Determination of a Rate Law

Experiment (11) Determination of a Rate Law Objectives: To determine the rate law for the redox reaction Apparatus and Chemicals: 250-mL Erlenmeyer flask (5) Thermometer 250-mL beakers (5) 0.1M H2O2 50-mL graduated cylinder 0.3M KI Pipet 0.02M Na2S2O3 Buret Starch Stirring rod Buffer containing 0.5M HC2H3O2 Stop watch and 0.5M NaC2H3O2 Graph papers Introduction: The rate of a chemical reaction describes how quickly the reactants disappear and the products form, and expressed as mol L-1 s-1. hypothetical reaction, For the the rate can be measured by observing the rate of disappearance of either A2 or B2, or the rate of appearance of AB2, depending on which species is the easiest to detect and monitor during the reaction. The rate of a reaction is related to the concentration of reactants by the rate expression, for the above reaction the rate expression or the rate law is: rate expression or rate law where k is the specific rate constant, varies with temperature but its value is independent of reactant concentrations. The superscripts p and q are the order of the reaction with respect to A2 and B2 respectively. p and q are always determined experimentally for a given reaction, and there is no relation between their values and the coefficients in the chemical equation. The overall reaction order is p + q. One of the objectives of chemical kinetics is to determine k, p, and q. In this experiment you will determine the rate law for the reaction between hydrogen peroxide and potassium iodide in aqueous solution: In this reaction H2O2 oxidizes I- to I2 in the presence of H+ ions, the rate law is; This reaction is very rapid to be detected when the concentration of H+ is greater than 10-3 M. However, if the concentration is less than 10-3M, the reaction occurs at a measureable rate. The concentration of H+ is maintained constant during the reaction at a concentration of 1 x 10-5M by using a buffer solution of sodium acetate / acetic acid solution. As [H+] is consumed during the reaction, the buffer equilibrium shifts to the right to supply additional [H+], Therefore, the reaction rate is independent of the H+, and the rate is: a. Determination of p: To determine p, [H2O2] is kept constant and [I-] is changed, and therefore, c is a constant = [H2O2]q Taking the logarithms of both sides: Combining all constants, we have an equation for a straight line: or Therefore, a plot of log (Rate) Versus log [I-] yields a straight line with a slope equal to p. b. Determination of q: To determine q, the concentration of [H2O2] varies, while [I-] is kept constant, and the logarithmic form of the rate law is: Where C' equals log k' + log[I-]p. Therefore, a plot of log(Rate) versus log [H2O2] yields a straight line with a slope equal to q. c. Measuring reaction rate: The rate of the reaction is measured as follows: (1) Solutions of (thiosulfate), and starch are prepared. reacts with I-, I2 is produced, which reacts rapidly with (2) When : the consumption of I2 proceeds until all (3) When is used. is gone, the produced I2 reacts with starch forming a deep blue color due to the formation of (I2. Starch) complex: (4) The reaction rate is then followed by observing the time required for a measured amount of to react completely, i.e the time required the rate of I2 production (mol I2 / time), which is affected only by the initial H2O2 and I- concentrations. (5) The number of moles of I2 = 1/2 the number of moles of indicated in the above equation between I2 and , as . Example: 25mL of a buffer solution and starch were added to 15mL of 0.2M KI. To this solution, 10 mL of 0.05M Na2S2O3 was added. When 25 mL of 0.1M H2O2 was added a blue color appeared after 90 seconds. Calculate the measured rate and the rate constant for this reaction? Assume a first order reaction with respect to both I- and H2O2, i.e rate = k[I-] [H2O2]. Rate = mol I2 / time (s) Total volume of solution aftr mixing = 25 + 15 + 10 + 25 = 75 mL The, initial concentrations of I- and H2O2 are calculated as follows: Using the rate law for the reaction we have: and Procedure: Use clean and dry glassware, do not dry glassware with paper towels. 1. Preparation of solutions of reactants: The following Table summarizes the preparation of the test solutions required for this experiment, Each test is performed by mixing two solutions A and B. Solution A contains the I- reactant in addition to buffer, starch, and Na2S2O3. Solution B contains the H2O2 reactant. Table (1): Composition of test solutions (All volumes are measured in mL) Solution A Test solution Distilled water Solution B Buffer 0.3M KI Starch 0.02M Na2S2O3 0.1M H2O2 1 61.5 12.5 5.0 1 5.0 15 2 56.5 12.5 10.0 1 5.0 15 3 51.5 12.5 15.0 1 5.0 15 4 41.5 12.5 15.0 1 5.0 25 5 31.5 12.5 15.0 1 5.0 35 1- In a clean and dry 250 Erlenmeyer flask or 250-mL beaker add components of solution A required for test 1. Measure the volumes of KI and Na2S2O3 using either a pipet or a buret. Stir thoroughly and record the temperature. 2- Measure 15mL of 0.1 M H2O2 (solution B) in a clean and dry 50-mL graduated cylinder and transfer quickly to the Erlenmeyer flask. The reaction begins where. H2O2 is added to solution. A. After mixing start timing immediately in seconds with a watch or clock until the blue color of I2. starch appears. Stop timing and record time elapsed to the nearest second. 3- Repeat steps 1 and 2 for test solutions 2, 3, 4 and 5 as described in Table 1. B. Calculations: Refer to the Report sheet and calculate the following: 1. Number of moles of 2. Number of moles of . . 3. The rate of reaction Since initial mol I2 = 0, and initial time = 0, we have: . 4. Calculate the initial concentrations of I- and H2O2. Note that the total volume of reaction's mixture = 100 mL. 5. Calculate the logarithms of rate, [I-]i, and [H2O2]i. C. Determination of reaction order: 1. Plot on linear graph paper log (mol I2/t), i.e log rate versus log [I-]i for test solutions 1, 2, and 3 (these solutions have constant concentrations of H2O2). Draw the best straight line through the three points. Calculate the slope of the straight line. The value of the slope is the order of the reaction, p, with respect to I-. 2. Plot on linear graph paper log (mol I2/t) vs log [H2O2] for test solution 3, 4, and 5 (these solutions have constant concentrations of I-). Draw the best straight line and calculate the slope which gives the order of reaction, q, with respect to H2O2. D. Determination of the rate constant, k': Use the values of p and q (obtained from part C) and the rate law equation: To determine k' for the five solution. Calculate the average values of k' and determine the average deviation and standard deviation. Report Sheet: Sec: Name: No.: Date: A. Data Solution Temp. 1 2 3 4 5 _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ _______ (mol/sec) _______ _______ _______ _______ _______ 4. log (mol I2/t) _______ _______ _______ _______ _______ 5. [I-]I (mol/L) _______ _______ _______ _______ _______ 6. log [I-]i _______ _______ _______ _______ _______ 7. [H2O2] (mol/L) _______ _______ _______ _______ _______ 8. log [H2O2]i _______ _______ _______ _______ _______ 1. Solution Temperature (oC) 2. Time elapsed for color change, t (seconds) B. Calculations 1. Mol reacted 2. Mol I2 produced 3. Mol I2/t, rate C. Determination of reaction order Plot of log (mol I2/t) vs log [I-]i: Slope = p = ______________ Plot of log (mol I2/t) vs log [H2O2] Slope = q = ______________ Rate law of the reaction: D. Determination of k' 1 2 3 4 5 1. Values of k' _______ _______ _______ _______ _______ 2. Average value of k' : ______________ 3. Average deviation : ______________ 4. Standard deviation : ______________ E. Show your calculations for solution 1