Transcript Ppt07c(Wk12)TM III-Basis for Color_v3

Basis for Color in Transition
Metal (TM) Complexes
Crystal (really “Ligand”) Field Theory
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Goals
• Learn that when ligands bind to a metal
complex, the d orbital sublevel becomes further
“split” into sub-sublevels
– Electronic transitions between these sub-sublevels
absorb visible light, causing “color”
• Learn what is meant by “strong field” and “weak
field” cases, and how these impact the electron
configuration (and paramagnetism).
• Use these ideas to explain why different
complexes of the same TM cation have different
colors (ruby vs emerald)
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Review, electron configurations for TM
cations (& # d electrons)
• Earlier (self study worksheet or exercise), you learned
how to determine the number of d
electrons in a transition metal (TM) cation
• Example, Ru3+
– Ru, 44 electrons. Kr, 36 electrons
 44-36 = 8  8 e-s “past” Kr
– Ru3+, remove 3 electrons  8-3 = 5 e– All electrons go into d sublevel (not s) [see next
slide for “explanation” if you wish]
 5 d electrons, [Kr] 5s0 4d5
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Creating Cations from Transition Metal
(TM) Atoms
• Recall that when filling up orbitals in ATOMS,
the ns fills before the (n-1)d (e.g., the 4s fills
before the 3d, etc.).
– We said that this is because the 3d orbital is “higher
in energy” than the 4s. (shapes/shielding issue)
• It turns out that once an electron (or more) is
removed (i.e., once you make a cation), the (n-1)d level
becomes lower than the ns! (see next slide) [You
are not responsible for knowing why this is so. Ask me in person if
you wish.]
– This means that the electron configurations for the TM
cations have no electrons in the ns subshell
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Effect of removing an electron on
relative energies of orbital sublevels
3d
remove electron(s)
4s
4s
3d
TM Cation
(I.e., once one or more
electrons have been removed)
Neutral TM Atom
This switching of positions of the ns and (n-1)d sublevels is why there are no “s”
electrons in TM cations. They would “fall down” to the (n-1)d sublevel which is
(now) lower in energy. Electrons are not “removed from the s sublevel” first!
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Effect of removing an electron on
relative energies of orbital sublevels
3d
remove electron
4s
3d
4s
4s
3d
Fe: [Ar] 4s2 3d6
Fe+: [Ar] 4s0 3d7
Neutral TM Atom
TM Cation
NOTE: This example is for illustrative purposes. Fe is usually in the +2 or +3
oxidation states in complexes. However, see Nature Chemistry, 5, pp 577–
581 (2013) for an interesting example of a linear iron(I) complex.
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TM Cation Configs--Examples
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Review, Orbital Diagrams
• You also learned to make an orbital diagram for
a TM cation:
• Ru3+: [Kr]
5s
OR
[Kr] __
4d
__ __ __ __ __
5s
4d
• NOTE: This orbital diagram is for a “free” metal
cation. I.e. With nothing bonded to it
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Ligand Field Theory
• Text refers to “crystal field theory”.
Simpler, but not “valid”.
• Ligand theory is better, but full treatment
beyond the scope of this course
• KEY IDEA: When ligands bind to a metal
cation, the ligand orbitals affect the energy
of the metal’s d orbitals--some d orbitals’
energies go up, and some go down. We
say that the d orbtial sublevel “splits” (into
sub sublevels). See diagram, next slide.
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When 6 Ligands Surround a TM cation (octahedral
environment), the d-orbital sublevel splits into two
“sub sublevels” (my terminology)
Refer back to Slide 6 to
see that this is what we
did “before”
This is called the d-orbital “splitting pattern”
for an octahedral complex. This is the only
splitting pattern you need to know for my
class (it would change for different geometries and
C.N.s). You do not need to know that the top
two orbitals are the z2 and x2-y2, etc. Just
know that there are “two up” and “three
down”.
Light & Color 10
Splitting Pattern and D, the splitting
energy
• The energy difference between the lower
“sub sublevel” and the higher one is called
the “splitting energy”, with symbol D
• D is just short for DE
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D varies in different complexes, but is
always small compared to “s to p” or “p
to d” gaps! (this is not obvious from text)
4d
4p
Free cation
4s
D
4d
4p
4s
Cation in octahedral complex
Light & Color 12
Recall (earlier PowerPoint):
Absorption at the Molecular Level
• Absorption of one photon of visible light
corresponds to the excitation of one electron from a
lower energy orbital to a higher energy one
• The bigger the DE (energy difference or gap)
between the orbitals, the greater the Ephoton
absorbed
• Different gaps yield different colors absorbed, and
thus different colors perceived
• Changing the DE in a metal complex (or other “dye”
molecule) will change the color of the complex (or
dye)
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The value of D, although it varies, is “just
right” to absorb photons of visible light!
(l  400 – 800 nm)
This is why transition metal complex are usually
colored! The small splitting that develops, results in
“gaps” that absorb some visible colors, leaving
others to reach our eyes! (More on this later.)
Light & Color 14
The greater the value of D, the
“stronger the (crystal or ligand) field”
• Recall that the reason for the splitting of
the d orbitals sublevel was the presence
of a (crystal or ligand) field.
– Think of this “field” as being like an
“environment”, imposed by the ligands.
• So hopefully it makes sense that the
“stronger the field” imposed by the ligands,
the greater the D.
– “strong field” means a bigger D
– “weak field” means a smaller D
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Different ligands tend to impose
different “field strengths”
• Ligands that tend to make a smaller D are called
“weak field” ligands
• Ligands that tend to make a larger D are called
“strong field” ligands
– You do not need to memorize which ligands are of
which type. You just need to know the concept (and
meaning of “strong / weak field”
– Recall the “Fun with colors” expt! Ni2+ with different
ligands had different colors!
• The type of metal cation also affects D (but again,
you just need to know that this is “so”, not “how” or “who does what”)
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Next task…
• Now let’s see how we can “populate” the
d-orbitals of a TM complex with electrons
– This is important because of some of the
properties of TM complexes that we’ll discuss
later (paramagnetism, color).
• This is similar to creating an orbital
diagram, except there is a “new”
consideration (as you will see)
• We’ll start with a review…
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Energy Considerations (reminder)
(briefly look back to slide 6 before looking below)
Why isn’t the config. for Ru3+ this one?
[Kr]
5p
4d
Ans: Because that configuration would have higher energy
than the one on the previous slide. This configuration
would describe an “excited state” of Ru3+. That
electron in the 5p orbital wouldn’t stay there. It would
“fall down” to the 4d.
 When we write electron configurations or orbital
diagrams, we assume ground state configurations.
We “want” the lowest overall energy possible.
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Hund’s Rule (revisited)
• OK, but what about this one? Why is this
is not the correct config?
[Kr]
4d
Ans: This also has a higher energy! This also represents
an “excited state” configuration!
Why? Because it takes a bit of energy to put two
electrons into the same orbital (electrons repel, right? See
next slide for more detail if you like). I call the amount of
“energy cost” to put two electrons in the same orbital
the pairing energy (P).
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Recap: It “costs” energy to pair up
electrons in the same orbital
• Electrons repel, so having them in the same
orbital makes the energy of the system a bit
higher
• The amount of energy it “costs” to pair up two
electrons is called the “pairing energy”
• P = “pairing energy”
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In the past, you could sort of ignore the
pairing energy
• In the past, once you had put one electron
into each sublevel of a “set”, you would
then pair up the electrons before
populating the next higher sublevel:
Next electron would go into 2p,
paired, rather than putting it up
“higher” into the 3s
3s
2p
This is because the price to pair up the electron is much less
than the energy needed to “get up to” that 3s orbital (the 2p – 3s
“gap”). This was never addressed in 1st semester, but please
take a moment to make sure you understand this now.
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Summary of last slide’s “concept”
• When “filling up” an orbital diagram, you don’t
pair up electrons unless the “price” to go up to
the next level is more than the pairing energy
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Preparation for next slide
• On the following slide, two complexes are
compared. Both involve Co3+, but in one case
(with F-’s), the ligands create a “weak field” and in
the other (with CN-’s), , they create a “strong field”
• In the “weak field” case, the D is so small that it
has become smaller than the pairing energy.
• As a result, the fourth electron placed into the
diagram goes “up” into the higher level rather
than pairing up in the lower one! (Click and look
below to see the “sequence” of filling)
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Low enough D results in electrons NOT
pairing right away
P (fixed value)
 weak-field
D<P
 strong-field
D>P
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Explanation of the “high spin” and “low
spin” designations on prior slide
• As will be discussed on the next couple of slides,
an unpaired electron has something called a
“spin” and has its own magnetic field.
• Thus, the more unpaired electrons, the greater
the total spin.
• The left complex had four unpaired electrons.
The right complex had none (which is fewer than
four!). So the left one is called “high spin” and
the right one “low spin”
• When there is a difference in total spin, the “high
spin” one is always the “weak field” one.
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Magnetic Properties are Related to
Electron “Spin”
• Electrons have a “spin” (up or down)
• “Spin” is a Magnetic property
– single electron (spin) will attract a magnetic
field (“paramagnetic”)
• If multiple electrons have the SAME spin in a
complex, the complex will be MORE attracted to a
magnetic field (more paramagnetic)
– Two paired electrons (one up, one down)
have no net “spin” (they cancel out)
• NOT attracted to a magnetic field
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Spin (continued)
• THUS:
• If a complex has ALL ITS ELECTRONS PAIRED
(“no spin”), then the complex:
– Will NOT attract a magnetic field
– Is called diamagnetic
• If a complex has ONE or MORE UNPAIRED
electrons, it:
– WILL attract a magnetic field
• And the more unpaired electrons (the “higher the [total]
spin”), the stronger the force of attraction
– Is called paramagnetic
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When is a species colored?
(Requirements)
• D must have an energy that falls in the range of
energies of visible light (photons)
– If too large, absorption band is in the UV (or higher)
range
– If too small, absorption is in the IR (or lower) range
(almost never is this small for electronic transitions)
• “bottom” energy level must contain at least one electron
(or else nothing to “excite”!)
• “top” energy level must have at least one “vacancy” (or
else nowhere for an electron to “go” to!)
• Result?
– If no d electrons, no color (Ti4+)
• Salts containing Al3+ or Gp I and Gp II cations
– If 10 d electrons, no color
• Complexes/compounds containing Zn2+ (or Gp III-Gp V metal cations)
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Calculating D (per electron) from l of
photon absorbed
If lmax for some absorption band is 795 nm (red end  low energy):
D(per e )  Ephoton
_
6.626 x 10


-34

J s 3.00 x 108 m/s
-9
(795 nm)(10 m/nm)
  2.500 x 10
-19
J (per photon)
(or electron)
 2.500 x 10-19 J   6.022 x 1023 e- 
5 J

1.506
x
10



mol
electron
mol
of
e-'s



1.506 x 105 J
mol
 1000 J/kJ  151 kJ/mol
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