Purpose: The purpose of this lab is to use microscale techniques to:

o find the volume of one drop of watero determine the reaction rate and the total rate law of a reaction

involving the oxidation of iodide ions by bromate ions in the presence of acid

o determine the order of each individual reactant by varying the concentration of each reactant individually

o determine the activation energy of a reaction involving the oxidation of iodide ions by bromate ions in the presence of acid by repeating the experiment at several different temperatures

o calculate the rate constants at the different temperatureso observe the effect of a catalyst on the rate

Background: This experiment in 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:

6I-(aq) + BrO3

-(aq) + 6H+

(aq) 3I2(aq) + Br-(aq) + 3H2O(l) Reaction 1

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 the reactants. If the rate of the reaction is expressed as the rate of decrease in concentration of bromate ion, the rate law has the form:

Rate = -∆[BrO3]/∆t = k[I-]x[BrO3-]y[H+]z Equation 1

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. -∆[BrO3

-], divided by the change in time for the reaction to occur, ∆t. 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 on e 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:

I2(aq) + 2S2O3-2

(aq) 2I-(aq) + S4O6

-2(aq) Reaction 2

Reaction 1 is somewhat slow. Reaction 2 is extremely rapid, so that as quickly as iodine is produced in reaction1, it is consumed in reaction 2. 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 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, Ea. 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 realted to the rate constant by the Arrhenius equation:

k = Ae-Ea/RT

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.

Materials: 5mL 0.1M cupric nitrate

solution, Cu(NO3) 5mL distilled water, 5mL 0.10M HCl solution 5mL 0.010M potassium

iodide, KI solution 5mL 0.040M potassium

bromate, KBrO3solution 5mL 2% starch solution 5mL sodium

thiosulfate,Na2S2O3 solution Cassette tape case Timer, seconds

Analytical balance (0.001g precision)

(1) 10ml beaker (7) Beral-type pipets with

microtip Cotton swaps Toothpicks Thermometer, 0◦-100◦C (2) 12-well reaction strips Trough 7 different colors of label

tape

Procedure:

Part 1 – Find the Volume of One Drop of Solution1) Obtain a microtip Beral-pipet. Fill the pipet with approx. 3mL of water.2) Mass a small beaker using an analytical balance. Record the mass.3) Holding the pipet vertically, deliver 5 drops of water into the beaker, nad find the total mass. Record the data.4) Add an additional 5 drops of water into the beaker, and again determine the mass. Record the value.5) Deliver 5 more drops and again find the mass. Record the data.

Part 2 – Determine the Reaction Rate and Calculate the Rate Law1) Obtain 6 microtip pipets and fold an adhesive label around the stem of each pipet. Label the pipets KI, H2O, Starch, Na2S2O3. 2) Fill each pipet with about 2mL of the appropriate liquid.3) Place the pipets in an opened cassette case for storage.4) Obtain 2 clean, 12-well reaction strips and arrage them so that the numbers can be read from left to right.5) Record the temperature of the one of the reaction solutions. 6) Using a provided table, 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 will be run in triplicate.

7) To wells 1-9 in the second reaction strip, add 2 drops of 0.040M KBrO3. The timer should start immediately when the first drop of KBrO3 is added.8) Record the time when the solution in each cell turns blue.9) 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.

Part 3 – Determine the Activation Energy1) Prepare a shallow warm water bath of about 40◦C.2) Using the provided table as a guide, fill each of the first 6 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 0.040M KBrO3 solution.5) Place this pipet in the warm temperature water bath for 5 minutes.6) With the reaction strip still in the warm temperature bath, add 2 drops of KBrO3 solution to the first well, stir, and immediately start the timer. 7) Record the time in seconds when the solution turns blue.8) Repeat steps 7-9 for the reaction solutions in wells #2 and #3.9) Remove he reaction strip and the pipet from the warm temperature bath.10) Add ice cubes and water to create a cold temperature water bath.11) Place both the reaction strip and Beral-type pipet containing the KBrO3

solution into the cold temperature water bath.12) Measure the temperature of the water bath with a thermometer and record the value.13) Repeat steps 7-9 for wells #4, #5, and #6. Record the time in seconds for each reaction.

Part 4 – Observe the Effect of a Catalyst on the Rate1) Repeat the procedure in Part 2 for Experiment 1 only, but this time add 1 drop of 0.1M 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. 2) Record the reaction times in the Part 4 data table.3) Dispose of the pipettes in a Ziploc bag.

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.

Observations:As thiosulfate ion was reacted, the solutions began turning blue. After

four minutes, the solutions reached a deep blue/indigo hue.

Analysis:The average masses of 1 drop of water was calculated by averaging the

average mass of 1 drop of water calculated in the 3 trials, in which the mass of a beaker plus 5 drops of water was subtracted from the mass of the same beaker when empty. The volume of one drop of solution was calculated through dimensional analysis by multiplying the average masses of 1 drop of water by the conversion factor of the weight of 1mL of water, 1.0g, and then dividing the product by 1000mL to calculate the volume in liters.

The mols of S2O3-2 was calculated by multiplying the mols of Na2S2O3

by 1, as there is a 1:1 stoichiometric ratio between mols of Na2S2O3 and mols of S2O3

-2. Mols of BrO3

- was calculated by dividing mols of mols of S2O3-2 by 2

to calculate the mols of I2 because there is a 2:1 stoichiometric ratio between mols of S2O3

-2 and mols of I2, and then dividing the quotient by 3 because there is a 3:1 stoichiometric ratio between mols of I2 and mols of BrO3

-. The value of -∆[BrO3] was calculated by dividing the mols of BrO3

- reacted by the volume of 12 drops. The rate of each reaction was found by dividing -∆[BrO3] by the number of seconds for the reaction to occur.

The initial concentration of each reactant for each experiment was calculated by using the proportion Vconcentrated X Mconcentrated = Vdilute X Mdilute, where Vconcentrated and Mconcentrated are the volume and molarity of the starting, concentrated solutions, and Vconcentrated and Mconcentrated 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.

To calculate the order of iodide ion, “x,” Experiments 1 and 2 were compared so that the concentrations of the species of the reaction were substituted in the differential rate law equation, R = k[I-]x[BrO3

-]y[H+]z

To calculate the order of bromate ion, “y,” Experiments 1 and 4 were compared so that the concentrations of the species of the reaction were substituted in the differential rate law equation, R = k[I-]x[BrO3

-]y[H+]z

To calculate the order of hydrogen ion, “z,” Experiments 1 and 6 were compared so that the concentrations of the species of the reaction were substituted in the differential rate law equation, R = k[I-]x[BrO3

-]y[H+]z

Once the order of each species involved in the reaction was determined, the rate constant of each experiment was calculated by substituting data from each respective experiment into the rate law equation, R = k[I-]x[BrO3

-]y[H+]z

The rate of each reaction was calculated by dividing -∆[BrO3] by the average reaction time.

The activation energy was calculated by graphing the natural log of the rate constant on the y-axis versus the inverse of the temperature in Kelvins on the x-axis to determine the slope of linear regression, and then equating the slope to –Ea/R, where Ea is the activation energy and R = 8.31 J/mol X K.

Post-Lab Questions:1) Why does the reaction rate change as concentrations of the reactants change?

The reaction rate changes as concentrations of the reactants change because of the direct relationship that exists between the concentrations of the reactants and the reaction rate through the differential rate law.

2) Explain the general procedure used to find the rate law.The general procedure to find the rate law is to first determine the

reaction order of each reactant (by conducting several trials in which the concentration of one particular reactant is varied while the other concentrations are kept constant), which will then allow the calculation of the rate constant. 3) Why does the reaction rate change as temperature changes?

The reaction rate changes as temperature changes because the kinetic energy of molecules and atoms is directly related to temperature.

4) Explain the general procedure used to determine the activation energy.The general procedure to determine the activation energy is graphing

the natural log of the rate constant, ln(k), on the vertical axis versus 1/T (temperature in Kelvins) on the horizontal axis using a graphing calculator and calculating the best-fit linear regression. The slope = –Ea/R, where Ea is the activation energy and R = 8.314J/mol X K, to calculate the activation energy for the reaction.

5) Differentiate between reaction rate and the specific rate constant.The reaction rate is the rate at which the moles of substance change

that varies with both temperature and concentration of the reactants. The

specific rate constant is a proportionality constant that will vary only with temperature.

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

The addition of a catalyst increased the rate of the reaction. This is because the function of catalysts in chemical reactions is to effectively lower the activation energy barrier that a reaction must overcome to take place.

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 order 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 collected was consistent, as the calculated order was close to integers, the check of the order gave the same value for the order, and the points of the graph were close to a straight line; this statement was made possible by impeccable lab technique.

Conclusion:The volume of one drop of water was experimentally determined to be

1.1 X 10-5 L. This value was achieved through the mathematical principle of

averaging multiple values by adding values and dividing the sum by the number of values.

[I-] was experimentally determined to be 1st order; [BrO3-] was

experimentally determined to be 1st order; and [H+] was experimentally determined to be 2nd order. Varying the concentration of one reactant individually at any particular time while keeping the concentration other two reactants constant enabled the mathematical calculation of the order of each reactant.

The average reaction rate was experimentally determined to be 2.0 X 10-7 M/s. This value was calculated by averaging the quotients of the -∆[BrO3] divided by the number of seconds for each of the 7 reactions to take place.

The activation energy of the reaction at was experimentally determined to be 28 kJ/mol. This was done by equating the slope of the best-fit linear regression (-3393K) to –Ea/R, where Ea is the activation energy and R = 8.31 J/mol X K. Repeating the experiment at several different temperatures allowed for the calculation of these values of energy; whereas at lower temperatures molecular particles collide with less energy than required to overcome the activation energy barrier and therefore bounce apart, resulting the lack of a reaction, or result in a reaction that does take place but proceeds relatively

slowly, higher temperatures cause molecular particles to collide with energies equal to or greater than the activation energy, which result in a rapid reaction. One of the postulates for the Kinetic Molecular Theory is that the average kinetic energy of the gas particles depends only on the temperature of the system; thus, temperature is directly related to the rate of the reaction.

The rate constant was experimentally determined to be 22 M-3s-1. This value was calculated using the differential rate law and equating the experimentally determined rate of the reaction to the product of k and the concentrations of each reactant raised to their respective orders, and then averaging the calculated rate constants for each trial.

The total rate law of a reaction involving the oxidation of iodide ions by bromate ions in the presence of acid was determined to be R = k[I-][BrO3

-][H+]2. This equation was determined by calculating the individual order of each reactant and the rate constant k of the reaction and relating them using the mathematical relationship stated by the differential rate law.

The addition of a catalyst increased the rate of the reaction; the reaction time decreased 36 seconds from 111 seconds uncatalyzed to 75 seconds catalyzed. This is because the function of catalysts in chemical reactions is to effectively lower the activation energy barrier that a reaction must overcome to to initiate.