EXPERIMENT 4 Le Chatelier s Principle INTRODUCTION Le Chatelier s Principle states: When a stress is applied to a chemical system at equilibrium, the equilibrium concentrations will shift in a direction that reduces the effect of the stress. This experiment will investigate this principle in several systems by shifting the equilibrium concentrations of reactants and products by applying a stress to the equilibrium. The stresses that will be applied in this experiment will consist of the following: Increasing the concentration of a species in the reaction by adding it to the equilibrium mixture Decreasing the concentration of a species in the reaction by the addition of a species that chemically reacts with either a reactant or product in the equilibrium mixture. Adding or removing heat from the system by changing the temperature of the equilibrium mixture For each equilibrium reaction, you will first need to decide how to set it up. A list of reagents available for each reaction will be given. To set up the equilibrium reaction, each species in the equilibrium must be present. For example, you may be asked to prepare the following equilibrium reaction: Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) and you are given only the following reagents: 0.1M Mg(NO 3 ) 2, 1 M NaOH, 1 M HCl. In order to set up the equilibrium reaction above, you must add Mg(NO 3 ) 2 to obtain Mg 2+ ions and NaOH to obtain OH - ions. The ions will react, producing Mg(OH) 2 solid, so all components in the equilibrium reaction will be present. The NO 3 - and Na + ions are spectator ions and do not participate in the reaction. If you are then asked to, first, shift the reaction to the left (towards the solid) and to, second, shift the reaction to the right (towards the aqueous ions), you will need to determine what of the listed reagents must be added to accomplish each. In this example, to shift the reaction to the left, the 1M NaOH can be added. To shift the reaction to the right, the 1M HCl can be added. You will record the observations verifying that each desired shift has occurred. NET IONIC EQUATIONS You will also be asked to write a net ionic equation (NIE) that demonstrates why the reactions shifted the way they did. NIEs show only the ions that participate in the reaction. In this lab the NIEs you write will follow one of two methods. Method 1 is used if the reagent added to the equilibrium system is already involved in the equilibrium, and Method 2 is used if the reagent added to the equilibrium system is not already involved in the equilibrium. 31
METHOD 1 If the stress applied to the equilibrium system is an increase in concentration of a reactant or product, then the NIE is given by the original equilibrium equation, with (1) the reaction arrow pointing in the direction of the spontaneous reaction, and (2) the added reagent underlined. For the stress that shifted the equilibrium to the left: Equilibrium System: Shift Caused by adding: Reason for shift: NIE: METHOD 2 Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) NaOH (aq) Adding sodium hydroxide increases the hydroxide concentration. This increases the rate of the reverse reaction, causing the reverse reaction to be spontaneous, (producing more solid magnesium hydroxide) until equilibrium is reestablished. Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) If the stress applied to the equilibrium system is caused by the addition of a species not in the equilibrium, then the NIE is given by the chemical reaction between the added species and either a reactant or product, lowering the concentration of that reactant or product. The equilibrium concentrations will then shift accordingly. To write the overall NIE several steps are involved. For the stress that shifted the equilibrium to the right: Equilibrium System: Shift Caused by adding: Reason for the shift: NIE: Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) HCl (aq) Adding hydrochloric acid produces hydronium ions, which react with the hydroxide ions, decreasing their concentration. This decreases the rate of the reverse reaction, causing the forward reaction to be spontaneous, (producing more of the dissolved ions) until equilibrium is reestablished. Step 1: Start with the original net ionic equilibrium reaction Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) Step 2: The H 3 O + ions from the strong acid react with the OH - ions in the original equilibrium system to form water. Add an equal number of H 3 O + ions to both sides of the equation, but enough to react away all of the OH - ions 2H 3 O + (aq) + Mg(OH) 2 (s) Mg 2+ (aq) + 2OH - (aq) + 2H 3 O + (aq) Step 3: combine the hydronium ions and hydroxide ions to form water 2H 3 O + (aq) + Mg(OH) 2 (s) Mg 2+ (aq) + 2H 2 O (l) Step 4: Cancel out any spectator ions (if there are any) and write the equation (1) with the reaction arrow pointing in the direction of the spontaneous reaction, and (2) with the added reagent underlined. 2H 3 O + (aq) + Mg(OH) 2 (s) Mg 2+ (aq) + 2H 2 O (l) PROCEDURE 1. Students will work individually for this experiment. Except for the laboratory handout, remove all books, purses, and such items from the laboratory bench top, and placed them in the storage area by the front door. For laboratory experiments you should be wearing closed-toe shoes. 32
2. Do not add more of a reagent than the minimum needed to produce a desired effect. After each reaction, discard all solutions into a waste beaker on your lab bench. At the end of the experiment, all waste should be disposed of in the waste bottle in the fume hood. REACTION 1 - DEMONSTRATION 3. Place 10 drops of 0.1 M cobalt (II) nitrate solution and 8 drops of 12 M hydrochloric acid into a small test tube. Mix with a vortex mixer. This will establish the following equilibrium showing the color of pink: Co 2+ (aq) + 4Cl - 2- (aq) CoCl 4 (aq) pink blue 4. Using one of the following reagents: 12 M HCI, 0.1 M AgNO 3 Add enough drops to force the equilibrium to shift to the right. Mix with a vortex mixer. Record in your Data Table the reagent used, as well as your observations verifying the shift to the right. 5. Using one of the following reagents: 0.1 M Co(NO 3 ) 2, 12 M HCI, 0.1 M AgNO 3 force the equilibrium to shift to the left. Record in your Data Table the reagent used, as well as your observations verifying the shifted to the left. 6. Prepare a new equilibrium solution as in step 1, and split it into 2 test tubes. Place one test tube in ice water and one in hot water. Wait 5 minutes. Record in your Data Table the direction of the equilibrium shift for each, as well as your observations verifying the direction of the equilibrium shift REACTION 2 7. Place 20 drops of 0.5 M potassium chromate solution in a small test tube, and add 1 drop of 0.1 M hydrochloric acid. This should establish the following equilibrium showing the color of yellow: 2-2CrO 4 (aq) + 2H + (aq) Cr 2 O 2-7 (aq) + H 2 O (l) yellow orange 8. Using one of the following reagents: 6 M HCI, 6 M NaOH Add enough drops to force the equilibrium to shift to the right. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the right. 9. Using one of the following reagents: 0.5 M K 2 CrO 4, 6 M HCI, 6 M NaOH force the equilibrium to shift to the left. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the left. REACTION 3 10. Place 10 drops of 0.1 M barium chloride solution and 1 drop of 0.1 M hydrochloric acid in a small test tube, and add 10 drops of 0.1 M ammonium oxalate. This should establish the following equilibrium showing the showing the white solid: BaC 2 O 4 (s) + H + (aq) Ba 2+ (aq) + HC 2 O 4 - (aq) white solid 11. Using one of the reagents: 0.1 M BaCl 2, 1.0 M HCl, 1 M NH 3 force the equilibrium to shift to the right. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the right. 33
12. Using one of the reagents: 1.0 M HCl, 1 M NH 3 force the equilibrium to shift to the left. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the left. REACTION 4 13. Place 20 drops of 0.01 M iron (III) chloride and 1 drop of 0.01 M ammonium thiocyanate in a small test tube. This should establish the following equilibrium showing the showing an orange color: Fe 3+ (aq) + SCN - (aq) FeSCN 2+ (aq) yellow red 14. Using one of the reagents: 0.01 M NH 4 SCN, 0.01 M HgCl 2 force the equilibrium to shift to the right. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the right. 15. Using one of the reagents: 0.01 M FeCl 3, 0.01 M NH 4 SCN, 0.01 M HgCl 2 force the equilibrium to shift to the left. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the left. REACTION 5 16. Place 10 drops of 0.1 M silver nitrate and 1 drop of 1 M ammonia in a small test tube, and add 10 drops of 0.1 M sodium chloride. This should establish the following equilibrium showing the showing the white solid: AgCl (s) + 2NH 3 (aq) Ag(NH 3 ) 2 + (aq) + Cl - (aq) white solid 17. Using one of the reagents: 0.1 M AgNO 3, 0.1 M NaCl, 1 M NH 3 force the equilibrium to shift to the right. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the right. 18. Using one of the reagents: 0.1 M NaCl, 1 M NH 3 force the equilibrium to shift to the left. Record in your Data Table the reagent used, as well as your observations verifying that the equilibrium shifted to the left. 19. Prepare a new equilibrium solution in a different test tube with 10 drops of 0.1 M silver nitrate, 10 drops of 0.1 M sodium chloride, and enough drops of 1 M ammonia until the precipitate completely dissolves. 20. Add drops of 0.1 M silver nitrate to force the equilibrium to the left. Record in your Data Table your observations verifying that the equilibrium shifted to the left. 21. All excess solutions should be disposed of in the Chem 1B Waste Container in the fume hood. 22. Clean and wipe dry your laboratory work area and all apparatus. When you have completed your lab report have the instructor inspect your working area. Once your working area has been checked your lab report can then be turned in to the instructor. 34
EXPERIMENT 4 LAB REPORT Name: Student Lab Score: Date/Lab Start Time: _ Lab Station Number: DATA TABLE REACTION 1 Observation for Step 3 Initial Equilibrium Reagent for Step 4 Right Shift Observation for Step 4 Right Shift Equilibrium Reagent for Step 5 Left Shift Observation for Step 5 Left Shift Equilibrium Observation for Step 6 Cold Temperature Observation for Step 6 Hot Temperature REACTION 2 Observation for Step 7 Initial Equilibrium Reagent for Step 8 Right Shift Observation for Step 8 Right Shift Equilibrium Reagent for Step 9 Left Shift Observation for Step 9 Left Shift Equilibrium 35
REACTION 3 Observation for Step 10 Initial Equilibrium Reagent for Step 11 Right Shift Observation for Step 11 Right Shift Equilibrium Reagent for Step 12 Left Shift Observation for Step 12 Left Shift Equilibrium REACTION 4 Observation for Step 13 Initial Equilibrium Reagent for Step 14 Right Shift Observation for Step 14 Right Shift Equilibrium Reagent for Step 15 Left Shift Observation for Step 15 Left Shift Equilibrium REACTION 5 Observation for Step 16 Initial Equilibrium Reagent for Step 17 Right Shift Observation for Step 17 Right Shift Equilibrium Reagent for Step 18 Left Shift Observation for Step 18 Left Shift Equilibrium Observation for Step 20 Left Shift Equilibrium with AgNO 3 36
QUESTIONS 1. For Reaction 1, Step 4, explain why the equilibrium shifted to the right. 2. For Reaction 1, Step 4, write the NIE for the right shift. 3. For Reaction 1, Step 5, explain why the equilibrium shifted to the left. 4. For Reaction 1, Step 5, write the NIE for the left shift. 5. For Reaction 1, rewrite the equilibrium by adding energy to the proper side of the equation. 6. Is Reaction 1 exothermic or endothermic? Explain based upon Le Chatelier s Principle. 7. For Reaction 2, Step 8, explain why the equilibrium shifted to the right. 37
8. For Reaction 2, Step 8, write the NIE for the right shift. 9. For Reaction 2, Step 9, explain why the equilibrium shifted to the left. 10. For Reaction 2, Step 9, write the NIE for the left shift. 11. For Reaction 3, Step 11, explain why the equilibrium shifted to the right. 12. For Reaction 3, Step 11, write the NIE for the right shift. 13. For Reaction 3, Step 12, explain why the equilibrium shifted to the left. 14. For Reaction 3, Step 12, write the NIE for the left shift. 38
15. For Reaction 4, Step 14, explain why the equilibrium shifted to the right. 16. For Reaction 4, Step 14, write the NIE for the right shift. 17. For Reaction 4, Step 15, explain why the equilibrium shifted to the left. 18. For Reaction 4, Step 15, write the NIE for the left shift. _ 19. For Reaction 5, Step 17, explain why the equilibrium shifted to the right. 20. For Reaction 5, Step 17, write the NIE for the right shift. 21. For Reaction 5, Step 18, explain why the equilibrium shifted to the left. 39
22. For Reaction 5, Step 18, write the NIE for the left shift. 23. For Reaction 5, Step 20, explain why the equilibrium shifted to the left. 24. For Reaction 5, Step 20, write the NIE for the left shift. 25. The following exothermic reaction is at equilibrium in a closed container: 2N 2 O (g) 2N 2 (g) + O 2 (g) (a) What will be the effect on the equilibrium position with each of the following disturbances: shift left (L), shift right (R), or no effect (NE)? The addition of a dinitrogen monoxide gas to the reaction container at constant volume The addition of oxygen gas to the reaction container at constant volume An increase in temperature A decrease in the volume of the reaction container Adding argon gas to the reaction container at constant volume (b) What will be the effect on the equilibrium concentration of N 2 (g) with each of the following disturbances: increases (I), decreases (D), or no effect (NE)? The addition of dinitrogen monoxide gas to the reaction container The addition of oxygen gas to the reaction container at constant volume The addition of nitrogen gas to the reaction container at constant volume An increase in temperature The addition of a catalyst to the reaction container 40