A basic kinetic study of a chemical reaction often involves conducting the reaction at varying concentrations of reactants. In this way, you can determine the order of the reaction in each species, and determine a rate law expression. Once you select a reaction to examine, you must decide how to follow the reaction by measuring some parameter that changes regularly as time passes, such as temperature, pH, pressure, conductance, or absorbance of light. In this experiment we will use color change to derive the initial rates of the reaction. Before coming to lab you should read in your text the section on initial rates and reaction order.
In this experiment you will conduct the reaction between solutions of potassium permanganate and oxalic acid. As this reaction proceeds, it undergoes a color change from purple to red to yellow. Although the reaction itself is quite complex, by carefully varying the concentrations of the reactants, you will be able to determine the effect each reactant has on the rate of the reaction, and consequently the order of the reaction. From this information, you will write a rate law expression for the reaction.
OBJECTIVES
In this experiment, you will
- Conduct the reaction of KMnO4 and H2C2O4 using various concentrations of reactants.
- Determine the order of the reaction with respect to KMnO4 and H2C2O4.
- Determine the rate law expression for the reaction.
MATERIALS
~60 mL 0.755 M H2C2O4 solution Glass stirring rod
~20 mL 0.130 M KMnO4 solution Timer
Three graduated 10 mL pipets ~60 mL distilled water
Three 100 mL beakers Test tube rack
Six clean, dry 20 × 150 mm test tubes
PROCEDURE
1. Obtain and wear goggles.
2. Obtain the stated volumes of solutions and distilled water into the three 100 mL beakers. These are your stock solutions. Make note of the actual concentrations of the solutions used.
3. Assemble the test tubes, pipets, stirring rod and timer in preparation for the experiment.
Determination 1
4. Dispense 5.0 mL of oxalic acid solution, using the pipet, into one test tube.
5. Into the same test tube add 6.0 mL of distilled water using another pipet and mix thoroughly with the stirring rod.
6. Dispense 1.0 mL of potassium permanganate solution, using the third pipet, into a separate test tube.
7. Quickly add the oxalic acid solution to the potassium permanganate solution, starting the timer when half the oxalic acid solution has been added. Thoroughly mix the solution.
8. Once all traces of red have disappeared stop the timer and make note of the time.
9. Repeat steps 4 – 8 until you can reproduce the time to within 10 seconds.
Determinations 2 and 3
10. Repeat steps 4 – 8 for determinations 2 and 3, but using the volumes indicated for determinations 2 and 3 in Table 1.
REPORT
For your lab report you will be required to:
1. Calculate the molar concentrations of KMnO4 and H2C2O4 for each reaction test tube and record the values in the table above. You should provide one example of your calculation.
For Determination 1,
[KMnO4] = [KMnO4]initial * Volume(KMnO4) / Volume(total) = 0.130 M * 1.0 mL / (5.0 + 1.0 + 6.0)mL = 1.083 * 10-2 M
[H2C2O4] = [H2C2O4]initial * Volume(H2C2O4) / Volume(total) = 0.755 M * 5.0 mL / (5.0 + 1.0 + 6.0) mL = 3.146 * 10-1 M
2. Calculate the average initial rate for each determination using:
For Determination 1,
[KMnO4]final = 0 M
(rate)initial = −([KMnO4]final − [KMnO4]initial)/t = -(0 - 1.083 * 10-2)M / 559 s = 1.938 * 10-5 M s-1
For Determination 2,
[KMnO4]final = 0 M
(rate)initial = −([KMnO4]final − [KMnO4]initial)/t = -(0 - 1.083 * 10-2)M / 351 s = 3.085 * 10-5 M s-1
For Determination 2,
[KMnO4]final = 0 M
(rate)initial = −([KMnO4]final − [KMnO4]initial)/t = -(0 - 2.167 * 10-2)M / 428 s = 5.063 * 10-5 M s-1
3. Determine the order of the reaction with respect to KMnO4 and H2C2O4, and derive the overall rate law expression for the reaction.
Comparing Determination 1 and 2, in which [H2C2O4] doubles in the latter while [KMnO4] remains the same,
Therefore, m ≈ 1, and the initial rate is proportional to [H2C2O4].
Comparing Determination 1 and 3, in which [KMnO4] doubles in the latter while [H2C2O4] remains the same,
Therefore, n ≈ 2, and the initial rate is proportional to [KMnO4]2.
The overall rate law expression for the reaction is: rate = k[KMnO4]2[H2C2O4]
4. Calculate the rate constant, k, for the reaction stating any assumptions made.
k = rate/[KMnO4]2[H2C2O4] = (rate)initial /[KMnO4]initial2[H2C2O4]initial
For Determination 1,
k = 1.938 * 10-5 M s-1 / (1.083 * 10-2 M)2 / 3.146 * 10-1 M = 0.525 M-2 s-1
For Determination 2,
k = 0.418 M-2 s-1
For Determination 3,
k = 0.343 M-2 s-1
The average k = 0.429 M-2 s-1
Reference
- Randall, Jack. Advanced Chemistry with Vernier. Vernier Software & Technology, 2006. Print.
- Blauch, David N. Method of Initial Rates. Chemical Kinetics, 2009. Web. 21 Apr. 2013.
