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Chemistry 324
Introduction to Transition Metal Chemistry
Welcome to the d-block!
Home to:
structural metals (Fe, Ti, Cr, W, Ni, etc)
highly active redox catalysts (Pd, Pt, Rh, etc)
polymerization catalysts (Ti, Zr, Mo)
magnetic materials (Fe, Ni, alloys)
electrical components (Cu, Ag, Au, Ni, Cd, Co)
and
lots of pretty colours!
Review: atomic electron configurations
What features would you include in a simple drawing of an atom?
Positive nucleus surrounded by a cloud of electrons
We know from 1st and 2nd year that electron positions are not
random but are in fact specifically defined or ‘quantized’
Think of e- sitting in orbitals with a specific energy, size,
shape and direction
But what do orbitals really represent?
Probability of finding an electron of specific energies at
specific distances from the nucleus (‘standard’ orbitals are
90% ‘probability’ surfaces)
Early 20th century physicists and chemists trying to
understand light and the quantized behaviour of e- in atoms
came to the conclusion that we must think of e- as a wave, as
well as a particle in order to understand its position. This
gives rise to the concept of phase.
Atomic orbitals are discrete mathematical solutions of the
Schrödinger wave equation: Ψ
each Ψ describes an electron wave in 3D space
orbitals are directly related to (Ψ)2: the probability of finding
an e- at a given point in space
Ψ is specified by three quantum numbers: n, l, ml while a 4th
quantum number, ms, specifies the electron spin
Best to use polar coordinates to describe the wave function:
Ψ(r, θ, φ) = R(r)Y(θ,φ)
R(r) Radial part describes wavefunction moving along the
radius vector and R2(r) corresponds to the probability of
finding an e- at a given distance from the nucleus.
Y(θ,φ)
space
Angular part describes the shape of the orbital in
Electron configuration: multi-electron atoms
Predict orbital occupancy according to:
a) Aufbau principle
lowest energy orbitals fill first
no distinction between p orbitals (or within any other set)
b) Pauli exclusion principle
each e- must have a unique set of quantum numbers
c) Hund’s rule of maximum multiplicity
electrons are placed in orbitals so there are the maximum
number of parallel spins
Why? Electron-electron repulsion effects are less for orbitals
with same spin because of ‘correlated motion’
Aufbau and many e- atoms: must take into account how orbital
energies change with:
a) increasing nuclear charge
all orbital energies are lowered to some extent but the effect
varies with l
b) electron-electron repulsion
dictates that e- will spread out among degenerate orbitals
c) shielding and effective nuclear charge (Z*)
filled orbitals do an incomplete job of shielding valence
(outer) electrons from nuclear charge but this is dependent on
the shape of the orbital (l). Filled d orbitals are particularly
poor at shielding so as we move across the d series, Z*
increases and orbitals energies decrease.
actual orbital fill order gets complicated after the 3p set fills
because of a-c above
Electron configurations: transition metals
Transition metals have use of (n-1)d, ns and np orbitals
For elemental transition metals, the ns orbitals fill first
exceptions are shown on handout for Cr (3d54s1) and Cu
(3d104s1). Better to have electrons unpaired in s when energies
are very close.
BUT, when positive metal ions form, the (n-1)d set drops below
the ns set
on average the (n-1)d set is closer to the nucleus than the ns
and is therefore more densely packed in space; removal of an
electron decreases e- repulsion more for the (n-1)d than the
ns, so the d set drops below the ns set and fills first. The effect
is even more pronounced for M2+ ions.
Transition metal ions have only d electrons in their outer (valence)
shells, not s electrons
dn configurations where n is the number of d electrons NOT
the principal quantum number
Eg. Ti+ is a d3 ion:
take group number as total d ‘potential’ electrons and then
subtract the charge to get the ‘d-count’
How does a partially filled d shell affect properties of the metal
ion?
d orbitals project out further than other orbitals and therefore
are more affected by the surrounding ligand environment than
are s or p electrons
as a result, transition metals properties are strongly influenced
by their dn electron count: egs. colour, magnetism, reactivity
Note that one definition of transition elements is to include only
those elements with partially filled d shells. This would exclude
group 12 ions Zn2+, Cd2+ and Hg2+ because they are all d10 without
exception
we usually still include them but they DO in fact have a lot of
similarities to the group 2 elements
So, why transition elements?
Ions of these elements represent a transition in bonding type
from ionic of the s-block to covalent in the p-block and the
‘transition’ is fairly smooth to increasing covalence moving
right (although oxidation state plays a role here too).