7# -NR,LXLXLXLXLXLfJLLL LLLL4M(xLM MM*NLXMM MMNDMMMMMMChemistry 142 LE CHATELIER'S PRINCIPLE Purpose The purpose of this lab is to study the shift in the equilibrium position of chemical reactions when an external stress is applied. The shift can be predicted by using Le Chatelier's Principle. Reading in Preparation for this Lab J. Olmstead III, G. M. Williams, Chemistry, The Molecular Science, Mosby, St. Louis, MO, 1994, pp. 745-748, 760. J. C. Kotz, K. F. Purcell, Chemistry & Chemical Reactivity, Second Edition, Saunders College Publishing, Philadelphia, PA, 1991, pp. 568-571 and 679-684. R. H. Petrucci, General Chemistry, Macmillan Publishing Company, New York, NY, 1989, pp. 581-585. Introduction No chemical reaction goes to completion. When the reaction stops, some of the reactants will remain. For example, even though we write 2 CO2(g) 2 CO(g) + O2(g) (1) as if it went to completion, at 2000 K only 2% of the CO2 decomposes. A chemical reaction reaches equilibrium when the concentrations of the reactants and products no longer change. The position of the equilibrium describes the relative amounts of reactants and products that remain at the end of a reaction. The position of the equilibrium for reaction (1) is said to lie with the reactants, or to the left, because at equilibrium very little of the reactants has reacted. On the other hand, in the reaction H2(g) + 1/2 O2(g)  H2O(g) (2) the equilibrium position lies very far to the right since only very small amounts of H2 and O2 remain after the reaction reaches equilibrium. Since we often wish to maximize the yield from a reaction, it is vital to determine the kind of control we can exercise over the position of the equilibrium. The equilibrium position of a reaction may shift if an external stress is applied. The stress may be in the form of a change in temperature, pressure, or the concentration of one of the reactants or products. For example, consider a flask with an equilibrium mixture of CO2, CO and O2, as in reaction (1). If a small amount of CO is then injected into the flask, we will find that the concentration of CO2 increases. Here the stress is the increase in concentration of CO. The system responds by reacting some of the added CO with O2 to yield an increased amount of CO2. That is, the position of equilibrium shifts to the left, yielding more reactant and less CO. A general statement of our observation is given by Le Chatelier's Principle: If a system at equilibrium is subjected to a stress the system will react in a way that tends to relieve the stress. Reaction (1) will also shift with changes in pressure. Starting with reaction (1) at equilibrium, an increase in pressure will cause the position of equilibrium to shift to the side of the reaction with the smaller number of moles of gas. That is, by shifting the equilibrium position to the left, the reaction will decrease the number of moles of gas, thereby decreasing the pressure in the flask. In so doing, some of the applied stress is relieved. The shift in equilibrium position of a reaction with changes in temperature will depend on its heat effect (enthalpy change). The decomposition of CO2, reaction (1), is endothermic in the forward direction. If the temperature of the reaction is increased, the reaction will shift in the forward direction to try to minimize the temperature increase. That is, the equilibrium will shift to the right for an increase in temperature. The formation of ammonia, N2(g) + 3H2(g)  2NH3(g) (3) on the other hand, is exothermic, so that for an increase in temperature, the reaction will shift in the reverse (endothermic) direction. That is, for an increase in temperature the equilibrium position shifts to the left. In summary, supplying heat always favors the endothermic process. In the following two experiments you will explore further applications of Le Chatelier's Principle. EXPERIMENT 1: Factors Influencing the Iron-Thiocyanate Equilibrium1 When ammonium thiocyanate, NH4NCS, is mixed with iron(III) nitrate, Fe(NO3)3, in solution, an equilibrium mixture of Fe3+, NCS-, and the complex ion FeNCS2+ is formed: Fe3+ + NCS-  FeNCS2+ yellow colorless red The solution also contains the ions NH4+ and NO3-, but these are merely spectator ions and do not participate actively in this reaction. The relative amounts of the various ions participating in the reaction can be judged from the color of the solution since in neutral or slightly acidic solutions, Fe3+ is light yellow, NCS- is colorless, and FeNCS2+ is blood red. If the solution is initially reddish, and the equilibrium shifts to the right (more FeNCS2+), the solution becomes darker red, while if the equilibrium shifts to the left (less FeNCS2+), the solution becomes lighter red or perhaps straw-yellow. In this experiment you will carry out various processes which will apply an external stress to this equilibrium. You should be able to explain the resulting shift of the equilibrium to the right or to the left in terms of Le Chatelier's Principle. Procedure Add one drop each of 1M Fe(NO3)3 and 1M NH4NCS to 25 ml of distilled water. Mix well. Add a few drops of this solution to each of seven wells of a spot plate. One well will serve as a color standard against which to judge color changes in the other wells. For each of the operations described, information is provided as necessary regarding the manner in which one or more of the chemical species is affected. Use your observation of the color change of the solution and the information provided to explain how the results are consistent with Le Chatelier's Principle. Use a different well for each of the operations described. As an example, if you added a drop of concentrated HCl to the solution, the blood red color would lighten or perhaps even disappear altogether. This indicates that the FeNCS2+ concentration has decreased. To explain this result, it is necessary to know that in the presence of a large excess of Cl-, Fe3+ forms complex ions: Fe3+ + 6 Cl-  FeClA(3-,6) This change reduces the Fe3+ concentration, so in accord with Le Chatelier's Principle, some FeNCS2+ dissociates to replace some of the Fe3+ removed by reaction with Cl-. Operations to Introduce an External Stress (a) Add one drop of 1 M Fe(NO3)3 and mix well. (b) Add one drop of 1 M NH4NCS and mix well. (c) Add one drop of 0.1M SnCl2 and mix well. Tin(II) ions reduce iron(III) ions to iron(II) ions: Sn2+ + 2 Fe3+ Sn4+ + 2 Fe2+ (d) Add one drop of 0.1 M AgNO3 and mix well. Silver ions react with thiocyanate ions to give a white precipitate of silver thiocyanate: Ag+ + NCS-  AgNCS(s) (e) Add one drop of 0.1 M Na2HPO4 and mix well. Hydrogen phosphate ions form a complex ion with iron (III) ions: Fe3+ + HPOA(2-,4 )  FeHPOA(+,4 ) (f) Add one drop of 1 M NH3 and mix well. Any base will form a precipitate or a colloidal suspension of iron (III) hydroxide when mixed with iron (III) ions: NH3 + H2O  NHA(+,4) + OH- Fe3+ + 3 OH-  Fe(OH)3(s) (g) Pour about 4-5 mL of the iron-thiocyanate solution into two test tubes. Set one tube aside as a color standard against which to judge color changes in the other tube. Gently warm the other tube in a hot water bath on a hot plate. Do not boil the solution. Considering the results and Le Chatelier's Principle, must the forward reaction be exothermic (release energy) or endothermic (consume heat energy)? h) Cool the tube in a beaker of ice water. Do these results indicate the forward reaction is exothermic or endothermic? EXPERIMENT 2: Hydrolysis of Sodium Acetate A buffer is a solution of a weak acid and its conjugate base. Buffer systems are common examples of reversible reactions in a state of equilibrium. Because of the reversible nature of this equilibrium, the position of equilibrium can be shifted in the appropriate direction to consume or release hydrogen ions as needed: A(CH3CO2H,(acetic acid)) A(CH3CO2- ,(acetate ion)) + H+ If a salt of a weak acid such as acetic acid is dissolved in water, the salt can undergo hydrolysis to set up a similar equilibrium state: CH3 COA(-,2) + H2O  CH3CO2H + OH- In this experiment, you will study the effect of temperature on the latter equilibrium. Procedure To 8-10 mL of 1 M sodium acetate, NaCH3CO2, in a medium size test tube, add 30 drops of phenolphthalein indicator. Mix well and describe the color. Transfer about half of this solution to a second test tube and set it aside as a color control. Now heat the solution in the original test tube almost to boiling in a hot water bath on a hot plate. What happens to the color of the solution? From what you know about the pH dependence of the colors of phenolphthalein (pink at pH> 9), in which direction did the equilibrium shift? Now cool the test tube first in air, then in an ice-water bath. What happens to the color? In which direction did the equilibrium shift this time? To convince yourself that this reaction is truly reversible, heat the test tube and then cool it again. Were the results the same? Considering Le Chatelier's Principle, is the forward reaction exothermic or endothermic? Report In your lab book give a short explanation of your observations for each reaction you performed, using Le Chatelier's Principle. Turn in a handwritten or photocopy of your observations and explanations. For example... Stress: Observation: Explanation: a) added 1 drop Sol'n became The equilibrium shifted to 1 M NH4NO3 yellower/or/redder the RIGHT/LEFT because... Footnote 1. These experiments are taken from a laboratory manual by James P. Birk at Arizona State University. Chemistry 142 LE CHATELIER'S PRINCIPLE PRELABORATORY ASSIGNMENT Name 1. Many home water softeners contain zeolites, which can exchange sodium ions for calcium and other ions in solution, in a reversible process: Na2Al2Si4O12(s) + Ca2+ (aq)  CaAl2Si4O12(s) + 2 Na+(aq) The zeolite can be recharged by washing with a concentrated solution of NaCl. Explain how the recharging process works in terms of Le Chatelier's Principle. 2. The proper maintenance of a swimming pool requires control of both the chlorine level and the pH. Chlorine is usually added in the form of a hypochlorite salt which forms chlorine by the reversible reaction: OCl-(aq) + Cl-(aq) + 2H+(aq)  Cl2(aq) + H2O(l) The chloride is generally present naturally in the water and is also formed by photochemical decomposition of chlorine. Predict the effect on the concentration of Cl2 if more chloride is added to the water. Predict the effect if the pH is raised. Explain your answers. 3. Champagne is poured cold into chilled glasses because it would go "flat" [lose dissolved CO2] if poured into a warm glass. The temperature dependence of the position of equilibrium is the same for all processes of the type: gas + liquid  solution. On the basis of the above observation and Le Chatelier's Principle, is the dissolving of a gas in a liquid an exothermic or an endothermic process? 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