CHEM 122L

General Chemistry Laboratory

Revision 2.1

Determination of the Rate Law for the Oxidation of Oxalic Acid

by Permanganate

To learn about the Kinetics of Chemical Reactions.

To learn about the Reaction Rate.

To learn about the Factors Affecting the Chemical Reaction Rate.

To learn about Rate Laws.

In this laboratory exercise we will determine the Rate Law for the reaction between

Permanganate (MnO42-

) and Oxalic Acid (H2C2O4).

3 H2C2O4(aq) + 2 MnO4-(aq) 6 CO2(g) + 2 MnO2(s) + 2 OH

-(aq) + 2 H2O

clear soln purple soln yellow colloid (Eq. 1)

It is expected the Rate Law for this reaction will have a simple form:

Rate = k [H2C2O4]m

[MnO4-]n (Eq. 2)

where m and n are the Orders of the Reaction with respect to Oxalic Acid and Permanganate,

respectively. (We must keep in mind the Rate Law does not have to have this simple form.) It is

these reaction Orders we wish to determine.

The Rate of this reaction is easily measured because the Permanganate 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 has reacted away when the purple color fades

and is replaced by a golden yellow. The faster the purple color of the Permanganate fades, the

faster the reaction is occurring.

The Rate of Reaction for any general reaction:

a A + b B c C + d D (Eq. 3)

can be defined in terms of any of the observable species involved in the reaction:

Rate of Reaction = (Eq. 4)

=

P a g e | 2

=

=

In practice it is difficult to determine instantaneous concentration changes, so instead we

approximate the above definitions by:

Rate of Reaction = in the limit t ~ small (Eq. 5)

= etc.

For our reaction, the observable species we wish to follow is Permanganate, so we define

the Reaction Rate as:

Rate of Reaction = - (1/2) (Eq. 6)

Because t is large, it is the time required for all the Permanganate Ion to disappear, we are actually measuring the Average Rate of Reaction over the reactions time course.

In general, the rate of reaction can depend on any number of factors, just as the rate of travel in

an automobile depends on the position of the accelerator or brake, hilliness of the ground, wind

speed, size of the engine, etc. In the case of a chemical reaction, the Rate can depend on:

Reactant Concentrations

Product Concentrations

Catalyst Concentration

Temperature

Amount of Light

Size of Heterogeneous Particles

etc.

These factors 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.

Typically, Reactant Concentration and Temperature are the major influences on the rate of

reaction. At a given temperature, the Rate of Reaction frequently depends on reactant

concentration according to a rather simple expression:

Rate = k [A]n [B]

m (Eq. 7)

where A and B are the reaction reactants. This form of the reaction rate dependence is known as

the Rate Law for the reaction and k is referred to as the Rate Constant. The reaction is said to

proceed according to Orders n and m. It should be emphasized the parameters k, n, and m must

be measured experimentally.

P a g e | 3

For our reaction, the Rate Law will have a simple form such as this and this form is given in

Equation 2 above. It is the goal of this laboratory exercise to measure the Reaction Orders n and

m and thus determine the Rate Law for the Oxidation of Oxalic Acid by Permanganate.

So, how are the reaction orders for a chemical reaction measured? There are a number of

methods available. One of the most common is the Method of Initial Rates. In this procedure,

we selectively change the Initial Concentration of each Species and see how this change affects

the rate of reaction. Suppose we double the Initial Concentration of Species A and find it causes

the reaction rate to quadruple. We then know n = 2; or the reaction is Second Order in A. We do

the same with each subsequent Species suspected of influencing the rate of the reaction to

determine each Reaction Order. (Initial Rates are used in this methodology because they provide

a convenient reference point for comparing one trial against another. As mentioned above, for

our case, we will be using the reactions average rate.)

For an example, consider the following reaction:

NO2(g) + CO(g) NO(g) + CO2(g)

We will assume the Rate of this Reaction is governed by a Rate Law of the following form:

Rate of Reaction = k [NO2]n [CO]

m

In order to determine the Reaction Orders n and m, the following Initial Rate data has been

obtained:

Trial Init. Rate (M/sec) Init. [NO2] (M) Init. [CO] (M)

1 0.005 0.10 0.10

2 0.080 0.40 0.10

3 0.005 0.10 0.20

In comparing Trials 1 and 2, we see a quadrupling of the initial NO2 concentration cause the

Reaction Rate to increase by a factor of 16; 0.005 x 16 = 0.080. This means the Reaction Order n

= 2. The reaction is Second Order with respect to NO2. Now we compare Trials 1 and 3. Here a

doubling of the initial CO concentration causes no change in the Reaction Rate. Hence, the

Reaction Order m = 0 and the reaction is Zeroth order with respect to CO. Thus, the Rate Law for

this reaction is:

Rate = k [NO2]2

Thus, for our reaction, we need to selectively vary the concentrations of Oxalic Acid and

Permanganate and determine how the Rate of Reaction responds to these changes. The key to

determining the reaction orders is the ability to measure the reaction rate for each variation in

initial reactant concentration. And, as was noted above, this is simply a matter of determining

how fast the purple color of the Permangante fades.

P a g e | 4

Pre-Lab Questions

1. To determine the Reaction Order of the decomposition of Formaldehyde:

CH3CHO(g) CH4(g) + CO(g)

which is assumed to proceed according to the following Rate Law:

Rate = k [CH3CHO]n

the following Initial Rate data has been obtained:

Trial Init. Rate (M/sec) Init. [CH3CHO] (M)

1 0.76 0.30

2 1.4 0.40

Use this data to determine the Reaction Order with respect to Formaldehyde.

Do this by comparing the Rate of Trial 2 to the Rate of Trial 1:

= n

Because the Reaction Order does not have to be an Integer value, we should solve for n by taking

the logarithm of both sides:

log = n log

and solving this algebraically for n.

Because the Reaction Order typically is an Integer value, we then round this experimentally

determined result to its closest Integer value. You should follow this procedure for this case.

What is the experimentally determined Reaction Order n for this decomposition reaction?

What is the approximate Integer value for n?

2. For our reaction, the last reagent to be added to the reaction mixture is the Permanganate;

step 6 of the procedure below. The timing of the reaction starts "when half of the

Permanganate solution has been added to the test tube." This choice is arbitrary and could

be problematic. Under what circumstances could it be problematic?

3. It is possible the Reaction Order is negative; for example consider the Rate Law for the

kinetics of the conversion of Ozone to Oxygen.

2 O3(g) 3 O2(g)

P a g e | 5

rate = k [O3]2 [O2]

-1

What are the implications for the reaction rate of a negative Reaction Order?

P a g e | 6

Procedure

Stock solutions of each of the reagents are provided at the following concentrations:

[H2C2O4] = 0.755 M

[MnO4-] = 0.130 M

We will vary the initial concentration of the Oxalic Acid and Permanganate by varying the

volume of each reagent added to the reaction mixture. The concentration (volume) of only one

species will be varied from one experiment to the next so as to determine the influence of only

the species being varied. Because the endpoint of the reaction is difficult to determine, each

experiment will be performed in triplicate and the results will be averaged. Each experiment will

be prepared as follows:

Experiment Oxalic Acid Permanganate Water

1 5.00 1.00 6.00

2 10.00 1.00 1.00

3 5.00 2.00 5.00

1. Label a clean 100mL beaker:

Oxalic Acid Obtain ~75 mL of Oxalic Acid solution and place it in the beaker. Oxalic Acid is

poisonous. Wear gloves when handling solutions containing Oxalic Acid. If some

spills on your skin, rinse it off with copious amounts of water. Large spills require

flushing with water for 15 minutes.

2. Do the same for a beaker for Water.

3. Label a clean 50 mL beaker:

Permanganate Ion Obtain ~25mL of Permanganate and place it in the beaker. Permanganate is a strong

oxidizing agent. Handle solutions of Permanganate with gloves.

4. Obtain a large test tube and clamp it in place over a stir plate. Add a small stir bar to the

test tube.

P a g e | 7

(The Reaction Is Pictured Above After It Has Reached Its Endpoint.)

5. Using pipets, add the Oxalic Acid solution and the Water to the test tube as specified for

Experiment #1 above. (Use a different pipet for each solution!) Turn on the stir plate and

allow the contents of the test tube to mix thoroughly. When adding the solutions to the

test tube, make sure the tip of the pipette is very close to the bottom of the tube.

Otherwise, some of the solution will stick to the sides of the tube and will not be a part

of the reaction mixture. Be sure to use good pipetting techniques as using poor

techniques will seriously degrade the quality of your data.

6. Pipet the Permanganate solution into the test tube. Start the stopwatch when half of the

Permanganate solution has been added to the test tube. This must be done in one

quick addition. Again, you must be careful to use good pipetting techniques. This is

particularly important for the Permanganate solution. A single misplaced drop can

seriously affect the quality of the kinetic data.

P a g e | 8

7. Observe the solution over time. Stop the stopwatch when the last trace of the

Permanganates purple color disappears. The final solution will be a golden yellow with some bubble formation due to the production of Carbon Dioxide gas by the reaction.

Determining the endpoint of the reaction is a bit subjective. (You may wish to perform a

sample trial so as to get some practice in determining when the purple color has been

replaced by the golden yellow.) Also, as the reaction approaches its endpoint, it is best not

to look at the time on the stopwatch. This will help you to remove any bias due to an

expectation from previous results.

7. Repeat this procedure for Experiment 1 twice more. All three trials should agree to within

10 seconds.

8. Now, perform triplicate trials for Experiments #2 and #3 as well.

P a g e | 9

Data Analysis

1. Calculate the Initial Concentration, in units of Molarity, for each reagent used in each of

the three trials in this experiment. Make an appropriate Table of these concentrations.

2. Calculate the Reaction Rate for each trial using an average of the Time required for the

reaction to go to completion. Add these to your above Table.

3. Calculate the Reaction Order for each Species; m and n. (See the note for Pre-Lab

Question 2 and follow the same procedure here.) Report two values for each Reaction

Order; the exact experimentally determined value and the nearest Integer value.

4. Report the Reaction Rate Law.

P a g e | 10

Post Lab Questions

1. Why is it necessary to keep the Total Volume of reagents in each of the four trials in this

experiment the same?

2. We have labeled all the Oxalic Acid bottles "Poison." What are the toxic effects of

ingesting Oxalic Acid? Write an appropriate chemical equation.

3. The reaction under study is an Oxidation-Reduction reaction. Assign an Oxidation State to

the Carbon and the Manganese in each relevant species. Identify the Oxidation and the

Reduction.