Rate of Reaction of Potassium Permanganate and Oxalic Acid

The purpose of this experiment was to determine the reaction order and write a rate equation with respect to changes in permanganate ion and oxalic acid concentrations and to examine the effect temperature has on the rate of the reaction 1. In part one, the reactants potassium permanganate (KMnO4) and oxalic acid (H2C2O4), three determinations were performed, each with different initial concentrations of the reactants. Each initial concentration resulted in a unique reaction rate; these rates were then examined using the method of initial rates to determine the order of the reaction with respect to both KMnO4 and H2C2O4. The rate constant (k), was then calculated, and the rate equation for the reaction written1. In part two; determinations were done at varying temperatures while keeping the initial concentrations for each reactant stayed constant to prove that a change in temperature results in a change of the reaction rate1. Introduction:

Chemical reactions occur when reactant ions or molecules collide with enough energy to break and form bonds; referred to as kinetics1. The rate is the velocity, or how quickly the reaction proceeds. Rates can be altered in a variety of ways; the increase of reactant concentration and increase reactant temperature results in a rise in the amount of collisions and thus a faster rate1. In part one of lab the initial concentrations of the reactants were varied to examine the effect on the reaction rate. The rate of reaction is affected differently by changes of concentrations in one reactant compared to another1. Changes in concentrations of individual reactants and the effect on the rate can be expressed mathematically through the rate equation (1): rate= k[H2C2O4]x [KMnO4]y (1)

In this equation, k is the rate constant which only varies with changes in temperature. Potassium permanganate and oxalic acid are both expressed as
molar concentrations (mol/L) and the superscripts x and y represent the reaction order of each adjacent concentration1. The reaction order is an exponential term that is useful in determining the relationship between an increase in reactant concentration and the resulting effect it has on the reaction rate whether being an increase or decrease1.The overall reaction order of a chemical reaction is the sum of the individual reaction orders (x and y) in equation (1), of each reactant; in this experiment, H2C2O4 and KMnO4. If both reaction orders equal a sum of 2, then the reaction is considered second order1. The reaction order with respect to the reactants was determined using the method of initial rates. Equation (2): (rate2)/(rate1)= ([H2C2O4]1x [KMnO4]1y)/([H2C2O4]2x [KMnO4]2y) (2) was used to solve for the reaction order (X). The same equation is used with determinations 1 and 3 to find (Y). If x or y equal 1, then the reaction is first order with respect to the corresponding reactant. A reaction order of one signifies that the change in concentration of the reactant is proportional to the product concentration.

When the concentration of the reactant doubles, the reaction rate will also double. When x and y are added together they equal the overall reaction order1. Once the reaction order for each concentration is established the rate constant (k) can be found by plugging data into the rate equation above (1) along with the overall reaction order. Reaction rates can be measured by knowing that the rate at which a reactant’s concentration decreases is proportional to the rate at which the products’ concentrations increase1. Using the balanced equation of potassium permanganate and oxalic acid (3) the reaction between reactants can be examined: 2 Mn04- + 6H+ + 5 H2C2O4 → 2 Mn2+ + 8 H2O + 10 CO2 (3) KMnO4 is purple and Mn +2 is yellow colored, so the decrease in reactant concentration or the progression of the chemical reaction can be visualized; the solution transforms from a purple color to yellow color. This is important because the initial presence of reactant KMnO4 can be visualized thus allowing the rate of the reaction to be measured due to a decrease in KMnO4 1. By measuring KMnO4, the reaction rate can be calculated by using equation (4): Rate=- Δ[KMnO4]/ Δt (4)

Δ[KMnO4] is the difference between initial and after concentrations when reaction is complete Δt= elapsed time1. In part two, determinations were performed between solutions of KMnO4 and H2C2O4 with constant initial concentrations and varying temperatures to examine the effect on reaction rates. The “Rule of Thumb1” states that when a reaction temperature is raised 10 °C, the reaction rate will double. To prove this theory, temperatures were increased roughly 10°C in every determination to prove that for every 10°C increase in temperature results in doubling the reaction rate1. The reaction rate was calculated using equation (4): Experimental:

Part 1:
To a test tube containing 6.00 ml of distilled water, 5.00 ml of 0.755M H2C2O4 was added and thoroughly mixed with a glass stirring rod. The solution was added to a test tube containing 1.00 ml of 0.130M KMnO4. A timer was started after half of the H2C2O4 solution was added. The solution was thoroughly mixed with a glass stirring rod and placed in a spectrometer to visualize the absorbency decrease. The solution was taken out of the spectrometer and placed in a test tube support once the solution turned yellow. Once the last hint of red had vanished, the timer was stopped and data was recorded in data table one. A total of three determinations were conducted and each determination was performed using the above procedure. Determination 2 was performed using the above procedure and 1.00 ml of distilled water, 10.00 ml of H2C2O4 and 1.00 ml of KMnO4. Determination 3 was performed using 2.00 ml of distilled water, 5.00 ml of H2C2O4 and 2.00 ml KMnO4. All reactions were performed at room temperature1. Part 2:

To a test tube containing 6.00 ml of distilled water, 5.00 ml of 0.755M H2C2O4 was added and thoroughly mixed with a glass stirring rod. To a test tube, 1.00 ml of 0.130M KMnO4 was pipetted. Reactants were placed in a warm water bath (31 °C) for 10 minutes. Once ten minutes elapsed, the H2C2O4 solution was quickly added to the KMnO4 solution. After half of the H2C2O4 solution was added, the timer was started and the solution was placed back into the warm water bath. The timer was stopped once the last trace of red disappeared from the solution. Determination 2 was performed using the above procedure and a water bath temperature of 38 °C. Determination 3 was performed using the same procedure and a water bath temperature of 53 °C1. All of the initial concentrations remained the same throughout part two.

ChemistryResults and Discussion:

Part 1:
With the data in table 2 and the method of initial rates (2), the orders with respect to each reactant were calculated. The order for each reactant was 1, making each reactant first order and the overall reaction, second order. This indicated that when the concentrations of each reactant is doubles independently, the reaction rare doubles. The rate constant was calculated to be .00929 mol/L using the data in determination 1 and equation (1); solving for k. With known reaction orders, the rate equation was written as; rate= k[H2C2O4][KMnO4]. Part 2:

Using the data collected in table 3 and the reaction rate equation (4), the reaction rate was calculated. With known reaction rates, the average factor for both time decrease and reaction rate increase were calculated. The average factor by which elapsed time is decreased with each 10 °C increase was 2. The average factor by which reaction rate increased with each 10 °C increase resulted in 2. This proved the “Rule of Thumb”1, stating that the reaction rate is expected to double for every 10 °C increase in temperature, thus making the amount of time cut in half. In conducting this experiment, there were several potential errors. Some errors could include failing to dispense the exact amount of reactant into the test tube, failing to mix the solution with a stirring rod after each reactant was added, and failing to properly time each reaction. For example, if the timer was not stopped when each solution reached the same color, it could alter the results because each reaction was carried out to a different point, not necessarily the end point. In part two, some potential errors include, failing to initially heat the reactants in the warm bath and failing to submerge all of the solution in the water. These errors could result in the reaction proceeding at a slower rate due to the lower temperature of the solution, affecting the proper results.