Chemistry 1b - Transition metals

WHAT IS A TRANSITION METAL COMPLEX?

A metal (Lewis acid) + ligands (Lewis bases)

Metal = accepts electrons

Ligand = donates electrons (lone pair)

the d block is special because

partially filled d orbitals

variable oxidation states

form complexes easily

The 5 d-ORBITALS

eg orbitals → point along axes

dz²

dx²−y²

t₂g orbitals → point between axes

dxy

dxz

dyz

Free atom orbital

4s is lower than 3d

In complexes: orbital

3d becomes LOWER than 4s

So electrons go into d orbitals first

Metal ion electron count

dⁿ = group number − oxidation state

exceptions

half filled (d^5) and full d^10 are execeptionally stable

oxidation states key trends

Early metals → high oxidation states

Middle → many options

End → prefer full d-shell

Lanthanide contraction

f-orbitals shield poorly

nuclear charge increases

atoms shrink

Ligand definition

molecule/ion that donates a lone pair

ligand types

Anionic (X ligands)

Neutral (L ligands)

Denticity

number of donor atoms

monodentate - 1 donor

bidentate - 2 donors

polydentate - many donors

COORDINATION NUMBER

Number of donor atoms bonded to metal

coordination number shapes

CN = 2 → Linear

CN = 4 → Tetrahedral most cases /Square planar -usually d⁸ metals

CN = 6 → Octahedral( most common geometry)

CN = 5 → trigonal bipyramidal/square pyramidal

Complex must be electrically neutral overall

using counterions - ions outside complex

ISOMERISM

Coordination isomerism

Linkage isomerism

Ionisation isomerism

Stereoisomerism

Coordination isomerism

ligands swap between metals

[Zn(NH₃)₄][CuCl₄] ↔ [Cu(NH₃)₄][ZnCl₄]

Linkage isomerism

ligand binds through different atoms

NO₂⁻ → nitro (N-bound)

nitrito (O-bound)

Ionisation isomerism

ligand swaps with counterion

[CrCl₂(H₂O)₄]Cl·2H₂O vs

[Cr(H₂O)₆]Cl₃

Stereoisomerism

Geometric (cis/trans) - exists for square planar/octahedral

fac / mer

fac = same ligands on one face

mer = spread around

Optical isomerism

mirror images (non-superimposable)

Ligands affect

the energy of the metal's d orbitals

This causes:

splitting of d orbitals

which determines:

-colour

-magnetism

-stability

d-ORBITAL SPLITTING (OCTAHEDRAL)

In a free atom all 5 d orbitals have same energy

In a complex (CN = 6) orbitals split into TWO groups:

t₂g (lower energy)

- dxy, dxz, dyz

eg (higher energy)

- dz², dx²−y²

Δo

energy gap between t₂g and eg

WHY SPLITTING HAPPENS

Ligands approach along axes

eg orbitals point directly at ligands → high repulsion → higher energy

t₂g orbitals point between ligands → lower repulsion → lower energy

ELECTRON FILLING

Two competing effects:

-Δ₀ (splitting energy)

-Pairing energy (P)

Weak field ligands

small Δ₀

electrons spread out

HIGH SPIN

Strong field ligands

large Δ₀

electrons pair up

LOW SPIN

paramagnetic

unpaired electrons

diamagnetic

no unpaired

Why complexes are coloured

electrons absorb light to jump:

t2g→eg

FACTORS AFFECTING Δ₀

Metal

Ligands

Geometry

in metals delta_o

higher oxidation state → larger Δ₀

In ligands delta_o

Spectrochemical series:

Weak → Strong:

I⁻ < Br⁻ < Cl⁻ < F⁻ < H₂O < NH₃ < CN⁻ < CO

stronger ligand → bigger Δ₀

Geometry effect on delta_o

octahedral > tetrahedral splitting

CFSE equation

(−0.4×n_t2g+0.6×n_eg)Δo

nₜ₂g = electrons in t₂g

nₑg = electrons in eg

Δ₀ = splitting energy

Ligands actually form bonds with orbitals (σ and π interactions)

This explains:

why ligands have different strengths

why Δ₀ changes

why some complexes are very stable

π-DONOR LIGANDS

Weak field

increases energy of t2g

smaller delta_o

high spin

π-ACCEPTOR LIGANDS

take electron density FROM metal

stabilise t₂g orbitals

larger Δ₀

strong field

LOW SPIN

ligands differ

σ only - medium

π-donor - decrease Δ₀

π-acceptor - increase Δ₀

Spectrochemical Series

I⁻ < Br⁻ < Cl⁻ < F⁻ < H₂O < NH₃ < CN⁻ < CO

Different ligands -

different Δ₀ → different colours

magnetism summary

weak field → more unpaired → paramagnetic

strong field → fewer unpaired → diamagnetic

strong field ligand summary

increase CFSE

make complexes more stable

visible light range

400-700nm

magnetic moment formula

μ=sqrt(n(n+2))

μ = magnetic moment

n = number of unpaired electrons

d orbitals

extend further so are easy to remove electrons

latimer diagrams show

reduction potentials between oxidation states

lowest point - most stable state

highest point - strongest oxidising agent

slope = E^O

Hard acids prefer

hard bases

hard species

Small

high charge

low polarizability

soft species

large

low charge

highly polarizable

bonding type

hard = ionic

soft = covalent

chelate effect

multidentate ligands form more stable complexes - due to entropy increase

macrocyclic effect

cyclic ligands even more stable

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