Kinetic Reaction

Abstract
This experiment is to study the effect of temperature on the rate of reaction between potassium permanganate with oxalic acid. We used 2cm3 of 0.02M potassium permanganate and 4cm3 of 1M sulphuric acid into a test tube. In another test tube, we placed 2cm3 of oxalic acid. We placed the test tubes in a water bath at 40, 45, 50, 55 and 60oC respectively. When the solutions have attained these temperatures pour the oxalic acid into the acidified permanganate solution and recorded the time taken for the permanganate to decolorize. At 400C , the time taken for the permanganate to decolorize is 66 seconds. At 600C, the time taken for the permanganate to decolorize is 10s. The higher temperature of the reaction, the faster the time taken for the permanganate to decolorize.

This is because the higher temperature implies higher average kinetic energy of molecules and more collisions per unit time. The rate of effective collision increases, the rate of reaction increases. As a result, the time taken for reaction decreases when temperature increasing. The graph shows that 1/T is decreasing when In 1/t is increasing. When the value of 1/T is bigger, the value of In 1/t will be smaller. In opposite situation, the value of 1/T smaller, the value of In 1/t will be bigger. Activation energy is the minimum energy is needed by particle to form product. It can calculated by using Arrhenius equation, k = Ae-Ea/RT .

Introduction
The purpose of this experiment is to determine the rate activation energy of the reaction between potassium permanganate with oxalic acid. Usually, an increase in temperature is accompanied by an increase in the reaction rate. Temperature is a measure of the kinetic energy of a system. The higher temperature implies higher average kinetic energy of molecules and more collisions per unit time. The rate of effective collision increases, the rate of (2013, 07). Ghost Mun. StudyMode.com. Retrieved 07, 2013, from

xperiment 17
Reaction Kinetics- Determination of the Activation Energy of the Reaction Between Oxalic Acid and Potassium Permanganate.

Objective

To determine the activation energy of the reaction between oxalic acid and potassium permanganate.

Theory and Background

Activation energy is the minimum amount of energy that is required to activate atoms or molecules to a condition in which they can undergo chemical transformation or physical transport. In terms of the transition-state theory, the activation energy is the difference in energy content between atoms or molecules in a transition-state configuration and the corresponding atoms and molecules in their initial configuration. The activation energy is usually represented by the symbol Ea in mathematical expressions for such quantities as the reaction-rate constant, k = Aexp(-Ea/RT).The Arrhenius equation gives the quantitative basis of the relationship between the activation energy and the rate at which a reaction proceeds. This equation suggests that the activation energy is dependent on temperature, but this effect is cancelled by the temperature dependence of the reaction rate coefficient. A substance that modifies the transition state to lower the activation energy is termed a catalyst.

A biological catalyst is an enzyme. It is important to note that a catalyst increases the rate of reaction, but isn’t consumed by it nor does it change the energies of the neither original reactants nor products. Instead, the reactant energy and the product energy remain the same while the activation energy is lowered. Arrhenius concept is about the acids and bases split into ions when they dissolved in water. The Arrhenius acid-base concept classifies a substance as an acid if it produces hydrogen ions H+ or hydronium ions in water. A substance is classified as a base if it produces hydroxide ions OH- in water.

Chemical kinetics, also known as reaction kinetics, is the study of rates of chemical processes.Chemical kinetics includes investigations of how different experimental conditions can influence thespeed of a chemical reaction and yield information about the reaction’s mechanism and transitionstates, as well as the construction of mathematical models that can describe the characteristics of achemical reaction (Pearson, 2004).The rate of reaction is the change in the amount of reactants or products over a time interval.The rate of reaction in chemistry, is usually expressed in moles/second (mol/s) or molarity/second( mol/L∙s ). There are two ways in which the rate can be expressed: average rate of a reaction orinstantaneous rate of reaction. The average rate of reaction shows how the concentration or molaritychanges over a specific time interval, while the instantaneous rate of reaction shows the rate of changeat a specific time. Differential calculus can often be used when calculating the instantaneous rate of change. An easier method would be to calculate the slope of the line at the point of tangency.

There are many ways in which chemists can measure reaction rates. Various methods include:monitoring mass, pH, conductivity, pressure, colour, and volume. When monitoring mass, the mass of the reactants can be measured over a time interval and if it releases gas, the mass will decrease. In areaction involving gasses, the pressure of the system can change as the reaction progresses. Anexample would be the decomposition of hydrogen peroxide. As it decomposes in a closed system, thepressure will increase as oxygen gas is produced. As well as pressure and mass, change in colour can beused to monitor the progress of a reaction. The absorbance of light is directly related to theconcentration of the compound, so by observing the change in absorbance, the rate of reaction ismonitored.The speed of a chemical reaction is affected by factors such as the temperature, concentration,volume, surface area, and orientation. These factors are sufficiently explained through collision theory.

When the temperature is greater, there is a greater fraction of particles that have more energy than theactivation energy, enabling them to collide and react. These particles also have more kinetic energy. Byincreasing concentration, while keeping volume and pressure constant, there is a greater chance thatthe particles will collide and react. Decreasing the volume is essentially another form of increasing theconcentration. With greater surface area, more collisions can occur; increasing the rate of reaction.Lastly, orientation is the key for a reaction to occur. If particles do not collide with the correctorientation, a reaction will not occur. As well, catalysts have the ability to increase the rate of reactionby lowering the activation energy barrier.

Involved in an elementary reaction. They can be either unimolecular (one), bimolecular (two), ortermolecular (three). The molecularity of the slowest step of the reaction mechanism is equal to theorders of reaction. The slowest step of the reaction mechanism is also called the rate determining step.This is because a reaction can only be as fast as its slowest step, so it has bearing on the order of reaction.

3. Discussion:
3.1 Determination of Reaction Orders and Rate Constant
The experiments are conducted based on the rate equation, R = k [I-]n[S2O82-]m, where k is the rate constant while n and m are the reaction orders of I- and S2O82- respectively. As reaction orders, n and m is defined as the power to which the concentration of that reactant is raised to in the experimentally determined rate equation.nand mcannot be found theoretically and are experimentally determined to be 1. This means that the reaction is first order with respect to [I-] and first order with respect to [S2O82-]. The overall rate order is 2.This reaction is said to be bimolecular since two reactant species are involved in the rate determining step. It was observed that the rate of reaction increases with increasing concentration. The Collision Theory explains the phenomenon by stating that for a chemical reaction to occur, reactant molecules must collide together in the proper orientation and the colliding molecules must possess a minimum energy known as the activation energy, EA, before products are formed.

An increase in the concentration of reactants leads to an increase in the number of reactant molecules having energy ≥ EA, hence increasing the collision frequency. The increase in the effective collision frequency leads to an increase in the reaction rate. When performing a chemical kinetics experiment, the procedures have to be conducted at a constant temperature. According to the Arrhenius equation, k=Ae-Ea/RT, a slight increase in temperature increases reaction rate significantly as the equation is exponential in nature. This is affirmed by the Maxwell-Boltzmann distribution curve (diagram on the right) as a slight increase in temperature increases the number of colliding particles with Ea and consequently, reaction rates, significantly. Hence, because slight deviations in temperature may affect reaction rates significantly, the temperature at which the experiment was carried out must be kept constant.

To prevent errors from occurring, all glassware used in this experiment must be kept clean and dry to prevent contamination by the previous batch of experimental products. The overall volume of the solution was also kept constant at 26mL by adding deionized water, to standardize the conditions of the reaction environment, thus increasing accuracy.Swirling of the conical flask contents for the same length of time must be done consistently so that results obtained will be fair. Instead of swirling with one’s hands, the conical flasks can be placed on an electronic swirl to ensure consistent swirling when conducting the experiment. Also, there is inaccuracy as the stopwatch was stopped only when an arbitrary colour intensity was observed. There should be a consensus between lab partners as to when the stopwatch should be stopped.

3.2 Temperature Effect on a Chemical Reaction
The results of this set of experiment show that the rate of reaction increases as temperature increases. Using the Arrhenius equation, k=Ae-Ea/RT, the activation energy, EA, can be determined by keeping the concentration of all the reactants constant while varying the temperature for each experiment. When performing a chemical kinetics experiment, the procedures have to be conducted at a constant temperature. According to the Arrhenius equation, k=Ae-Ea/RT, a slight deviation in temperature changes reaction rate significantly. This is affirmed by the Maxwell-Boltzmann distribution curve (diagram on the right) as a slight increase in temperature increases the number of colliding particles with Eaand consequently, reaction rates, significantly. Hence, since slight deviations in temperature may affect reaction rates significantly, the temperature at which the experiment was carried out must be kept constant.

This is especially important for experiments being conducted at 10oC and 20oC, the conical flasks were placed in an ice bath to maintain the reaction temperature. There were several fluctuations above and below the desired temperatures. Moreover, the time taken for the blue solution to turn colourless is relatively longer for these 2 lower temperatures which creates a greater room for error. Keeping temperatures constant can be done by conducting the experiments in a thermostatic water bath. Reactants were poured imprecisely into the conical flask. There may be leftover reactants in the test tubes and some reactants may stain the sides of the conical flask during the addition. This reduces the concentration of the reactants in the conical flask. Pipetting the reactants into the conical flask would ensure that the reactants are added in the requisite quantities and that the eventual results are accurate. Swirling of the conical flask contents for the same length of time must be done consistently so that results obtained will be fair.

Instead of swirling with one’s hands, the conical flasks can be placed on an electronic swirl to ensure consistent swirling when conducting the experiment. Also, there is inaccuracy as the stopwatch was stopped only when an arbitrary colour intensity was observed. There should be a consensus between lab partners as to when the stopwatch should be stopped. The reaction is autocatalysed as the product of the reaction acts as a catalyst for the reaction. An autocatalysed reaction is slow at first and then becomes more rapidly as the catalyst is produced in the reaction. For the reaction, Mn2+ is the autocatalyst. This accounts for why vigorous effervescence of CO2 is not observed immediately when the reactants were added but only observed after a little while when Mn2+ is produced. 2MnO42- + 5C2O42- + 16H+ -> 2Mn2++10 CO2 + 8H2O

4. Conclusion:
The rate equation of the chemical reaction between I- and S2O82- to produce I2 and SO42- has been found to be: Rate = k[I-][S2O82-], where rate constant k =5.055 × 10-3 mol-1Ls-1 The reaction is first order with respect to [I-] and the reaction is first order with respect to [S2O82-]. The overall order of reaction is 2. This reaction is said to be bimolecular since two reactant species are involved in the rate determining step. Using the Arrhenius equation, k=Ae-Ea/RT, the activation energy, EA, of the oxidation reaction of oxalic acid by permanganate was determined to be 76.01KJmol-1. This means that the minimum amount of energy that reactant particles must possess in order to react successfully is experimentally determined to be 76.01KJmol-1.

5. References:

1)http://www.shodor.org/unchem/advanced/kin/arrhenius.html
2)http://www.webchem.net/notes/how_far/kinetics/maxwell_boltzmann.htm 3)http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA3/MAIN/AUTOCAT/PAGE1.HTM – See more at: http://webcache.googleusercontent.com/search?q=cache:http://www.art-xy.com/2011/10/lab-report-on-chemical-kinetics-initial.html#sthash.GOpENJx2.dpuf