Marc DeeleyCarl Caughell, Kristi Kulig
CHEM 126/Section 01
Dates of Experimentation: 10/12/10; 10/19/10
Title: Studying the Rate of Reaction of Potassium Permanganate and Oxalic Acid Abstract:
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...
References: Michael Stranz, Signature Labs Series, Cengage Learning: Mason, Ohio, 2008; pgs. 23-32.
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