1-50 ml beaker stirring rod 1-10 ml beaker 0.10 M NaOH (1 ml) calibrated plastic dropper (1 ml) 50 ml dispensing burette (for Crystal Violet)

Transcription

1 Exercise 2 Page 1 Illinois Central College CHEMISTRY 132 Name: Kinetics, Part II. Equipment Objectives ml beaker stirring rod 1-10 ml beaker 0.10 M NaOH (1 ml) calibrated plastic dropper (1 ml) 1.5x10-5 M Crystal Violet 50 ml dispensing burette (for Crystal Violet) The objectives of this experiment are to: - study the rate of reaction of crystal violet with NaOH using the LabWorks interface colorimeter. - determine the order of reaction with respect to NaOH. - calculate the rate constant for the reaction at room temperature. Background As shown in last week's experiment, the rate of the reaction of crystal violet with NaOH is given by the generalized rate expression: Rate = k[oh -1 ] x [CV] y (1) In Equation (1), k is the rate constant for the reaction, CV is an abbreviation for crystal violet, C 25 H 30 N +1, x is the order of reaction with respect to OH -1, and y is the order of reaction with respect to CV. In that experiment, the [OH -1 ] was always much greater than [CV]. Thus the change in [OH -1 ] had a negligible effect on the initial [OH -1 ]. For this reason, [OH -1 ] x was treated as a constant and Equation (1) was rewritten: Rate = k' [CV] y where k' = k[oh -1 ] x. k' was termed a pseudo rate constant. So the value of this pseudo rate constant k' depends on the concentration of OH -1. In this part of the experiment we will re-run the reaction with a higher concentration of OH -1 in order to determine "x", the order of the reaction with respect to OH -1 and k, the true rate constant. (1b)

2 Exercise 2 Page 2 Connecting the Colorimeter Locate the Logger Pro icon and double-click on it, or use the Start menu. Connect the Vernier Colorimeter to the GoLink USB interface and connect the GoLink to the USB input on your computer. From the Menu Bar select File/Open and click on the folder Chemistry with Computers. Open the file Xtalviol.cmbl. You should now see the window displayed here. Use the arrow buttons on the colorimeter to select the 565 nm LED. Select a single cuvette to use for both your blank and your samples for this experiment. Safety Precautions Crystal violet solutions may cause skin and eye irritation. Sodium hydroxide solutions are caustic and will cause skin burns if not immediately washed with copious amounts of water. Safety goggles must be worn at all times. As usual, wash hands with soap and water before leaving the lab. PROCEDURE - Part II Recall that the k' obtained in last week's experiment is a pseudo rate constant, whose value depends upon the concentration of OH -1, i.e. k' = k[oh -1 ] x. In Part I of the experiment, 9.00 ml of M crystal violet and 1.0 ml of 0.05 M NaOH were combined to form the reaction mixture. A second kinetic run will now be made in exactly the same way except that the concentration of NaOH will now be doubled to 0.10 M. 1. Fill the cuvette with distilled water to serve as a "blank". The blank contains all the constituents used in the analysis except the substance to be measured. We can assume then that the difference in the color between the blank and the sample is due only to the substance to be measured. Distilled water is the reference blank for this experiment. 2. Insert the cuvette containing the distilled water into the opening of the colorimeter. Note that the cuvette is "ribbed" on two sides. IMPORTANT: Be certain that the light path is passing through the CLEAR sides of the cuvette facing the arrow at the top of the cuvette slot. Close the lid of the colorimeter (to keep out stray light) and press the "CAL" button on the colorimeter to calibrate it. Release the CAL button when the red LED begins to flash. When the LED stops flashing, the calibration is complete and your unit is ready to collect data.

3 Exercise 2 Page 3 3. Empty the solution cell and dry it thoroughly inside and out. 4. Using the burette provided, dispense 9.00 ml of M crystal violet solution into a clean, dry 50 ml beaker. 5. Using the calibrated plastic dropper provided, add 1.0 ml of 0.05 M NaOH to the CV solution as rapidly as possible without splashing. At the same instant, click the Collect button on the Toolbar to start the timer. (you should see the statement "waiting for data" displayed on your graph). Hustle, because you only have 60 seconds before the first data point is collected. 6. Thoroughly mix the CV/NaOH solution with a stirring rod and then fill the cuvette 3/4 full. Position the cuvette in the colorimeter in exactly the same manner as was used for pure water. Shut the lid, and wait. At the 1 minute mark, the colorimeter should take its first reading. If so, you have 30 minutes to kill. 7. The program will take absorbance readings at one minute intervals for a period of 30 minutes and then automatically stop. If there is a need to stop data collection prior to the end of 30 minutes, click the Stop button on the Toolbar and the program will terminate. Data Analysis - Part II 1. In our last experiment, we determined that the this reaction was first order with respect to the crystal violet. So there is no need to plot the zero order or second order data. We will plot only the first order -ln(absorbance) vs time. 2. Select Data from the Menu Bar and click on New Calculated Column. In the New Calculated Column dialogue box enter -ln(absorbance) for the name (-ln(abs) for the short name) and leave the units blank. For the equation in our dialogue box, simply type -ln("absorbance"). Click "Done". 3. A plot of Absorbance versus time should currently be displayed. Simply click on the "Absorbance" label on the x-axis of your graph. A box containg column choices should appear. Choose "-ln(absorbance)" to be plotted. Click <Ctrl>J to autoscale your graph. 4. Select Analyze from the Menu Bar, and choose Linear Fit. (Or click on the Linear Fit icon found on the Toolbar.) Prepare and print a carefully labeled linear regression graph for this plot. The slope of the straight line displayed in the Linear Fit dialogue box is the best value of k'. Record this value with proper units and to correct number of significant figures on the Report Sheet.. 5. From the ratio of the k' value obtained today (using 0.10 M NaOH) and the k' value from last week (using 0.05 M NaOH), determine the order of reaction with respect to OH -1 (the value of "x"). Clearly indicate your reasoning in evaluating x.

4 Exercise 2 Page 4 6. Calculate the value of the true rate constant k using each of the k' values. In the calculations, the concentrations of OH -1 will have to be adjusted to account for the dilutions which occurred when the NaOH and crystal violet solutions were mixed. Finally, average the two k values obtained. Again, be sure to watch significant figures and use proper units. 7. Select File/Print Graph to obtain a print out of your regression plot. Also select File/Print Data Table to obtain the printout of your data. Attach your graph and data and to your Report Sheet.

5 Exercise 2 Page 5 Chemical Equilibria Equipment Objectives ml beaker solid ammonium chloride (NH 4 Cl) 4-test tubes solid CoCl. 2 6H 2 O Iron (III) chloride (FeCl 3 ) solution concentrated HCl Ammonium thiocyanate (NH 4 CNS) solution The objectives of this experiment are to: - illustrate systems in equilibrium. - demonstrate Le Chatelier's Principle. Background When compounds react, they sometimes form a mixture of products and (unreacted) reactants, in a dynamic equilibrium. Much like water in a U-shape tube, the water constantly mixes back and forth through the lower portion of the tube, as if the forward and reverse "reactions" were occurring at the same rate. This makes the system appear to be "static" (or, stationary), when in reality, it is "dynamic" (in constant motion). Reactants Products Once the water seeks the proper level, that is, reaches equilibrium it continues to migrate back and forth across the "arrow" but at the same rate making it appear to be static. For example, The Haber process for producing ammonia from nitrogen and hydrogen gas does not go to completion, but instead reaches an equilibrium state where all three participants are present N 2(g) + 3 H 2(g) 2 NH 3(g) Chemical Equilibrium is the state reached by a reaction mixture when the rates of the forward and reverse reactions have become equal.

6 Exercise 2 Page 6 Le Chatelier's Principle Obtaining the maximum amount of product from a reaction depends on the proper selection of reaction conditions. By changing the reaction conditions, one can increase or decrease the amount of products. Le Chatelier's Principle states that an equilibrium will shift under conditions of "stress" in such a way as to remove that stress. This applied "stress can be achieved in several ways. 1. Changing the concentrations by removing products or adding more reactants to the reaction vessel. 2. Changing the partial pressure of gaseous reactants and products by changing the volume of the reaction vessel. 3. Changing the temperature. Removing Products or Adding Reactants. Let's refer back to the illustration of the U-tube. It's a simple concept to see that if we were to remove products (analogous to dipping water out of the right side of the tube) the reaction would shift to the right until equilibrium was re-established. Likewise, if more reactant is added (analogous to pouring more water in the left side of the tube) the reaction would also shift to the right..again until equilibrium is re-established. Conversely, if more product were to be added, or reactants removed, the reaction would shift to the left until equilibrium is re-established. Effects of Pressure Change A pressure change obtained by changing the volume of the reaction vessel can affect the yield of products in a gaseous reaction only if the reaction involves a change in the total moles of gas present. If the products in a gaseous reaction contain fewer moles of gas than the reactants, then it is logical that they would require less space. So reducing the volume of the reaction vessel would favor the products. Conversely, if the reactants require less volume (that is, fewer moles of gas on the left side of the arrows) then decreasing the volume of the reaction vessel would shift the equilibrium to the left. So literally, "squeezing" the reaction will cause a shift in the equilibrium toward the fewer moles of gas. Its a simple step, then, to see that reducing the pressure in the reaction vessel by increasing its volume would have the opposite effect. In the event that the number of moles of gaseous products equals the number of moles of gaseous reactants...vessel volume will have no effect on the position of the equilibrium.

7 Exercise 2 Page 7 Effect of Temperature Change Temperature has a significant effect on most reactions. Reaction rates generally increase with an increase in temperature, consequently, equilibrium is established sooner. In addition, the "position" of the equilibrium shifts. If we look at heat as if it were a product in exothermic reactions...and a reactant in endothermic reactions, we can see that increasing the temperature of a reaction is analogous to adding more product (in the case of exothermic) or adding more reactant (in the case of endothermic). This ultimately has the same effect as if this added heat were a physical substance. For example, consider the following generic exothermic reaction. a A + b B c C + d D + HEAT ( H is negative) Increasing the temperature would be analogous to adding more product, consequently the equilibrium would now shift to the left. For an endothermic reaction, the opposite is true. HEAT + a A + b B c C + d D ( H is positive) Increasing the temperature in this case would cause the equilibrium position to shift toward the products. Effect of a Catalyst A catalyst is a substance that increases the rate of a reaction but is not consumed by it. It is important to understand that a catalyst has no effect on the equilibrium composition of a reaction mixture. A catalyst merely speeds up the attainment of equilibrium. Although a catalyst cannot affect the composition at true equilibrium, in some cases it can affect the product in a reaction because it affects the rate of one reaction out of several possible reactions.

8 Exercise 2 Page 8 Procedure Pour 1 ml of iron(iii) chloride, FeCl 3, and 1 ml of ammonium thiocyanate, NH 4 CNS, into a beaker containing 40 ml of water. The equilibrium established is illustrated in the equation; FeCl 3(aq) + 3 NH 4 CNS (aq) 3 NH 4 Cl (aq) + Fe(CNS) 3(aq) The color is due to iron (III) thiocyanate or, more precisely, to the formation of complex ions containing iron(iii) and thiocyanate ions. Consequently, a change in the concentration of this ion can be noted by the change in the depth of color. The color of the solution should be light enough that any deepening can be detected. If the solution is too dark, dilute it with more water. Fill each of four test tubes about one-third full with this solution. Add 2 ml of iron(iii)chloride to the first tube, 2 ml of the ammonium thiocyanate solution to the second, and 3 grams of solid ammonium chloride to the third (shake until the solid dissolves). Keep the fourth tube for comparison. Based on your observations, answer the following questions on your report sheet. 1. What color change occurred in the first tube? Explain 2. What color change occurred in the second tube? Explain. 3. What color change occurred in the third tube? Explain. Aqueous solutions of simple cobalt(ii) salts are pink due to the presence of hydrated Co +2 ions. In the presence of excess chloride ions, the solution changes color. The equilibrium that is established is illustrated below. [Co(H 2 O) 6 ] +2 (aq) + 4 Cl-1 (aq) [Co(H 2 O) 2 Cl 4 ] (aq) + 4 H 2 O (l) Put about 1 gram of CoCl 2. 6H 2 O into a test tube, add 4 ml of distilled water, and shake to dissolve the solid. 1. What is the color of the solution? Add concentrated HCl dropwise until the color just changes. 2. What is the color of the solution now? Place the test tube in a hot water bath for about ten minutes. 3. What happens? Based on this observation, do you believe the forward reaction is exothermic or endothermic?

9 Exercise 2 Page 9 Illinois Central College CHEMISTRY 132 Name: REPORT SHEET A Kinetic Study: Part II SHOW YOUR WORK 1. What is the value of k' when using 0.10 M NaOH? Include units and use appropriate significant figures. 2. Determine the ratio of the two k' values. Round to the nearest integer. Use this ratio to determine "x", (the order of the reaction with respect to [OH -1 ]). Clearly state your reasoning. 3. Attach a printout of your labeled graph and data points to this report sheet. 4. Calculate k using each k' value. Adjust [OH -1 ] for dilution. 5. calculate k average

10 Exercise 2 Page 10 Chemical Equilibria FeCl 3(aq) + 3 NH 4 CNS (aq) 3 NH 4 Cl (aq) + Fe(CNS) 3(aq) 1. What color change occurred in the first tube? Explain 2. What color change occurred in the second tube? Explain. 3. What color change occurred in the third tube? Explain. [Co(H 2 O) 6 ] +2 (aq) + Cl-1 (aq) [Co(H 2 O) 2 Cl 4 ] (aq) + 4 H 2 O (l) 1. What is the color of the solution? 2. What is the color of the solution now? 3. What happens? Based on this observation, do you believe the forward reaction is exothermic or endothermic?

11 Exercise 2 Page 11 Illinois Central College CHEMISTRY 132 Name: PRELAB: Exp. 2 Chemical Kinetics and Chemical Equilibria 1. Describe four factors that affect the rate of a reaction. 2. A reaction has the experimental Rate Law Rate = k[a] 2 [B] If the concentration of A is doubled and the concentration of B is halved, what happens to the reaction rate? 3. If you plot 1/[reactant] versus time and the plot is linear, what is the order of the reaction. 4. Define "chemical equilibrium". 5. State Le Chatelier's Principle. 6. Tell the direction of "shift" in the equilibrium, N 2(g) + 3 H 2(g) 2 NH 3(g) + 22 kj a) add N 2 caused by the following. b) Add NH 3 c) Remove H 2 d) Increase pressure e) Lower temperature

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