Materials Included In Kit

Additional Materials Required

Prelab Preparation

Potassium iodide solution, 0.010 M: To prepare 100 mL of 0.010 M potassium iodide, KI, dissolve 0.17 grams of KI in approximately 50 mL of distilled water. Dilute to 100 mL with distilled water and mix. Five milliliters is sufficient for a team of two students. Preparing 100 mL will give enough solution for 20 students working in pairs. Potassium iodide solution does not keep well and should be prepared fresh.

Starch solution, 2%: Prepare 100 mL of starch solution by making a smooth paste of 2 g soluble starch and 10 mL distilled water. Pour the paste into 90 mL of boiling water while stirring. Cool to room temperature before using. Starch solution has a poor shelf life. Dispose of it after the experiment, as it will form mold if kept too long. Prepare fresh for use. Five milliliters is sufficient for a team of two students. Preparing 100 mL will give enough solution for 20 students working in pairs.

Safety Precautions

Dilute hydrochloric acid solution is severely irritating to skin and eyes and is slightly toxic by ingestion and inhalation. Dilute copper(II) nitrate solution is irritating to skin, eyes and mucous membranes and slightly toxic by ingestion. Dilute potassium bromate solution is irritating to body tissue and slightly toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please consult current Safety Data Sheets for additional safety 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. All reaction solutions, the starch solution, the potassium iodide solution, the copper(II) nitrate solution and the potassium bromate solution may be disposed of according to Flinn Suggested Disposal Method #26b. Any solid starch and potassium iodide may be disposed of according to Flinn Suggested Disposal Method #26a. The hydrochloric acid solution may be disposed of according to Flinn Suggested Disposal Method #24b. The sodium thiosulfate solution may be disposed of according to Flinn Suggested Disposal Method #12b.

Lab Hints

• Two to three 50-minute periods are needed to carry out the experiment. Parts 1 and 2 can be completed in the first 50-minute period, with Parts 3 and 4 completed in the second 50-minute period.
• Two hours are required to prepare the solutions and sets of reagents.
• The use of microscale techniques in this experiment has many advantages—setup is easy, different concentrations can be tested simultaneously, reaction times are short, and multipletrials can be performed within a typical 50-minute laboratory period. There are, however, some disadvantages to the use of microscale techniques in kinetics, and these disadvantages should be accounted for in the instructional process. Quantities of reactants must be carefully measured using microtip pipets. Temperatures of reactants must also be controlled. Since only twelve drops of reactants are used in each experimental measurement, care must be used in adding and mixing the reagents. Students can quickly repeat any questionable measurements. They must evaluate the data and determine if additional measurements are needed.
• Consider using Experiment 1 as a practice run to allow students to become comfortable with the procedure in Part 2. The instructor should demonstrate the “shakedown” technique for the students before they begin the lab.
• The time of the reaction is the time from when the potassium bromate solution, KBrO₃, is introduced until the first tinge of blue color appears. It is helpful to observe the wells against a white background.

Teacher Tips

• Students must have an understanding of the purpose of the various steps. They should recognize that Experiment 2 has double the potassium iodide concentration while all others remain constant, while Experiment 3 has triple the potassium iodide concentration. Since reaction orders are generally whole numbers, it is to be expected that Experiment 2 should take one-half as long as Experiment 1 if the reaction is first order in potassium iodide, and should take one-fourth as long if the reaction is second order. Similar relationships occur with the other dilutions. If values do not follow the expected trends, experiments need to be repeated until reproducible values are obtained.

• This experiment may work better with students working in groups of three rather than in pairs. The “iodine clock” changes color in quick succession in each series of wells, making it easier for a team of three to measure and record the reaction times for 9 or all 12 wells.
• It is convenient, but not necessary, to dispense the solutions in Beral microtip pipets. Sets of the reagents can be stored in a cassette tape box. Store with the pipet tips pointing up. A cassette box of the pipets containing the potassium iodide, potassium bromate, sodium thiosulfate, hydrochloric acid, starch, copper nitrate, and distilled water is most convenient for storing the solutions and for use by the students.
• Sensitive balance. A balance which is capable of reading ±0.001 g or better is needed to get accurate measurements of the mass of the small drops delivered with the Beral pipets. If only centigram balances are available, change the procedure which calls for measuring the mass of 5 drops of water to measure the mass of 50 drops of water.
• Trough for hot and cold water baths. A low, flat container is needed to heat or cool the 12-well strips in Part 3. Troughs work well because they heat or cool the small quantities of reactants fairly quickly. Petri dishes may also work well.

Answers to Prelab Questions

Another version of the iodine clock reaction involves reaction of iodide ions with persulfate ions (Reaction 3).

The following rate data were collected by measuring the time required for the appearance of the blue color due to the iodine-starch complex.

1. In each trial, the blue color appeared after 0.0020 M iodine (I₂) had been produced. Calculate the reaction rate for each trial by dividing the concentration of iodine formed by the reaction time.

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 to 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 for 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 this version of 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.

Sample Data

Part 1. Find the Volume of One Drop of Solution

Part 2. Determine the Reaction Rate and Calculate the Rate Law

Part 3. Determine the Activation Energy

Part 4. Observe the Effect of a Catalyst on the Rate

To observe the effect of the catalyst, compare the reaction times for the uncatalyzed experiment (Experiment 1 from Part 2) and the catalyzed reaction. The uncatalyzed time was 173 seconds, while the catalyzed time was 52 seconds, indicating that the reaction occurred about 3.3 times faster with the catalyst.

Calculations

Part 1. Calculate the Volume of One Drop of Solution

Part 2A. Calculate the Rate

The rate will be expressed as –Δ[BrO₃^–]/Δt. In each reaction there is one drop of 0.0010 M Na₂S₂O₃ solution. Calculate the number of moles of S₂O₃^2– present in one drop:

The blue color begins to appear when all the thiosulfate ion is consumed. Examination of reactions (1) and (2) allows us to calculate the moles of BrO₃^– which react as all of the S₂O₃^2– ion is used up:

The value of –Δ[BrO₃^–] in all reactions, since all experiments have a total volume of 12 drops, is:

The rate of each reaction can be found by dividing –Δ[BrO₃^–] by the number of seconds for the reaction to take place.

Calculate the rate of reaction in each experiment and enter into the data table which follows. Use the average data for each experiment.

Moles of S₂O₃^2–

moles S₂O₃^2– = 1.9 x 10^–8 moles S₂O₃^2–

Moles of BrO₃^–

moles BrO₃^– reacted = 3.2 x 10^–9 moles BrO₃^–

–Δ[BrO₃^–]

Rate

Find the rate of each reaction by dividing –Δ[BrO₃^–] by the number of seconds it took for the reaction to occur. A sample calculation for Experiment 1 is given below. For each calculation, the average time for each experiment from Data Table 2 was used.

Part 2B. Calculate Initial Concentrations

Calculate the initial concentration of each reactant for each experiment. This will not be the same as the concentration of the starting solution because combining the reactants dilutes all of the solutions. On dilution, the number of moles of reactant stays the same, Therefore:

no. moles = V_concentrated x M_concentrated = V_dilute x M_dilute

where V_concentrated and M_concentrated are the volume and molarity of the starting, concentrated solutions, and V_dilute and M_dilute are the volume and molarity of the diluted reaction mixtures. Since volumes will be proportional to the number of drops of solution used, the number of drops substitute for volumes.

For example, in Experiment 1 the initial [I^–] is found in this way:

Find the initial concentration of each reactant and record in the data table.

Part 2C. Calculate the Order of Each Reactant

Next, the values for the exponents x, y and z need to be determined. The experiment is designed so that the concentration of one ion changes while the others remain constant. Comparing values in Experiments 1, 2 and 3, we see that Experiment 2 has double the I^– concentration as Experiment 1, and Experiment 3 has triple the I^– concentration as Experiment 1.

Substitute the concentration values for Experiments 1 and 2 into the equation:

Rate = k[I^–]^x[BrO₃^–]^y[H⁺]^z

Exp. 1: Rate₁ ______ = k[    ]^x [    ]^y [    ]^z

Exp. 2: Rate₂ ______ = k[    ]^x [    ]^y [    ]^z

Divide the first equation by the second. Notice that most of the terms will cancel out and the ratio reduces to:

Divide and solve for x. Report the value of x to the nearest integer. Repeat the calculations using Experiments 1 and 3 to confirm the value for x.

Note

To solve for an exponential value, take the logarithm of both sides of the equation.

For example:     8 = 2ⁿ     log 8 = n log 2      

Next use the same procedure with Experiments 1, 4 and 5 to find the value of y. Lastly, use Experiments 1, 6 and 7 to find the value of z. Show how the calculations are carried out.

1. To calculate the order of iodide ion, “x”, first compare Experiments 1 and 2.

Rate = k[I^–]^x[BrO₃^–]^y[H⁺]^z

0.48 = 0.50^x

log 0.48 = x log 0.50

Comparison of Experiments 1 and 3 yields x = 1.0.

2. To calculate the order of bromate ion, “y”, first compare Experiments 1 and 4.

Rate = k[I^–]^x[BrO₃^–]^y[H⁺]^z

0.51 = 0.52^y

log 0.51 = y log 0.52

Comparison of Experiments 1 and 5 yields y = 0.94.

3. To calculate the order of hydrogen ion, “z”, first compare Experiments 1 and 6.

Rate = k[I^–]^x[BrO₃^–]^y[H⁺]^z

0.25 = 0.52^z

log 0.25 = z log 0.52

Comparison of Experiments 1 and 7 yields z = 2.0.

Part 2D. Find the Rate Constant

Substitute data from each experiment into the rate law equation to find the value of k. Report the average value of k. Don’t forget to include proper units for k. A sample calculation for Experiment 1 is given.

Experiment 1: Rate = k[I^–]^x[BrO₃^–]^y[H⁺]²

Average value of k = 25 M^–3s^–1

Part 3. Calculate the Activation Energy, E_a

Using the data from Part 3, calculate the values listed in the table below for each measured temperature.

Part 3A. Calculation of the Rates

Determine the rate of each reaction dividing –Δ[BrO₃^–] by the average reaction time, just as in Part 2A. From Part 2A, –Δ[BrO₃^–] = 1.4 x 10^–5 M for each reaction. A sample calculation for the experiment at approximately 0 °C is given.

Rate = 4.8 x 10^–8 M/s

Part 3B. Calculation of the Rate Constant

Calculate the rate constant for each reaction as in Part 2D, using the rates calculated above, the initial concentrations listed in Part 2B, and the orders for each reactant determined in Part 2C. A sample calculation for the experiment at approximately 0 °C is given.

Rate = k[I^–][BrO₃^–][H⁺]²

Part 3C. Calculation of the Activation Energy

Graph the natural logarithm of the rate constant, ln k, on the vertical axis versus 1/T (temperature in the Kelvin scale) on the horizontal axis. Draw a straight line that is closest to the most points, and determine the slope of the line. Use the data points at (1/277 K) and (1/303 K) to calculate the slope because they lie closest to the line. The slope = –E_a/R, where E_a is the activation energy and R = 8.314 J/mol•K, to calculate the activation energy for the reaction.

Slope = –E_a/R
E_a = –slope x R = –(–3200 K x 8.314 J/mol•K) = 27000 J/mol = 27 kJ/mol

Answers to Questions

1. Why does the reaction rate change as concentrations of the reactants change?

Reactions occur because of collisions of particles. As the concentrations of particles increase, the number of collisions also increases, and so the reaction occurs more rapidly. This is only true if the particles occur in the rate law.

2. Explain the general procedure used to find the rate law.

In the experiment the concentration of each individual reactant is varied, and the rate of the reaction determined. This allows the calculation of the order of each reactant. After the orders of reactants are known, the values are inserted into the rate law to calculate the rate constant.

3. Why does reaction rate change as temperature changes?

Reaction rates increase as the temperature rises because molecules are moving faster at higher temperatures. This means that molecules collide more frequently, and more importantly, a greater number of molecules possess the necessary kinetic energy to overcome the activation energy barrier and react to form products.

4. Explain the general procedure used to determine the activation energy.

The determination of activation energy is based on the Arrhenius equation: ln k =(–E_a/RT) + ln A. This equation corresponds to the straight line equation: y = mx + b. If a graph is made plotting ln k on the y-axis and 1/T on the x-axis, a straight line should result with –E_a/R as the slope and ln A as the y intercept.

5. Differentiate between reaction rate and specific rate constant.

The reaction rate is the rate of change in concentration of a reactant versus the time the reaction has proceeded. Reaction rates almost always change as reactions proceed. The rate constant, k, is a constant which allows the calculation of the rate if the concentrations of reactants are known and the rate law is known.

6. Comment on the effect of the catalyst. Predict the activation energy changes when a catalyst is added to the reaction.

The catalyst decreased the time of reaction from 173 seconds to 52 seconds. A catalyst works by providing a different reaction mechanism in which the activation energy is lowered, so that more molecules are able to collide with sufficient kinetic energy to overcome the activation energy barrier and react to form products.

7. Make a general statement about the consistency of the data as shown by calculating the orders of reactants, and by the graphical analysis which leads to activation energy. Were the calculated orders close to integers? Did the check of the order give the same value for the order? Were the points on the graph close to a straight line?

The data were fairly consistent, giving values close to integers for the orders of the reactants. The check of the reaction orders with Experiment 8 agreed within 7%. There was more scatter in the graphical analysis calculating activation energy. The difficulty in being sure of a constant temperature is probably responsible for this difference.

Discussion

AP Standards: This lab fulfills the requirements for the College Board Recommended AP Experiment 12: Determination of the Rate of Reaction and Its Order.

Introduction

How fast will a chemical reaction occur? If a reaction is too slow, it may not be practical. If the reaction is too fast, it may explode. Measuring and controlling reaction rates makes it possible for chemists and engineers to make a variety of products, everything from antibiotics to fertilizers, in a safe and economical manner. The purpose of this experiment is to investigate how the rate of a reaction can be measured and how reaction conditions affect reaction rates.

Concepts

• Rate law
• Activation energy
• Reaction rates
• Rate constant
• Catalyst

Background

This experiment is designed to study the kinetics of a chemical reaction. The reaction involves the oxidation of iodide ions by bromate ions in the presence of acid:

The reaction is somewhat slow at room temperature. The reaction rate depends on the concentration of the reactants and on the temperature. The rate law for the reaction is a mathematical expression that relates the reaction rate to the concentrations of reactants. If the rate of reaction is expressed as the rate of decrease in concentration of bromate ion, the rate law has the form:

where the square brackets refer to the molar concentration of the indicated species. The rate is equal to the change in concentration of the bromate ion, –Δ[BrO₃^–], divided by the change in time for the reaction to occur, Δt. The term “k” is the rate constant for the equation, which changes as the temperature changes. The exponents x, y and z are called the “orders” of the reaction with respect to the indicated substance, and show how the concentration of each substance affects the rate of reaction.

The total rate law for the process is determined by measuring the rate, evaluating the rate constant, k, and determining the order of the reaction for each reactant (the values of x, y and z).

To find the rate of the reaction a method is needed to measure the rate at which one of the reactants is used up, or the rate at which one of the products is formed. In this experiment, the rate of reaction will be measured based on the rate at which iodine forms. The reaction will be carried out in the presence of thiosulfate ions, which will react with iodine as it forms:

Reaction 1 is somewhat slow. Reaction 2 is extremely rapid, so that as quickly as iodine is produced in reaction 1, it is consumed in reaction 2. Reaction 2 continues until all of the added thiosulfate has been used up. After that, iodine begins to increase in concentration in solution. If some starch is present, iodine reacts with the starch to form a deep blue–colored complex that is readily apparent.

Carrying out reaction 1 in the presence of thiosulfate ion and starch produces a chemical “clock.”When the thiosulfate is consumed, the solution turns blue almost instantly.

In this laboratory procedure, all of the reactions use the same quantity of thiosulfate ion. The blue color appears when all the thiosulfate is consumed. An examination of reactions 1 and 2 shows that six moles of S₂O₃^2– are needed to react with the three moles of I₂ formed from the reaction of one mole of BrO₃^–. Knowing the amount of thiosulfate used, it is possible to calculate both the amount of I₂ that is formed and the amount of BrO₃^– that has reacted at the time of the color change. The reaction rate is expressed as the decrease in concentration of BrO₃^– ion divided by the time it takes for the blue color to appear.

There is an energy barrier that all reactants must surmount for a reaction to take place. This energy can range from almost zero to many hundreds of kJ/mol. This energy barrier is called the activation energy, E_a.

Reactants need to possess this amount of energy both to overcome the repulsive electron cloud forces between approaching molecules and to break the existing bonds in the reacting molecules. In general, the higher the activation energy, the slower the reaction.

The activation energy is related to the rate constant by the Arrhenius equation:

Where A is the frequency constant and is related to the frequency of collisions, R is the universal gas constant, and T is the temperature in K.

Catalysts are substances that speed up a reaction, but are not consumed in the reaction. Catalysts work by lowering the overall activation energy of the reaction, thus increasing the rate of the reaction.

The experiment is designed so that the amounts of the reactants that are consumed are small in comparison with the total quantities present. This means that the concentration of reactants is almost unchanged during the reaction, and therefore the reaction rate is almost a constant during this time.

The experiment utilizes a microscale procedure. Only 12 drops of reactants delivered from capillary droppers are used for each measurement. A special microscale “shakedown” technique is used to mix the reactants. The steps involved are as follows:

Part 1. Measure the volume of a drop of solution. This must be done to determine the number of moles of thiosulfate ion in one drop. This will allow the moles of bromate ions that react to be calculated.

Part 2. Determine the reaction rate and calculate the rate law. This is done by carrying out an experiment at specific concentrations of each of the reactants and measuring the reaction rate. The concentration of one reactant is then changed and the reaction rate change is observed. This is repeated for each reactant. This data allows the calculation of the order of each reactant. Once the orders are known, the value of the rate constant can be determined.

Part 3. Determine the activation energy. Reaction rates generally increase as the temperature goes up. By measuring how the rate changes as the temperature is varied, the activation energy, Ea, for the reaction can be calculated. The natural log of the Arrhenius equation is:

where ln k is the natural logarithm of the rate constant, E_a is the activation energy, R is the gas constant, 8.314 J/mol•K, and T is the temperature on the kelvin scale. A is the frequency factor.

This equation follows the straight line relationship: y = mx + b. A plot of the natural logarithm of k versus 1/T will give a straight line graph. The slope of the graph is –E_a/R. By determining the slope, the activation energy can be calculated.

Part 4. Observe the effect of a catalyst on the rate of the reaction. The catalyst used is copper(II) nitrate solution.

Experiment Overview

The purpose of this experiment is to utilize a microscale technique to determine the total rate law for the oxidation of iodide ions by bromate ions in the presence of acid:

There are several steps in the experiment. First, the order for each of the reactants is found by varying the concentration of each reactant individually. Once the orders are known, the rate constant is calculated. Second, the activation energy is found by repeating the experiment at several different temperatures, measuring the rate, and calculating the rate constants at the different temperatures. A graph of the reciprocal of absolute temperature versus the natural logarithm of the rate constant allows the calculation of the activation energy. Last, a catalyst is added and the change in reaction rate is observed.

Materials

Prelab Questions

Another version of the iodine clock reaction involves reaction of iodide ions with persulfate ions (Reaction 3).

The following rate data was collected by measuring the time required for the appearance of the blue color due to the iodine–starch complex.

1. In each trial, the blue color appeared after 0.0020 M iodine (I₂) had been produced. Calculate the reaction rate for each trial by dividing the concentration of iodine formed by the reaction time.
2. Compare trials 1 and 2 to 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 for 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 this version of the iodine clock reaction. Could the rate law have been predicted using the coefficients in the balanced chemical equation? Explain.

Safety Precautions

Dilute hydrochloric acid solution is severely irritating to skin and eyes and is slightly toxic by ingestion and inhalation. Dilute copper(II) nitrate solution is irritating to skin, eyes and mucous membranes and slightly toxic by ingestion. Dilute potassium bromate solution is irritating to body tissue and slightly toxic by ingestion. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Find the Volume of One Drop of Solution

1. Obtain a microtip Beral-type pipet. Fill the pipet with approximately 3 mL of deionized water.
2. Mass a small beaker using an analytical balance. Record the mass in the Part 1 Data Table.
3. Holding the pipet vertically, deliver five drops of water into the beaker, and find the total mass. Record your data in the Part 1 Data Table.
4. Add an additional five drops of water into the beaker, and again determine the mass. Record this value in the Part 1 Data Table.
5. Deliver five more drops and again find the mass. Record the data in the Part 1 Data Table.

Part 2. Determine the Reaction Rate and Calculate the Rate Law

It is necessary to use consistently good technique to obtain reproducible data. Hold pipets vertically and be sure no air bubbles are introduced. Since such small quantities of reagents are used, it is very easy to repeat measurements. Calculation of the orders of reactants are all based on the values obtained for the first experiment, so be sure to get reproducible data from the beginning. All other experiments should be carried out at least twice.

Table 1 shows the reagent quantities to be used in carrying out the reactions needed. It is important to use care in measuring out the solutions. Since the total solution volume is quite small, even one extra drop can cause a substantial change in concentrations.

A study of Table 1 shows that all experiments contain the same total number of drops of solution. Only one drop of sodium thiosulfate, Na₂S₂O₃, and one drop of starch solution are added to each well. In Experiments 1, 2 and 3, the concentration of potassium iodide, KI, is gradually increased while all other solutions volumes remain constant. Experiments 1, 4 and 5 have an increasing concentration of potassium bromate, KBrO₃. Experiments 1, 6 and 7 show an increase in the concentration of hydrochloric acid, HCl.

Read the entire procedure before beginning the experiment.

1. Obtain six microtip pipets and fold an adhesive label around the stem of each pipet (see Figure 3). Label the pipets KI, H₂O, HCl, Starch, Na₂S₂O₃ and KBrO₃.
2. Fill each pipet with about 2 mL of the appropriate liquid.
3. Place the pipets in an opened cassette case for storage (see Figure 1).
4. Obtain two clean, 12-well reaction strips and arrange them so the numbers can be read from left to right.
5. The first determination will vary the concentration of only the KI. Using the following table as a guide, fill each numbered well in the first reaction strip with the appropriate number of drops of the reagent listed. Mix the solution in each well with a new toothpick. Each experiment in Table 1 will be run in triplicate.

Reaction Strip 1

6. To wells 1–9 in the second reaction strip, add 2 drops of 0.040 M KBrO₃.

7. Turn the second reaction strip upside down and place it on top of the first strip so that the numbered wells are lined up on top of each other (see Figure 2). Note: It seems strange at first, but the liquid will not flow out of the wells until the reaction strips are “snapped” downward in step 8. Surface tension prevents the liquid from flowing out of the wells.

8. Holding the aligned plates firmly together “open end to open end” as shown in Figure 2, shake them downward once vigorously with a sharp downward motion. This is done by dropping the hands as fast as possible and stopping abruptly. There is no upward motion—this is a “shakedown” technique. The lab partner should start timing immediately when the strips are snapped.
9. In the Part 2 Data Table, record the time when the solution in each cell turns blue.
10. When all of the cells have turned blue, take the temperature of one of the reaction solutions. Record this temperature for all the reactions in the Part 2 Data Table.
11. Rinse the contents of the well strips with warm water. Use a cotton swab to dry the inside of each well.
12. Repeat the entire process (steps 4–11) for the following combinations that cover experiments 4 and 5 and experiments 6 and 7, respectively.

Reaction Strip 1 (Experiments 4 and 5)

Reaction Strip 2

Reaction Strip 1 (Experiments 6 and 7)

Reaction Strip 2

Part 3. Determine the Activation Energy

In this part of the experiment, the reaction will be carried out at several different temperatures using the concentrations given in Part 2 for Experiment 1. The temperatures will be about 40 °C, 20 °C and 0 °C. Use data from Experiment 1 at room temperature for the second measurement.

1. Prepare a shallow warm water bath of about 40 °C.
2. Using the following table as a guide, fill each of the first six wells in the reaction strip with the appropriate number of drops of the reagent listed. Mix the solutions well with a new toothpick.

3. Place the reaction strip in the warm temperature bath.
4. Fill the Beral pipet labeled KBrO₃ half-full with 0.040 M KBrO₃ solution.
5. Place this pipet in the warm temperature water bath for at least five minutes.
6. Measure the temperature of the water bath with a thermometer and record the value in the Part 3 Data Table.
7. Take the pipet out of the water bath and dry the outside of the pipet.
8. With the reaction strip still in the warm temperature bath, add two drops of KBrO₃ solution to the first well, stir, and immediately start the timer. Place the pipet back in the warm temperature bath.
9. Record the time, in seconds, when the first blue color appears.
10. Repeat steps 7–9 for the reaction solutions in wells 2 and 3.
11. Remove the reaction strip and the pipet from the warm temperature bath.
12. Add ice cubes and water to create a cold temperature water bath.
13. Place both the reaction strip and Beral-type pipet containing the KBrO₃ solution into the cold temperature water bath.
14. Measure the temperature of the water bath with a thermometer and record the value in the Part 3 Data Table.
15. Repeat steps 7–9 for wells 4, 5 and 6. Record the time, in seconds, for each reaction in the Part 3 Data Table.

Part 4. Observe the Effect of a Catalyst on the Rate

Repeat the procedure given in Part 2 for Experiment 1 only, but this time add 1 drop of 0.1 M cupric nitrate solution, Cu(NO₃)₂, and only 3 drops of distilled water to the mixture. Fill only the first reaction wells. The total volume will still be 12 drops. Record the reaction times in the Part 4 Data Table.

Student Worksheet PDF

10533_Student1.pdf