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## Using Spectrophotometry to Determine the Rate Law for the Reaction of KI with FeCl3

### Introduction

The differential rate law of a chemical reaction relates the rate of the reaction in M/s to the concentration of the reactants raised to various exponents called orders of reaction. For a generalized reaction of A with B the rate law will have the form:

Rate = k[A]n[B]m

In this equation k is the rate law constant which makes the proportion into an equation. The rate law constant varies with temperature and all experiments in this investigation will be carried out at constant temperature. The exponents n and m must be determined experimentally. In this investigation you will determine the values of these exponents and therefore define the rate law for the following chemical reaction:

2I(aq) + 2Fe3+(aq) --> I2(aq) + 2Fe2+(aq)

By varying the concentrations of iodide and iron(III) ions in solution and measuring the effect on the rate you will be able to determine the order of reaction with respect to each reactant and determine a and b in the following equation:

Rate = k[I]a[ Fe3+]b

The formation of elemental iodine in this reaction makes it possible to measure the progress of the reaction spectrophotometrically. The iodine absorbs light at 450 nm so when the absorbance of the solution is measured as a function of time the rate of the production of iodine can be indirectly measured. By graphing absorbance vs. time a straight line is produced, the slope of which is proportional to the rate of the reaction. Recall that when Beer’s Law applies, concentration is directly proportional to absorbance. By plotting absorbance on the y-axis and time on the x-axis the slope will have units of absorbance over time. Since we do not have measurements of the proportion between concentration and absorbance we cannot measure the rate of reaction directly in this lab.

Objective

Determine the rate law for the reaction under investigation by collecting rate vs. initial concentration data.

### Materials

• spectrophotometer capable of providing absorbance at 450 nm
• 4 cuvets
• distilled water
• 0.020 M potassium iodide (KI) solution
• 0.020 M iron(III) chloride (FeCl3) solution
• syringes for measuring solutions

### Safety

The following list does not cover all possible hazards, just the ones that can be anticipated. Move slowly and carefully in the lab: haste and impatience have caused more than one accident.

• Wear chemical splash goggles, gloves, and a chemical-resistant apron.
• Iron(III) chloride is a skin and tissue irritant; it is corrosive and toxic LD50 = 260 mg/kg. Use care in handling the solution.
• Potassium iodide is considered non-toxic.
• All solutions used in this lab must be collected for hazardous waste disposal. Neither the elemental iodine nor the iron(II) or iron(III) ions may be dumped into the sewer system.
 Experiment Number Volume Distilled H2O (mL) Volume 0.020 M KI (mL) Volume 0.020 M FeCl3 (mL) 1 1.0 1.0 1.0 2 0 1.0 2.0 3 0 2.0 1.0

### Procedure

1. Prepare a blank solution by mixing 1 mL KI solution with 2 mL distilled water in a cuvet.
2. Plug the SpectroVis Plus unit into the USB port of the computer and start the Logger Lite or Logger Pro software.
3. Click on the “Experiment” menu, select Calibrate and click the word Spectrometer.
4. A dialog box will pop up to inform you that the lamp is warming up. Do not skip this step, it only requires 90 seconds.
5. Pick up a cuvet by touching the grooved sides. Handling the smooth sides will put fingerprints on it which could have a small effect on the data collection. Fill a cuvet with the blank solution and place a cap on it. Once warm-up is complete place the cuvet into the SpectroVis device so that the smooth sides face the white circle and white arrow. Click “Finish Calibration”.
6. Once Calibration is complete, click OK.
7. Click the ‘Wave’ button. Set up the data collection to be ‘Absorbance vs. Time’ for an individual wavelength. Select the wavelength listed that is nearest to 450 nm. As the iodine is produced it has an absorbance maximum at this wavelength.
8. The default is to collect data at one data point per second for 200 seconds. This can be changed in the ‘Experiment Menu’ under the ‘Data Collection…’ menu item. Collect one data point per second for 120 seconds. Or just halt data collection once it gets to 120 s.
9. Using the quantities in the table at right perform three trials for each set of quantities.
10. Mix water and KI solution in the cuvet. Measure out the FeCl3 solution using the syringe. When ready, add the iron(III) chloride solution, cap the cuvet, shake vigorously, and place the cuvet in the spectrophotometer.
11. Immediately begin data collection by pressing the ’Collect‘ button. Try to perform these actions as repeatably as possible so that data collection always begins after the same delay to measure the FeCl3 solution, the same amount of shaking, and the same time before placing the cuvet in the device.
12. Immediately enter the data into a spreadsheet program in order to generate a graph and determine the slope of the line.
13. The graph of all of the data will show a curve from a steep upward slope over toward a horizontal line. This is because the production of iodine (I2) will slow down as the reaction progresses. Use the spreadsheet software to make a linear trendline which displays the slope of the graph for the interval from 10 s to 70 s. This span may be adjusted if it seems necessary so use your judgement. This will be an average rate over the time span but if you always use the same interval then you can compare trials and experiments.
14. Calculate a slope for all three trials for each set of reactant quantities. Find an average slope for each set of trials with the same quantities of reactants. Only use this average in your analysis.
15. If a trial produces a slope that is clearly different from the other two then either repeat the trial or discard it from your analysis.
16. Continue data collection and analysis until three trials for each set of quantities have been performed and you have a reasonable average slope for the absorbance vs. time graph.

### Analysis

Perform the following calculations and answer the following questions. Some questions are specific to your lab data and results and some are more general and probe your knowledge of the concepts central to a deep understanding of chemical kinetics. You may not have collected data relevant to the answers to all questions.

1. Provide one sample graph for each experiment number.
2. Provide a data table showing the values of the slope for each trial and the average.
3. In your own words, briefly describe the lab techniques you used in this lab to collect data to determine the order of reaction of each reactant in order to be able to write the differential rate law.
4. Why in this lab did we work on finding the differential rate law and not the integrated rate law?
5. Calculate the order of each reactant: Show one sample calculation for each reactant.
6. Using initial concentrations calculate the rate constant for this reaction. Take the slope of your graphs to be the rate in M/s. (It isn’t but it is proportional to the rate). Is the rate constant the same (or at least close) for all your experiments?
7. From the data in the lab it is not possible to calculate the value of the rate constant directly. Explain what other data you would need.
8. What is the differential rate law for this reaction? Write it in the form: Rate = k[I]a[Fe3+]b.
9. Comment on the reliability of your data. Were your experiments consistent for a given set of initial conditions? If not, why not? Discuss.
10. When you found the slope of the graph were you calculating an instantaneous rate or an average rate? Justify your answer.
11. How could you change the experiments to improve the data?
The following questions are generalized questions about chemical kinetics and do not depend on data collected in the course of carrying out your lab work:
1. Mathematically speaking, how does the reaction rate depend on the concentrations of the reactants? In other words, as concentration increases, what happens to the rate?
2. Why does the reaction rate change as you alter the experimental conditions by changing the initial concentrations? Give an answer by referring to events at the molecular level.
3. What is the difference between the reaction rate and the rate constant?
4. What is the effect of a catalyst on the measured rate of the reaction? How does a catalyst affect the activation energy of the reaction? Why does it have this effect?
5. What effect does a catalyst have on the rate constant?
6. Why is the reaction rate dependent on the temperature at which the reaction takes place?

There are notes available to help students understand important background information, data analysis procedures, and calculations.

Based on a lab provided in the book Flinn Scientific Spectrophotometer Laboratory Manual ©1996 ISBN 1-877991-33-3

Last updated: Mar 08, 2019 Home