Potassium Iodide and Iron (III) Chloride Chemical Reaction Report

Abstract

The rate of a chemical reaction can be defined as the rate at which the reactants are used up in the reaction or the rate at which new products are made. Certain factors such as temperature and concentrations of reactants influence the velocity of chemical reactions. The investigation of the progress of chemical reactions is called chemical kinetics. Such studies involve following up certain parameters, for example, absorbance to monitor the progress of the reaction. The experimental parameters are then used to compute the rate and rate law of the reaction. This experiment investigated the reaction between potassium iodide and iron (III) chloride. It aimed at finding the order and rate law for the reaction. The logger Pro program was used to determine the changes in concentration of the products formed in the reaction between potassium iodide and iron (III) chloride. 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 that rate laws and order of reactions could be established only by experimental means.

Introduction

The study of chemical reactions as concerns the velocities of the reaction and the outcome of various variables on the velocity is called chemical kinetics. Such studies involve performing reactions at different concentrations of reactants to establish how changing these parameters affects the overall rate of the reaction, which is the quantity of products yielded by the reaction as a function of time. Various reactions have varying velocities. In reactions with high velocities, the amount of reactant diminishes at a faster rate. Conversely, the amounts of reactants decrease at a lower rate in reactions with slow velocities.

Rate law, therefore, is an equation that shows the contribution of reactant concentrations and catalysts to the overall rate of a chemical reaction (“Kinetics: Rates and Mechanisms of Chemical Reactions” 627). For a reaction such as eE + fF →gG + hH, the rate law is given by Rate= k [E]^a + [F]^b. The key components of a rate law are the concentration of the reactants, the rate constant and reaction orders. In the above example of a rate law, k is the rate constant while ‘a’ and ‘b’ are the reaction orders. [E] and [F] are the concentrations of the reactants E and F respectively.

The power to which the concentrations of reactants are raised in a rate law is referred to as the order of the reaction. Zero order with regard to a given reactant in a reaction implies that altering the concentration of the reactant does not affect the rate of the reaction. Therefore, the reaction rate is independent of the concentration of the reactant. The first order in connection with a particular reactant in a chemical reaction means that increasing or reducing the concentration of the reactant doubles the overall velocity of the reaction. In second order reactions, on the other hand, doubling the concentration of a given reactant causes a fourfold escalation in the rate of the reaction.

Experimental methods are often used to determine the rate of chemical reactions because it is only possible to establish the velocities of chemical reactions by experimental runs. One such method entails varying concentrations of the reactants and obtaining the resultant velocity. In this technique, the concentration of one reactant is changed while the concentration of the other is kept constant. The advancement of the reaction with time is then determined by spectrophotometric techniques and making a plot of concentration versus time. The gradient of the plot usually gives the initial rate of the reaction. The effects of the changes in reactant concentrations are then evaluated using the method of isolation to determine the rate of the reaction.

Time-dependent concentrations can also be used to determine the rate law of reactions. In this method, the change in the concentration of a given reactant with time is monitored. Three different plots can be made to determine the order of the reaction. A zero order reaction is established when the changes in concentration with time are constant. A first order reaction is proved when the plot of the natural logarithms of the concentrations against time yields a straight line. Conversely, a second order reaction is determined when graphing the inverse of the concentrations against time produces a straight line.

The reaction between potassium iodide and iron (III) chloride at changing concentrations of the reactants was investigated in this experiment. The reaction was given by the chemical equation 2Fe³⁺_(aq) + 2I^–_(aq) → I_2(aq) + 2Fe²⁺_(aq). The experiment intended to establish the rate of the reaction as well as the rate law expression for the reaction using the method of initial rates.

Experimental

Goggles were worn to ensure the safety of the eyes before commencing the experiment. A calorimeter was linked to Channel one of the Vernier computer interface, which was then connected to the computer using the appropriate cable. The file ‘25 rate and Order’ was opened under the Logger Pro program that was used to perform the experiment (Grossie and Underwood 38). The calorimeter was calibrated, and its wavelength was set to 430 nm. Thereafter, all the materials that were required for the procedure were assembled. They included three graduated measuring cylinders whose capacities were 25 ml, 100 ml of 0.020 M potassium iodide solution in a 100 ml beaker, 100 ml of 0.020 m iron (III) chloride solution that had been dissolved in 0.1 M hydrochloric acid, and 60 ml of distilled water.

Twenty milliliters of iron (III) chloride and 20 ml of KI were measured into two 100 ml beakers. The cuvette containing water was removed from the calorimeter, and the water that was held in it was discarded. Twenty milliliters of iron (III) chloride were added to the beaker containing the potassium iodide solution, and the beaker was swirled gently for the reactants to mix. The mixture of the two chemicals was used to rinse the cuvette, which was then filled three quarter way with the mixture. Subsequently, the cuvette was placed inside the calorimeter after its outer surface had been wiped clean. The ‘collect’ button was clicked to start the collection of absorbance data. The advancement of the reaction was examined for two minutes. The contents of the cuvette were then discarded after which the cuvette was rinsed with distilled water in preparation for the next trial. The graph of the initial run was carefully observed, and the linear region within the first minute was selected and analyzed to obtain the rate of the reaction. The linear regression button was used to produce a straight line from which the gradient of the line was obtained. The slope was noted down as the initial rate of reaction for the first trial, and the linear regression box was closed.

The above procedure was repeated for trial 2 and 3. Twenty milliliters of iron (III) chloride and ten milliliters of potassium iodide were used in the second trial, whereas the third trial used ten milliliters of iron (III) chloride and twenty milliliters of potassium iodide. The data were recorded in a table and were used to determine the order and rate of the reaction.

Data and Results

Table 1: The rates of reaction of the reaction between iron (III) chloride and potassium iodide at varying concentrations of the reactants

Table one above showed that increasing the concentration of KI did not alter the rate of the reaction, whereas reducing the concentration of FeCl3 increased the rate of the reaction twofold.

Discussion and Conclusion

The concentrations of the reactants were calculated from the volume used and the concentration of the stock solutions in moles per liter as follows. The original FeCl₃ solution had a concentration of 0.020 M, which meant that 1000 ml of FeCl₃ contained 0.020 moles. Therefore, the number of moles that were present in 20 ml of FeCl₃ was given by the equation:

• 0.02 M = number of moles / 0.005 L
• Number of moles = 0.0001
• Concentration = 0.0001 Mol / 0.01 L
• Concentration = 0.01 M

The reaction between FeCl₃ and KI was a first order reaction. The order of the reaction was obtained from the experimental values using the method of isolation. From table 1 above, a pair of experimental runs where the concentration of only one reactant varied was selected and used to determine the order of the reactant. In runs 1 and 2, for example, it was seen that the concentration of potassium iodide (KI) was reduced by half the concentration of FeCl₃. Reducing the concentration of KI did not have any impact on the initial rate of the reaction, which was given by the slope at that concentration. Therefore, the order of KI was zero^th. In runs 1 and 3, it was seen that lowering the concentration of FeCl₃ doubled the rate of the reaction implying that the reaction was a first order regarding FeCl₃. The overall order for the reaction was then given by the sum of the order of FeCl₃ (1) and the order of KI (0). The sum of zero and one gave one as the overall order of the reaction.

The rate law expression for the reaction was given by rate=k [FeCl₃]¹[KI]⁰, which was also given by rate=k[FeCl3]¹ since [KI]⁰ was equivalent to zero. It was possible to compute the rate constant from the data because the order of the reaction had been determined and the concentrations of the reactants were known. The rate constant was calculated from the rate equation rate=k [FeCl₃]¹[KI]⁰. The concentrations of reactants and the rate of reaction in one of the runs (run 3) were then substituted into the equation.

• 0.0024373=k [0.005]¹[0.01]⁰
• K=0.0024373 Ms^-1/0.005 M
• K=0.48746 s^-1.

The complete rate equation was, therefore, given as rate= 0.48746 [FeCl₃]. It was concluded that reaction rates and rate laws could be obtained from experimental values only since initial reaction rates could only be found by performing experiments.

Works Cited

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

Kinetics: Rates and Mechanisms of Chemical Reactions n.d. Web. 2014.