Download Experiment 18 Lecture

Experiment 18 Lecture
Note: I post my lecture notes for these reasons: in case I forgot to mention some crucial points
during my actual lecture, in case students did not take notes, and to give students ideas if they
are stumped on some part of their lab reports. Please use these notes for ideas, but use your
own words when you are writing your discussion section. Do not take a large part of this
document, simply change a few words, and use it as your discussion!
1st week lecture - Coordination complex concepts and terminology
Two models that have been used to describe coordination complexes:
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Lewis model – In this model, coordination complexes are seen as the result of a Lewis
acid-base reaction. Ligands are said to attach to the metal center (the Lewis acid) with
special covalent bonds in which both of the electrons in the bonds are donated by the
ligand (the Lewis base).
Crystal field model – In this model, the metal-ligand bonds are seen as ionic bonds
between the positively-charged metal center and the negative charge from the ligands’
non-bonding electrons. Differences in spectroscopic and magnetic properties can be
explained by the energy differences on various d or f orbitals on the metal in the complex
due to interactions between d or f electrons and electric fields around the ligands.
Other important terms in this experiment:
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Coordination compound – a compound that contains a transition metal complex
Ligands – ions or molecules that bond (“coordinate”) to a metal center
Complex ion – a charged complex of a metal center and ligands
Counter ion – ion in a coordination compound that is not directly bonded to the metal
center; it serves to balance the charge of the complex ion to make a neutral compound
Coordination number – number of bonds to the metal center. Remember that some
ligands can occupy more than one coordination site, such as ethylene diamine (“en”)
which can form 2 bonds to the metal.
2nd week lecture – More theory and examples
What causes different colors?
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In Part 1 of this experiment, you synthesize the coordination compound [Co(NH3)6]Cl3,
which is orange. In Part, you synthesize [Co(NH3)5Cl]Cl2, which is violet.
What causes the color difference? One model we can use to understand this phenomenon
is the crystal field theory. This theory says that the difference in light absorption is due to
the energy splitting between d-orbitals in the complex, which is caused by differences in
the interaction between its d-electrons and the non-bonding electrons of the ligands due
to orbital shape.
D-orbital diagram (see board notes)
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A good aid to help you think about the model is the d-orbital diagram of the complex.
Let’s look at the diagram for [Co(NH3)6]Cl3.
You can start by writing out the valence electron configuration of the neutral metal atom.
For cobalt, if you look at the periodic table, you can see the configuration is [Ar]4s23d7.
Then, find out what the oxidation state of the metal is in the complex. In [Co(NH3)6]Cl3,
there are 3 Cl- counter-ions, so the complex has a +3 charge. NH3 ligands are neutral, so
the oxidation state of the Co in the complex is +3.
Thus, we must remove 3 electrons from the neutral Co. The 4s electrons come off first,
then one 3d electron. So the valence configuration will be 3d6.
The free Co+3 ion has d-orbitals that all have the same energy (degenerate)
But in the complex ion, according to the crystal field model, certain d-orbitals are higher
in energy than others, due to unfavorable interactions between the ligands and the delectrons. D-orbitals that have shapes that overlap with the electron density of the
ligands in the complex will have higher energy than those that have shapes that allow the
ligands to avoid the d-electrons. (See the figure below- the axes topped with green
negative charges represent the ligand electron density, and the wireframe shapes
represent the d-orbitals).
Image © 2002, 2007 by Stephen Lower - Simon Fraser University - Burnaby/Vancouver Canada
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The splitting between the higher energy levels (dx2-y2 and dz2) and the lower energy levels
(dxy, dyz and dxz) is the crystal field splitting energy, called Δ.
The color of a coordination complex comes from the absorbance of light by the complex,
which promotes electrons from the lower levels to the higher levels of d-orbitals. The
crystal field splitting energy is what dictates the wavelength of light absorbance of the
complex and therefore its color.
Looking at the relationship between energy and wavelength,
Δ =E=hv=hc/λ
we see that a large crystal field splitting is indicated by a shorter wavelength of light absorbed,
since the numerator is made of the constants h=Plank’s constant (6.626068 × 10-34 m2 kg / s) and
c=the speed of light in vacuum (3 X 108 m/s), and wavelength is in the denominator.
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Ligands differ in their ligand field strength, that is, their ability to increase Δ. Greater
ligand field strength gives rise to a greater crystal field spitting energy.
A general, empirically-based list that can help you predict the extent of crystal field
splitting is the spectrochemical series (see board notes or page 230 in the lab manual).
Also influenced by ligand field strength –the “spin” of the complex
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Complexes with larger crystal field splitting energies are more likely to be low spin
complexes, in which electrons will pair up in the lower level d-orbitals before they
traverse the large energy gap to get to the higher d-orbitals. A low spin complex with 6
d-electrons will be diamagnetic, because all of its d-electrons will be paired up in the
lower levels.
Complexes with small crystal field splitting energies are more likely to be high spin
complexes, in which electrons will jump up to the higher level empty d-orbitals before
they start pairing up. A high spin complex with 6 d-electrons will be paramagnetic,
because 4 of its d-electrons will be unpaired.
You have to find out if a complex is diamagnetic or paramagnetic via experiment, so the
crystal field model can only help you predict if a complex is more likely to be high spin
or low spin.
Part 1 and Part 2 Example
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[Co(NH3)6]Cl3 is orange and it has all NH3 ligands
[Co(NH3)5Cl]Cl2 is violet and it has five NH3 ligands and one Cl- ligand
NH3 has a higher ligand field strength than Cl-.
The light absorbed by an object is complementary in color to the light the object is
observed to be. (Look at the color wheel diagram on the board notes)
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The complementary color of orange is blue, and the complementary color of violet is
yellow. Blue (~450 nm) is shorter wavelength light than yellow (~580 nm)
It makes sense for the complex with stronger field ligands to absorb a shorter wavelength
of light than the complex with one weaker field ligand, because it has a greater crystal
field splitting energy.
Part 3 and Part 4 – Geometric isomers
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Geometric isomers are compounds that have the same chemical composition but a
different geometry (different arrangement of ligands)
In Part 3, you synthesize trans-[Co(en)2Cl2]Cl and in Part 4, you convert the trans isomer
to cis-[Co(en)2Cl2]Cl isomer (see lab manual for structures). In general, a trans isomer
has bulky groups located as far as possible from one another, and a cis isomer has bulky
groups adjacent to one another.
You will see a color difference between these two isomers. Obviously, the difference in
color is not due to ligand field strength, because the isomers have the same ligands.
The shape of the complex causes the difference in energy—it is highly unfavorable to
have bulky groups close to each other, so the cis isomer is higher in energy relative to the
trans isomer.
o The colors make sense here. trans-[Co(en)2Cl2]Cl appears green (absorbs red
light ~650 nm) and cis-[Co(en)2Cl2]Cl is violet (absorbs yellow light). Red light
has a longer wavelength than yellow light, so the lower energy complex absorbs
the longer wavelength of light, as expected.
This is why it is important to heat slowly in Part 4, Step 4, where you are supplying heat
energy to the trans-[Co(en)2Cl2]Cl in order to impart enough energy to overcome the
activation energy between the two isomers and convert to the cis isomer.
o If you increase the temperature too quickly, you run the risk of supplying too
much heat, in which case the complex will have enough energy to revert back to
the lower energy trans isomer. The idea is to supply just enough energy to “trap”
the complex into the cis conformation.
Absorbance spectra as a test
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After you synthesize each compound in this experiment, you will take an absorbance
spectrum to see if the compound is absorbing in the correct wavelength range. So, for
example, trans-[Co(en)2Cl2]Cl should have its max absorbance in the red range of light
(around 650 nm) and cis-[Co(en)2Cl2]Cl should have its max absorbance in the yellow
range of light (around 580 nm). You may have contaminants that absorb in other ranges.
It is also common, especially in the absorbance spectrum for your cis-[Co(en)2Cl2]Cl, to
see evidence of a mixture of the cis and trans isomers, which shows up as 2 peaks (one in
the red and one in the yellow) or a big, fat peak ( 2 or more peaks that are so close
together that it is difficult to tell them apart).