Catalytic decomposition of hydrogen peroxide Experiment 4 APPARATUS - 2 Erlenmayer flasks (100cm 3 ) - Pipette (25cm 3 ) - Pipette (5cm 3 ) - Biuret with stand - Stop watch REAGENTS - 5% solution of H 2 O 2-0.05M solution of FeCl 3 + 0.4M HCl (catalyst of the reaction) - 2M solution of H 2 SO 4-0.1M solution of HCl - 0.1M solution of KMnO 4-0.1M solution of MnCl 2 PURPOSE OF THE EXPERIMENT The goal of the experiment is to establish reaction rate constant in catalytic decomposition of H 2 O 2 in the presence of various amounts of catalyst (MnO 2 ) and without catalyst. INTRODUCTION The reaction rate for a given chemical reaction may be expressed as the change in concentration of the reactants or the change in the concentration of the products per unit of time. Therefore it can be defined as: (1) where: c concentration of the reactants; x concentration of products; t time. Catalyst is the chemical compound which accelerates chemical reaction. The role of the catalyst is to decrease activation energy via creating an alternative reaction route, which eliminates the slowest reaction step. It results in total increase of reaction rate at given temperature. Catalysts performance may be very effective e.g. activation energy of H 2 O 2 decomposition in aqueous solution and at room temperature is equal to 76kJ/mol, thus this reaction proceeds slowly at room temperature. Addition of small amounts of I - ions into H 2 O 2 aqueous solution will result in decrease of activation energy to the value of 57kJ/mol, increasing at the same time reaction rate ca. 2000 times. 1
Catalysis is divided into two main groups: Homogenous catalysis when catalyst and reactants are in the same phase. Heterogeneous catalysis when catalyst is present in a separate phase, most commonly these type of reactions proceed on the solid surface of catalyst and reactants are present in the gaseous phase. Indirect group of catalytic processes is microheterogenous catalysis, when catalyst is present as a separate phase, but it is highly dispersed. Into that group of reactions biocatalysts may be included, which describes biological reactions proceeding in the presence of catalyst (enzyme). Catalytic reaction may be described in few consecutive steps: 1. Transport of reactants from gaseous or liquid phase into the surface of catalyst. This step is controlled by diffusion rate, which might be regulated e.g. change in the velocity of stirring. 2. Adsorption of the reactants on the catalyst surface. This step is controlled by the rate of adsorption. 3. Reaction between adsorbed species on the catalyst surface surface reaction. 4. Desorption of the products of reaction. 5. Transport of products form catalyst surface to gaseous or liquid phase. Total reaction rate is influenced by the rate of the slowest step. Reaction rate constant or reaction rate coefficient is increasing with the temperature (for each step) according to Arrhenius Law: (2) where: A constant; E a activation energy; T temperature; R gas constant; The process of adsorption is especially important in heterogeneous catalysis. Adsorption occurs on the most active spots on catalyst surfaced called active centres. If in catalytic process gaseous reagents are involved, the adsorption process is usually described by Langumir isotherm. In case of reaction proceeding in the liquid phase, the Freundlich isotherm is used to described adsorption process. According to Ostwald reaction rate in homogenous or biocatalysis may be described by the equation: (3) where: k 1 and k 2 constants; a starting concentration of reactant; b concentration of catalyst; x concentration of product after time t; n order of reaction; Right side of the equation (3) may be expressed as: (4) 2
It may be concluded from the equation (4) that rate of catalytic reaction is the sum of the rates of two processes: (i) process without the presence of catalyst and (ii) process dependent on the catalyst presence. If the sum (k 1 + k 2 b) will be described by k, then equation (3) becomes kinetic equation of n-ordered reaction: (5) The final effect of catalyst performance may be thus expressed as the change in value of reaction rate constant. An example of microheterogenous catalysis is decomposition of H 2 O 2 in the presence of MnO 2. The progress of the reaction may be measured by titration of H 2 O 2 concentration with KMnO 4. Decomposition of H 2 O 2 id first order reaction, thus it is described by equation: (6) where: c 0 starting concentration of reactant; c concentration of reactant after time t; In case of hydrogen peroxide decomposition c 0 describes starting concentration of H 2 O 2, and c concentration of H 2 O 2 after time t. Both concentrations are proportional to volume of KMnO 4 used upon titration, thus: (7) where: V 0 volume of H 2 O 2 solution; V t volume of titrant By combining equations (6) and (7) we obtain: (8) or (9) The graph of lnv t =f(t) is straight line, and its slope (angular coefficient) is equal to minus reaction rate constant (-k). EXPERIMENTAL Pour into Elenmayer flask 25cm 3 of H 2 O 2 solution and 5cm 3 of solution containing 0.05 mol/dm 3 FeCl 3 and 0.4 mol/dm 3 HCl. In the prepared mixture the following reaction will proceed: H 2 O 2 H 2 O + ½O 2 (10) This reaction is accelerated in the presence of Fe 3+ ions. From the prepared solution measure 5cm 3 and move it to Erlenmayer flask containing 25cm 3 2M H 2 SO 4 and 1 cm 3 of 0.1M MnCl 2. So prepared solution should be titrated by 0.1 M KMnO 4 until pale 3
pink colour of solution will be obtained. Next to titrations should be performed after 30min and 60min after the first titration. The results of measurements should be put in Table (according to the example presented below). Volume of the solution containing catalyst [cm 3 ] time[s] V KMnO4 [cm 3 ] lnv KMnO4 MEASURING CONCENTRATION OF H 2 O 2 BY TITRATION WITH KMnO 4 Hydrogen peroxide in the presence of potassium permanganate is behaving as a reducer and in acidic environment it reduces KMnO 4 to Mn 2+ according to the reaction: 5H 2 O 2 + 2 MnO - 4 + 6 H + 5O 2 + 2 Mn 2+ + 8H 2 O (10) 5H 2 O 2 + 2KMnO 4 + 3H 2 SO 4 5O 2 + 2MnSO 4 + K 2 SO 4 + 8H 2 O (11) Reaction is catalysed by Mn 2+ ions; First drops of KMnO 4 are discoloured very slowly, however, when concentration of Mn 2+ is increased, reaction rate becomes fast. Durability of H 2 O 2 decreases with the decrease of concentration. Therefore, titration should be performed immediately after addition of tested solution. Simple stoichiometric calculations are needed in order to calculate concentration of H 2 O 2 in tested solutions (2 moles of KMnO 4 react with 5 moles of H 2 O 2 ). RESULTS 1. Draw graph presenting the lnv KMnO4 as a function of time and basing on that graph establish reaction rate constant. 2. Results of calculations should be presented in Table. 3. Study the influence of catalyst concentration on reaction rate. According to Ostwald the reaction rate constant is linear function of catalyst concentration. Colloidal MnO 2 is a catalyst of hydrogen peroxide decomposition, which is formed upon addition of KMnO 4. Thus the catalyst concentration is proportional to the volume of added KMnO 4. The Ostwald equation is true when the graph of k=f(v kat ) is linear. 4. Discuss obtained results. THEORETICAL ISSUES: 1. Definition of reaction rate, reaction rate constant and order of reaction. 2. Kinetic equation of first and second order reactions. Methods of measuring reaction rate constant. 3. Kinetics of catalytic reactions. 4. Catalysis and its division. 4
5. Applications of catalysis in industry. 6. Manganometry (properties of KMnO 4, examples of manganometry application). LITERATURE 1. Peter Atkins, Physical Chemistry 2. Jens Hagen, Industrial Catalysis: A Practical Approach, Wiley 3. I. Chorkendorff, J.W. Niemantsverdriet, Concepts of Modern Catalysis and Kinetics, Wiley 4. James T. Richardson, Principles of Catalyst Development, Springer 5. Guy Marin, Gregory S. Yablonsky, Kinetics of Chemical Reactions, WIley 5