Chemical Equilibria: Coordination Compounds

Transcription

1 E11 Chemical Equilibria: Coordination Compounds Objective Ø Illustrate the tendency of metal ions to form metal coordination complexes with ions and neutral polar molecules that act as electron-pair donors (Lewis bases) toward metal ions. Ø Study metal ion complexes with ammonia, chloride ion, and hydroxide ion. Ø Determine the dissociation constant, K diss, of the diamminesilver(i) complex ion, Ag(NH 3 ) 2. Discussion In Chapter 23 we have studied the most common class of coordination compounds. We have noted that all ions in aqueous solution attract polar water molecules to them to form ion dipole (Lewis acid-base) bonds. The cupric ion, Cu 2, for example, exist as the hexaaquocopper(ii) ion, Cu(H 2 O) 6 2. The negative side of the polar water molecule is strongly attracted to metal cations. In general, the greater the charge on the cation and the larger the cation, the greater the number of coordinated water molecules. Examples are H(H 2 O), Be(H 2 O) 4 2, Cu(H 2 O) 6 2, Al(H 2 O) 6 3, Fe(H 2 O) 6 3. Anions attract the positive side of the polar water molecule and are also hydrated. However, the attraction is not as large, the hydrates are less stable, and fewer water molecules are coordinated. Other neutral, but polar, molecules such as ammonia, NH 3, and also a number of anions, such as OH -, Cl -, CN -, S 2-, S 2 O 3 2-, and C 2 O 4 2-, likewise can form similar very stable coordination groupings about a central metal ion. Such coordination compounds result from the replacement of the water molecule from the hydrated ion by other molecules or ions when they are present in a solution at high concentration, forming a still more stable bond. The resulting coordination compound may be a positively or negatively charged ion (a complex ion), or it may be a neutral molecule, depending on the number and kind of coordinating ligands (groups) attached to the central ion. In this experiment we will focus on ammonia, hydroxide, and chloride complexes of Cu 2, Zn 2, and Ag. Ammonia Complex Ions Some of the important ammonia complexes are Co(NH 3 ) Ni(NH 3 ) 4 2 Ni(NH 3 ) 6 Cu(NH 3 ) 2 2 Cu(NH 3 ) 4 Ag(NH 3 ) 2 Au(NH 3 ) 2 2 Zn(NH 3 ) 4 2 Cd(NH 3 ) 4 These complexes are formed by adding ammonia to a solution containing the hydrated cation. Ammonia molecules are bound by the cation one at a time as the concentration of ammonia increases. At low concentrations of ammonia, smaller numbers of ammonia molecules may be bound as ligands. For instance, the two NH 3 molecules that bond to the Ag bind in successive steps. The equilibrium constant is known for each step of the reaction sequence Ag NH 3 Ag(NH 3 ) Ag(NH 3 ) NH 3 Ag(NH 3 ) 2 K 1 = [Ag(NH 3 ) ] =1.6 [Ag x103 ][NH 3 ] K 2 = [Ag(NH 3 ) 2 ] = 6.8 [Ag(NH 3 ) x103 ][NH 3 ]

2 Equilibria of Coordination Compounds Ag 2 NH 3 Ag(NH 3 ) 2 K formation = K 1 K 2 = 1.1 x 10 7 The formation constant, K f, represents the overall equilibrium constant given above. We will exam complex ion equilibria in Chapter 17. In the last part of this experiment we will determine the dissociation constant, K diss = 1/K formation, for the Ag(NH 3 ) 2 complex ion. 1 Amphoteric Hydroxides - The Hydroxide Complex Ions The hydroxides of most metals are relatively insoluble in water. Thus, when a strong base like sodium hydroxide is added to a metal ion in solution, such as chromium ion, a precipitate is formed. Cr(H 2 O) OH - Cr(H 2 O) 3 (OH) 3 (s) 3 H 2 O or, using the unhydrated metal ion formula it simplifies to, Cr 3 3 OH - Cr(OH) 3 By Le Châtelier s principle, it is expected that excess hydroxide ion would give more complete precipitation. Instead, the precipitate dissolves! This is explained by the tendency of chromium ion to form a more stable coordination compound with excess hydroxide ion: Cr(H 2 O) 3 (OH) 3 (s) OH - Cr(H 2 O) 2 (OH) 4 - H 2 O or Cr(OH) 3 OH - Cr(OH) 4 - For simplicity we shall use the unhydrated formulas except where it is important to emphasize the hydrated structure. It is important to note that the reactions that form these hydroxide complex ions are entirely reversible. The addition of acid to the above strongly basic Cr(OH) - 4 solution reacts first to reprecipitate the hydroxide, - Cr(OH) 4 H Cr(OH) 3 (s) H 2 O and then, with excess acid, Cr(OH) 3 (s) 3 H Cr 3 3 H 2 O Metal hydroxides, like these, that may be dissolved by an excess of either a strong acid or a strong base, are called amphoteric hydroxides, Table-1. Table-1 Some Important Amphoteric Hydroxides Simple Ion (acidic solution) Precipitate Hydroxide Complex ion (strongly basic solutions Al 3 Al(OH) 3 Al(OH) - 4, tetrahydroxoaluminate ion Cr 3 Cr(OH) 3 Al(OH) - 4, tetrahydroxochromate(iii) ion Pb 2 Pb(OH) 2 Pb(OH) - 3, trihydroxoplumbate(ii) ion Sn 2 Sn(OH) 2 Sn(OH) - 3, trihydroxostannate(ii) ion Sn 4 Sn(OH) 4 Sn(OH) 2-6, hexahydroxostannate(iv) ion Zn 2 Zn(OH) 2 Zn(OH) - 4, tetrahydroxozincate(ii) ion 1 In the experimental procedure, the [NH3 ] is so high that the [Ag(NH 3 ) ] may be neglected. -2-

3 Equilibria of Coordination Compounds Experimental Procedure Chemicals: 1 M ammonium chloride, NH 4 Cl; 1 M ammonia, NH 3 ; 15 M (conc.) ammonia, NH 3 ; 6 M ammonia, NH 3 ; 6 M sodium hydroxide, NaOH; 6 M nitric acid, HNO 3 ; 6 M hydrochloric acid, HCl; 12 M (conc.) hydrochloric acid, HCl; copper(ii) sulfate pentahydrate, CuSO 4 5H 2 O(s); 0.1 M copper(ii) sulfate, CuSO 4 ; 0.1 M sodium chloride, NaCl; 0.1 M silver nitrate, AgNO 3 ; 0.1 M zinc nitrate, Zn(NO 3 ) 2 ; 0.1% phenolphthalein, alizarin yellow R, and indigo carmine indicators. SAFETY PRECAUTIONS: Fuming concentrated (12 M) HCl and concentrated (15 M) NH 3 are lung irritants. These solutions and 6 M HCl, 6 M HNO 3, 6 M NH 3, and 6 M NaOH are hazardous to the skin and eyes. If they contact your skin, wash them off immediately. Dispense the solutions in a wellventilated fume hood. Clean up any spills immediately. Waste Collection: Waste containers are provided under the fume hood for the copper, zinc, and silver compounds formed in this experiment. 1. The Formation of Complex Ions with Ammonia To 3 ml of 0.10 M CuSO 4, add a drop of 6 M NH 3. Mix this well. (Record your observations and write the equation for the reaction.) Continue to add NH 3 a little at a time, with thorough mixing, until a distinct change occurs. Save this solution. Is this result contrary to Le Châtelier s principle? Obviously the OH - concentration was increasing while the Cu(OH) 2 dissolved. How must have the Cu 2 concentration changed? Did it increase or decrease? To learn which of the substances present in an ammonia solution (NH 4, OH -, NH 3, H 2 O) is responsible for the change you noted, try the following tests: (a) To a 15x100 mm test tube add 1 ml of 1 M NH 4 Cl to 1 ml of 0.10 M CuSO 4, and mix well. (b) To another 15x100 mm test tube add 2 drops (an excess) 2 of 6 M NaOH to 2 ml of 0.10 M CuSO 4, and mix well. (c) Add ammonia gas by placing several crystals of CuSO 4 5H 2 O(s) in a small dry 50 ml beaker as follows: At one side in the beaker, place a small piece of filter paper moistened with concentrated (15 M) NH 3. Cover with a watch glass and observe any changes. From this evidence, write an equation to show the formation of this new substance when excess NH 3 is added to Cu 2. To 1 ml of the cupric complex ion solution, you saved above, add 6 M HNO 3 in excess. Explain the result and write the equation for the reaction. Dispose of the solution in the copper waste container. 2. The Formation of Amphoteric Hydroxides To 5 ml of 0.10 M Zn(NO 3 ) 2, in a 15x100 mm test tube add 6 M NaOH drop by drop, with mixing, until the precipitate that first forms just redissolves. Avoid undue excess of NaOH. Divide the solution into two portions; test one portion with 2 drops of alizarin yellow R and the other with 2 drops of indigo carmine indicator. Estimate the OH - concentration (for later comparison in part 3), using the information on the color changes and ph intervals of the indicators given in the endnotes (end of experimental procedure). Now to one portion, add 6 M HCl, drop by drop with mixing, until a precipitate forms (what is it?), then add more HCl mix well to redissolves the precipitate. Interpret all these changes as related to Le Châtelier s principle and to the relative concentration of the various constituents (the zinc in its various forms, H, and OH - ), both in words and in net ionic equations. 2 This provides an excess of OH -, a much stronger base than NH 3. The strong base OH! shows some amphoteric effect (see part 2) with copper(ii) salts but is far from complete. -3-

4 Equilibria of Coordination Compounds 3. The Reaction of Zinc Ion with Ammonia When ammonia is added gradually to Zn 2, does the precipitate of zinc hydroxide that first forms redissolve as zincate ion, Zn(OH) 4 2-, owing to the excess base added, or does it redissolve as Zn(NH 3 ) 4 2, owing to the NH 3 molecules added? To answer this question, to 3 ml of 0.10 M Zn(NO 3 ) 2 in a 15x100 mm test tube add 6 M NH 3 drop by drop, with mixing, until the precipitate that first forms just redissolves. Divide this solution into two portions, test one portion with 2 drops of phenolphthalein and the other portion with 2 drops of alizarin yellow R. Estimate the approximate OH - concentration and compare this with the corresponding situation in part 2, where NaOH was used (see endnotes). What can you conclude about the possibility of forming Zn(OH) 4 2- by adding NH 3 to a zinc salt solution? Explain. Write the equation for the equilibrium that you have verified. 4. Chloride Complex Ions (a) To 2 ml of 0.10 M CuSO 4 in a 15x150 or 20x150 mm test tube, add 2 ml of 12 M (concentrated) HCl; then dilute the solution with about 5 ml of water. Write equations and interpret the color changes you observed, assuming that the complex formed is CuCl (b) To a 15x150 or 20x150 mm test tube add 1 ml of 0.10 M AgNO 3, plus 3 ml of 12 M (concentrated) HCl; then agitate the solution well for several minutes to redissolve the precipitated AgCl. Now dilute this mixture with about 5 ml of distilled water. Write equations for the reactions and interpret the changes you observed, assuming that the complex formed is AgCl 2 -. Dispose of the tube contents in the silver waste container. 5. The Equilibrium Constant of an Ammonia Complex Ion The dissociation of diamminesilver(i) complex ion is represented by the equilibrium and the corresponding equilibrium constant expression Ag(NH 3 ) 2 Ag 2 NH 3 (1) K diss = [Ag ][NH 3 ] 2 [Ag(NH 3 ) 2 ] (2) If you add sufficient Cl - gradually to an equilibrium mixture of Ag and NH 3 {represented by Reaction (1)} so that you just barely begin precipitation of AgCl(s), a second equilibrium is established simultaneously without appreciably disturbing the first equilibrium. This may be represented by the combined reactions Ag(NH 3 ) 2 Ag 2 NH 3 Cl - AgCl(s) (3) By using a large excess of NH 3, you can shift Reaction (1) far to the left, with reasonable assurance that the Ag is converted almost completely to Ag(NH 3 ) 2 rather than to the first step only, Ag(NH 3 ). From the measured volumes of NH 3, Ag, and Cl - solutions used, you can determine the concentrations of the species in Reaction (1) and calculate the value of K diss. -4-

5 Equilibria of Coordination Compounds To prepare the solution, place 3.0 ml of 0.10 M AgNO 3 (measure it accurately in a 10-mL graduate) into a 15 x 150 or 20x150 mm test tube. Add 3.0 ml (also carefully measured) of 1.0 M NH 3. Obtain 10.0 ml of M NaCl in a 10-mL graduate cylinder, and record exact volume. {If the M NaCl is not available, prepare some M NaCl by diluting 2.0 ml of 0.10 M NaCl to 10.0 ml in your 10-mL graduate cylinder. Mix this thoroughly by pouring it into a 15x150 mm test tube mix well and pour it back into the 10-mL graduated cylinder and record the exact volume.} Then, using a medicine dropper or small pipet, start by adding about 1 to 1.5 ml of the M NaCl to the mixture of AgNO 3 and NH 3, mix well, then add the NaCl drop by drop mixing well until a very faint, permanent milky precipitate of AgCl remains. Return any excess NaCl from the medicine dropper/pipet to the graduate cylinder and record the exact volume used. From these data, K diss can be calculated. Dispose of the solution in the silver waste container. Acid Base indicators for approximate ph measurements. Indicator ph interval color change alizarin yellow R yellow - red indigo Carmen blue yellow phenolphthalein colorless - red -5-

6 E11 Chemical Equilibria: Coordination Compounds Report Form Name: Partner s Name: (if any)_lab Section: MW/TTH/M-TH (circle) Data and Observations -4-

7 Chemical Equilibria: K sp of Calcium Iodate Report Form -5-

8 Chemical Equilibria: K sp of Calcium Iodate Report Form -6-

9 Chemical Equilibria: K sp of Calcium Iodate Report Form 2. When ammonia is added to Zn(NO 3 ) 2 solution, a white precipitate forms, which dissolves upon addition of the excess ammonia. But when ammonia is added to a mixture of Zn(NO 3 ) 2 and NH 4 NO 3, no precipitate forms at anytime. Suggest an explanation for this difference in behavior. {Hint: Consider the common ion affect and the solubility products constant, K sp, of the white precipitate.} -7-