Chapter 11 Experiment: Transition Metal Chemistry OBJECTIVES: To observe the various colors associated with transition metal ions To determine the relative strength of a ligand To compare the stability of complexes To synthesize a coordination compound TECHNIQUES: Cleaning Glassware Handling Chemicals Disposing Chemicals Measuring Mass Test Tubes for Small Volumes Decanting a Liquid or Solution from a Solid Flushing a Precipitate from the Beaker Vacuum Filtration Venting Gases Erlenmeyer Flask over a Cool Flame INTRODUCTION: One of the most intriguing features of the transition metal ions is their vast array of colors. The blues, greens, and reds that we associate with chemicals are oftentimes due to the presence of transition metal ions. In previous experiments, we have already noted some of the characteristic colors of certain hydrated transition metal ions; for example, hydrated Cu 2+ salts are blue, Ni 2+ salts are green, and Fe 2+ salts are rust colored. A second interesting feature of the transition metal ions is that subtle, and on occasion very significant, color changes occur when molecules or ions other than water bond to the metal ion to form a complex. These molecules or ions other than water, called ligands, are Lewis bases (electron pair donors) which bond directly to the metal ion, producing a change in the electronic energy levels of the metal ion. As a result, the energy (and also the wavelength) of light absorbed by the electrons in the transition metal ion and, consequently, the energy (and wavelength) of light transmitted also change. The solution has a new color. The complex has a number of ligands bonded to the transition metal ion which form a coordination sphere. The complex along with its neutralizing ion is called a coordination compound. For example, K 4 [Fe(CN) 6 ] is a coordination compound: the six CN 1 ions are the ligands, Fe(CN) 6 is the coordination sphere, and [Fe(CN) 6 ] 4 is the complex. More than the intrigue of the color and color changes are the varied uses that these complexes have in chemistry. A few examples of the applicability of the complex formation follow: For photographic film development, sodium thiosulfate, Na 2 S 2 O 3, called hypo, removes the unsensitized silver ion from the film in the form of the soluble silver complex, [Ag(S 2 O 3 ) 2 ] 3 : AgBr (s) + 2 S 2 O 3 2 (aq) [Ag(S 2 O 3 ) 2 ] 3 (aq) + Br 1 (aq)
For the removal of calcium ion hardness from water, soluble polyphosphates, such as sodium tripolyphosphate, Na 5 P 3 O 10, are added to detergents to form a soluble calcium complex: Ca 2+ (aq) + P 3 O 10 5 (aq) [CaP 3 O 10 ] 3 (aq) The bond strength between the transition metal ion and its ligands varies, depending on the electron pair donor (Lewis base) strength of the ligand and the electron pair acceptor (Lewis acid) strength of the transition metal ion. Ligands may be neutral (e.g., H 2 O, NH 3, H 2 NCH 2 CH 2 NH 2 ) or anionic (e.g., CN 1, SCN 1, Cl 1 ). A single ligand may form one bond to the metal ion (a monodentate ligand), two bonds to the metal ion (a bidentate ligand), three bonds to the metal ion (a tridentate ligand), and so on. Ligands that form two or more bonds to a transition metal ion are also called chelating agents and sequestering agents. Note that a nonbonding electron pair (a Lewis pair) is positioned on each atom of the ligand that serves as a bonding site to the transition metal ion. The complex formed between a chelating agent (a polydentate ligand) and a metal ion is generally more stable than that formed between a monodentate ligand and a metal ion. The explanation is that the several ligand-metal bonds between a polydentate ligand and a metal ion are much more difficult to break than a single bond between a monodentate ligand and a metal ion. The stability of complexes having polydentate ligands will be compared with the stability of those having only monodentate ligands in this experiment. The number of bonds between a metal ion and its ligands is called the coordination number of the complex. If four monodentate ligands or if two bidentate ligands bond to a metal ion, the coordination number for the complex is 4, and so on. Coordination numbers of 2, 4, and 6 are most common among the transition metal ions. The stability of a complex can be determined by mixing it with an anion known to form a precipitate with the cation. The anion most commonly used to measure the stability of a complex is the hydroxide ion. When hydroxide ion is added to a system containing the complex [CuX 4 ] 2 (aq), copper (II) ion has a choice of now combining with the anion that forms the stronger bond in other words a competition for copper (II) ion between the two anions exists in solution. If the ligand forms the stronger bond to copper (II) ion, the complex remains in solution and no change is observed; if on the other hand the hydroxide forms a stronger bond to copper (II) ion, Cu(OH) 2 precipitates from solution. Therefore a measure of the stability of the complex is determined. To write the formulas of the complexes in this experiment we will assume that the coordination number of copper (II) ion is always 4 and that of nickel (II) and cobalt (II). Later stages of the experiment outline the syntheses of several coordination compounds. Tetraamminecopper(II) sulfate monohydrate [Cu(NH 3 ) 4 ]SO 4 H 2 O Hexamminenickel(II) chloride [Ni(NH 3 ) 6 ]Cl 2 Tris (ethylenediamine) nickel (II) Chloride, dihydrate [Ni(en) 3 ]Cl 2 2H 2 O Tetraamminecopper (II) Sulfate Monohydrate: Ammonia is added to an aqueous solution of copper (II) sulfate pentahydrate. Ammonia displaces the four water ligands from the coordination sphere: [Cu(H 2 O) 4 ]SO 4 H 2 O + 4 NH 3 [Cu(NH 3 ) 4 ]SO 4 H 2 O + 4H 2 O The solution is cooled and 95% ethanol is added to reduce the solubility of the blue tetraamminecopper (II) sulfate salt.
Hexaamminenickel (II) Chloride: For the preparation of Hexaamminenickel (II) Chloride, six ammonia molecules [Ni(H 2 O) 6 ]Cl 2 displace the six water ligands in [Ni(H 2 O) 6 ]Cl 2 + 6NH 3 [Ni(NH 3 ) 6 ]Cl 2 + 6 H 2 O [Ni(NH 3 ) 6 ]Cl 2 coordination compound is cooled and precipitated with 95% ethanol; the lower polarity of the 95% ethanol causes the lavender salt to become less soluble Tris(ethylenediamine)nickel(II) Chloride Dihydrate: Three ethylenediamine (en) molecules displace the six water ligands of [Ni(H 2 O) 6 ]Cl 2 for the preparation of the violet tris (ethylenediamine)nickel(ii) chloride dihydrate. [Ni(H 2 O) 6 ]Cl 2 + 3 en [Ni(en) 3 ]Cl 2 2H 2 O + 4 H 2 O PROCEDURE: Several complexes of copper (II) ion, nickel (II) ion, and cobalt (II) ion are formed and studied. The observations of color change that result from the addition of a ligand are used to understand the relative stability of the various complexes that form. One or more coordination compounds are synthesized and isolated. Some colors may be difficult to distinguish. If a precipitate initially forms when a solution of the ligand is added, add more of the solution. Always compare the test solution with the original aqueous solution. On occasion you may need to discard some of the test solution if too much of the original solution was used for testing at the outset. If a color change occurs (not a change in color intensity) then a new complex has formed. Also, after each addition of the ligandcontaining solution, tap the test tube to agitate the mixture and view the solution at various angles to note color change. Caution: Several of the ligand-containing solutions should be handled with care. The conc HCl, conc NH 3, and the ethylenediamine reagents produce characteristic odors that are strong irritants of the respiratory system. Use drops as suggested and avoid inhalation and skin contact. Most complex ion formation reactions are completed in small test tubes. While volumes of solutions need not be exact, a volume close to the suggested amount should be used. A 75-mm test tube has a volume of 3 ml proportional fractions should be used in estimating the volume suggested in the procedure. From six to twelve small test tubes are used for each part of the experiment. Plan to keep them clean, rinsing thoroughly after each use. However do not discard any test solutions until the entire subsection of each part of the Experimental Procedure has been completed. Consult with your laboratory instructor to determiner which of the coordination compounds outlined in the latter parts of the experiment you are allowed to synthesize. Chloro Complexes of the Copper (II), Nickel (II), and Cobalt (II) ions: Form the complexes. Place about 0.5 ml of 0.1 CuSO 4, 0.1 M Ni(NO 3 ) 2, and 0.1 M CoCl 2 into each of three separate small test tubes. Add 1 ml (20 drops) of conc HCl to each. (CAUTION: Do not allow conc HCl to contact skin or clothing. Flush the affected area with water immediately.) Tap the test tube to agitate.
Dilute the Complexes. Slowly add 1-2 ml of water to each test tube. Compare the colors of the solutions to about 2.5 ml of the original solutions. Does the original color return? Complexes of the Copper (II) Ion: Form the Complexes. Place about 0.5 ml of 0.1 M CuSO 4 in each of five test tubes and transfer them to the fume hood. Add 5 drops of conc NH 3 to the first test tube, 5 drops of ethylenediamine to the second, 5 drops of 0.1 M KSCN to the third test tube, and nothing more to a fourth test tube. The fifth test tube can be used to form a complex with a ligand selected by your instructor the thiosulfate anion, the oxalate anion, the tartrate anion, and the ethylenediamine anion, readily form complexes and are suggested. If a precipitate forms in any of the solutions, add an excess of the ligand-containing solution. Compare the appearance of the test solutions with 1 ml of the original CuSO 4 solution in the fourth test tube. Test for Stability. Add 3-5 drops of 1 M NaOH to each of the five solutions. Account for your observations. Complexes of the Nickel (II) Ion: Form the Complexes. Place about 0.5 ml of 0.1 M NiCl 2 in each of five test tubes and transfer them to the fume hood. Add 5 drops of conc NH 3 to the first test tube, 5 drops of ethylenediamine to the second, 5 drops of 0.1 M KSCN to the third test tube, and nothing more to a fourth test tube. The fifth tube will be used to form a complex, as in the previous procedure. Test for Stability. Add 3-5 drops of 1 M NaOH to each of the five solutions. Account for your observations. Complexes of the Cobalt (II) Ion: Form the Complexes. Place about 0.5 ml of 0.1 M Co(NO 3 ) 2 in each of five test tubes and transfer them to the fume hood. Add 5 drops of conc NH 3 to the first test tube, 5 drops of ethylenediamine to the second, 5 drops of 0.1 M KSCN to the third test tube, and nothing more to a fourth test tube. The fifth tube will be used to form a complex, as in the previous procedure. Test for Stability. Add 3-5 drops of 1 M NaOH to each of the five solutions. Account for your observations. Rinse the test tubes twice with tap water and twice with deionized water. Discard the rinses in the sink, followed by a generous amount of tap water. Synthesis of Tetraamminecopper (II) Sulfate Monohydrate: Dissolve the Starting Material. Measure 6 g of copper(ii) sulfate pentahydrate in a clean, dry, 150-mL beaker. Dissolve the sample in 15 ml of deionized water. Heating may be necessary. Transfer the beaker to the fume hood. Precipitate the Complex Ion. Add conc NH 3 until the precipitate that initially forms has dissolved. Cool the deep blue solution in an ice bath. Cool 20 ml of 95% ethanol to ice bath temperature and then slowly add it to the solution. The blue, solid complex should form. Isolate the Product. Premeasure the mass of a piece of filter paper and fit it in a Buchner funnel. Vacuum filter the solution. Wash the solid with two 5-mL portions of cold 95% ethanol and 5 ml of acetone. Place the filter and the
sample on a watchglass and allow the sample to air dry. Determine the mass of the filter paper and the product. Calculate the percent yield. Obtain the Instructor s Approval. Transfer your product to a clean, dry test tube, stopper and properly label the test tube, and submit it to your instructor for approval. Synthesis of Hexaamminenickel (II) Chloride: Form the Complex Ion. In a 125-mL Erlenmeyer flask, dissolve 6 g of nickel(ii) chloride hexahydrate in no more than 10 ml of warm deionized water. In a fume hood, slowly add 20 ml of conc NH 3. Precipitate the Complex Ion. Cool the mixture in an ice bath. Cool 15 ml of 95% ethanol to ice bath temperature and then slowly add it to the solution. Allow the mixture to settle for complete precipitation of the lavender product. The supernatant should be nearly colorless. Isolate the Product. Premeasure the mass of a piece of filter paper, fitted for a Buchner funnel. Vacuum filter the product; wash with two 5-mL volumes of cold 95% ethanol and 5 ml of acetone. Air dry the product and filter paper on a watchglass. Determine the mass of the filter paper and product. Calculate the percent yield. Obtain the Instructor s Approval. Transfer your product to a clean, dry test tube, stopper and properly label the test tube, and submit it to your instructor for approval. Synthesis of Tris(ethylenediamine)nickel (II) Chloride Dihydrate: Form the Complex Ion. In a 125-mL Erlenmeyer flask, dissolve 6 g of nickel(ii) chloride hexahydrate in 10 ml of warm deionized water. Cool the mixture in an ice bath. In a fume hood, slowly add 10 ml ethylenediamine. Precipitate the Complex Ion. Cool the mixture in an ice bath. Cool 15 ml of 95% ethanol to ice bath temperature and then slowly add it to the solution. Allow the mixture to settle for complete precipitation of the lavender product. The supernatant should be nearly colorless. Isolate the Product. Premeasure the mass of a piece of filter paper, fitted for a Buchner funnel. Vacuum filter the product; wash with two 5-mL volumes of cold 95% ethanol and 5 ml of acetone. Air dry the product and filter paper on a watchglass. Determine the mass of the filter paper and product. Calculate the percent yield. Obtain the Instructor s Approval. Transfer your product to a clean, dry test tube, stopper and properly label the test tube, and submit it to your instructor for approval.
Prelaboratory Assignment NAME: Answer the following questions. 1. Consider the coordination compound, [Cr(NH 3 ) 2 Cl 3 (H 2 O)]Br. Use the definitions in the introduction to identify the following with the formula and charge as applicable. a. the ligand b. the complex c. the coordination sphere d. the coordination number of chromium 2. Ethylenediamine is a bidentate ligand and the ethylenediaminetetraacetate ion, EDTA 4, is a hexadentate ligand, both of which are chelating agents. a. What is the difference between a bidentate and a hexadentate ligand? Explain. b. What is a chelating agent? c. If the coordination number of a metal ion is 6, how many ethylenediamine molecules and how many ethylenediaminetetraacetate ions will bind the metal ion in a complex? Explain. 3. Ammonia is a stronger ligand than is the chloride ion. What does this statement mean?
Prelaboratory Assignment (continued) NAME: 4. Write the formula and the charge of the complex formed between the following. a. chromium(iii) and ammonia, assuming a coordination number of 6 b. iron(iii) and cyanide ion, assuming a coordination number of 6 c. iron(iii) and ethylenediamine, assuming a coordination number of 6 5. What is the coordination number of the metal ion in each of the following complexes? a. [Ag(S 2 O 3 ) 2 ] 3 b. [Pt(NH 3 ) 2 Cl 2 ]