Purpose: To determine the experimental rate law for a chemical reaction.
1. Find the Volume of a Drop of Solution
For each of the three droppers you are using, perform the following procedure. Fill a graduated cylinder to approximately half full with tap water. Using your dropper, count the number of drops it takes to increase the volume of water in the cylinder by 1.00 mL. Repeat this three times for each dropper. You will use this data later to determine the volume contained in 1 drop of liquid.
2. Determine the Reaction Rate and Calculate the Rate Law
The table that follows shows the reactant quantities to be used in carrying out the reactions needed. Because we don't want the reaction to start until we are ready, be sure the I2 solution is the last solution added. It is important to use care in measuring out the solutions. Since the total volume of the solution is quite small, even one extra drop can cause a substantial change in concentrations. Try to hold the droppers vertically and be sure no air bubbles are introduced. Carry out each experiment number in the table three times and strive to get consistent results.
Place your dropper plate on a sheet of white paper. For
each trial, add all of your volumes of solution for each experiment into
1 well in your dropper plate. You must begin timing at the addition
of the first drop of I2 solution. After all of the I2 has been added,
stir the solution thoroughly with a toothpick. Record the amount
of time it takes for the yellow color to completely disappear. For
your first trial, fill one well in your dropper plate with distilled water
so that you can compare the color of your solution to it.
|Experiment No.||4.0 M Acetone||1.0 M HCl||Distilled H2O||.005 M I2|
|1||10 drops||10 drops||20 drops||10 drops|
|2||20 drops||10 drops||10 drops||10 drops|
|3||10 drops||20 drops||10 drops||10 drops|
|4||10 drops||10 drops||10 drops||20 drops|
1. You have enough information from your experiment to calculate the exponents of the experimental rate law for this reaction. To determine the rate of each reaction, divide the total number of moles of iodine added by the time it took for the iodine to disappear. In your calculations, find the nearest whole-number exponent for each reactant. In your determination of the rate constant, k, average the value of the rate constant that you get for each trial.
2. Compare your results to the actual values which your instructor will give you.
1. Sketch a curve of concentration versus time for the reactant of a reaction which is 1st order. What information does the slope of this curve give you at any point in the reaction?
2. Sketch a curve of concentration versus time for the reactant of a reaction which is 0 order.
3. In the reaction used for this lab, you should have discovered that the reaction was 0 order with respect to iodine, the reactant used to determine the rate. Judging from the graphs you sketched above, would it have been correct to measure the rate of the reaction as you did in this experiment if the reaction had been 1st order with respect to iodine? Explain.