The Rate of a Chemical Reaction: Chemical Kinetics A.P. Chemistry
Purpose: To determine the rate law for the reaction: - + H2O2 + 2I + 2H I2 + 2H2O The rate law will take the form: - + Rate = k[H2O2][I ][H ]
Background: Two important questions may be asked about a chemical reaction: 1. How far or how completely do the reactants interact to yield products? 2. How fast is the reaction? The first question, how far, is a question of chemical equilibrium which will be studied later on in the year. The second question, how fast, is the realm of chemical kinetics, the subject of this experiment.
In this experiment we will see how changing the concentrations of each of the reactants affects the rate law for the reaction. We will also measure the effect of both temperature and catalytic activity on the rate.
The Experimental Method In this experiment we will study the rate of oxidation of iodine ion by hydrogen peroxide, which proceeds according to the equation: - + H2O2 + 2I + 2H I2 + 2H2O By varying the concentrations of each of the three reactants we will determine the order of the reaction with respect to each reactant.
To measure the reaction rate we will employ a clever variation of the initial rate method. A small amount of sodium thiosulfate and starch indicator is added to the reaction mixture. Thiosulfate ion does not react at an appreciable rate with any of the reactants, but it does react rapidly with iodine, I2 according the equation: 2- - 2- 2S2O3 + I2 2I + S4O6 As a result, iodine is reduced back to iodide ion as fast as it is formed until all of the thiosulfate is used up. At this point the solution suddenly turns blue because the iodine concentration rapidly increases to the point where the iodine forms an intense blue complex with the starch indicator.
When thiosulfate ion is present in solution, both reactions are happening at the same time. The reactions can be added together for the following result: - + H2O2 + 2I + 2H I2 + 2H2O 2- - 2- 2S2O3 + I2 2I + S4O6 ______2- + 2- H2O2 + 2S2O3 + 2H 2H2O + S4O6
Adapted from: General Chemistry in the Laboratory: Roberts, Hollenberg, Postma Jon Bergmann Page 1 4/4/2018 You should be careful to note that the hydrogen peroxide is not directly reacting with the thiosulfate, but the result is the same as if it did, and 2 moles of thiosulfate are consumed per mole of hydrogen peroxide. Therefore the rate the reaction can be written as:
[ H O ] 1 [ S O 2 ] r a t e 2 2 2 3 t 2 t Experimental Procedure
Supplies 1-ml graduated pipet 5-ml graduated pipet 10-mL graduated pipet thermometer timer with a second hand Chemicals 1. 3% H2O2 stabilized with 0.001-M H2SO4: 1 bottle per 2 groups 2. 0.050-M sodium thiosulfate: 150-mL per group 3. 0.050-M KI: 250-mL per group 4. 0.050-M acetic acid-sodium acetate buffer (0.050-M for both): 500-mL per group 5. 0.1% starch solution (freshly made): 50-mL per group 6. 0.010-M Mo(VI) catalyst (1.76-g/L of ammonium heptamolybate ((NH4)6Mo7O24•4H2O): 10-mL per group 7. 2-M H2SO4
A. Standardization of the Hydrogen Peroxide Solution
1. Fill the buret with 0.050-M Na2S2O3. Then pipet exactly 1.00-mL of the approximately 0.8-M H2O2 solution into a 125-mL Erlenmeyer flask containing 25-mL of H2O. 2. Add 10-mL of 2-M H2SO4, 1-g of solid KI and 3 drops of 3% ammonium molybdate catalyst. 3. Swirl until the KI dissolves and immediately titrate the brown iodine solution that forms with 0.050-M sodium thiosulfate until a brown color begins to fade to yellow. 4. At this point add 2-mL of 0.1% starch indicator and titrate till the disappearance of the blue color. The end point is very sharp and should require a total of 32 to 36-mL of the 0.050-M Na2S2O3. Calculate the exact concentration of the H2O2.
Adapted from: General Chemistry in the Laboratory: Roberts, Hollenberg, Postma Jon Bergmann Page 2 4/4/2018 B. Reaction Rate Measurements
Six reaction mixtures will provide the information necessary to determine the effects of - + the concentration of H2O2, I , and H on the rate of the reaction.
Make up each reaction mixture in a 250-mL beaker, adding the reactants for each reaction mixture in the order that they appear in the table (reading from left to right). Read and record the temperature to the nearest 0.1˚ C. It is advisable to set up your work area in an organized fashion. Have all of the equipment you need to dispense the chemicals. CAUTION: At no time should you allow cross contamination of your glassware. Use one pipet for the H2O2, one graduated cylinder for the water, etc. Put labels on the glassware so you do not confuse them.
Reaction Temp H2O 0.05-M 0.3-M 05M .1% .05M .01 0.8-M Mixture ˚C HC2H3O2, HC2H3O2 KI starch Na2S2O3 Mo H2O2 NaC2H3O2 cat bufffer 1 R.T. 75 30 0 25 5 5 0 10 2 R.T. 80 30 0 25 5 5 0 5 3 R.T. 50 30 0 50 5 5 0 10 4 R.T. 30 30 45 25 5 5 0 10 5 R.T. 75 30 0 25 5 5 0 10 + 10˚ 6 R.T. 70 30 0 25 5 5 5 10
To start each reaction add exactly 10 or 5-mL of the H2O2. Start the timer when half of the H2O2 has been added to the mixture. Watch the solution carefully for the sudden appearance of the blue color and stop the timer when it appears.
C. Interpretation of the Data
1. Calculating the reaction orders and the rate constant
a. Calculate the concentrations of each chemical in the table above. Remember that each chemical has been diluted by all of the others. b. Determine the initial rate of each reaction. [ H O ] 1 [ S O 2 ] r a t e 2 2 2 3 t 2 t c. Use the method of initial rates to determine the order of the reaction for each reactant. (This will only apply for mixtures 1-4 d. Once the orders have been determined calculate the value of the rate constant, k. This should be done 4 times and an average taken.
2. The Effect of Temperature
Adapted from: General Chemistry in the Laboratory: Roberts, Hollenberg, Postma Jon Bergmann Page 3 4/4/2018 a. Using reaction mixtures 5 and 1 and determine the ratio of the rates. r a t e t k T b. 5 1 5 r a t e 1 t 5 k T 1 c. Determine the rate constant at the elevated temperature
3. Effect of a Catalyst a. Using the measured reaction times, we can calculate the ratio of the reaction rates from the data for reaction mixtures 6 and 1 r a t e 6 ( c a t a l y z e d ) t 1 b. r a t e 1 ( u n c a t a l y z e d ) t 6
Adapted from: General Chemistry in the Laboratory: Roberts, Hollenberg, Postma Jon Bergmann Page 4 4/4/2018 Reaction Kinetics Lab (#26): Expectations
1) Formal Lab write-up 2) Data and Calculations 3) Data in neat data tables: 4) Calculations
a) Show how [ ] was determined b) Show how a, b, c was determined: c) Show how 2 values of k were determined d) Determine ratio of rate of R.T. vs R.T. + 10˚ 5) Class Data (Including averages and standard deviations) a) Value of k (room temperature b) Value of k (+ 10˚ C)
6) Error Discussion a) In cause-effect data table
7) Questions: a) Describe why reactions at elevated temperatures happen more rapidly than reactions at lower temperatures: Use the kinetic molecular theory in your explanation. b) Describe the effect of a catalyst on a reaction: Include a discussion about what changes in the rate law. 8) Conclusion a) Include the final answer: Rate law and2 values of k b) Comment on the effect of temperature upon reaction rate.
Adapted from: General Chemistry in the Laboratory: Roberts, Hollenberg, Postma Jon Bergmann Page 5 4/4/2018