactant concentration is doubled, and when n = 2, the rate will occur four (2²) times as fast when the reactant concentration is doubled. The values of the exponents must be determined by experiment—they cannot be predicted simply by looking at the balanced chemical equation.

Consider, for example, the reaction of hydrogen peroxide with iodide ions to give water and iodine (Equation 1). The general rate law for this reaction is given by Equation 2. The order of reaction with respect to hydrogen peroxide—the value of n—can be determined by measuring the rate of the reaction for several different initial concentrations of hydrogen peroxide. If the concentrations of the other reactants are not changed, the rate will depend only on the concentration of H₂O₂ and the value of n. The general rate law for the reaction will reduce to the form where the constant k' includes the [I^–]^m and [H⁺]^p terms. Calculating the average rate of a reaction may be accomplished by using Equation 4, Since reactants are converted to products, the average reaction rate is expressed as a negative quantity. However, rates are always expressed as positive or as absolute values.

Experiment Overview

The purpose of this activity is to complete a series of kinetics questions in the homework assignment prior to lab day. The homework guides you through a few distinct kinetics problems that will prepare you to determine the rate of the iodination of acetone by performing a series of simple reactions. The reactions will proceed from yellow to colorless on a time scale that can be measured with a stopwatch or cellphone. Build your own data table and conduct post-lab analysis. Plan ahead to maintain chemical economy; your instructor will only give you enough chemicals for five runs.

Prelab Questions

Complete the following homework set and write a lab procedure to be approved by your instructor prior to performing the lab. Along with your procedure, you will turn in any graphs or figures you were asked to create in this homework set, and answers to the questions on a separate sheet of paper, if needed.

1. A student designed an iodine clock experiment by mixing two colorless solutions and immediately measured the time required for a blue color to suddenly appear. Refer to the following data table to answer a–d.

   2I^–(aq) + S₂O₈^2–(aq) → I₂(aq) + 2SO₄^2–(aq)

   {12311_PreLab_Table_1}
   1. In each trial, the blue color appeared after 0.0020 M iodine (I₂) had been produced. Calculate the reaction rate for each trial.
   2. Compare trials 1 and 2 and determine the order of reaction with respect to iodide ions. How did the concentration of iodide ions change in these two trials, and how did the rate change accordingly? What is the reaction order with respect to iodide?
   3. Which two trials should be compared to determine the order of reaction with respect to persulfate ions? What is the reaction order for persulfate?
   4. Write the rate law for the iodine clock reaction. Could the rate law have been predicted using the coefficients in the balanced chemical equation? Explain.
2. In a separate kinetics experiment, the student analyzed a color fading reaction spectroscopically. In sodium hydroxide solution, purple crystal violet (CV⁺) cation slowly combines with hydroxide ions to form the neutral product, CVOH, which is colorless. The rate of this reaction is slower than typical acid–base proton transfer reactions and depends on the initial concentration of both crystal violet and hydroxide ions. Answer a–d.
   {12311_PreLab_Equation_5}
   1. Assume that the reaction of CV⁺ with OH⁻ ions (Equation 5) proceeds to completion, that is, the solution turns colorless. What percentage of OH⁻ ions will remain at the end of the reaction if the initial crystal violet to sodium hydroxide mole ratio is 1:1? What if the initial ratio is 1:1000?
   2. In Equation 5, the rate law has the general form of rate = k[CV⁺]ⁿ[OH^–]^m. If the reaction is carried out under certain conditions, then the rate will reduce to the form of rate = kʹ[CV⁺]ⁿ, where kʹ = k[OH^–]^m. The constant kʹ is a new “pseudo” rate constant incorporating both the “true” rate constant k and the [OH^–]^m term. The pseudo-rate law is valid when the concentration of OH^– ions is much greater than the concentration of CV⁺ ions. From 2a, which mole ratio can be used to ensure that the reaction between CV⁺ and OH^– ions can be treated using a pseudo-rate law to determine the reaction order n with respect to [CV⁺]?
   3. Based on observations from 2a and 2b, the student mixed crystal violet and sodium hydroxide solutions in a beaker, transferred to a sample cell, placed into the spectrophotometer, and collected absorbance readings every 20 seconds for 15 minutes. Graph the student’s data and generate four plots: A vs. time, [CV] vs. time, ln[CV] vs. time, and 1/[CV] vs. time.
      {12311_PreLab_Table_2}
   4. Select the appropriate linear graph from 2c that provides the best fit straight line through the data points. Consult your textbook and determine the order of the reaction (Equation 5). What is the value of the pseudo-rate constant kʹ?
3. On lab day, you will determine the rate of the iodination of acetone reaction (Equation 6). The iodination of acetone is a simple and measurable reaction, it proceeds from yellow to colorless. The observable color change is slow enough to measure in minutes, but not too slow that it will exceed a lab class period.
   {12311_PreLab_Equation_6}

   We can use the following equation to calculate the rate of this reaction.

   {12311_PreLab_Equation_7}

   Because iodine is the limiting reagent and the solutions of acetone and hydrochloric acid are in excess, we can use the initial and final concentrations of iodine to calculate the rate of the reaction. Look at the chemical formula (Equation 6). What is the initial concentration of iodine? What is the final?
   
   You will be provided with 10-mL, each of three solutions: 4 M acetone, 0.005 M iodine, and 1 M hydrochloric acid, in addition to distilled (or deionized) water. Run five reactions, four at room temperature and one at 0 °C to 10 °C. Write your own procedure to be pre-approved by your instructor prior to lab day. Helpful Tips:

   1. Think safety, first. Make sure you have the proper PPE available to perform this lab (i.e. goggles, apron and gloves).
   2. Make a list of the equipment and glassware needed for this lab.
   3. Number the steps in your procedure; remember to be as detailed as possible, from set-up to clean-up.
   4. Draw necessary data tables in your notebook for data collection during the lab.
   5. You have a total chemical volume of 40-mL (10-mL of each solution). You should plan to run at least one standard reaction, and three reactions where at least one reactant concentration changes. Pre-plan to maintain chemical economy. Hint: attempt to run a 1:1:1:1 ratio (a total of 4-mL) to observe the standard color change.
   6. Run the cold reaction first using the same concentrations as the standard. Why run the cold reaction first?

Safety Precautions

The iodine solution causes skin and eye irritation. Acetone is a highly flammable liquid and vapor, causes serious eye irritation and may cause drowsiness or dizziness. Keep acetone away from heat, sparks, open flames and hot surfaces. Hydrochloric acid solution causes severe skin burns and eye damage and respiratory irritation. Hydrochloric acid solution may be harmful if swallowed. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

Post-Lab Analysis
(to be completed in laboratory notebook)

1. Determine the rate law of the iodination of acetone reaction and the order of the reaction with respect to acetone, iodine and hydrochloric acid.
2. Calculate the rate constant, k.
3. The iodination of acetone occurs in a series of elementary reactions or reaction mechanism. Seek education resources and draw the reaction mechanism of the iodination of acetone.
4. What is the purpose of the hydrochloric acid in the reaction?