Preparation and Characterization of Nickel and Cobalt Complexes 1

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1 Preparation and Characterization of Nickel and Cobalt Complexes 1 Frank Walmsley, Department of Chemistry, Trinity University, San Antonio, TX Nickel(II) commonly forms complexes with three different geometries: octahedral, tetrahedral, and square planar. Some five-coordinate complexes are known but are rare. Nickel(II) is a 3d 8 system so octahedral and tetrahedral complexes will have 2 unpaired electrons and square planar complexes usually will have none. Octahedral complexes can be prepared from both strong field and weak field ligands (or a mixture of both). Among four-coordinate nickel(ii) complexes, those with strong field ligands tend to be square planar and those with weak field ligands tend to be tetrahedral. Cobalt(III) complexes are almost always low-spin octahedral. Cobalt(II) forms both tetrahedral and octahedral complexes; square planar complexes are very rare. As a 3d 7 system, cobalt(ii) complexes have 3 unpaired electrons. With moderate to strong-field ligands, cobalt(ii) complexes tend to be easily oxidized to cobalt(iii). When the ligands can bind in only one way, the coordination number of the complex can be determined from the formula. NH 3 and similar ligands can only be monodentate, o-phenanthroline can only be bidentate, and so forth. But some ligands can bond in more than one way. For example, NO 3 - can be uncoordinated (present as an ion), monodentate (coordinating through one oxygen), or bidentate (coordinating through two oxygens). Thus when nitrate is present in a complex, its contribution to the coordination number can be 0, 1, or 2 for each NO 3 - and all nitrates do not need to be the same. In such instances, information other than the formula must be available in order to determine the complex s coordination number. Similar to nitrate (NO 3 - ), carbonate (CO 3 2- ) can be ionic, monodentate or bidentate. Some ligands can bond through different atoms; they are ambidentate. For example, a sulfoxide (R 2 SO) can coordinate either through the sulfur or through the oxygen but not through both. The NO 2 - ion can bond through the nitrogen or through an oxygen but not both to the same metal center. The thiocyanate ion, SCN -, can coordinate through the sulfur or through the nitrogen but only through both when it bridges between two metal centers. The complexes to be studied in this experiment are of known formula: Ni((C 6 H 5 ) 3 P) 2 (SCN) 2, Ni((C 6 H 5 ) 3 P) 2 (NO 3 ) 2, Ni(NH 2 CH 2 CH 2 NH 2 ) 2 (NO 3 ) 2, Ni(NH 2 CH 2 CH 2 NH 2 ) 2 (NO 3 )I, Co((C 6 H 5 ) 3 PO) 2 (NO 3 ) 2, Co(NH 3 ) 4 (CO 3 )(NO 3 ), Ni(C 5 H 5 N) 3 (NO 3 ) 2, and Co(C 5 H 5 N) 3 (NO 3 ) 2. In the case of the thiocyanate complex, the questions are (1) whether the thiocyanate ligand is bonded through the sulfur or through the nitrogen, and (2) whether the nickel geometry is tetrahedral or square planar or octahedral. Since there are four monodentate ligands present, two (C 6 H 5 ) 3 P and two SCN -, the coordination number is fixed. In the nitrate complexes, the number of ligands is fixed but the coordination number is not known since the nitrate ion can be present in three possible ways. Thus we need to know how the nitrate coordinates in order to determine the coordination number. In the complex with the carbonate, the question is similar to that for the nitrate since it, too, can coordinate in the same 1

2 three ways. How can the structure of a coordination compound be determined in the absence of X-ray crystal data? Enough examples have been studied to permit determination in many cases using (1) infrared spectra, (2) visible-near IR spectra, (3) magnetic susceptibility, and (4) conductance. In this experiment you will prepare two complexes and determine the structures by these indirect means: infrared spectra, color, magnetic susceptibility, and conductance. The following discussion shows how these techniques are applied to the type of nickel(ii) and cobalt complexes being prepared. 1. Conductance. When a coordination compound contains "anions" (such as NO 3 - or Cl - ), those anions may or may not be coordinated. If they are coordinated, they will remain coordinated when the complex 0is dissolved in a noncoordinating solvent. If they are not coordinated, they will be free solvated ions when the complex is dissolved. In the complexes you will prepare, there are these possibilities. If the ions are coordinated, the complexes will be nonelectrolytes. If any of the ions are not coordinated, the complexes will be composed of ions and will be electrolytes. A table of values in several solvents is given in Table 3. In the nomenclature of Electrolyte Type, NaCl would be 1:1, CaCl2 or Na 2 SO4 would be 2:1, etc. Nonelectrolytes do not have zero conductance; the conductance is low. Table 3. Molar conductance ranges for electrolytes at 1 x 10-3 M in some common solvents at 25 o C. Solvent Electrolyte Type 1:1 2:1 3:1 4:1 Nitromethane Acetonitrile Methanol Acetone Water ~ Magnetic Susceptibility. The magnetic moments due to electron spin of unpaired electrons may be calculated from the equation μ = (n(n+2)) ½. where n is the number of unpaired electrons. For some metal ion systems, there may be an orbital contribution to the magnetic moment, which makes the measured value somewhat larger. Nickel(II). Square planar nickel(ii) complexes usually have no unpaired electrons, are diamagnetic and thus have a magnetic moment of zero. Nickel(II) with 2 unpaired electrons would have a calculated spin- 2

3 only magnetic moment of 2.83 μ B. Octahedral nickel(ii) complexes have magnetic moments between 2.9 and 3.4 μ B which indicates a small but definite orbital contribution to the magnetic moment. Tetrahedral complexes, however, range from 3.0 to 4.0 μ B ; the larger the distortion from a regular tetrahedron, the lower the magnetic moment. Thus if a magnetic moment is found between 3.0 and 3.4 μ B, this information alone is not sufficient to distinguish between octahedral and tetrahedral nickel(ii). If the coordination number is known from other data to be four, then the magnetic moment can be used to distinguish between tetrahedral and square planar. Cobalt(II). Cobalt(II) with 3 unpaired electrons would have a calculated spin-only magnetic moment of 3.87 μ B. Octahedral complexes typically have magnetic moments between 4.7 and 5.2 μ B. Tetrahedral complexes are typically below 4.7 μ B but may be higher for stronger field ligands. Thus magnetic moments give some indication of geometry but are not completely reliable. 3. Color. The characteristic absorptions of the complexes give colors that may be indicative of structure. Square planar nickel complexes are usually red or yellow or brown. Tetrahedral nickel complexes are usually blue or green. Octahedral nickel complexes are typically blue or green or yellow and are less intensely colored than either square planar or tetrahedral nickel complexes. Tetrahedral cobalt(ii) complexes are usually blue but may be violet. Octahedral cobalt(ii) complexes are usually pink but may be violet or even blue. Cobalt(III) complexes have a wide variation in color. 4. Infrared Spectra. When the fundamental IR bands of the NO 3 - group can be identified, it is possible to tell if the NO 3 - group is ionic or coordinated as shown in Table 1. If the NO 3 - group is coordinated, these bands can not distinguish between monodentate and bidentate coordination. However, the combination bands in the cm -1 region are able to make these distinctions. Ionic nitrates have a single absorption in this region while coordinated nitrate has two absorptions. For nickel(ii) and cobalt(ii), if the NO 3 - group is monodentate the two absorptions are separated by about 24 cm -1 while for bidentate NO 3 - the separation is about 52 cm -1. These absorptions are very weak and can only be used for diagnosis when there are no other absorptions in this region. An FTIR instrument is able to expand this portion of the spectrum to make identification of the bands possible. For examples of the use of this technique see references 2, 3, and 4. Figure 1 shows the IR spectra of two nickel complexes taken from Reference 3. The left spectrum is of an ionic nitrate and shows only 1 absorption in the cm -1 region. The right spectrum is of a monodentate nitrate and shows 2 absorptions separated by 23 cm -1. There may be other bands in this region; look for the strongest bands among these weak bands and look for a single strong one, two of about equal intensity separated by about 24 cm -1 or two of about equal intensity separated by about 52 cm -1 Figure 1. IR absorptions in the cm -1 region. a. [Ni(phen)(H 2 O) 4 ](NO 3 ) 2?H 2 O and 3 b. [Ni(bipy)(H 2 O) 3 (NO 3 )](NO 3 )

4 Table 1 Infrared absorptions of ionic and coordinated nitrate. Ionic NO 3 - Assignment NO stretch Out-of-plane NO 2 stretch NO 2 bend Frequency/cm Coordinated NO 3 - Assignment NO stretch Out-of-plane NO 2 stretch NO 2 stretch asymmetric NO 2 bend sym asym Frequency/cm ~739 ~713 The carbonate ion is isoelectronic with the nitrate ion and can coordinate in the same ways. The mode of binding of the carbonate ion can be determined from the fundamental vibrations; that is, it is not necessary to resort to looking for weak combination bands. Table 2 lists the fundamental bands for ionic, monodentate, and bidentate carbonate. Table 2. Infrared absorptions for ionic and coordinated carbonate in cm Vibration Assignment Ionic Monodentate Bidentate , , , The thiocyanate group can be bonded through the sulfur, the nitrogen, or both. The CN stretching frequency can be used for characterization. N-bonded complexes have this absorption near and below 2050 cm -1 while S-bonded complexes are near 2100 cm -1. Bridging SCN - exhibits this CN band well above 2100 cm -1. This absorption is quite strong so it is usually easy to identify. The NCS bending (deformation) absorption is sometimes helpful. For N-bonding, there is a single weak, sharp peak near 480 cm -1. For S-bonding, there are several weak bands near 420 cm -1. The CS stretching band is weak and often obscured by other bands so is not usually useful. 4

5 5. Visible-Near IR Spectra. Different geometries of nickel(ii) have quite different spectra in the visible and near infrared portions of the spectrum. The absorptions in this region are due to d-d transitions and are quite weak compared to typical ultraviolet spectra. Octahedral nickel(ii) has three absorptions: nm (9,000-11,000 cm -1 ); nm (14,000-18,000 cm -1 ); and nm (25,000-30,000 cm -1 ). The molar absorptivities of these absorptions are typically less than 10. Tetrahedral nickel(ii) typically has a similar spectrum but shifted to lower energies. A characteristic absorption is found at about 670 nm (15,000 cm -1 ) with a molar absorptivity of about 100. Square planar nickel(ii) has a more complex spectrum but typically has an absorption between 450 and 600 nm (16,500-22,000 cm -1 with a molar absorptivity near 60. In order to determine molar absorptivities, the complex must be sufficiently soluble in a noncoordinating solvent to measure both the band positions and absorbances. Octahedral cobalt(ii) complexes typically have a band with its maximum between 500 and 600 nm (16,500-20,000 cm -1 with a molar absorptivity of about 10. This band may have a slight shoulder on the high energy side. Tetrahedral cobalt(ii) typically has a band with its maximum between 600 and 700 nm 14,300-16,500 cm -1 with a molar absorptivity between 500 and This band also has some fine structure usually with 3 visible maxima on the large envelope. PROCEDURE Synthesis Ni((C 6 H 5 ) 3 P) 2 (NO 3 ) 2. 6 Work in a hood for this synthesis. Dissolve 1.48 g of Ni(NO 3 ) 2. 6H 2 O in about 1 ml of distilled water. The dissolution may be slow but do not use more water. When the salt is dissolved, add 25 ml of glacial acetic acid [CAUTION: skin irritant]. In another beaker dissolve 2.68 g of (C 6 H 5 ) 3 P [CAUTION: do not inhale or ingest] in about 15 ml of glacial acetic acid. The solution may be warmed to speed dissolution; if you do warm it, allow it to cool somewhat before proceeding. Add the triphenylphosphine solution to the nickel nitrate solution, mix well, and warm near 50 o C (50 o C is not too hot to touch) for about 5 minutes. If after cooling, a white precipitate forms, reheat the solution for another 5 minutes. Cover the container with Parafilm and allow it to stand in the hood at least overnight, preferably for several days. Label your beaker with your name, the course number, and the comment: "CAUTION: glacial acetic acid; skin irritant". Suction filter the solid on a Buchner funnel with coarse filter paper. Wash with about 20 ml of glacial acetic acid [CAUTION: skin irritant] and draw air through the crystals for several minutes. Wash with one 10 ml portion of ethanol followed by two 5 ml portions of acetone. Air dry (thoroughly) the crystals. Ni((C 6 H 5 ) 3 P) 2 (SCN) 2. 6 Work in a hood for this synthesis. Place 75 ml of n-butanol in a 250 ml round bottom flask. Add 1.42 g of Ni(NO 3 ). 2 6H 2 O and 1.47 g of KSCN. Reflux the mixture for 2 hours. Allow the solution to cool; some solid will form. Gravity filter the solution through coarse filter paper discarding the solid. Reheat the solution and add 2.67 g of (C 6 H 5 ) 3 P [CAUTION: do not inhale or ingest]. Heat for 5 minutes then allow to cool using an ice bath if desired. The color of the solution will be pale when precipitation is complete. Clean the filter flask that you will use. Filter off the crystals using a Buchner funnel and medium filter paper. If some of the crystals pass through the filter paper, 5

6 refilter through the same piece of filter paper. Allow the crystals to air dry in your drawer until the next period or dry them in a 70 o C oven. It is important that the crystals be thoroughly dry. Ni(NH 2 CH 2 CH 2 NH 2 ) 2 (NO 3 ) 2. 7 In a 150 ml or 250 ml beaker, dissolve 5.8 g of Ni(NO 3 ) 2. 6H 2 O in 35 ml of methanol. Heat the solution to boiling to aid dissolution. To the hot solution, add 2.7 ml of ethylenediamine dropwise. Do this in a hood since ethylenediamine is similar to concentrated aqueous ammonia. Add the liquid SLOWLY - the reaction is exothermic and the methanol can overheat and boil over. After all the ethylenediamine has been added, cool the solution in ice. Filter the crystals on a Buchner funnel. Wash the crystals with 5 ml of cold methanol. Dry in the air. Ni(NH 2 CH 2 CH 2 NH 2 ) 2 (NO 3 )I. 7 In a 150 ml or 250 ml beaker, dissolve 5.8 g of Ni(NO 3 ) 2. 6H 2 O in 35 ml of methanol. Heat the solution to boiling to aid dissolution. To the hot solution, add 2.7 ml of ethylenediamine dropwise. Do this in a hood since ethylenediamine is similar to concentrated aqueous ammonia. Add the liquid SLOWLY - the reaction is exothermic and the methanol can overheat and boil over. After all the ethylenediamine has been added, add 3.0 g of NaI to the hot solution. Keep the mixture hot until all the NaI has dissolved. Allow the solution to cool to room temperature; do not cool in ice. Collect the crystals on a Buchner funnel and wash with 5 ml of methanol. Dry in the air. Co(NH 3 ) 4 (CO 3 )(NO 3 ). 8 Dissolve 10.0 g of (NH 4 ) 2 CO 3 in 30 ml H 2 O. Add 30 ml of concentrated NH 3 to this solution; the NH 3 (aq) can be added before all the solid is dissolved. In a separate beaker, dissolve 7.5 g Co(NO 3 ) 2. 6H 2 O in 15 ml H 2 O. Add the (NH 4 ) 2 CO 3 solution to the Co(NO 3 ) 2 solution with stirring. Oxidize the cobalt(ii) to cobalt(iii) by slowly adding, with stirring, 12 ml of 10% H 2 O 2. Heat the resulting solution on a hot plate to decompose the excess H 2 O 2. Then evaporate the solution to about 50 ml without boiling. Cool the solution in an ice bath. Vacuum filter the crystals and wash them once with 5 ml cold H 2 O and twice with 10 ml 95% ethanol. Dry the crystals in air. Co((C 6 H 5 ) 3 PO) 2 (NO 3 ) 2. 9,10 Dissolve 2.78 g of triphenylphosphine oxide in about 40 ml of acetone. In a separate beaker, dissolve 1.46 g of Co(NO 3 ) 2 2 O in about 20 ml of acetone. Pour the (C 6 H 5 ) 3 PO solution into the Co(NO 3 ) 2 solution with stirring. Allow the solution to stand for about 1 hour for precipitation to proceed. Vacuum filter the crystals and wash them once with 20 ml acetone. Ni(C 5 H 5 N) 3 (NO 3 ) 2 and Co(C 5 H 5 N) 3 (NO 3 ) 2. 2, Place 1.15 g of Ni(NO 3 ) 2 6H 2 O or 1.15 g of Co(NO 3 ) 2 6H 2 O in a 50 ml beaker with a stir bar. Add 10 ml of 2,2-dimethoxypropane and stir for 40 minutes. 2,2-Dimethoxypropane is a dehydrating agent and will form a light colored precipitate of the anhydrous metal nitrate. Add 5 ml of pyridine (in a hood; wear gloves) and continue stirring for an additional 15 minutes. A precipitate of the desired complex will form. Vacuum filter the crystals and wash them once with 5 ml of acetone. Dry the crystals in the air (preferably in a hood) for at least 3 days. 6

7 Characterization For each compound, (1) Measure the melting point. (2) Measure the magnetic moment as directed by your instructor. (3) Record the infrared spectrum using the Attenuated Total Reflection (ATR) attachment for the infrared spectrometer. For best results, the sample must be finely ground and will require at least 200 scans. Identify the positions of all relevant bands. Note: you need to know what regions to examine before you mark the bands in your spectrum! (4) Measure the conductance as a 1.00 x 10-3 M solution in an appropriate solvent. (5) Note the color. In the IR spectrum of the thiocyanate complex, note particularly the positions of the CN stretch and the NCS deformation. In the IR spectrum of the nitrate complex, expand the cm -1 region. You may also want to try to identify the NO 3 - bands according to Table 1 which is taken from Reference 16. Look for the CO 3 2- bands according to Table 2. For the phosphine oxide complex, look for the P=O stretching bands. Conductance x 10-3 M solutions will be prepared for you. Note the solvent used for each. Set the scale to 200. Turn the instrument to the conductance (micromho) setting. Place the probe into the solution to be measured and gently stir the solution with the probe or swirl the solution with the probe clamped in place. It is important to continually refresh the solution at the electrode surfaces while taking a reading. Read the meter. Remove the probe from the solution and gently dry the probe with Table 4. Conductance Solvents Complex Solvent Ni(en) 2 (NO 3 ) 2 nitromethane Ni(en) 2 (NO 3 )I insoluble Co(OPPh 3 ) 2 (NO 3 ) 2 nitromethane Co(NH 3 ) 4 (NO 3 )(CO 3 ) water Ni(PPh 3 ) 2 (NO 3 ) 2 nitromethane Ni(PPh 3 ) 2 (SCN) 2 Ni(py) 3 (NO 3 ) 2 Co(py) 3 (NO 3 ) 2 acetonitrile nitromethane nitromethane a KimWipe. Since the solution is 1.00 x 10-3 M in compound, the cell constant is 1.000, and since the meter measures in micromho (1 mho = 1 ohm -1 ), the reading is equal to the molar conductance. The conductance of pure acetonitrile is 4.4 micromho. A 1:1 electrolyte in acetonitrile under the same conditions as this experiment gave a reading of micromho. Magnetic Moment. Turn the instrument on and choose the X1 scale. Adjust the reading to zero. If the reading drifts from zero during the experiment, rezero the instrument. Place a clean empty tube into the instrument and record the reading, R o. Weigh the empty tube. Grind the sample lightly to crush any lumps. Pack the tube with sample. Add a small amount of solid and tap the tube gently on the bench. Repeat this procedure until the sample height in the tube is at least 2 cm but no more than 3.5 cm. Measure the height of the sample, h, in centimeters as accurately as possible. Weigh the sample and tube. Calculate the mass of the sample, m, in grams. Place the packed tube into the instrument and record the reading, R. Measure the room temperature. Empty the tube with gentle tapping. Rinse the remaining sample 7

8 from the tube with a stream of water from a wash bottle followed by rinsing with acetone. The mass susceptibility is calculated as followed; the instrument constant, C, has been previously determined and is posted on the magnetic susceptibility instrument. C x h x (R - R o ) χ g = (1 x 10 9 ) x m Calculate the molar susceptibility, χ M. χ M = χ g x Formula Weight The molar susceptibility includes both the paramagnetism of the unpaired electrons and the diamagnetism of the atoms. Correct for this diamagnetism using Pascal's constants from Table 4. Calculate the corrected molar susceptibility, χ ' M. χ ' M = χ M - correction. The correction is the sum of the Pascal's constants for all components of the formula. When there is more than one of a component, the correction must be applied once for each one; if there are two nitrates then multiply the nitrate correction by two. The magnetic moment (sometimes called the effective magnetic moment) is calculated from: μ eff = 2.84 x ( χ ' M x T) ½ where T is the kelvin temperature. Table 5. Molar Diamagnetic Corrections (Pascal Constants) 16 Ni x 10-6 (C 6 H 5 ) 3 P -167 x 10-6 Co x 10-6 NO x 10-6 (C 6 H 5 ) 3 PO -172 x 10-6 Co x 10-6 en x 10-6 SCN x 10-6 NH 3-18 x 10-6 I x 10-6 CO x 10-6 Py -49 x 10-6 Results. Obtain all necessary data. Calculate the theoretical yield and the percent yield for each of the two syntheses. On the basis of the data obtained, determine the structure of each compound. A data sheet will be provided to guide you through the assignment and calculations. Be sure to attach calculations, spectra, and relevant IR band assignment information to the data sheet. References 1. Eagle, C.T.; Walmsley, F. J. Chem. Educ. 1991, 68,

9 2. Lever, A.B.P.; Mantovani, E.; Ramaswamy, B.S. Can. J. Chem. 1971, 49, Walmsley, F.; Pinkerton, A.A.; Walmsley, J.A. Polyhedron 1989, 8, Wheeler, M.T.; Walmsley, F. Thermochim. Acta 1986, 108, Nakamoto, K.; Fujita, J.; Tanaka, S.; Kobayashi, M. J. Am. Chem. Soc. 1957, 79, Venanzi, L.M. J. Chem. Soc. 1958, Curtis, N.F.; Curtis, Y.M. Inorg. Chem. 1965, 4, Schlessinger, G. Inorg. Syn. 1960, 6, Bannister, E.; Cotton, F.A. J. Chem. Soc. 1960, Dias Rodrigues, A.M.G.; Francisco, R.H.P.; Lechat, J.R. Cryst. Struct. Comm. 1982, 11, Rosenthal, M.R.; Drago, R.S. Inorg. Chem. 1965, 4, Biagetti, R.V.; Haendler, H.M. Inorg. Chem. 1966, 5, Atwood, J.L.; Bott, S.G.; Vincent, R.L. J. Crystallogr. Spectrosc. Res. 1990, 20, Choca, M.; Ferraro J.R.; Nakamoto, K. J. Chem. Soc., Dalton Trans. 1972, Herlocker, D.W.; Rosenthal, M.R. Inorg. Chim. Acta 1970, 4, Bain, G.A.; Berry, J.F. J. Chem. Educ. 2008, 85, Addison, C.C.; Gatehouse, B.M. J. Chem. Soc. 1960, 613. Frank Walmsley