Experiment 4 Studying the Spectrochemical Series: Crystal Fields of ...

Experiment 4
Studying the Spectrochemical Series: Crystal Fields of Cr(III)
Introduction
Coordination compounds of transition metals are often highly colored. The color results
from absorption of light at specific wavelengths of visible light associated with electronic
transitions within the d-orbitals. Thus, these d-d transitions give many transition metal ions their
characteristic color (eg: cobalt blue).
The d orbitals of a metal ion in an octahedral crystal field (surrounded by an octahedral
array of ligands) are split into a higher energy eg set and a lower energy t2g set (Figure 1). This is
due to electron clouds around each ligand (L) destabilizing those d-orbitals that lie along the X, Y,
and Z axes. The energy difference between the upper and lower energy levels is designated as Δo
(pronounced del-oh) or 10Dq.
Figure 1. d-orbitals split by an octahedral crystal field.
The degree of splitting of the d orbitals and hence the magnitude of Δo depends on several
factors, but the most important are the charge on the metal and the identity of the ligand.
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Understanding the trends in ligand-field splitting is simplified considerably by considering a series
of complexes with the same metal in a given oxidation state; the only major variable in this case is
the ligand identity. From a large number of studies it is known that ligands can be arranged in a
sequence according to their ability to cause d-orbital splitting. This series is known as the
spectrochemical series:
halides < NCS¯ < OH¯ < oxalate < H2O < NCS¯ < pyridine < NH3 < en < NO2¯ <
CN¯ < CO
(Underlined atom is the one coordinated to the metal).
Note the general trend shown here: halide ligands are weaker-field ligands than sulfur
donors, which are weaker field than oxygen donor ligands, which are weaker-field than nitrogen
donors, which are weaker-field than multiply-bonded ligands. The stronger the interaction of the
ligand with the metal's d-orbitals, the stronger the field strength of that ligand.
The magnitude of ligand field strength increases by a factor of about two as one moves from
halide to CN¯ in the spectrochemical series.
The structure of the oxalate ion, C2O4 2- , is shown below. Note the resonance forms of this
compound.
O
O
C C
-O O- Figure 2. Structure of the oxalate ion
- O
O
C C
O O- - O
O
C
- O
C
O
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The objective of this experiment is to quantify Δo for a series of Cr(III) complexes by
electronic absorption spectroscopy. Cr(III) compounds are d 3 , and their electronic spectral
characteristics are reasonably easy to interpret.
Because of electronic selection rules (see lecture), the absorption band corresponding to the
energy of the crystal field strength (Δo) will be the one at the longest wavelength (lowest energy) in
the spectrum, and it should be more intense than any other nearby transition.
Ordering the octahedral Cr III compounds from longest to shortest wavelength will place the
ligands in order of increasing crystal field strength, as # " 1 ! o , and will allow you to build your
own spectrochemical series.
In mixed-ligand complexes, the “Rule of Average Environments” states that the observed
value of Δo in such complexes is the weighted average of Δo for each of the homoleptic (single type
of ligand) complexes. The first equation below is general, the second equation is for the specific
example of [Cr(H2O)4Cl2] + .
If the Δo's of [Cr(H2O)4Cl2] + and [Cr(H2O)6] 3+ are known, then by rearranging the equation,
you can solve for the Δo of [CrCl6] 3+ . You can use this Δo to find the Δo of other unknown
complexes.
Syntheses
ΔoMAnBm = (1/6) {nΔoMA6 + mΔoMB6} (1)
Δo[Cr(H2O)4Cl2] + = (1/6) {4Δo[Cr(H2O)6] 3+ + 2Δo[CrCl6] 3- } (2)
In this experiment, the bidentate ligand acetylacetonate (acac¯) will be generated via the
deprotonation of acetylacetone (acacH) by ammonia. The ammonia is generated by hydrolysis of
urea (Figure 3); subsequently, ammonia acts as a base to deprotonate acacH. Note the resonance
forms of acac-.
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O
H 2N NH 2
O
H3C C CH3 H2 + H 2O 2 NH 3 + CO 2
O
H
O
O
H3C C CH3 H
O
O
H3C C CH3 H
Figure 3. Top: Hydrolysis of urea. Bottom: Deprotonation of acetylacetone forms the bidentate
ligand acetylacetonate (acac-). Systematic name is 2,4-pentanedionate.
Experimental Procedure
You will work in pairs in this lab to prepare two compounds: Cr(acac)3, and
[Cr(en)3]Cl3•2H2O. One of the pair will synthesize one compound, and the other partner will
prepare the other compound. The procedure for the syntheses for *both* compounds must be
written in your prelab. You will then share experimental and spectroscopic data (for all the
compounds, not just for the two that you prepare).
Another compound, [Cr(H2O)6](NO3)3•3H2O, will be provided to you; you will
spectroscopically analyze both this complex and your starting material, CrCl3•6H2O (this compound
is actually [CrCl2(H2O)4]Cl•2H2O]).
Finally, the electronic spectrum of [Cr(NH3)5Cl]Cl2, will be provided on the class web page
for your use in the Data Analysis section of your lab report.
How to Heat a Reaction with a Sand Bath/Magnetic Stirrer:
In this experiment, both syntheses require that you heat your reaction mixtures to boiling (or
to "reflux") in a sand bath placed on top of a magnetic stirring plate. Getting the reaction mixture to
stir well while heating can be problematic, so use the following procedure:
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Place an empty beaker containing a stir bar on the magnetic stirrer, and turn on the stirrer.
Check to see that the bar can stir quickly and well. If it cannot, use a different magnetic stirrer.
Place your sand bath on top of the magnetic stirrer and bury your flask containing the stir
bar and the solution to be heated deep into the sand. Turn on the magnetic stirrer very slowly, and
try to get the stir bar to stir vigorously. There is often a problem in accomplishing this, because the
stir bar is too far away from the magnetic stirrer to respond well to it. To correct this, take sand out
of your sand bath (place it in a beaker and RETURN IT to the sand bath when you are finished)
and move your flask down towards the bottom of the bath, until it almost touches the metal bottom
of the sand bath. The closer the bottom of the flask comes to the bottom of the sand bath, the better
the stir bar will stir. Do not however, let the bottom of your flask come into direct contact with the
metal of the sand bath; the flask may melt.
Again, turn on the magnetic stirrer very slowly, and keep removing sand and moving the
flask down into the bath until the bar is stirring vigorously and well.
Now heap the extra sand back into the bath underneath and around the shoulders of the
flask, to help transfer heat and encourage quick heating of the solution. When your reaction
mixture is vigorously stirring, and the extra sand has been replaced around the shoulders of the
flask, turn on the sand bath and heat your reaction mixture.
Make sure, if possible, that your reaction mixture is stirring vigorously throughout the
heating period. If the stirring stops, repeat the above procedure until you can get efficient stirring
again.
A. Cr(acac)3 – the preparation of Tris(2,4-pentanedionate)chromium(III)
Dissolve 260 mg of CrCl3•6H2O in 4.0 mL of distilled water in a 25 mL round-bottomed
flask, and then add 1 g of urea to the flask. Measure out 0.8 mL of acetylacetone using the 1-mL
graduated pipette provided, and add the acetylacetone to the flask also. Next place a reflux
condenser on the flask, after first applying a thin layer of stopcock grease to the male ground glass
joint of the reflux condenser. Reflux your reaction, with stirring, for one hour using a sand bath. As
the urea releases ammonia and the solution becomes basic, deep maroon crystals begin to form.
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After one hour, cool the flask thoroughly in an ice bath. Very thorough cooling is needed. If
no crystals form, carefully make a small scratch on the bottom of the flask using a glass rod or
metal spatula, and cool the flask again. Collect the crystals by suction filtration and wash them with
three 0.3 mL portions of distilled water. Dry the crystals as much as possible on the filter using the
aspirator, and then collect them and spread them on a piece of filter paper to air dry. Determine the
percentage yield, and transfer the Cr(acac)3 to a labeled vial.
B. [Cr(en)3]Cl3•2H2O – the preparation of Tris(ethylenediamine)chromium(III)
Weigh out 100 mg of mossy zinc into a 25 mL round-bottomed flask. Remove the surface
layer of ZnO and generate a clean and active Zn 0 surface by washing the Zn with HCl immediately
prior to use. To accomplish this, add enough 6M HCl to cover the mossy zinc, swirl the flask briefly
(the mossy zinc may completely dissolve upon prolonged contact with HCl), and pipette off the
acid. Then wash the mossy zinc three times with 10 mL portions of methanol, to remove as much
acid as possible.
Weigh out 266 mg of CrCl3•6H2O, and add it and 1 mL of methanol to the round-bottomed
flask. In the hood, add 1 mL of ethylenediamine. Next place a reflux condenser on the flask, after
first applying a thin layer of stopcock grease to the male ground glass joint of the reflux condenser.
Reflux your reaction, with stirring, for one hour using a sand bath.
Take care not to turn up the sand bath power too high: methanol refluxes at a lower
temperature (65 o C) than water (100 o C).
Cool the purple solution in an ice bath. Collect the yellow crystalline product (it will appear
purple, because it is wet with the solution) by suction filtration using a Hirsch funnel. Remove any
un-reacted zinc with tweezers.
Wash the filtered product with 0.5 mL portions of 10% ethylenediamine in methanol until
the purple solution is gone, the product is yellow, and the washings are colorless. Follow this with a
0.5 mL rinse with ether, to rinse away the methanol and help dry the product. Dry the crystals as
much as possible on the filter using the aspirator, and then collect them and spread them on a piece
of filter paper to air dry. Determine the percentage yield, and transfer the [Cr(en)3]Cl3•2H2O to a
labeled vial.
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C. Spectroscopy of the Cr(III) Complexes.
Prepare aqueous solutions of [Cr(en)3]Cl3•2H2O and [Cr(H2O)6](NO3)3•3H2O by dissolving
approximately 8 mg of each complex in about 5 ml of distilled water. Because the analysis of Δo,
the objective of this experiment, requires only that the longest-wavelength λmax of each complex be
identified, you do not need to know the exact concentration of the solution, as you do when
calculating extinction coefficients using Beer's law.
Also prepare a similar ethanol solution of Cr(acac)3.
Make a similar solution of CrCl3•6H2O in water, but note that this complex, consisting of
[CrCl2(H2O)4] + ions in water, slowly substitutes to the hexaaquo species, [Cr(H2O)6] 3+ +, so it should
be analyzed immediately after its preparation. Do not prepare this solution until your name is called
for use of the UV-vis instrument. At that point, quickly add already-prepared portions of
CrCl3•6H2O and water to a cuvette , mix the solution a few times by drawing it up in a Pasteur
pipette, and then acquire the spectrum of this solution first.
Obtain the electronic absorption spectrum of each complex. Record the wavelengths and
absorbance values for all absorbances for each sample in your notebook. Determine and record the
longest wavelength of the absorbance peaks for each complex in units of nanometers. This is the
wavelength you will use to calculate the Δo's of the complexes in the Data Analysis section of your
lab report.
Data Analysis
1. (3 points). Draw the crystal field energy levels and electron occupancies for octahedral Cr 3+ ions.
Indicate the energy gap that corresponds to the transition that is being investigated in this lab.
2. (10 points) Prepare a table of your data, including columns for λmax in nm and ΔO in cm -1 , eV, and
kJ/mol.
a) Convert the wavelengths which correspond to Δo into wavenumbers (cm -1 ) using the
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following relationship:
Δo = [1/λ (nm)] (1 x 10 7 ) cm -1
b) Other energy units for Δo may be obtained using the following conversion factors:
1 cm -1 = 1.24 x 10 -4 eV = 0.01196 kJ/mol
c) Arrange the entries in order of increasing gap energy.
Show your complete set of calculations for one of the complexes. For the rest of the
complexes, just record the calculated values in the table.
3. (12 points) List the five ligands in order of increasing ligand field strength.
Use the Rule of Average Environments to calculate the true ligand field strengths of two of
the ligands (to be included in the list above). Show your complete set of calculations.
Discussion
1. (10 points) Discuss the order of the ligand field strengths you obtained, comparing them with the
spectrochemical series given in the introduction. Does the order of ligands obtained by this
experiment correspond to the established order of the spectrochemical series? Explain any
deviations.
2. (5 points) List the five ligands in order of the increasing strength of their interaction with Cr +3 ,
and give the reasoning behind your answer.
3. (5 points) The ligand acac¯ is not included in the spectrochemical series given in the Introduction.
Examining the chemical structure of acac - and of those of related ligands shown in the
spectrochemical series, explain where you would expect to find acac - in the spectrochemical series.
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Does your calculated ligand field strength for acac - agree with this prediction?
Conclusions (10 points): Discuss what knowledge has been gained by this experiment. Discuss the
importance to inorganic chemistry of being able to ascertain the crystal field energy gap of
complexes, and the ligand field strength of ligands. What information can be obtained, or what
advantages in other experiments can be gained, about inorganic complexes using these two
quantities?
Questions
1. (5 points) What is a significant difference that is seen between the UV-visible spectrum of
Cr(acac)3 and the spectra of the other complexes? What is the reason for this difference?
2. (5 points) High-spin Mn(II) and Fe(III) complexes are much less intensely colored than those of
Cr(III). Why are they so weakly colored?
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