Transition metal complexes
A complex is a central
metal atom or ion surrounded
by coordinately bonded ligands
A coordinate bond is the same as a dative covalent bond (both e- come from same ligand)
Ligands donate a pair of
electrons to the central
metal ion to form a
coordinate bond
They must have at least
one lone pair to donate
Ligands with ONE LONE PAIR = MONODENTATE
Ligands with TWO LONE PAIRS = BIDENTATE
Ligands with MORE THAN TWO LONE PAIRS = POLYDENTATE
Coordination number = number of coordinate bonds formed
Number of lone pairs = coordination number
Shapes of complexes
2 COORDINATE BONDS = LINEAR
SIX COORDINATE BONDS = OCTAHEDRAL
FOUR COORDINATE BONDS = TETRAHEDRAL (109.5 bond angle) or SQUARE PLANAR (90 bond angle)
Transition metals can form complexes
Normally the 3d orbitals of transition metal ion have the same energy, however when ligands bond to them, the 3d orbital is split into two different energy levels
For electrons to jump from the ground state (lower orbital) to an excited state (higher orbital), they must absorb energy (visible light) equal to the energy gap
The frequency of visible light
absorbed by electrons depends
on the size of the energy gap
Only the frequencies with this
energy are absorbed, the rest
are transmitted or reflected
The combination of these
transmitted/reflected
frequencies is the
complimentary colour
of the frequency absorbed
If there are no 3d electrons or the
3d sub-level is full, no electrons will
jump, no energy will be absorbed
and the compound will look white/colourless
The energy needed for electrons to make the jump depends on the central metal ion + its oxidation number, and the ligands + their coordination number
The colour of the complex depends
on the number of e- in the d orbital,
the arrangement of ligands and the
nature of these ligands
Colours of common
transition metal ions
+3 ions
+2 ions
Fe+2 = pale green
Mn+2 = pale pink
Cu+2 = blue
Fe+3 = orange/brown
Ti+3 = purple
V+3 = green
Cr+3/Mn+3 = violet