Teacher Notes
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Teacher Notes![]() Kinetics of a ReactionStudent Laboratory KitMaterials Included In Kit
Copper( II ) nitrate solution, Cu(NO3)2, 0.1 M, 75 mL
Hydrochloric acid solution, HCl, 0.10 M, 75 mL Potassium bromate solution, KBrO3, 0.040 M, 75 mL Potassium iodide, KI, 5 g Sodium thiosulfate solution, Na2S2O3, 0.0010 M, 75 mL Soluble starch, 2 g Cotton swabs for cleaning well plates Pipets, Beral-type, microtip, 100 Toothpicks for stirring Additional Materials Required
Water, distilled or deionized, H20, 75 mL
Balance, 0.001-g precision Beakers, 10- or 50-mL, 12 Cassette tape cases, 12 Label tape for pipets Marking pens, 12 Reaction strips, 12-well, 24 Thermometers, microscale, 0–100 °C, 12 Timers, 12 Troughs for hot & cold water baths, shared Prelab PreparationPotassium 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. Safety PrecautionsDilute 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. DisposalPlease 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
Teacher Tips
Answers to Prelab QuestionsAnother version of the iodine clock reaction involves reaction of iodide ions with persulfate ions (Reaction 3). {13816_PreLabAnswers_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.
{13816_PreLabAnswers_Table_1}
Sample DataPart 1. Measure the Volume of One Drop of Solution {13816_Data_Table_8}
Calculate the volume of one drop of solution. Assume the denstity of water to be 1.00 g/mL.
{13816_Data_Equation_3}
Part 2. Determine the Reaction Rate and Calculate the Rate Law {13816_Data_Table_9}
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 S2O32– present in one drop: {13816_Data_Equation_4}
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 S2O32– ion is used up:
{13816_Data_Equation_5}
The value of –Δ[BrO3–] in all reactions, since all experiments have a total volume of 12 drops, is:
{13816_Data_Equation_6}
The rate of each reaction can be found by dividing –Δ[BrO3–] by the number of seconds for the reaction to take place.
{13816_Data_Equation_7}
Calculate the rate of reaction in each experiment and enter the results into the following table. Use the average time for each experiment. Moles of S2O32–
{13816_Data_Equation_8}
Moles of BrO3–
{13816_Data_Equation_9}
–Δ[BrO3–]
{13816_Data_Equation_10}
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. For each calculation, the average time for each experiment from Data Table 2 was used. {13816_Data_Equation_11}
{13816_Data_Table_10}
Calculate Initial Concentrations Calculate the initial concentration of each reactant for each experiment. These are the concentrations of each reactant after all the reactants have been mixed, but before any reaction has taken place. 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 = Vconcentrated x Mconcentrated = Vdilute x Mdilute where Vconcentrated and Mconcentrated are the volume and molarity of the starting, concentrated solutions, and Vdilute and Mdilute 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 as follows: {13816_Data_Equation_12}
Find the initial concentration of each reactant and record in the following data table.
{13816_Data_Table_11}
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: {13816_Data_Equation_13}
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
{13816_Data_Equation_14}
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. To calculate the order of iodide ion, “x”, first compare Experiments 1 and 2.
Substitute data from each experiment into the rate law equation to find the value of k. Report the average value of k. Do not forget to include proper units for k. A sample calculation for Experiment 1 is given. Experiment 1: Rate = k[I–][BrO3–] [H+]2 {13816_Data_Equation_21}
{13816_Data_Table_12}
Average value of k = 25 M–3s–1 Part 3. Determine the Activation Energy {13816_Data_Table_13}
Calculate the Activation Energy, Ea Using the data from Part 3, calculate the values listed in the following table for each measured temperature. {13816_Data_Table_14}
Calculations 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. {13816_Data_Equation_22}
Rate = 4.8 x 10–8 M/s Calculate the rate constant for each reaction, using the rates calculated above, the initial concentrations, and the orders for each reactant. A sample calculation for the experiment at approximately 0 °C is given. Rate = k[I–] [BrO3–] [H+]2 {13816_Data_Equation_23}
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 = –Ea/R, where Ea is the activation energy and R = 8.314 J/mol•K, to calulate the activation energy for the reaction. 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. {13816_Data_Figure_3}
{13816_Data_Equation_24}
Part 4. Observe the Effect of a Catalyst on the Rate of Reaction
{13816_Data_Table_15}
Answers to QuestionsPost-Laboratory Review
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Student Pages
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Student Pages![]() Kinetics of a ReactionIntroductionHow 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
BackgroundThis 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: {13816_Background_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 reactants. If the rate of reaction is expressed as the rate of decrease in concentration of bromate ion, the rate law has the form:
{13816_Background_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, –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: {13816_Background_Reaction_2}
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 S2O32– 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, Ea. {13816_Background_Figure_1}
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. 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: k = Ae–Ea/RT 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: {13816_Background_Equation_2}
where ln k is the natural logarithm of the rate constant, Ea 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 –Ea/R. By determining the slope, the activation energy can be calculated. Part 4. Observe the effect of a catalyst on the rate of reaction. The catalyst used is copper(II) nitrate solution. Experiment OverviewThe 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: {13816_Overview_Reaction_1}
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
Copper( II ) nitrate solution, Cu(NO3)2, 0.1 M, 5 mL
Hydrochloric acid solution, HCl, 0.10 M, 5 mL Potassium bromate solution, KBrO3, 0.040 M, 5 mL Potassium iodide solution, KI, 0.010 M, 5 mL Sodium thiosulfate solution, Na2S2O3, 0.0010 M, 5 mL Starch solution, 2%, 5 mL Water, distilled or deionized, H2O, 5 mL Balance, 0.001-g precision Beaker, 10-mL or 50-mL Cassette tape case Cotton swabs for cleaning well plate Label tape, for pipets Marking pen Pipets, Beral-type, microtip, 7 Reaction strips, 12-well, 2 Thermometer, microscale, 0–100 °C Timer, seconds Toothpicks for stirring Trough for hot and cold water baths Prelab QuestionsAnother version of the iodine clock reaction involves reaction of iodide ions with persulfate ions (Reaction 3). {13816_PreLabAnswers_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.
{13816_PreLabAnswers_Table_1}
Safety PrecautionsDilute 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. ProcedurePart 1. Measure the Volume of One Drop of Solution
Part 2. Determine the Reaction Rate and Calculate the Rate Law {13816_Procedure_Table_2}
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.
Part 3. Determine the Activation Energy
Part 4. Observe the Effect of a Catalyst on the Rate Student Worksheet PDF |