Basis for Color in Transition Metal (TM) Complexes

1
Basis for Color in Transition Metal (TM) Complexes

• Crystal (really Ligand) Field Theory

2
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)

3
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

4
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

5
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!
6
Effect of removing an electron on relative
energies of orbital sublevels
3d
remove electron
4s
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 577581 (2013) for an interesting example of a
linear iron(I) complex.
7
TM Cation Configs--Examples
8
Review, Orbital Diagrams

• You also learned to make an orbital diagram for a
  TM cation
• Ru3 Kr
• OR Kr __ __ __ __ __ __
• NOTE This orbital diagram is for a free metal
  cation. I.e. With nothing bonded to it

5s
4d
5s
4d
9
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
  metals 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.

10
When 6 Ligands Surround a TM cation (octahedral
environment), the d-orbital sublevel splits into
two sub sublevels (my terminology)
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.
11
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

12
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
Cation in octahedral complex
4s
13
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)

14
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.)
15
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

16
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)

17
Next task

• Now lets 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 well discuss
  later (paramagnetism, color).
• This is similar to creating an orbital diagram,
  except there is a new consideration (as you
  will see)
• Well start with a review

18
Energy Considerations (reminder)(briefly look
back to slide 6 before looking below)

• Why isnt 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
wouldnt 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.
19
Hunds 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).
20
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

21
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.
22
Summary of last slides concept

• When filling up an orbital diagram, you dont
  pair up electrons unless the price to go up to
  the next level is more than the pairing energy

23
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)

24
Low enough D results in electrons NOT pairing
right away
P (fixed value)
? weak-field
? strong-field
? D lt P
? D gt P
25
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)

29
Calculating D (per electron) from l of photon
absorbed
If lmax for some absorption band is 795 nm (red
end ? low energy)