Transition Metal Chemistry

From the electronic configuration point of view, filling the 3d, 4d, and 5d, sub-shells results in the 1st, 2nd, and 3rd transition metal series. The 3rd and 4th series include the "inner" transition elements called lanthanides and actinides, or f-block elements. The transition elements are typical metals, and when they are oxidized, their oxidation states have a range with +2 to +3 being most common. These ions have incomplete d subshell and the number of d electrons affect their coordination chemistry.

Transition Metals and Their Ions

The lower-energy 4s atomic orbitals are filled before the 3d orbitals in neutral atoms. However, when oxidized, the 4s orbitals of all ions have a slightly higher energy than the 3d orbitals. Thus, the +2 and +3 ions have no 4s electrons. The electronic configurations of the metals and their ions and some other physical properties of transition elements are given below:
Z21222324252627282930
ElementScTiVCrMnFeCoNiCuZn
e config
[Ar]+
3d14s2 3d24s23d34s23d54s13d54s23d64s23d74s23d84s23d104s13d104s2
M2+3d13d23d33d43d53d63d73d83d93d10

M3+3d03d13d23d33d43d53d63d73d83d9

Ionicradii (Shannon)
M2+-1.000.930.940.970.920.890.830.870.88
M3+0.890.810.780.760.790.790.750.74--
IP(kJ/mol)
I1633658650653718759759737745906
I21235131014141592150915621648175319581733
I32389265328282987324829583232340835543838
E0 (V)
M2+(aq)
M(s)
--1.63-1.18-0.91-1.18-0.44-0.28-0.250.34-0.76
M3+(aq)
M(s)
-2.03-1.21-0.87-0.74-0.28-0.040.46---
Color of
M2+(aq)
--violetbluepinkgreenpinkgrenbluecolorless

Some points of interest about the data given above are:

The photon energy is calculated by Planck’s equation:

DE = h v Where h = 6.626e-34 J s, the Planck constant, and v frequency of light.

Transition Metal Complexes

Complexes are metal ions surrounded by neutral molecules or negative ions. For example, Ag(NH3)2+, Cu(H2O)62+, Fe(CN)6, Au(Cl)4, [Ti(H2O)2(NH3)2Cl2]+, [V(en)2BrCl]+ and FeF6 are complexes.

In the above example, you can easily identify the central ion as transition metals, and NH3, H2O, CN–, Cl–, Br–, en (= NH2-CH2- CH2-NH2, ethylenediamine) are ligands.

Here are some of the activities that may help you learn the coordination chemistry.

Nomenclature

Prefixes: mono, di, bis, tri, tetra, penta, and hexa are used to represent 1, 2, 3, 4, 5, and 6 respectively.

Anionic ligands usually end in o), for example O oxo, Cl– Chloro, CN– cyano, OH–, hydroxo, C2O4 Oxalato.

Neutral ligands' names are not changed, for example NH3 ammine, H2O aquo, CO carbonyl , NO nitrosyl.

Oxidation states of central ion are indicated by Roman numerials: silver(I), zinc(II), iron(III) etc.

If the complex is positively charged, they are named ions; whereas negative complexes end in ate

Learn the names as you encounter them.

Bonding Consideration of Complexes

Why complexes form?

Electrostatic, sigma, or pi?

 Crystal Field Theory: Purely electrostatic consideration. The field depends on the spatial arrangement of the ligands. In general, octahedral, tetrahedral, distorted octahedral, and square planar are some of the geometric arrangements.

 In the absence of ligands, the d orbitals are degenerate.

 The shapes and orientations of the degenerate d orbitals described in lecture.

 Crystal field implied by octahedral coordination of ligands described.

 Interaction between crystal field and d orbitals splits the degeneracy into eg and t2g levels.

eg

– –

 

eg

– –


3/5
D 0

– – – – – - -

degenerate d –2/5D 0

– – –

t2g

– – –

t2g

 

The spectrochemical series in increasing order of D 0 are: (memorize)

 Br– < Cl– < F– < OH– < C2O4 < H2O < NH3 < NO2– < CN–

 Approximate crystal-field splitting parameters D 0 (kJ/mol) of H2O, NH3, and CN–.

No. of d electron. H2O NH3 CN–.

Ti3+ 1 277

V3+ 2 222

Cr3+ 3 208 258 319

Fe3+ 5 170

Co3+ 6 217 274 401

Mn2+ 5 100

Fe2+ 6 112 376

Co2+ 7 98 121

Ni2+ 8 102 129

Cu2+ 9 137 180

 

The electronic configurations in terms of eg and t2g are due to crystal filed spliting D 0. Generally speaking, low spin or high spin configurations can be differentiated from the magnetic properties of the complexes. Individual cases are given.

Electronic Configuration of Transition Metal Complexes

Pairing energy P and energy gap of splitting D E: Suppose we have only two orbitals separated by an energy gap D E. If D E is higher than P, the system will be at a lower energy state when both electrons go into the same orbital. If the two energy level gap is small compared to the energy required to form a pair with opposite spin, placing the two electrons at different levels will result in a system with lower total energy. The idea is illustrate by the diagram below:

 

–– ­ –

­

| D E. 2E0+D E. 2E0+P <= energy

–– E0 ­ ­

 

Electronic Configurations of Octahedral Complexes:

 

eg

– – – – ­ ­ – – ­ ­ – – ­ ­


3/5
D 0

- –2/5D 0 –4/5D 0 –6/5D 0

–2/5D 0

– – – ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­
t2g

d1 d9 d2 d8 d3 d10

t2g1 eg t2g6 eg3 t2g2 eg0 t2g6 eg2 t2g3 eg0 t2g6 eg4

 

Possible electronic configurations with d4, 5, 6, & 7

high spin low spin high spin low spin

­ – – – ­ ­ – –

 

 

­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­

d4 d5

–3/5D 0 –8/5D 0 + P 0 -2D 0 + 2P

t2g3 eg1 t2g4 eg0 t2g3 eg2 t2g5 eg0

 

 

high spin low spin high spin low spin

­ ­ – – ­ ­ ­ –

 

 

­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­ ­

d6 d7

–2/5D 0 + P –12/5D 0 + 3P –4/5D 0 + 2P –9/5D 0 + 3P

 

t2g4 eg2 t2g6 eg0 t2g5 eg2 t2g6 eg1

Crystal Field Splitting by Tetrahedral Field

Splitting of the degenerated d orbitals by a tetrahedral field gives a reverse to the orbitals. This point will be demonstrated during the lecture, because a lot of writing is required to make the picture clear. First of all, you should recall the shapes of the d orbitals. Imaging the coordination ligands come in from the direction away from the x, y, z; -x, -y, z; -x, y, -z; and x, -y, -z directions. The energy levels of the d orbitals are as shown:

 

t2

– – –

t2

– – –

2/5D t

degenerate d
– – – – – - -

–3/5D t

– –

e – –

e

 

In general, the splitting due to a tetrahedral field is much smaller than that of an octahedral field. Thus, low spin complexes are usually formed.

 

Tetragonal Distortion and Square Planar Complexes

Square planar complexes are likely formed when the central ions have 8 or 9 d electrons. In the case of Cu2+ (a d9 ion), tetragonal distorted complexes are formed due to lowering the energy of the complexes.

 

–– dx2- y2

eg –– ––

–– dxy

–– dz2

t2g –– –– ––

dxz –– –– dyz

 

Ligand Field Theory

The basic ligand field theory considers the central ion forms sigma bonds with the ligands. Hybrid atomic orbitals (sp 3 or d 2sp 3) of the central ion are made available to form sigma bonds with atomic orbitals of the ligands. For octahedral complexes, six sigma bonds are formed.

For first transition metal ions, we consider one 4s (a1g) three 4p (t1u) and two 3d orbitals hybrid to give a set of atomic orbitals (d 2sp 3). When these are combined with six atomic orbitals of the ligands, six sigma molecular orbitals and six antibonding MO orbitals are generated. An approximate energy level diagram is shown below.

. – s *a1g

. – – – s *t1u

. 4p (t1u) – – –

 

. 4s (a1g) –

. – – s *eg

 

. 3d (eg) – –


. 3d (t2g) – – – – – – t2g

. – – – – –(a1g, eg, t1u)

. – – s eg

. – – – s t1u

. – s a1g

Note: This approach is essentially the same as the MO for three atomic molecules in the handout earlier (check file MO3atoms.doc)

Charge transfer spectra (in the text).

Metal Carbonyls

Carbon monoxide, CO, forms many compounds with transition metals. This neutral molecule reacts with metal to form molecular compounds:

Ni(s) + 4 CO ® Ni(CO)4(g)

Fe(s) + 5 CO ® Fe(CO)5(g)

And carbonyls react with other compounds to form carbonyls,

WCl6 + 3 Fe(CO)5 ® W(CO)6 + 3 FeCl2 + 9 CO

The structures of Ni(CO)4 and Fe(CO)5 are tetrahedral and trigonal bipyramidal. Many metal carbonyls are dimers, Mn2(CO)10, with metal-metal bonds whereas some such as (CO)3Co(m Χ CO)2Co(CO)3 have CO as bridges binding two metal atoms together in a molecule. The chemistry of metal carbonyls is interesting for the preparation of pure metals.

The stability and formation of metal carbonyls can be explained as due to a interesting sigma and back pi bonding system between the metal and the carbonyl group. Bonding between metals and carbon monoxide is particularly interesting. This part is described in the Text on page 825, explained during the lecture.