Experiment 13: The Decomposition of Hydrogen Peroxide, Lab Reports of Chemistry

Download Experiment 13: The Decomposition of Hydrogen Peroxide and more Lab Reports Chemistry in PDF only on Docsity! 65   E x p e r i m e n t 13 The Decomposition of Hydrogen Peroxide Objectives • To determine the general rate law of a reaction. • To determine the rate constant for a reaction. • To determine the activation of a reaction from two trials or multiple trials. • To understand the functions and properties of catalysts. • To understand dilution calculations. • To convert between mass percent and molarity. In the Lab • Students will work in pairs. Waste • Solutions can be washed down the drain with excess water. Safety • The stopper in the test tube may pop out if the pressure raises enough. Do not point it at yourself or others.  66  E x p e r i m e n t 1 3 • The Decomposition of Hydrogen Peroxide Expt. 13 Any molecule in motion possesses kinetic energy, and the faster it moves, the greater the kinetic energy it has. When molecules collide, some of the energy is converted into vibrational energy. If the vibrational energy is large, it may cause some of the chemical bonds in the molecule to break. Breaking bonds is the first step towards product formation. If the initial kinetic energies are small, the molecules will merely bounce off each other without breaking any bonds. In order to react, the molecules must have a total kinetic energy equal to, or greater than, the activation energy, Ea. The activation energy is the minimum amount of energy required to initiate a chemical reaction. The energy change in a reaction is given on an energy- level diagram, or potential energy profile, as shown in Figure 13.1. The vertical axis gives the potential energy for the reaction, while the horizontal axis is a relative (i.e., time) scale that shows the progress of the reaction. The diagram indicates that there is a “hill” or energy barrier that needs to be overcome before any products can be formed. If the collision between the molecules produces enough energy to overcome the barrier, the reactant molecules are in a temporary transition state, forming an activated complex at the height of the barrier before forming the product molecules. ©Hayden-McNeil, LLC Progress of Reaction R el at iv e E ne rg y Reactants Products Ea Activated complex Hrxn Figure 13.1. Energy diagram of a reaction. Factors Influencing the Reaction Rate The rate at which the reaction occurs depends on several factors: 1. the nature of the reaction 2. the concentration of the reactants 3. the temperature of the reaction 4. the presence of a catalyst The change in concentration and temperature affect the rate of a chemical reaction by influencing the collisions among the molecules. An increase in concentration in- creases the number of molecular collisions. An increase in temperature increases the rate because the molecules move faster, increasing the kinetic energy and collisions are more frequent. A catalyst is a compound that affects the rate of chemical reactions by lowering the activation energy, Ea. With a lower activation energy, less energy is needed to overcome the energy barrier to form the product(s). A catalyst takes part in the reaction, but is not consumed. It always returns to its original composition at the end of the reaction. Catalysts are defined as homogeneous or heterogeneous where the catalyst is either in the same phase as the reactants or in a different phase, respectively. Reaction Rate and Rate Laws Consider a chemical system in which there are two re- actants, A and B, with stoichiometric coefficients a and b, forming c moles of product C. aA  bB → cC The rate of the reaction can be defined as the change in concentration of the product as a function of the change in time: [ ] rate t C  D D 69  E x p e r i m e n t 1 3 • The Decomposition of Hydrogen Peroxide  Expt. 13 so the slope could be used to calculate the activation energy. With only two data points to use for the tem- perature comparison (only Parts I and IV have the same concentrations of reactants), graphing is a little excessive so instead we can use another form of the Arrhenius equation, shown below. ln R E T 1 T 1 k k   a 1 21 2 = G where k1 and k2 are the rate constants at two different temperatures, T1 and T2 (in Kelvin). The value of the activation energy, Ea, is in joules per mole. Pay attention to the units and value of the gas constant, R. This is not the same value of R used in the previous calculation. Tips for Procedure • Record exact volumes used in each step to one digit beyond the markings on the glassware. Volumes given in the table are approximate values. • Allow the solutions to sit before mixing for at least two minutes for their temperatures to equilibrate. • Be prepared to start collecting data immediately after the solutions are mixed. • The stopper may pop out of the test tube near the end of the experiment. This is to be expected as the pressure inside the test tube increases. Use the data up to the point of the pop when completing the data analysis. • If you don’t see an increase in pressure, repeat the trial after checking for leaks in your setup. • T2 should be approximately 10 degrees higher than T1. • Always use distilled water. • Look at the rate law problems in your textbook (example: Problem 13.70 in Chang, 10th edition on page 607). You should set up a similar table with your data. Suggested Materials MeasureNet gas pressure probe kit thermometer 3% hydrogen peroxide, (H2O2), solution 0.5 M potassium iodide, (KI), solution two 6″ test tubes pipets pipet bulb 400 mL beaker other glassware and equipment, as needed Procedure You must write your procedure and prepare your lab notebook to record data before coming to lab to do the experiment. Make sure that you will collect the data necessary to complete the data analysis questions. It’s better to have too much information and not need it, than to need something and not have it. Instructions on using graduated pipets can be found in Chapter 3. Steps for using the gas pressure probe with the MeasureNet workstation are in Chapter 4. Table 13.1. Approximate volumes for each trial. Trial Volume of 3.00% H2O2 (mL) Volume of 0.500 M KI (mL) Volume of Water (mL) Temperature File Name 1 4 1 0 T1 001 2 4 0.5 0.5 T1 002 3 2 1 2 T1 003 4 4 1 0 T2 004  70  E x p e r i m e n t 1 3 • The Decomposition of Hydrogen Peroxide Expt. 13 Data Analysis 1. What is the concentration of the 3% H2O2 solution in units of molarity? Assume the density of H2O2 is 1.00 g/mL. 2. Find the concentration for the H2O2 and KI solutions for each trial after the solutions and/or water are mixed together. 3. Determine the initial rate in torr/s of each reaction using a graph of pressure vs. time for each part. Look at the units of the data being graphed to determine what information is needed from the graph to de- termine this. 4. Convert the initial rate for each part to M/s using the ideal gas law. 5. Comparing the parts done at the same temperature, look at the data for the initial concentrations of the reactants (found in question 2) and the initial rates determined in question 4 to determine the order of the reaction with respect to each reactant. Write the general rate law for the reaction. 6. Determine the value of the rate constant, k, for each part using the general rate law equation from question 5. For the parts at the same temperature, report the value of k as an average with its standard deviation. 7. Use the Arrhenius equation to determine the activa- tion energy, Ea, for this reaction by comparing trials with the same concentrations of reactants at two different temperatures.