Chemical Equilibria: Coordination Compounds

E11

Objective

Ø Illustrate the tendency of metal 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 complexes with ammonia, chloride ion, and 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-) 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 - The Hydroxide Complex Ions The hydroxides of most metals are relatively insoluble in water. Thus, when a strong base like 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 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 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. 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 2- you conclude about the possibility of forming Zn(OH)4 by adding NH3 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 CuSO4 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 2- you observed, assuming that the complex formed is CuCl4 . (b) To a 15x150 or 20x150 mm test tube add 1 mL of 0.10 M AgNO3, 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 AgCl2 . 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

+ + Ag(NH3)2 ⇐⇒ Ag + 2 NH3 (1) and the corresponding equilibrium constant expression

[Ag+ ][NH ]2 3 (2) Kdiss = + [Ag(NH3 )2 ]

- + If you add sufficient Cl gradually to an equilibrium mixture of Ag and NH3 {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(NH3)2 ⇐⇒ Ag + 2 NH3 + Cl- ⇑ ⇓ AgCl(s) (3)

By using a large excess of NH3, you can shift Reaction (1) far to the left, with reasonable assurance that + + + the Ag is converted almost completely to Ag(NH3)2 rather than to the first step only, Ag(NH3) . From + - the measured volumes of NH3, Ag , and Cl solutions used, you can determine the concentrations of the species in Reaction (1) and calculate the value of Kdiss.

-4- Equilibria of Coordination Compounds

To prepare the solution, place 3.0 mL of 0.10 M AgNO3 (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 NH3. Obtain 10.0 mL of 0.020 M NaCl in a 10-mL graduate cylinder, and record exact volume. {If the 0.020 M NaCl is not available, prepare some 0.020 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 0.020 M NaCl to the mixture of AgNO3 and NH3, 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, Kdiss 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 10.0 – 12.0 yellow - red indigo Carmen 11.4 – 13.0 blue – yellow phenolphthalein 8.3 - 10.0 colorless - red

-5- Chemical Equilibria: Coordination Compounds E11 Report Form

Name: ______Partner’s Name: ______(if any)_Lab Section: MW/TTH/M-TH (circle)

Data and Observations

-4- Chemical Equilibria: Ksp of Calcium Iodate Report Form

-5- Chemical Equilibria: Ksp of Calcium Iodate Report Form

-6- Chemical Equilibria: Ksp of Calcium Iodate Report Form

2. When ammonia is added to Zn(NO3)2 solution, a white precipitate forms, which dissolves upon addition of the excess ammonia. But when ammonia is added to a mixture of Zn(NO3)2 and NH4NO3, no precipitate forms at anytime. Suggest an explanation for this difference in behavior. {Hint: Consider the common ion affect and the solubility products constant, Ksp, of the white precipitate.}

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