Exp 4-rate law

EXPERIMENT 4: DETERMINING RATE OF REACTION BY THE EFFECT OF CONCENTRATION AND CATALYST

Objectives: 

To determine the rate of reaction (reaction order) of KMnO4 and H2C2O4 by the effect of concentration and catalyst.



To learn about the kinetics of chemical reaction.



To learn about rate law.

Introduction: There are many factors that affect the speed of a chemical reaction, some of the more common ones being concentration effects, temperature effects, surface area and catalysts. In this experiment, two variables that will affect the reaction rate will be studied that are concentration effects and catalysts. An increase in the concentration of the reactants often increases the speed of a reaction. If two molecules, A and B must collide in order to react, anything that increases the frequency of those collisions increases the rate of the reaction. Increasing the concentration of a reactant most often results in an increased reaction rate as the molecules are closer together so collisions can occur more frequently. Catalysts dramatically speed up a chemical reaction by allowing the reaction to proceed via a much easier (energy wise) pathway or by giving the reaction a sort of short cut to follow.

In this laboratory the Rate Law for the reaction between potassium permanganate (KMnO4) and oxalic acid (H2C2O4). The chemical equation of this reaction is as following: 3 H2C2O4(aq) + 2 KMnO4(aq)  6 CO2(g) + 2 MnO2(s) + 2 KOH(aq) + 2 H2O The rate of this reaction was easily measured because the permanganate ion is bright purple and the only other colored species in the system is manganese dioxide (MnO2), which is a golden yellow in solution. Thus, all the permanganate ion had reacted away when the purple color fades and was replaced by a golden yellow. The faster the purple color of the permanganate fades, the faster the reaction was occurring.

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The Rate of Reaction for any general reaction: a A + b B  c C + d D can be defined in terms of any of the observable species involved in the reaction: Rate of reaction = Because ∆t is large, it was the time required for all the permanganate ion to disappear, so actually the average rate of reaction over the reaction’s time course was measured. Rate of reaction =

. The factors such as concentration and catalysts will not

affect the rate of all chemical reactions, but at least some will affect any given reaction. And, occasionally, odd factors, such as such the amount of dust present in the reaction vessel, will also affect the rate.

While knowing the rate of a reaction after the reaction was completed, it provides some information, in order to predict how altering a variable will change the reaction rate, the rate law for the reaction need to be determined. A Rate Law is an equation that directly relates the concentration of the reactants to the speed of the reaction when all other conditions, such as temperature, pressure, etc were held constant. For a reaction A+B  C, the general rate law would be: Rate = k[A]m[B]n. These exponential variables are called reaction orders and were found experimentally through a series of experiment called the method of initial rates. In this experiment, the Rate Law for the reaction will have a simple form: Rate = k [H2C2O4]m[KMnO4]n where m and n are the orders of the reaction with respect to H2C2O4 and KMnO4, respectively. A brief overview of this experiment is oxalic acid was added into the conical flask containing KMnO4 which was mix thoroughly using the magnetic bar and magnetic stirring hotplate. The time taken for the solution to turn yellowish was recorded. In part I, the concentration was set differently by diluting with different amount of water whereas in part II different amount of catalyst H2SO4 was used. The rate constant, and order with respect to each reactants and overall order were then calculated.

Materials & Apparatus:

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KMnO4 (0.02M), H2C2O4 (0.5M), H2SO4, conical flask, magnetic bar, magnetic stirring hotplate, stopwatch

Procedures: Part I: Determining the rate of a chemical reaction effect of concentration 1) The reagents were prepared as tabulated in table below: Reagents

Experiment 1

Experiment 2

Experiment 3

H2C2O4 (mL)

150

150

100

KMnO4 (mL)

10

5

30

H2O (mL)

0

5

30

2) The required amount of KMnO4 was placed in a thoroughly washed and dried conical flask. The magnetic stirring bar was placed into the conical flask. 3) The oxalic acid was added into the conical flask containing KMnO4 and the timing was started once the beaker containing oxalic acid was emptied. 4) The mixture was mixed thoroughly by swirling the conical flask or magnetic stirring hotplate. 5) The swirling or stirring was continued until the solution turned a light yellow or brown color. 6) The timing was stopped and the time it actually took for the reaction to take place was recorded. 7) Steps 2-6 were repeated for experiments 2 and 3. 8) The rate for each of the 3 experiments was determined. 9) The full rate equation for each experiment was written out.

Part II: Determining the rate of a chemical reaction effect of catalyst 1) The reagents were prepared as tabulated in table below: Reagents

Experiment 1

Experiment 2

H2C2O4 (mL)

100

100

KMnO4 (mL)

30

30

H2SO4 (mL)

10

2 3

2) The required amount of KMnO4 was placed in a thoroughly washed and dried conical flask. The magnetic stirring bar was placed into the conical flask. 3) The oxalic acid and sulphuric acid were added into the conical flask containing KMnO4 and the timing was started once the beaker containing oxalic acid was emptied. 4) The mixture was mixed thoroughly by swirling the conical flask or magnetic stirring hotplate. 5) The swirling or stirring was continued until the solution turned a light yellow or brown color. 6) The timing was stopped and the time it actually took for the reaction to take place was recorded. 7) Steps 2-6 were repeated for experiments 2 and 3. 8) The rate for each of the 3 experiments was determined. Remember this was just the [KMnO4] / t average 9) The full rate equation for each experiment was written out.

Results & Discussion: Part I: Data collection: Experiment 1

Experiment 2

Experiment 3

Molarity of KMnO4 (M)

0.02

0.02

0.02

Molarity of H2C2O4 (M)

0.5

0.5

0.5

Volume of KMnO4 (mL)

10

5

30

Volume of H2C2O4 (mL)

150

150

100

Volume of H2O (mL)

0

5

30

Time taken (s)

195

215.4

235.8

Experiment 1

Experiment 2

Experiment 3

1.25 x 10-3

6.25 x 10-4

3.75 x 10-3

Calculations:

Concentration of KMnO4 (M)

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Concentration of

0.4688

0.4688

0.3125

Average time (s)

195

215.4

235.8

Initial rate for

6.1403 x 10-6

2.9016 x 10-6

1.5903 x 10-5

H2C2O4 (M)

KMnO4 (M/s)

Sample calculation: 1) Concentration of KMnO4 M1V1 = M2V2 M2 = M1V1 / V2 M2 = (0.02)(10) / 160 = 1.25 x 10-3 M # 2) Concentration of H2C2O4 M1V1 = M2V2 M2 = M1V1 / V2 M2 = (0.5)(150) / 160 = 0.4688 M # 3) Average time = Time taken (since only one trial was conducted for each experiment) = 195 s # 4) Initial rate for KMnO4 r = [KMnO4] / taverage = 1.25 x 10-3 M / 195 s = 6.4103 x 10-6 M/s #

5) Find the order with respect to KMnO4 and H2C2O4 (I) Order with respect to KMnO4 5

[

] [

]

0.4526 = 0.5 m m log 0.5 = log 0.4526 m = 1.1437

#

(II) Order with respect to H2C2O4

[

] [

5.4808= ( ) (

]

)

0.6666n = 0.9135 n log 0.6666 = log 0.9135 n = 0.2231

#

6) Find the overall order for this reaction, Overall order = m + n =1+0 =1# 7) Find the rate constant, k for this reaction, By using experiment 1, =k( k = 5.1282 x 10-3 s-1

) (

)

#

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In part I, based on the result and calculation, it was shown that from experiment 1 to experiment 3, as the concentration of reactants decreases, the time taken for the reaction increases and the rate of reaction also decreases from experiment 1 to experiment 3. Therefore, it can be summarized that the rate of reaction increases with concentration of reactants. Amongst the experiments, the reaction during experiment 1 (highest concentration of KMnO4 that is 1.25 x 10-3 M) had the highest rate of reaction that was 6.4103 x 10-6 M/s and it taken the shortest time for reaction that was 195s. According to the collision theory, the rate of a reaction is directly proportional to the number of effective collisions per second between the reactant molecules. Effective collision is the fraction of total collisions that actually result in the formation of the product(s). If the concentration of the reactants increases (i.e. particles per given volume) the greater the number of total collisions. The greater the frequency of total collisions, the greater the frequency of effective collisions. If the frequency of effective collisions increases, so does the reaction rate. (N.a., n.d.) By comparing two experiments using initial rate and concentration of both KMnO4 and H2C2O4, it was determined that this reaction is first order with respect to KMnO4 and zero order with respect to H2C2O4. Hence, the overall order of reaction is first order. The value of k constant also being determined that was 5.1282 x 10 -3 s-1. Part II: Reagents

Experiment 1

Experiment 2

H2C2O4 (mL)

100

100

KMnO4 (mL)

30

30

H2SO4 (mL)

10

2

Time taken (s)

3.05

3.12

1) Why does it take so much longer for the color of the solution in the beaker on the right (experiment 2) to change? 

The time taken for the color of the solution in the beaker on the right to change is longer because the amount of catalyst used (H2SO4) is lesser. Catalysts are 7

substances used to speed up the rate of reaction without being consumed. As the amount of catalysts used is lesser, the rate of reaction of experiment 2 is lower, thus it takes a longer time to change color.

2) Write a balanced chemical equation for the reaction of permanganate ion and oxalic acid to give Mn2+ and carbon dioxide. The reaction takes place in an acidic solution. 

2MnO4- + 5 H2C2O4 + 6H+  2Mn2+ + 10CO2 + 8H2O

3) The reaction in the beaker on the right eventually proceeded at a reasonably rapid rate. Explain how this observation supports a reaction mechanism involving autocatalysis. 

In autocatalysis, the reaction is catalyzed by one of its products. Initially, the reaction is very slow at room temperature. The reaction is then catalyzed by manganese (II) ions. There obviously aren't any of those present before the reaction starts, and so it starts off extremely slowly at room temperature. However, from the equation, it was found that manganese (II) ions amongst the products. More and more catalyst is produced as the reaction proceeds and so the reaction speeds up. (Jim.C., 2010)

Conclusion: The rate of reaction increases with increasing concentration of reactant. The rate of reaction also increases with increasing amount of catalysts used. The reaction between H2C2O4 and KMnO4 is first order with respect to KMnO4, zero order with respect to H2C2O4. The overall order for the reaction is one and the rate constant, k is 5.1282 x 10-3 s-1.

Safety Precaution: 

Oxalic Acid is poisonous. Gloves must be wearing when handling solutions containing oxalic acid. If some spills on skin, rinse it off with copious amounts of water. Large spills require flushing with water for 15 minutes.



Permanganate is a strong oxidizing agent. The solutions of permanganate must be handled with gloves. 8

Limitations of experiment: 

There was a human limitation to make sure that the final color of solution was same for all 3 experiments. In some samples the color of the reaction had gone past the color of what the comparison was so then the time recorded of the reaction would be recorded as taking longer.

References: 

Jim.C. (2010) Autocatalysis. [Online]. Available from: http://www.chemguide. co.uk/physical/catalysis/introduction.html. [Accessed: 30 September 2013].



N.a, (n.d.) Determination of the Rate Law for the Oxidation of Oxalic Acid by Permanganate. [Online]. Available from: http://infohost.nmt.edu/~jaltig/OxalicAcid Kinetics.pdf. [Accessed: 30 September 2013].



N.a. (2013) Determining the rate law. [Online]. Available from: http://www.chm.uri.edu/dfreeman/chm192/expt2_2013.pdf. [Accessed: 30 September 2013].



N.a. (n.d.) Chemical Kinetics---Determining the Rate Equation. [Online]. Available from: http://www.chem.umass.edu/genchem/chem112/112_Experiment_3.htm. [Accessed: 30 September 2013].



N.a. (n.d.) Kinetics: Factors affecting reaction rates. [Online]. Available from: http://www.chem.tamu.edu/class/majors/tutorialnotefiles/factors.htm [Accessed: 3 October 2013].

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