Determining the Rate of Reaction for the Oxidation of Oxalic Acid by Permanganate

Date:  2021-05-28 07:18:51
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Harvey Mudd College
Type of paper: 
Research paper
This essay has been submitted by a student. This is not an example of the work written by our professional essay writers.

A redox reaction is one in which both reduction and oxidation take place. Oxidation is the loss of electrons while reduction is the gain of electrons. Species such as potassium permanganate are said to be strong oxidizing agents while those like oxalic acid are said to reducing agents (Duke 2887). The reaction in this paper will involve titration of oxalic acid and potassium permanganate. The reaction requires no use of a catalyst as it is self indicating due to the color change from purple to yellow.

Background Information

Chemical kinetics states that for a reaction to occur, particles must collide with each other as prescribed in the collision theory (Duke 2886). This collision of particles reads to the increase in the activation energy of the reacting particles. To start with, activation energy is the minimum amount of energy that is required for particles of the reacting molecules to react. This energy is essentially used in breaking the bonds of attraction that hold particles together hence facilitating reactions to take place as new compound is formed (Wariishi et al. 23690). Once this energy is increased, the particles of different materials can then react with each other to form a new product with new sets of bonds. This is, however, dependent on the temperature at which the reaction is being undertaken. If the temperature of the reacting species is increased, then the average speed of the reacting particles is also increased which causes an increase in the collision (Kovacs et al. 11029). Temperature is therefore directly proportional to the reaction rate of any given reaction due to the increase in the kinetic energy of the reacting particles. Borrowing from the kinetic theory, we get that an increase in the collision rate frequency will lead to an increase in the reaction rate due breaking of the molecular forces holding the particles together. In addition, an increase in temperature also causes an increase in the internal energy and the total potential energy of the substance. Consequently, the proportion of particles that have energy is larger than the activation energy it would not generate if the temperature of the reacting conditions were nor increased.

From experiments done by Maxwell- Boltzmann, there is a direct relation between the temperature of the reacting species and the amount of the products formed (Dollimore 58). Borrowing from the curve of Maxwell-Boltzmann it is evident that the higher the temperature of the reacting particles the more the possibility of the particles exceeding the activation energy of the same particles (Popp 126). The experiment can be used to form an equation that relates the number of particles and the temperature of the reacting particles. An increase in the activation energy Ea is proved to be directly proportional to an increase in the temperature of the reacting species as shown in the Arrhenius equation shown below:


The equation is called Maxwell-Boltzmann equation where K is a constant rate, A is the Arrhenius constant representing the frequency of the orientation of the particles undergoing collisions, R represents the universal gas rate and Ea is the activation energy that is required for different species to undertake a reaction process (Dollimore 58). Finally T is the temperature of the reacting particles and is measured in the Kelvin scale.

In this question, the following reaction will be used to investigate the comparison between activation energy and the temperature of the system.

5KBr(aq) + KBrO3(aq) + 3H2SO4 (aq) 3K2SO4 + H2O(l) + 3Br2(aq)

The above equation can be manipulated by simple transformation so as to conform with the Arrhenius equation, as and can be used to calculate the activation energy of the reacting particles. However, before undertaking any calculation it is important that we consider the rate constant in the terms of reaction rate. This can be done by considering the rate law which relates the rate of reaction which is directly proportional to the products of the reaction based on the concentration of the reacting particles (Duke 2889).

Reaction between KBr and KBrO3

This reaction between the bromide and the bromated ions can only take place in an acidified solution. The ionic equation of the reaction is shown below:

5Br + BrO3 + 5H + 3Br2 + 3H2O

The above reaction represents a redox reaction in which the Br- is being oxidized and BrO3- is being reduced. Potassium and the sulfate ions are being used as the spectator ions in this redox reaction and hence they do not take part in this overall reaction. The time taken for the reaction to take place is taken at different temperatures where phenolphthalein indicator is added into the solution to identify the end point (Popp 126). The bromide ions are formed in this reaction and react vigorously with the indicator hence bleaching it hence being convenient to be used to show the end point of the reaction. However, the problem in using this method is to determine the rate of reaction at different temperatures including the room temperature. This is because the reaction is very fast at room temperature hence recording of the temperature and practicability of the same becomes impossible (Duke 2888).

To solve the above problem, a small amount of phenol is added to the reacting species which acts as an inhibitor in reducing the reaction rate at room temperature. The purpose of phenol is therefore to reduce the rate of reaction by providing a suitable environment before bromide ions are formed. This is made possible since phenol consumes part of the bromine in the reaction which reduces its concentration thereby reducing the amount of collisions and overall reduction in the reaction rate.

The above reaction can be used as the bases on which we can use to calculate the reaction rate between the oxidation of oxalic acid by potassium permanganate. The following chemical equation will be used:

3 H2C2O4(aq) + 2 MnO4 - (aq) 6 CO2(g) + 2 MnO2(s) + 2 OH- (aq) + 2 H2O

It is expected that the reaction will produce using the following rate law as seen in the earlier example of bromide ions:

Rate = k [H2C2O4] m [MnO4 -] n

where n and m represent the order of the reaction with respect to permanganate and the oxalic acid solutions respectively (Popp 126). In undertaking this experiment and consequently using it to determine the reaction orders m and n, then we must always keep in mind the general law that governs chemical kinetics and not just simple form of the above equation. It is therefore important to assume that the above equation remains sufficiently reliable and valid for our case.

The rate of this reaction can be easily measured since potassium permanganate solution is bright purple in color and in fact the only colored species in the system (Banerjee et al. 25). When this species is oxidized by oxalic acid, it forms manganese dioxide which is a golden yellow solution thereby providing an end point for the reaction. It is therefore assumed that when the purple color of permanganate ion fades from purple to golden yellow, an oxidation reaction has taken place. The faster the rate at which the color fades, then the higher the rate of reaction and the higher the order levels of n and m (Kovacs et al. 59).

Reaction can thus be defined in terms of the observed concentration of potassium permanganate solution and oxalic acid:

Rate of Reaction = - d([MnO4 -]dt - when dt represents a very little change in the chemical kinetics.

However, since we are taking the time taken for potassium permanganate to have its color faded, then the value of dt will certainly be large than when we take the average of the reaction rate directly over the course of time.

The equation for calculating the average rate of reaction is given as:

Av. Rate of Reaction = - d([MnO4 of completion- intial manganate ions conc.]Time taken to completion-0

For the purpose of this reaction, it is important that we measure the average reaction rate since we shall be determining how the rate of reaction is influenced by the reaction rate of the species. Using the earlier equation, we can calculate the rate of reaction of permanganate solution by determining the initial and the final concentration of permanganate ions and the time taken to consume those (Banerjee et al. 25).

Once the reaction rate of permanganate solution is determined, it is important that we determine the order of the reaction. There are various methods in which the order of the reaction can be found. In this article we are going to use initial rate method to calculate the values of m and n. this method involves the change in the concentration of the reacting species and carefully observing how these changes affect the reaction rate (Yu et al. 236). Similarly, we can double the concentration of oxalic acid species which increases the order of the reaction from order 2 to order 4. An advantage of using initial rate method is because it provides a convenient reference point to be used by a chemist when comparing one experimental trial against another one where the concentration changes are at maximum.

In conclusion, we need to selectively vary the concentration of the oxalic acid and potassium permanganate solution and to determine how reaction rate responds to changes in the concentration. Finally we shall infer these changes to the reaction order in respect to oxalic acid and potassium permanganate solutions thereby solving for the actual values of m and n.

Experimental Materials and Procedure

Stock solutions of the acid and the potassium permanganate solutions were provided as 0.755 M and 0.0130 M concentration respectively. The concentration of the acid and the permanganate was varied by using different volumes of reagents to be used in the mixture. The volume of the species will be varied separately form one experiment to another so as to determine the effect of that species added. Due to the difficulty in determining the end point of the reaction, this experiment will involve three trials after which the average of the three experiments will be determined.

The table below represents how the experiments were carried out.

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

A clean 100ml beaker is ladled oxalic acid where 75ml of the acid is carefully added into the empty beaker. It should be noted that oxalic acid is very poisonous therefore care must be taken when handling it. The person performing the experiment must always were gloves and incase some acid spill on the skin, it should be rinsed with little amounts of copious water. The same procedure was then repeated for water where the results were recorded as shown in the table above.

Potassium permanganate solution was prepared in a clean 50ml beaker labeled Permanganate ion. 25 ml of this solution was obtained and placed in another clean beaker. Due to the nature of the solution as a strong oxidizing agent it is important to handle it with care by wearing hand gloves. A large test tube was then obtained and placed on a clamp over a stir plate. A small stir bar was then put in the test tube to act as a stirrer.

The Titration Procedure

0.15g of copper oxalate powder was weighed onto a piece of paper and transferred into a clean 250ml Erlenmeyer flask

10ml of 3.0M ammonia solution was added to the flask containing the copper oxalate and the contents stirred using a stirring rod until all the powder had dissolved

In a separate 250ml conical flask, 100ml of 0.8M H2SO4 was heated to temperatures of 60-700C

The already hot H2SO4 acid was added to the dissolved copper oxalate in a separate beaker. The contents were stirred using a stirring rod to make sure no big c...


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