The Rate and Order of a Chemical Reaction

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Table of Contents

Abstract
Introduction
Experimental
Data and Results
Discussion and Conclusion
Works Cited

Abstract

The rate of a chemical reaction is the speed at which the reactants are consumed or the speed at which new products are made. To determine the progress of chemical reactions, it is necessary to monitor parameters such as absorbance, temperature and pH, which are used to establish the rate and the rate law of chemical reactions. This experiment investigated the reaction between potassium iodide and iron (III) chloride. It intended to find the order and rate law for the reaction. The Logger Pro program was used to find out the changes in concentration of the products produced in the reaction between potassium iodide and iron (III) chloride, which were measured as alterations in absorbance. The method of initial rates was employed in the determination of the order and rate law for the reaction. It was established that the reaction between potassium iodide and iron (III) chloride was a first order reaction. It was concluded experiments were the only reliable means of determining rate laws and order of chemical reactions.

Introduction

The analysis of chemical reactions regarding velocities and the outcome of certain variables on the velocity is known as chemical kinetics. These studies require that reactions are done at varying dilutions of reactants. This ascertains how altering these factors influences the entire reaction.

Rate law, therefore, is an equation that illustrates the role of reactant concentrations and catalysts in the overall rate of a chemical reaction (Saxena 97). For a reaction such as bB + cC →dD + eE, the rate law is given by rate= k [B]^a + [C]^b where k is the rate constant while ‘a’ and ‘b’ are the reaction orders. [B] and [C] are the concentrations of the reactants B and C respectively.

Experimental methods are employed in establishing the rate of chemical reactions since the velocities of chemical reactions can only be found experimentally. In this experiment, the advancement of the reaction between KI and FeCl₃ was determined through spectrophotometry by quantifying the changes in absorbance with time.

The rate of the reaction and the rate law expression for the reaction were then established using the method of initial rates.

Experimental

A calorimeter was coupled to Channel one of the Vernier computer interface that was subsequently linked to the computer using the appropriate cable. The file ‘25 rate and Order’ in the Logger Pro program was used to perform the experiment. The calorimeter was calibrated, and its wavelength was set to 430 nm (Grossie and Underwood 38). All the necessary materials were assembled in readiness for the first trial. Twenty milliliters of 0.020M FeCl₃ and 20 ml of 0.020M KI were measured into separate 100 ml beakers. Thereafter, the two reagents were mixed and swirled gently. The cuvette containing water was emptied and rinsed with the mixture of the two chemicals. The cuvette was then filled three quarter way with the mixture and placed inside the calorimeter. The ‘collect’ button was clicked to gather absorbance data. The progression of the reaction was assessed for two minutes after which the cuvette was emptied and rinsed with distilled water in readiness for the subsequent trials. The linear region within the first minute in the graph of the initial run was selected and analyzed to obtain the reaction rate. A straight line from which the gradient of the line was obtained was made by clicking the linear regression button. The slope was recorded as the initial rate of reaction for the first trial.

The above procedure was reiterated for the second and third trials. Twenty milliliters of FeCl₃ and 10 ml of KI were used in the second trial, whereas the third trial used 10 ml of FeCl₃ and 20 ml of KI.

Data and Results

It was realized that raising the concentration of KI did not affect the rate of the reaction significantly while lowering the concentration of FeCl₃ raised the rate of the reaction twice. The results are indicated in table 1.

Table 1: The rates of reaction of the reaction between FeCl₃ and KI at varying concentrations of the reactants

Trial	FeCl₃	KI	Slope
1	0.01	0.01	0.00272959
2	0.01	0.005	0.00189996
3	0.005	0.01	0.0030891

Discussion and Conclusion

The reaction between FeCl₃ and KI was a first order reaction. A pair of experimental runs where the concentration of only one reactant altered was used to find the order of the reactant. In the first and second runs, the concentration of KI was reduced by half without affecting the initial reaction rate. For that reason, the order of KI was zero^th. In the first and third runs, reducing the concentration of FeCl₃ increased twofold the rate of the reaction meaning that the reaction was a first order regarding FeCl₃. The overall order was given by the sum of the order of FeCl₃ (1) and the order of KI (0), which was 1.

The rate law expression for the reaction was rate=k [FeCl₃]¹[KI]⁰, which was equivalent to rate=k[FeCl3]¹. Calculating rate constant from the data was possible because the order of the reaction and the concentrations of the reactants were known. The rate constant was calculated by substituting the rate of one of the runs into the rate law expression.

0.0030891=k [0.005]¹[0.01]⁰
K=0.0030891Ms^-1/0.005 M
K=0.61782 s^-1.

The full rate equation was given by rate= 0.61782 [FeCl₃]. The results were consistent with other experiments in the literature concerning the rate and order of the reaction between FeCl₃ and KI. It was concluded that experiments were the only reliable means of determining rate laws and reaction rates of chemical reactions.

Works Cited

Grossie, A David and Kirby Underwood. Laboratory Guide for Chemistry. 6^th ed. 2013. Plymouth, MI: Hayden-McNeil. Print.

Saxena, P. B. I. I. T Chemistry. Meerut, India: Krishna Prakashan Media, 2012. Print.

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