The rates at which reactions occur depend on the composition and the temperature of the reaction mixture. Usually the rate of reaction is found to be proportional to the concentrations of the reactants raised to a power.1 There are many reactions that have a rate law in the form of: (1) v = k[A]a[B]b
According to reference1 the power to which the concentration of a species (product or reactant) is raised in a rate law of this nature is the order of the reaction with respect to that species. In equation (1) first order with respect to [A] and first order with respect to [B]; however, the overall reaction is the sum of the individual orders. Thus we have a second order reaction. In this experiment a hexacyanoferrate(III) ion ([Fe(CN)6]3-) oxidizes ascorbic acid (C6H8O6) by the following reaction: (2) 2[Fe(CN)6]3- + C6H8O6 = 2[Fe(CN)6]4- + C6H6O6 + 2H+
The reaction above is of a first order reaction at room temperature with respect to individual reactants; therefore the reaction stoichiometry and rate law at time t are: (3) aA + bBproducts and (4) -d[A] = k[A] [B]
where [A] represents the concentration of ascorbic acid and [B] represents the concentrations of [Fe(CN)6]3- at time t. For this experiment we will use an integrated rate law in the form of: (5) ln [A] = b [A]0 – a [B]0 kt + ln [A]0
Rates of Reaction BACKGROUND INFORMATION What affects the rate of reaction? 1) The surface area of the magnesium. 2) The temperature of the reaction. 3) Concentration of the hydrochloric acid. 4) Presence of a catalyst. In the experiment we use hydrochloric acid which reacts with the magnesium to form magnesium chloride. The hydrogen ions give hydrochloric acid its acidic properties, so that all ...
where [A]0 and [B]0 are the initial concentrations of C6H8O6 and [Fe(CN)6]3- and a=1 and b=2. From equation (5), it is possible to calculate the second-order rate constant k by plotting ln [A]/[B] against time (find slope of line where b=2 and a=1).
EDTA in this experiment is used as a masking agent to hide metal ions that would normally interfere with the analysis in this reaction. Thus the absorbance of [Fe(CN)6]3- at time t is given by: (6) Absorbance = 1012 [Fe(CN)6]3-
The oxidation of C6H8O6 by [Fe(CN)6]3- involves a mechanism that consists of 3 steps.2 In the first step, the ascorbate ion (AH-) is rapidly formed by ionization of the ascorbic acid. (7) AH2 AH – + H+
Following the ionization is the slow rate-determining step, the oxidation of the ascorbate ion to an ascorbate free radical (AH∙): (8) [Fe(CN)6]3- + AH-[Fe(CN)6]4- + AH∙ During the final step, an electron is rapidly transferred from the ascorbate free radical to the hexacyanoferrate(III) anion, producing dehydroascorbic acid (A): (9) [Fe(CN)6]3- + AH- [Fe(CN)6]4- + A + H+
The slow rate-determining step is an ionic reaction between [Fe(CN)6]3- and AH-. According to reference3, the specific rate constant of an ionic reaction in aqueous solution depends on two factors: the ionic strength I of the solution and on the charges ZA and ZB of the ionic species reacting to for the activated complex. (10) log k = log k0 + 1.02ZAZB I1/2
All reagents in this experiment were of reagent grade. Mass measurements were taken on a Shimadzu Libror AEG-120 analytical scale with an uncertainty of ±0.0001. Manual data acquisition was taken with a Barnstead/Turner SP-830 spectrophotometer and a stopwatch. The computerized data acquisition was completed by a Cary 50 Bio. The experiment began by preparing four solutions of 1 x 10-3 M of K3Fe(CN)6 with varied concentrations of NaNO3: 0.025 M, 0.05 M, 0.1 M and 0.2 M.
This was completed by dissolving 0.0329245 (±0.001) g of K3Fe(CN)6 with the specified concentrations of NaNO3 and deionized water in a 100 mL volumetric flask. A 25 mL aliquot of each solution was transferred into a 250 mL Erlenmeyer flask and the temperature of the aliquot was recorded. Next, a 500 mL 2.5 x 10-4 M solution of ascorbic acid was prepared by using a standardized 0.01 M HNO3 solution dissolved in 0.005 g of EDTA and deionized water. A 25 mL aliquot was transferred into each of the four 100 mL beakers by using a 25 mL pipet.
Introduction: In 1909 S. P. L. Sorensen published a paper in Biochem Z in which he discussed the effect of H 1+ ions on the activity of enzymes. In the paper he invented the term pH to describe this effect and defined it as the -log[H 1+ ]. In 1924 Sorensen realized that the pH of a solution is a function of the 'activity' of the H 1+ ion not the concentration and published a second paper on the ...
The spectrophotometer was set to 418 nm and the absorbance reading was zeroed by using deionized water as a standard. The ascorbic acid in the beaker was poured into the K3Fe(CN)6 solution and the timer was immediately started. The Erlenmeyer flask was swirled for 2-3 seconds before pouring the reacting mixture into a 1-cm cuvette. The cuvette was conditioned with the reacting solution 4 times before being placed into the sample holder of the spectrophotometer. An absorbance reading was taken at 30 seconds and every 30 seconds thereafter for a total of 6 minutes. The same process was implemented with the Cary 50 Bio except that each sample was analyzed by the computer for 7 minutes and 53 seconds. Data/Results