Chapter 20 Experiment: LeChâtelier s Principle: Buffers OBJECTIVES: Study the effects of concentration and temperature changes on the position of equilibrium in a chemical system. Study the effect of strong acid and strong base addition on the ph of buffered and unbuffered systems. Observe the common-ion effect on dynamic equilibria. TECHNIQUES: Cleaning Glassware Disposing of Chemicals Handling Chemicals Hot Water Bath or Hot Plate Test Tubes for Small Volumes Well Plates for Small Volumes INTRODUCTION: Most chemical reaction do not produce a 100% yield of product, not because of experimental technique or design, but rather because of the chemical characteristics of the reaction. The reactants initially produce the expected products, but after a period of time the concentrations of the reactants and products stop changing. This apparent cessation of the reaction before a 100% yield is obtained implies that the chemical system has reached a state where the reactants combine to form the products at a rate equal to that of the products re-forming the reactants. This condition is a state of dynamic equilibrium, characteristic of all reversible reactions. For the reaction: 2 NO 2(g) N 2 O 4(g) + 58 kj chemical equilibrium is established when the rate at which two NO 2 molecules react equals the rate at which one N 2 O 4 molecule dissociates. If the concentration of one of the species in the equilibrium system changes, or if the temperature changes, the equilibrium tends to shift in a way that compensates for the change. For example, assuming the system represented by the above equation is in a state of dynamic equilibrium, if more NO 2 is added, the probability of its reaction with other NO 2 molecules increases, reducing the amount of NO 2. As a result, the reaction shifts to the right, producing more N 2 O 4 until equilibrium is re-established. A general statement governing all systems in a state of dynamic equilibrium follows: If an external stress (change in concentration, temperature, etc.) is applied to a system in a state of dynamic equilibrium, the equilibrium shifts in the direction that minimizes the effect of that stress. This is LeChatelier s principle, proposed by Henri Louis LeChatelier in 1888. Often the equilibrium concentrations of all species in the system can be determined. From this information, an equilibrium constant can be calculated; its magnitude indicates the relative position of the equilibrium. Two factors affecting equilibrium position are studied in this experiment: changes in concentration and changes in temperature.
Changes in Concentration: Metal-Ammonia Ions. Aqueous solutions of copper ions and nickel ions appear sky blue and green, respectively. The colors of the solutions change, however, in the presence of added ammonia, NH 3. Because the metal-ammonia bond is stronger than the metal-water bond, ammonia substitution occurs and the following equilibria shift right, forming the metal-ammonia complex ions: [Cu(H 2 O) 4 ] 2+ + 4NH 3 [Cu(NH 3 ) 4 ] 2+ + 4H 2 O (l) [Ni(H 2 O) 6 ] 2+ + 6NH 3 [Ni(NH 3 ) 6 ] 2+ + 6H 2 O (l) Addition of strong acid, H +, affects these equilibria by its reaction with ammonia (a base) on the left sides of the equations. NH 3 + H 1+ NH 4 1+ The ammonia is removed from the equilibria, and the reactions shift left to relieve the stress caused by the removal of the ammonia, re-forming the aqueous Cu 2+ (sky blue) and Ni 2+ (green) solutions. For copper ions, this equilibrium shift may be represented as Multiple Equilibria with the Silver Ion: [Cu(H 2 O) 4 ] 2+ + 4NH 3 [Cu(NH 3 ) 4 ] 2+ + 4H 2 O (l) 4H + 4NH 4 + Many salts are only slightly soluble in water. Silver ion, Ag 1+, forms a number of these salts. Several equilibria involving the relative solubilities of the silver salts of the carbonate, CO 3 2, chloride, Cl, iodide, I, and sulfide, S 2, anions are investigated in this experiment. Silver Carbonate Equilibrium. The first of the silver salt equilibria is a saturated solution of silver carbonate, Ag 2 CO 3, in dynamic equilibrium with its silver and carbonate ions in solution. Ag 2 CO 3(s) 2Ag 1+ + CO 3 2 Nitric acid, HNO 3, dissolves silver carbonate: H + ions react with (and remove) the CO 3 2 ions on the right; the system, in trying to replace the CO 3 2 ions, shifts to the right. The Ag 2 CO 3 dissolves, and carbonic acid, H 2 CO 3, forms. Ag 2 CO 3(s) 2Ag 1+ + CO 3 2 2H + H 2 CO 3 H 2 O (l) + CO 2(g) The carbonic acid, being unstable at room temperature and pressure, decomposes to water and carbon dioxide. The silver ion and nitrate ion (from HNO 3 ) remain in solution. Silver Chloride Equilibrium. Chloride ion precipitates silver ion as AgCl. Addition of chloride ion (from HCl) to the above solution, containing Ag 1+ and NO 3, causes the formation of a silver chloride, AgCl, precipitate, now in dynamic equilibrium with its Ag 1+ and Cl ions.
Ag 1+ + Cl AgCl (s) Aqueous ammonia, NH 3, forms a complex ion with silver ion, producing the soluble diamminesilver(i) ion, [Ag(NH 3 ) 2 ] 1+. The addition of NH 3 removes silver ion from the equilibrium in the above equation, shifting its equilibrium position to the left and causing AgCl to dissolve. 2NH 3 [Ag(NH 3 ) 2 ] 1+ Ag 1+ + Cl AgCl (s) Adding acid, H +, to the solution again frees silver ion to recombine with chloride ion and re-form solid silver chloride. This occurs because H + reacts with the NH 3 to yield ammonium ion, NH 4 1+, as detailed earlier. Ag 1+ + Cl AgCl (s) 2NH 3 + 2H + 1+ 2NH 4 [Ag(NH 3 ) 2 ] 1+ Silver Iodide Equilibrium. Iodide ion, I (from KI), added to the Ag 1+ + 2NH 3 Ag(NH 3 ) 2+ equilibrium from above results in the formation of solid silver iodide, AgI I AgI (s) Ag 1+ + 2NH 3 [Ag(NH 3 ) 2 ] 1+ The iodide ion removes the silver ion, causing a dissociation of the [Ag(NH 3 ) 2 ] 1+ ion and a shift to the left. Silver Sulfide Equilibrium. Silver sulfide, Ag 2 S, is less soluble than silver iodide, AgI. Addition of sulfide ion (from Na 2 S) to the AgI (s) Ag 1+ + I dynamic equilibrium removes silver ions. AgI dissolves but solid silver sulfide forms. Buffers: AgI (s) Ag 1+ + I ½ S 2 ½ Ag 2 S (s) In many areas of research, chemists need an aqueous solution that resists a ph change when doses of hydrogen ion, from a strong acid, or doses of hydroxide ion, from a strong base, are added. Biologists often grow cultures that are very susceptible to changes in ph and therefore a buffered medium is required. A buffer solution must be able to consume small additions of H 3 O 1+ and OH without undergoing large ph changes. Therefore, it must have present a basic component that can react with the H 3 O 1+ and and acidic component that can react with OH. Such a buffer solution consists of a weak acid and its conjugate base (or weak base and its conjugate acid). This experiment shows that an acetic acid-acetate buffer system resists large ph changes. CH 3 COOH + H 2 O (l) H 3 O 1+ + CH 3 CO 2 The addition of OH shifts the buffer equilibrium, according to LeChatelier s principle, to the right, because of its reaction with H 3 O 1+, forming H 2 O. The shift right is by an amount that is essentially equal to the moles of OH added to the buffer system. Thus, the amount of CH 3 CO 2 increases and the amount of CH 3 COOH decreases by an amount equal to the moles of strong base added.
CH 3 COOH + H 2 O (l) H 3 O 1+ + CH 3 CO 2 OH 2 H 2 O (l) Conversely the addition of H 3 O 1+ from a strong acid to the buffer system causes the equilibrium to shift left: the H 3 O 1+ combines with the acetate ion (a base) to form more acetic acid, an amount (moles) equal to the amount of strong acid added to the system. CH 3 COOH + H 2 O (l) H 3 O 1+ + CH 3 CO 2 H 3 O 1+ As a consequence of the addition of strong acid, the amount of CH 3 COOH increases and the amount of CH 3 CO 2 decreases by an amount equal to the moles of strong acid added to the buffer system. This experiment compares the ph changes of a buffered solution to those of an unbuffered solution when varying amounts of strong acid or base are added to each. Common-Ion Effect: The effect of adding an ion or ions common to those already present in a system at a state of dynamic equilibrium is called the common-ion effect. The effect is observed in this experiment for the following equilibria: 4Cl + [Co(H 2 O) 6 ] 2+ [CoCl 4 ] 2 + 6H 2 O (l) 4Br + [Cu(H 2 O) 4 ] 2+ [CuBr 4 ] 2 + 4H 2 O (l) These equations represent equilibria of ligands bonded to metal ions. These equilibria are shifted by changes in the concentrations of the anions. Changes in Temperature: Referring again to the equation: 2NO 2(g) N 2 O 4(g) + 58 kj The reaction for the formation of colorless dinitrogen tetraoxide is exothermic by 58 kj. To favor the formation of N 2 O 4, the reaction vessel should be kept cool; removing heat from the system causes the equilibrium to replace the removed heat and the equilibrium therefore shifts right. Added heat shifts the equilibrium in the direction that absorbs heat. For this reaction, a shift to the left occurs with addition of heat. This experiment examines the effect of temperature on the systems described in the equations: 4Cl + [Co(H 2 O) 6 ] 2+ [CoCl 4 ] 2 + 6H 2 O (l) 4Br + [Cu(H 2 O) 4 ] 2+ [CuBr 4 ] 2 + 4H 2 O (l)
These systems involve an equilibrium between the coordination spheres, the water versus the halide ion about the cobalt(ii) and copper(ii) ions; both equilibria are concentration and temperature dependent. The tetrachlorocobaltate(ii) ion, [CoCl 4 ] 2, and the tetraaquacopper(ii) ion, [Cu(H 2 O) 4 ] 2+, are more stable at higher temperatures. CHEMICALS AND EQUIPMENT: AgNO 3, 0.01 M CoCl 2, 1 M HCl, 0.1 M, 1 M KI, 0.1 M Na 2 CO 3, 0.1 M Na 2 S, 0.1 M Universal Indicator CH 3 COOH, 0.1 M CuSO 4, 0.1 M HNO 3, 6 M NaCH 3 CO 2, 0.1 M NaOH, 1 M NH 3, concentrated Test Tubes PROCEDURE: Procedure Overview: A large number of qualitative tests and observations are performed. The effects that concentration changes and temperature changes have on a system at equilibrium are observed and interpreted using LeChatelier s principle. The functioning of a buffer system and the effect of a common ion on equilibria are observed. Perform this experiment with a partner. At each asterisk (*) in the procedure, stop, and record your observations in your laboratory notebook. Discuss your observations with your lab partner and instructor. Account for the changes in appearance of the solution after each addition in terms of LeChatelier s principle. Formation of Metal-Ammonia Ions. Place 1 ml (<20 drops) of 0.1 M CuSO 4 in a small, clean test tube. * Add drops of concentrated NH 3 until a color change occurs and the solution is clear. * Shift of Equilibrium. Add drops of 1 M HCl until the color again changes. * Silver Carbonate Equilibrium. In a 150-mm test tube, add ½ ml of 0.01 M AgNO 3 to ½ ml of 0.1 M Na 2 CO 3. * Add drops of 6 M HNO 3 to the precipitate until evidence of a chemical change occurs. * Silver Chloride Equilibrium. To the clear solution, add ~ 5 drops of 0.1 M HCl. * Add drops of concentrated NH 3 until evidence of a chemical change. * Reacidify the solution with 6 M HNO 3 and record your observations. * What happens if excess concentrated ammonia is again added? Try it. Silver Iodide Equilibrium. After trying it, add drops of 0.1 M KI. * Silver Sulfide Equilibrium. To the mixture, add drops of 0.1 M Na 2 S until evidence of chemical change has occurred. * Rinse the test tube twice with tap water and discard in the Waste Silver Salts container. Rinse twice with deionized water and discard in the sink. A Buffer System: Preparation of Buffered and Unbuffered Systems. Transfer 10 drops of 0.1 M CH 3 COOH to wells A1 and A2 of a 24-well plate, add 3 drops of universal indicator, and note the color. * Now add 10 drops of 0.1 M NaCH 3 CO 2 to each well.
Compare the color of the solution with the ph color chart for the universal indicator. * Place 20 drops of deionized water into wells B1 and B2 and add 3 drops of universal indicator. Effect of Strong Acid. Transfer 5 drops of 0.10 M HCl to wells A1 and B1, estimate the ph, and record each ph change. * Effect of Strong Base. Transfer 5 drops of 0.10 M NaOH to wells A2 and B2, estimate the ph, and record each ph change. * Effect of a Buffer System. Explain the effect a buffered system (as compared with an unbuffered system) has on ph change when a strong acid or strong base is added to it. * Equilibrium (Common-Ion Effect): Effect of Concentrated HCl. Place about 5 drops of 1 M CoCl 2 in well Cl of the 24-well plate. * Add drops of concentrated HCl until a color change occurs. * Slowly add water to the system and swirl. * Equilibrium (Temperature Effect): Add Heat from a Boiling Water Bath. Place about 1 ml of 1 M CoCl 2 in a 75-mm test tube into the boiling water bath. Compare the color of the hot solution with that of the original cool solution. *
Prelaboratory Assignment NAME: Answer the following questions. 1. Explain the meaning of the phrase, a system in a state of dynamic equilibrium. 2. A state of dynamic equilibrium exists between the solid sucrose molecules and sucrose molecules dissolved in aqueous solution for a saturated sucrose solution. a. Write an equation to represent the dynamic equilibrium. b. What change occurs in the concentration of dissolved sucrose if more solid sucrose is added to the saturated system? Explain. c. If water is added to the saturated solution and equilibrium is re-established, what change occurs in the amount of dissolved sucrose? Explain. d. If water is added to the saturated solution and equilibrium is re-established, what change occurs in the concentration of dissolved sucrose?
Prelaboratory Assignment (continued) NAME: 3. Consider the following aqueous chemical equilibrium of benzoic acid, a weak acid: C 6 H 5 COOH + H 2 O (l) H 3 O 1+ + C 6 H 5 CO 2 a. Addition of H 3 O 1+ to the chemical equilibrium the amount of C 6 H 5 CO 2 in the system. b. Addition of OH to the chemical equilibrium the amount of C 6 H 5 COOH in the system. c. Increasing the ph of the solution the amount of C 6 H 5 COOH in the system. 4. The carbon dioxide of the atmosphere, being a nonmetallic oxide and acid anhydride, has a low solubility in rainwater but produces a slightly acidic solution. Write an equilibrium equation for the dissolution of carbon dioxide in water and identify the acid that is formed. 5. Calcium carbonate, the major component of limestone, is slightly soluble in water. Write an equilibrium equation showing the slight solubility of calcium carbonate in water.