Preparation and Characterization of Nickel and Cobalt Complexes1
Frank Walmsley, Department of Chemistry, Trinity University, San Antonio, TX 78212.
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 3d8 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 3d7 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. NH3 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, NO3- 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 NO3- 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 (NO3-), carbonate (CO32-) can be ionic, monodentate or bidentate.
Some ligands can bond through different atoms; they are ambidentate. For example, a sulfoxide
(R2SO) can coordinate either through the sulfur or through the oxygen but not through both. The NO2ion 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((C6H5)3P)2(SCN)2,
Ni((C6H5)3P)2(NO3)2, Ni(NH2CH2CH2NH2)2(NO3)2, Ni(NH2CH2CH2NH2)2(NO3)I,
Co((C6H5)3PO)2(NO3)2, Co(NH3)4(CO3)(NO3), Ni(C5H5N)3(NO3)2, and Co(C5H5N)3(NO3)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 (C6H5)3P 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
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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 NO3- 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 Na2SO4 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 25oC.
Solvent
Electrolyte Type
1:1
2:1
3:1
4:1
Nitromethane
75-95
150-180
220-260
290-330
Acetonitrile
120-160
220-300
340-420
Methanol
80-115
160-220
Acetone
100-140
160-200
Water
118-130
235-275
410-435
~560
2. 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 spin2
only magnetic moment of 2.83 μB. Octahedral nickel(II) complexes have magnetic moments between 2.9
and 3.4 μB which indicates a small but definite orbital contribution to the magnetic moment. Tetrahedral
complexes, however, range from 3.0 to 4.0 μB; 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 μB, 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 μB. Octahedral complexes typically have magnetic moments between 4.7 and 5.2 μB. Tetrahedral
complexes are typically below 4.7 μB 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 NO3- group can be identified, it is possible
to tell if the NO3- group is ionic or coordinated as shown in Table 1. If the NO3- group is coordinated,
these bands can not distinguish between monodentate and bidentate coordination. However, the
combination bands in the 1700 - 1800 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 NO3- group is monodentate the two absorptions are separated by about 24 cm-1 while for
bidentate NO3- 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
1730-1800 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
Figure 1. IR absorptions in the 1700 - 1800 cm-1 region.
about 52 cm-1
a. [Ni(phen)(H2O)4](NO3)2? H2O and
b. [Ni(bipy)(H2O)3 (NO3)](NO3)
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Table 1 Infrared absorptions of ionic and coordinated nitrate.
Ionic NO3Assignment
NO stretch
Frequency/cm-1
1050
Out-of-plane NO2 stretch
831
1390
NO2 bend
720
Coordinated NO3Assignment
NO stretch
Frequency/cm-1
1034-970
Out-of-plane NO2 stretch
800-781
1290-1253
NO2 stretch
asymmetric
1531-1481
NO2 bend
sym
asym
~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-1.5
Vibration
Assignment
Ionic
Monodentate
Bidentate
1
-
1070
1020
2
870
850
835
3
1390,1370
1450,1370
1590,1260
4
690
750
755
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.
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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: 9001100 nm (9,000-11,000 cm-1); 550-700 nm (14,000-18,000 cm-1); and 325-400 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 1000. This band also has some fine
structure usually with 3 visible maxima on the large envelope.
PROCEDURE
Synthesis
Ni((C6H5)3P)2(NO3)2.6 Work in a hood for this synthesis. Dissolve 1.48 g of Ni(NO3)2.6H2O 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 (C6H5)3P [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 50oC (50oC 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((C6H5)3P)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(NO3)2.6H2O 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 (C6H5)3P [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,
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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 70oC oven. It is important that the crystals be thoroughly dry.
Ni(NH2CH2CH2NH2)2(NO3)2.7 In a 150 mL or 250 mL beaker, dissolve 5.8 g of Ni(NO3)2. 6H2O 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(NH2CH2CH2NH2)2(NO3)I.7 In a 150 mL or 250 mL beaker, dissolve 5.8 g of Ni(NO3)2. 6H2O 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(NH3)4(CO3)(NO3).8 Dissolve 10.0 g of (NH4)2CO3 in 30 mL H2O. Add 30 mL of concentrated NH3 to
this solution; the NH3 (aq) can be added before all the solid is dissolved. In a separate beaker, dissolve
7.5 g Co(NO3)2.6H2O in 15 mL H2O. Add the (NH4)2CO3 solution to the Co(NO3)2 solution with stirring.
Oxidize the cobalt(II) to cobalt(III) by slowly adding, with stirring, 12 mL of 10% H2O2. Heat the
resulting solution on a hot plate to decompose the excess H2O2. 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 H2O and twice with 10 mL 95% ethanol. Dry the crystals in air.
Co((C6H5)3PO)2(NO3)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(NO3)2
2O in about 20 mL of acetone. Pour the (C6H5)3PO
solution into the Co(NO3)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(C5H5N)3(NO3)2 and Co(C5H5N)3(NO3)2.2, 11-15 Place 1.15 g of Ni(NO3)2∙6H2O or 1.15 g of
Co(NO3)2∙6H2O 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.
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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 1700 - 1800 cm-1 region. You
may also want to try to identify the NO3- bands according to Table 1 which is taken from Reference 16.
Look for the CO32- bands according to Table 2. For the phosphine oxide complex, look for the P=O
stretching bands.
Table 4. Conductance Solvents
Conductance. 1.00 x 10-3 M solutions will be prepared
Complex
Solvent
for you. Note the solvent used for each. Set the scale to
Ni(en)2(NO3)2
nitromethane
200. Turn the instrument to the conductance (micromho)
Ni(en)2(NO3)I
insoluble
Co(OPPh
)
(NO
)
nitromethane
setting. Place the probe into the solution to be measured
3 2
3 2
Co(NH3)4(NO3)(CO3)
water
and gently stir the solution with the probe or swirl the
Ni(PPh
)
(NO
)
nitromethane
3 2
3 2
solution with the probe clamped in place. It is important
Ni(PPh
)
(SCN)
acetonitrile
3
2
2
to continually refresh the solution at the electrode
Ni(py)3(NO3)2
nitromethane
surfaces while taking a reading. Read the meter. Remove
Co(py)3(NO3)2
nitromethane
the probe from the solution and gently dry the probe with
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 119.5 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, Ro.
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
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 - Ro)
χg = ----------------------(1 x 109) x m
Calculate the molar susceptibility, χM.
χM = χg 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 = χM - 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:
μeff = 2.84 x ( χ'M x T)½
where T is the kelvin temperature.
Table 5. Molar Diamagnetic Corrections (Pascal Constants)16
Ni2+
-12 x 10-6
(C6H5)3P
-167 x 10-6
Co3+
-12 x 10-6
NO3-
-18.9 x 10-6
(C6H5)3PO
-172 x 10-6
Co2+
-12 x 10-6
en
-46.5 x 10-6
SCN-
-31 x 10-6
NH3
-18 x 10-6
I-
-50.6 x 10-6
CO32-
-28.0 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, 336.
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2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
Lever, A.B.P.; Mantovani, E.; Ramaswamy, B.S. Can. J. Chem. 1971, 49, 1957.
Walmsley, F.; Pinkerton, A.A.; Walmsley, J.A. Polyhedron 1989, 8, 689.
Wheeler, M.T.; Walmsley, F. Thermochim. Acta 1986, 108, 325.
Nakamoto, K.; Fujita, J.; Tanaka, S.; Kobayashi, M. J. Am. Chem. Soc. 1957, 79, 4904.
Venanzi, L.M. J. Chem. Soc. 1958, 719.
Curtis, N.F.; Curtis, Y.M. Inorg. Chem. 1965, 4, 804.
Schlessinger, G. Inorg. Syn. 1960, 6, 173.
Bannister, E.; Cotton, F.A. J. Chem. Soc. 1960, 2276.
Dias Rodrigues, A.M.G.; Francisco, R.H.P.; Lechat, J.R. Cryst. Struct. Comm. 1982, 11, 847.
Rosenthal, M.R.; Drago, R.S. Inorg. Chem. 1965, 4, 840.
Biagetti, R.V.; Haendler, H.M. Inorg. Chem. 1966, 5, 383.
Atwood, J.L.; Bott, S.G.; Vincent, R.L. J. Crystallogr. Spectrosc. Res. 1990, 20, 631.
Choca, M.; Ferraro J.R.; Nakamoto, K. J. Chem. Soc., Dalton Trans. 1972, 2297.
Herlocker, D.W.; Rosenthal, M.R. Inorg. Chim. Acta 1970, 4, 501.
Bain, G.A.; Berry, J.F. J. Chem. Educ. 2008, 85, 532.
Addison, C.C.; Gatehouse, B.M. J. Chem. Soc. 1960, 613.
©Frank Walmsley 2010
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