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Chapter 18: Coordination Complexes
Oxtoby: pages 659-689.
The three series of elements arising from the filling of the 3d, 4d, and 5d shells, and situated in the periodic table following the alkaline earth metals, are commonly described as "transition elements," though this term is sometimes also extended to include the lanthanide and actinide (or inner transition) elements. They exhibit a number of characteristic properties which together distinguish them from other groups of elements:
(i) They are all metals, and as such are lustrous and deformable and have high electrical and thermal conductivities. In addition, their melting and boiling points tend to be high and they are generally hard and strong.
(ii) Most of them display numerous oxidation states which vary by steps of 1 rather than 2 as is usually the case with those main-group elements which exhibit more than one oxidation state.
(iii) They have a strong propensity for forming coordination compounds with Lewis bases.
Coordination complexes have been known for ca. 100 years. They were initially called "complex compounds" because of the consternation they caused for early chemists. These materials are stable (more or less) and stoichiometrically reliable, but they are often made up
of combinations of other compounds that are independently stable. For example, Cu 2+^ ions combine
with H 2 O to give Cu[H2O] 4 2+. The Cu 2+^ ion coordinates four water molecules. The water
molecules are called ligands.
It is very important to realize that the interaction between the cation (Lewis acid) and the ligand (Lewis base) is actually a chemical bond. The bond has associated with it a highly variable but substantial bond enthalpy.
Remember that water is just another ligand. Much of coordination chemistry is carried out in water, and the competition for metal binding between water and various ligands is a central theme. Review your notes from Chapter 11 on complex ion equilibria.
Types of Ligand
Ligands are commonly classified according to the number of donor atoms which they
contain. Unidentate or monodentate ligands include H2O, NH3, HS-^ , RO -^ , etc.
Bidentate ligands are very common. Examples include ethylenediamine, 1,10- phenanthroline and oxalate. These are often "chelating" ligands (from Greek χηλ η´ , crab's claw).
CH 2
HN
CH 2
CH 2
CH 2 NH 2
NH 2
diethylenetriamine, dien
OC
OC O
O
Oxalate
CH 2 CH (^2) H 2 N NH 2
ethylenediamine, en
N N
1,10-phenanthroline
As
As
Me (^2)
Me 2 o -phenylenebis(dimethylarsine), diars [1,2-bis(dimethylarsino)benzene]
HC C O
C O
R
R
β-diketonates (e.g., R=Me: acetylacetonate, acac)
O
O
Tropolonate
N
N
N
terpyridine, terpy
N
N
N (^) N
N
N
N
N
phthalocyanine
N N
CH 2
C
O O
CH 2 C
O
O
CH (^2)
CH (^2)
O
O
C O
O
(HO 2 CCH 2 ) 2 N(CH 2 ) 2 N(CH 2 CO 2 H) 2
EDTA
- Anionic ligands are given an o suffix. Neutral ligands retain their usual name. Coordinated water is called aqua.
Examples: chloro, Cl - bromo, Br - sulfato, SO 4 2- methylamine, CH 3 NH (^2) ammine, NH 3 (the double m distinguishes NH 3 from alkyl amines) aqua, H2O
- The number of ligands of one kind is given by the following prefixes. If the ligand name includes these prefixes or is complicated, it is set off in parentheses and the second set of prefixes is used.
2 di bis 3 tri tris 4 tetra tetrakis 5 penta pentakis 6 hexa hexakis 7 hepta heptakis 8 octa octakis 9 nona nonakis 10 deca decakis
Examples: Simple ligands are given above. dichlorobis(ethylenediamine)cobalt(III), [Co(NH (^) 2CH (^) 2CH (^) 2NH (^) 2) (^) 2Cl 2 ] + tris(bipyridine)iron(II), [Fe(C (^) 5H (^) 4N–C (^) 5H (^) 4N) (^) 3] 2+
- The prefixes cis- and trans- designate adjacent and opposite geometric locations. Examples are given in the figure below. Other prefixes are used as well and will be introduced as needed in the text.
Examples: cis- and trans- diamminedichloroplatinum(II), [PtCl2(NH3) (^) 2] cis- and trans- tetraamminedichlorocobalt(III), [CoCl (^) 2(NH (^) 3) (^) 4] +
Pt
H 3 N Cl
H 3 N Cl Pt
Cl NH^3
H 3 N Cl
Co
NH 3
NH 3
H 3 N Cl
H 3 N Cl
Co
Cl
Cl
H 3 N NH (^3)
H 3 N NH (^3)
cis- and trans- dichlorodiammineplatinum(II) [PtCl 2 (NH 3 ) 2 ]
cis- and trans- dichlorotetraamminecobalt(III) [CoCl 2 (NH 3 ) 4 ] +
- Bridging ligands between two metal ions as in the figure below have the prefix μ-.
Examples: tris(tetraammine-μ-dihydroxocobalt)cobalt(6+), [Co(Co(NH (^) 3) 4 (OH) (^) 2) (^) 3] 6+
μ-amido-μ-hydroxobis(tetraamminecobalt)(4+), [(NH (^) 3) (^) 4Co(OH)(NH (^) 2)Co(NH (^) 3) (^) 4] 4+
Co
HO OH
Co
OH
NH 3
NH 3
H 3 N NH 3
HO Co
NH 3
H 3 N NH 3 NH 3
O
O Co
NH 3
NH 3
H 3 N
H 3 N
H
H
Co H 3 N
H 3 N
NH (^2)
O
NH (^3)
NH (^3)
Co NH 3
NH 3
NH 3
NH 3
tris(tetraammine-μ-dihydroxocobalt)cobalt(6+) [Co(Co(NH 3 ) 4 (OH) 2 ) 3 ] 6+
μ-amido-μ-hydroxodi-(tetraamminecobalt)(4+) [(NH 3 ) 4 Co(OH)(NH 2 )Co(NH 3 ) 4 ] 4+
Demonstration
A solution of Co 2+^ in water can be prepared by dissolving various Co salts in water.
The solution is pink due to the presence of the octahedral ion Co(H (^) 2O) 6 2+.
Addition of chloride (Cl-) in the form of concentrated HCl forms the blue tetrachloride dianion:
[Co(H (^) 2O) (^) 6] 2+^ + 4 Cl- [CoCl (^) 4] 2-^ + 6 H2O
The equilibrium can also be shifted to the left by addition of solid ZnSO 4 , since Zn2+
binds Cl -^ more strongly than Co 2+^.
Structures of Coordination Complexes
We will discuss structures first, and then move on to consideration of the bonding theories to attempt to explain the observed structures.
The coordination number of the central metal in a complex is simply equal to the number of ligands attached. Coordination numbers from 2 to 12 are known. Six is the most common. We will focus on coordination numbers 4, 5, and 6 in this class.
Early observations of four-coordinate platinum(II) complexes led to the surprising conclusion that two isomers exist for complexes such as Pt(NH3) (^) 2Cl2. This shows that their
structure cannot be tetrahedral. Werner suggested a square planar structure:
Pt
H 3 N Cl
H 3 N Cl
Pt
H 3 N Cl
Cl NH (^3) cis trans
Similar observations on six-coordinate complexes such as [Co(NH (^) 3) (^) 4Cl (^) 2] +^ showed that
there are two isomers, consistent with an octahedral geometry.
Co
NH 3
NH 3
H 3 N Cl
H 3 N Cl
Co
Cl
Cl
H 3 N NH (^3)
H 3 N NH (^3)
cis purple
trans green
An additional type of isomerism is noted for a complex closely related to cis-
[Co(NH 3 ) 4 Cl 2 ] +^. If the four NH 3 ligands are replaced with two ethylenediamine ligands (see
18.13), we now have the possibility of forming two enantiomers of the complex.
Co
N
N
Cl
N Cl
N Co
N
N
N
Cl N
Cl
Imagine a mirror plane here
More on Optical Isomers
(see p. 673 in Oxtoby.)
Complexes such as [Pt(en) (^) 3] 4+^ and [Fe(ox) 3 ] 3-^ (ox = oxalate) can exist as two optical
isomers. All complexes of this general type are called tris chelate complexes, and their structures can be represented schematically:
These two molecules are enantiomers. Official nomenclature calls them lambda (Λ) and delta (∆).
B. Complexes
In transition metal complexes, the n d orbitals end up lower in energy than the ( n + 1) s orbitals.
e.g., in Vanadium complexes:
V(0) (^) 3d^5 not 4s 2 3d 3 , as in free atoms
V(I) (^) 3d^4
V(II) (^) 3d^3
V(III) (^) 3d^2
V(IV) (^) 3d^1
General electron counting rules for number of d electrons:
- Count total number of valence electrons from the periodic table.
- Deduce oxidation state.
- Subtract.
Examples: Fe2+^ is d^6 , Fe3+^ is d^5 , Co2+^ is d^7 , Co3+^ is d^6.
The dxy , dxz, and d (^) yz orbitals are very little changed in energy because they are in
between the axes. The d (^) z2 and dx2-y2 orbitals are greatly destabilized. It turns out that typical
values of ∆ 0 are equivalent to the energy of visible light photons, so the promotion of
electrons from the lower to higher orbitals leads to absorption of visible light (color!).
The tetrahedral case is best approciated by inscribing the tetrahedron within a cube.
M m+
Now the dxy , dxz, and d (^) yz orbitals are destabilized the most, while the dz2 and d (^) x2-y
orbitals are affected less.
Another common geometry is the square planar geometry, which we can derive from the octahedral case by removing the 2 ligands on the z axis.
Ligand Field Model
It has been observed that the ∆ 0 values observed for various ligands follows the order:
I-^ < Br-^ < Cl-^ < F-^ ≈ OH -^ < H2O < :NCS -^ < NH 3 < en < CO, CN -
weak field intermediate strong field
This would be hard to explain on our purely electrostatic model, since ligands with full negative charges might be expected to be capable of giving large ∆ 0 values.
Looking at our correlation diagram allows us to make a more complete analysis. The ∆ 0 energy basically depends on how high in energy the antibonding (σ* (^) d) level ends up. This
will go up as the corresponding σd orbital goes down in energy. So ligands that form the
strongest σ bonds to M will give the highest ∆ values.
There is also a secondary but important interaction between the non-bonding d levels
and the filled p orbitals of ligands like Cl -^ , Br -^ , etc. These repulsive interactions will raise the energy of t (^) 2g slightly, decreasing ∆. This effect decreases down the halogens, so we have the
order I -^ < Br-^ < Cl-^ < F-^. Similar reasoning indicates that NH 3 (one "lone pair") should be a
better ligand than H2O (two "lone pairs").
What about CO and CN -^? We need to invoke π bonding to understand this.
Thermodynamic Effects of Ligand Field Splittings
- The magnitude of ligand field effects is large enough to have a considerable effect on the chemistry of the complexes.
- For example, a d^2 ion (in an octahedral ligand set) will have each of its two d electrons stabilized by 2 ∆ 0 /5, for a total stabilization of 4 ∆ 0 /5. Since ∆ 0 (measured by spectroscopy) is 10,000 to 20,000 cm -1^ for di- and tri-valent ions of the first transition series, the stabilization due to ligand field (LFSE) is 100 to 200 kJ mol-^.
LFSE for Octahedrally and Tetrahedrally Coordinated Ions (High Spin)
of d electrons oct. tetra. Diff* (oct.-tet.)
1, 6 2 ∆ 0 /5 3 ∆t /5 ∆ 0 /
2, 7 4 ∆ 0 /5 6 ∆t /5 2 ∆ 0 /
3, 8 6 ∆ 0 /5 4 ∆t /5 8 ∆ 0 /
4, 9 3 ∆ 0 /5 2 ∆t /5 4 ∆ 0 /
0, 5, 10 0 0 0
*assuming that ∆ 0 ≈ 2 ∆t
Hydration Energies
Consider the enthalpy of the process:
M2+ (gas) + H2O → [M(H2O)6] 2+ (aq)
Recalling the trend in ionic radii, we would anticipate a steady increase from Ca2+^ to Zn2+^ as the ions get smaller and the 2+ charges are concentrated in a smaller volume.
Hydration energies of some divalent ions. Solid circles are experimentally derived hydration energies. Open circles are energies corrected for LFSE.
Magnetic Properties of Complexes
(See Oxtoby, pp. 623-624.)
- Paramagnetic substances are strongly attracted or drawn into a strong magnetic field.
- Diamagnetic species are weakly repelled by the same field.
Demonstration: MnSO 4 • nH2O
Mn+2, d^5 , high spin, five unpaired electrons.
Ligand Substitution Rates
Complexes are found to undergo ligand exchange reactions at very widely variable rates; e.g.:
[Cu(H 2 O) (^) 6] 2+^ + 4 NH 3 → [Cu(NH3) (^) 4(H2O) (^) 2] 2+^ + 4 H2O
[Fe(H (^) 2O) (^) 6] 3+^ + SCN–^ → [Fe(H (^) 2O) (^) 5(SCN)]2+^ + H2O
These reactions are complete in the time that it takes to mix the solutions.
In contrast:
[Co(NH (^) 3) (^) 6] 3+^ + 6 H2O H^
→ [Co(H2O) (^) 6] 3+^ + 6 NH 4 +
occurs very slowly (weeks!), in spite of ∆G° < 0.
Definitions:
slow ( t 12 ≥ 1 minute) ≡ inert
fast ( t 12 ≤ 1 minute) ≡ labile
Labile complexes have electrons in high energy (antibonding) orbitals.
