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Chemical Equilibria: Coordination Compounds

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Chemical Equilibria: Coordination Compounds Chemical Equilibria: Coordination Compounds E11 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, Kdiss, of the diamminesilver(I) complex ion, Ag(NH3)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) 2+ 2+ bonds. The cupric ion, Cu , for example, exist as the hexaaquocopper(II) ion, Cu(H2O)6 . 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 + 2+ 2+ 3+ 3+ molecules. Examples are H(H2O) , Be(H2O)4 , Cu(H2O)6 , Al(H2O)6 , Fe(H2O)6 . 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, NH3, and also a number of anions, such as OH , - - 2-, 2- 2- Cl , CN , S S2O3 , and C2O4 , 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 Cu2+, Zn2+, and Ag+. Ammonia Complex Ions Some of the important ammonia complexes are 2+ 2+ + 2+ Co(NH3)6 Ni(NH3)4 Cu(NH3)2 Zn(NH3)4 2+ 2+ 2+ Ni(NH3)6 Cu(NH3)4 Cd(NH3)4 + Ag(NH3)2 + Au(NH3)2 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 NH3 molecules that bond to the Ag bind in successive steps. The equilibrium constant is known for each step of the reaction sequence + + + Ag + NH3 ⇐⇒ Ag(NH3) [Ag(NH ) ] K = 3 =1.6 x103 1 + [Ag ][NH3 ] + + + [Ag(NH3 )2 ] 3 Ag(NH3) + NH3 ⇐⇒ Ag(NH3)2 K2 = + = 6.8 x10 [Ag(NH ) ][NH ] 3 3 € € Equilibria of Coordination Compounds + + 7 Ag + 2 NH3 ⇐⇒ Ag(NH3)2 Kformation = K1·K2 = 1.1 x 10 The formation constant, Kf, 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 + 1 the dissociation constant, Kdiss = 1/Kformation, for the Ag(NH3)2 complex ion. 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. 3+ - Cr(H2O)6 + 3 OH ⇐⇒ Cr(H2O)3(OH)3(s) + 3 H2O or, using the unhydrated metal ion formula it simplifies to, 3+ - Cr + 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(H2O)3(OH)3(s) + OH ⇐⇒ Cr(H2O)2(OH)4 + H2O 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) + H2O and then, with excess acid, + 3+ Cr(OH)3(s) + 3 H ⇐⇒ Cr + 3 H2O 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 3+ - Al Al(OH)3 Al(OH)4 , tetrahydroxoaluminate ion 3+ - Cr Cr(OH)3 Al(OH)4 , tetrahydroxochromate(III) ion 2+ - Pb Pb(OH)2 Pb(OH)3 , trihydroxoplumbate(II) ion 2+ - Sn Sn(OH)2 Sn(OH)3 , trihydroxostannate(II) ion 4+ 2- Sn Sn(OH)4 Sn(OH)6 , hexahydroxostannate(IV) ion 2+ - Zn Zn(OH)2 Zn(OH)4 , tetrahydroxozincate(II) ion 1 + In the experimental procedure, the [NH3] is so high that the [Ag(NH3) ] may be neglected. -2- Equilibria of Coordination Compounds Experimental Procedure Chemicals: 1 M ammonium chloride, NH4Cl; 1 M ammonia, NH3; 15 M (conc.) ammonia, NH3; 6 M ammonia, NH3; 6 M sodium hydroxide, NaOH; 6 M nitric acid, HNO3; 6 M hydrochloric acid, HCl; 12 M (conc.) hydrochloric acid, HCl; copper(II) sulfate pentahydrate, CuSO4·5H2O(s); 0.1 M copper(II) sulfate, CuSO4; 0.1 M sodium chloride, NaCl; 0.1 M silver nitrate, AgNO3; 0.1 M zinc nitrate, Zn(NO3)2; 0.1% phenolphthalein, alizarin yellow R, and indigo carmine indicators. SAFETY PRECAUTIONS: Fuming concentrated (12 M) HCl and concentrated (15 M) NH3 are lung irritants. These solutions and 6 M HCl, 6 M HNO3, 6 M NH3, 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 well- ventilated 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 CuSO4, add a drop of 6 M NH3. Mix this well. (Record your observations and write the equation for the reaction.) Continue to add NH3 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 - 2+ the OH concentration was increasing while the Cu(OH)2 dissolved. How must have the Cu concentration changed? Did it increase or decrease? + - To learn which of the substances present in an ammonia solution (NH4 , OH , NH3, H2O) is responsible for the change you noted, try the following tests: (a) To a 15x100 mm test tube add 1 mL of 1 M NH4Cl to 1 mL of 0.10 M CuSO4, and mix well. 2 (b) To another 15x100 mm test tube add 2 drops (an excess) of 6 M NaOH to 2 mL of 0.10 M CuSO4, and mix well. (c) Add ammonia gas by placing several crystals of CuSO4·5H2O(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) NH3. Cover with a watch glass and observe any changes. From this evidence, write 2+ an equation to show the formation of this new substance when excess NH3 is added to Cu . To 1 mL of the cupric complex ion solution, you saved above, add 6 M HNO3 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(NO3)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 NH3. The strong base OH! shows some amphoteric effect (see part 2) with copper(II) salts but is far from complete. -3- Equilibria of Coordination Compounds 3. The Reaction of Zinc Ion with Ammonia When ammonia is added gradually to Zn2+, does the precipitate of zinc hydroxide that first forms 2- redissolve as zincate ion, Zn(OH)4 , owing to the excess base added, or does it redissolve as 2+ Zn(NH3)4 , owing to the NH3 molecules added? To answer this question, to 3 mL of 0.10 M Zn(NO3)2 in a 15x100 mm test tube add 6 M NH3 drop by drop, with mixing, until the precipitate that first forms just redissolves.
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