Determination of a Rate Law by the Method of Initial Rates | CHEM 220, Lab Reports of Chemistry

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Determination of a Rate Law by the Method of Initial Rates

Determination of a Rate Law by the Method of Initial Rates

PROCEDURE

Table 1 summarizes the volumes required for trials 1-12. All trials will be carried out in 150-mm test tubes.

1. Preparation of solution A: Pipette the listed volumes for the various solutions into a clean, dry test tube. Add the starch from a dropper. Prepare all of solution A for trials 1-6 at the same time.
2. Preparation of solution B: Pipette the indicated volume of 0.10 M H 2 O 2 into a small container (10-mL beaker). Prepare just before use in each trial.
3. Prepare to time the reaction. Place a piece of white paper behind the test tube rack so that you can easily see the color change.
4. Add solution B rapidly to the test tube with solution A for the trial. START TIME immediately. Stopper the tube and invert twice (only 2 times!). Remove the stopper and place into the test tube rack.
5. The change to a deep purple color will be sudden. Be prepared! STOP TIME when the blue color appears.
6. Record the time in seconds. Measure and record the temperature of the reaction mixture.
7. Repeat steps 2-6 for the additional kinetic trials.
8. Repeat steps 1-7 above for trials 7-
9. Run a Trial “Zero” that is just like trial #1, but excludes the thiosulfate solution. This should help to illustrate dramatically the effect of the thiosulfate on the reaction mixture.

DATA ANALYSIS

10. Prepare a results table to summarize the data below.
11. Calculate the moles of S 2 O 32 −^ ions consumed in each trial. From the stoichiometry of the reaction, calculate the moles of I 2 consumed by the reaction with the thiosulfate. We will designate this ∆(mole I 2 ). These values are the same for each trial and serve as the “clock” - when the moles of I 2 produced exceeds the stoichiometric ratio to the thiosulfate, it will complex with the starch to give the deep blue color of the starch-iodine complex.
12. Calculate the initial rate of the reaction for each trial with the equation: where ∆t is the elapsed time.
13. Calculate the log (rate) for each trial.
14. For each trial, calculate the initial concentrations of hydrogen peroxide [H 2 O 2 ] 0 and iodide [I−] 0 as well as log [H 2 O 2 ] 0 and log [I−] 0. Note: Consider the volume of the drops in the calculation of the initial volumes of the reactants. 1 drop ≈ 0.05 mL

∆t

∆mol rate

( I 2 )

Prepare a data/analysis table that summarizes, for each trial, all of the measured and calculated values describe above.

Preliminary order determinations:

15. Select two trials in which the [I−] 0 is constant between trials. Analyze the difference in the [H 2 O 2 ] 0 and the rate of reaction to make a preliminary determination of the order of the reaction for [H 2 O 2 ].
16. Select two trials in which the [H 2 O 2 ] 0 is constant between trials. Analyze the difference in the [I−] 0 and the rate of reaction to make a preliminary determination of the order of the reaction for [I−].

Graphical order determinations:

17. Graph #1: Plot log [I−] 0 (x-axis) vs. log (rate) (y-axis) for trials 1-6. Determine the slope of the line. The slope should be a good approximation of the order of the reaction for I−.
18. Graph #2: Plot log [H 2 O 2 ] 0 vs. log (rate) for trials 7-12. Determine the slope of the line. The slope should be a good approximation of the order of the reaction for H 2 O 2.

RESULTS

19. State the differential rate law expression for the reaction of the iodide ion and hydrogen peroxide.
20. Calculate k´ for each of the seven trials and determine the average value. Calculate the standard deviation.

Table 1. Composition of Reaction Mixtures

Solution A Solution B

Trial

Buffer (0.5 M CH 3 CO 2 H / 0.5 M NaCH 3 CO 2 )

0.020 M

Na 2 S 2 O 3 Starch DI Water 0.30 M KI 0.10 M H 2 O 2

1 1.0 mL 1.0 mL 5 drops 6.0 mL 1.0 mL 3.0 mL 2 1.0 mL 1.0 mL 5 drops 5.0 mL 2.0 mL 3.0 mL 3 1.0 mL 1.0 mL 5 drops 4.0 mL 3.0 mL 3.0 mL 4 1.0 mL 1.0 mL 5 drops 3.0 mL 4.0 mL 3.0 mL 5 1.0 mL 1.0 mL 5 drops 2.0 mL 5.0 mL 3.0 mL 6 1.0 mL 1.0 mL 5 drops 1.0 mL 6.0 mL 3.0 mL 7 1.0 mL 1.0 mL 5 drops 7.0 mL 2.0 mL 1.0 mL 8 1.0 mL 1.0 mL 5 drops 5.5 mL 2.0 mL 2.5 mL 9 1.0 mL 1.0 mL 5 drops 4.0 mL 2.0 mL 4.0 mL 10 1.0 mL 1.0 mL 5 drops 3.0 mL 2.0 mL 5.0 mL 11 1.0 mL 1.0 mL 5 drops 1.5 mL 2.0 mL 6.5 mL 12 1.0 mL 1.0 mL 5 drops none 2.0 mL 8.0 mL 0 1.0 mL 1.0 mL none 6.0 mL 1.0 mL 3.0 mL

2. State the purpose of each of the following solutions in this experiment.

A) Buffer solution

B) Starch solution

C) Sodium thiosulfate solution

D) Deionized water

3. Consider the reaction studied in this experiment (Equation 24.2). It is an oxidation-reduction reaction. Write the ½-reactions for this overall reaction.
4. Write the equation for the reaction that serves as the “clock” for this reaction.