Chemistry of Transition Metals Bonding in transition metal compounds - - PowerPoint PPT Presentation

Chemistry of Transition Metals

Theories: (i) Werner Coordination Theory (ii) 18 electron rule/ EAN (iii) Valence Bond Theory (iv) Crystal field theory (v) Molecular orbital approach

Bonding in transition metal compounds

Consequences: (i) High spin - low spin complexes (ii) Spectrochemical series (iii) Crystal Field Stabilization Energy (CFSE) (iv) Jahn-Teller distortions (v) Spinels

• Cobalt(III) complexes: CoCl2 + NH3 (aq) and then oxidized by air.

CoCl3 6 NH3

• range-yellow

CoCl3 5 NH3 H2O red CoCl3 5 NH3 purple CoCl3 4 NH3 green

• NH3 (aq) + HCl(aq)

 NH4+(aq) + Cl–(aq) However, CoCl3 6 NH3(aq) + HCl(aq)  no reaction

• Ag+(aq) + Cl–(aq)

 AgCl (s) (white)

Expected: CoCl3 6 NH3 + Ag+  3 AgCl (s) Observed: CoCl3 6 NH3 + Ag+  3 AgCl (s) Expected: CoCl3 5 NH3 + Ag+  3 AgCl (s) Observed: CoCl3 5 NH3 + Ag+  2 AgCl (s) Expected: CoCl3 4 NH3 + Ag+  3 AgCl (s) Observed: CoCl3 4 NH3 + Ag+  1 AgCl (s)

Werner Coordination Theory

• The secondary valence is the number of ions of molecules that

are coordinated to the metal ion. Werner assumed that the secondary valence of the transition metal in these cobalt(III) complexes is six. The formulas of these compounds can therefore be written as follows.

• Measurements of the conductivity:

CoCl3 6 NH3 four ions CoCl3 5 NH3 H2O four ions CoCl3 5 NH3 three ions CoCl3 4 NH3 two ions [Co(NH3)6]3+ 3Cl–

• range-yellow

[Co(NH3)5(H2O)]3+ 3Cl– red [Co(NH3)5Cl]2+ 2Cl– purple [Co(NH3)4Cl2]+ Cl– green

Werner : Nobel prize in 1913.

An example of a metal complex

Metal ion: Central & is a Lewis Acid Ligand: Is Attached & is a Lewis Base

How strong are the Co-ordinate bond?

Fe4[Fe(CN)6]3· x H2O. Prussian blue Prussian blue is used for certain heavy metal poisons. KCN is a poison by itself.... CH3CN Modest toxicity. Can be metabolised to produce HCN, which is very toxic

[Co(NH3)6]Cl3 and [CoCl(NH3)5]Cl2

Two compounds made of the same chemicals, yet they look different, and react differently with silver(I) salt.

Colors & How We perceive it

Black and White

If a sample absorbs all wavelength

• f visible light, none reaches our

eyes from that sample. Consequently, it appears black. If the sample absorbs no visible light, it is white

• r colorless.

When a sample absorbs light, what we see is the sum of the remaining colors that strikes our eyes.

Absorption and reflection

If the sample absorbs all but orange, the sample appears

• range.

Further, we also perceive orange color when visible light of all colors except blue strikes our eyes. In a complementary fashion, if the sample absorbed only orange, it would appear blue; blue and

• range are said to be complementary colors.

Linear (2) Trigonal plane (3) Square planar (4) Tetrahedral (4) Number Geometry Polyhedron

Examples: [Ag(NH3)2]+ [Zn(NH3)4]2+, [FeCl4]- [Ni(CN)4]2- [HgI3]-

Square pyramid (5) Trigonal bipyramid (5)

Number Geometry Polyhedron

Coordination No. 5

Example: [Ni(CN)5]3- [Cu(Cl)5]3- Example:

Triagonal prism Octahedral (6) Coordination No. 6

Singly capped

• ctahedron (7)

Number Geometry Polyhedron

Coordination No. 7 [Mo(CN)7]2-

Pentagonal bipyramidal (7) [Re(CN)7]2-

Doubly capped

• ctahedral (8)

[W(CN)8]3-

Coordination No. 8

Ligands Ligands are species (neutral

• r

anionic) bonded to the central metal ion They may be attached to the metal center through a single atom (monodentate) or two or three atoms or higher (bidentate, tridentate, etc.) Such polydentate (bidentate or higher) ligands are called chelating ligands

Some very common & simple ligands

Some very common but chelating ligands

Self Study

Isomerism in metal complexes

Isomers are compounds with the same chemical formula but different structures

• --Note that as they have different

structures, they will have different physical and chemical properties.

Self Study

Variable oxidation state

• 1. Increase in the number
• f oxidation states from

Sc to Mn. All are possible

• nly in case of Mn.
• 2. Decrease in the number
• f oxidation states from

Mn to Zn, due to the pairing of d-e’s after Mn

• 3. Stability of higher oxidn

states decreases along Sc to Zn. Mn(VII) and Fe(VI) are powerful oxidizers.

• 4. Down the group, the stability of high oxidation states

increases (easier availability of both d and s electrons for ionization).

Why does Scandium only have one oxidation state? Because if it had two - it would be Scandalous!

Stable metal complex at low oxidation state: metal electrons + lone pairs from ligands = 18 Ni(CO)4 - 4s23d8 and 4 lone pairs = 18

18 electron rule (earlier EAN: Sidgwick)

18 electron rule: explained by MO theory. (Filling of all the molecular bonding orbitals and none

• f the antibonding orbitals).

Drawback: not applicable for all complexes. Fe(CO)5 - 4s23d6 and 5 lone pairs = 18 Cr(CO)6 - 4s23d4 and 6 lone pairs = 18

EAN: Uses and Limitations

• G. N. Lewis (1902): atoms form covalent bonds by

sharing electron pair.

• W. Heitler and F. London (1927): showed how the

sharing of pairs of electrons holds a covalent molecule

• together. The Heitler-London model of covalent bonds

was the basis of the VBT.

• L. Pauling: atomic orbitals are mixed to form

hybrid orbitals, such as sp, sp2, sp3, dsp2, dsp3, and d2sp3 orbitals.

Valence Bond Theory (VBT)

(i) Ligands form covalent-coordinate bonds to the metal. (ii) Ligands must have lone pair of electrons. (iii) Available empty orbital of suitable energy for metal for bonding. (iv) Atomic (hybrid) orbitals are used for bonding (rather than molecular orbitals)

VBT – Assumptions / Features

Can explain: shape and stability of the metal complex. Can not explain: (i) Color (ii) Temperature dependence of magnetic properties

Outer sphere complex Reactive/Labile High spin Paramagnetic Inner sphere complex Less reactive Low Spin Diamagnetic

VBT – Examples

Ref: C. J. Ballhausen, J. Chem. Ed. 1979 56 194-197, 215-218, 357-361. J D Lee; pp. 204-222.

CFT largely replaces VB Theory for interpreting the chemistry

• f coordination compounds.

Crystal Field Theory (CFT)

CFT: Assumptions

• Pure electrostatic interaction
• Ligands: point charges
• Negative ligand: ion-ion

interaction, neutral ligand: ion dipole interaction

• The electrons on the metal are

under repulsive from those on the ligands

• The electrons on metal occupy

those d-orbitals farthest away from the direction of approach of ligands

Symmetric field

Degenerate 5–d orbitals in an isolated gaseous metal atom. Spherically symmetric field of negative charges: All degenerate 5–d orbitals are raised in energy due to repulsion from ligand.

dz2 dx2-y2 dxy dxz dyz

dz2 dx2-y2 dxy dxz dyz

dz2 dx2-y2 dxy dxz dyz

dz2 dx2-y2 dxy dxz dyz

dz2 dx2-y2 dxy dxz dyz

Octahedral Field

• Interaction of d-orbitals with six point charges

at +x, -x, +y, -y, +z and -z axes are not same.

• The orbitals lying along the axes (i.e. x2-y2, z2)

will be destabilized more than the orbitals lying in-between the axes (i.e. xy, xz, yz).

CFT- Octahedral Field

CFT-Octahedral Complexes

For Oh point group x2-y2, z2 orbitals: eg xy, xz, yz orbitals: t2g

• Difference between t2g and eg = Δ0 or 10 Dq.
• Conservation of barycenter from a spherical field to
• ctahedral field indicates t2g set must be stabilized as much

as the eg set is destabilized.

Illustration of CFSE

Maxima (UV-vis absorption spectrum): 20300 cm-1 So, Δo = 243 kJ/mol. (1000 cm-1 = 11.96 kJ/mol or 2.86 kcal/mol or 0.124 eV) Typical Δ0 values ~ energy of a chemical bond.

[Ti(H2O)6]3+ : 3d1 complex e− in lowest energy t2g orbitals. t2g

1 eg

t2g

0eg 1: Purple color

• For d1-d3 systems: t2g

3 (Hund's rule of maximum multiplicity)

• For d4-d7 systems:

(i) t2g

4 eg 0 (low spin case or strong field situation)

(ii) t2g

3 eg 1 (high spin case or weak field situation)

More than one d electrons

Parameters (for HS and LS case): (i) CFSE (value of Δ0, 10Dq) (ii) Pairing energy of e- (repulsive)

• CFSE > P.E.: LS complex
• CFSE < P.E.: HS complex

d4: Option I d4: Option II

High spin complexes Low spin complexes

So which one? Decided by (i)Δ0 (ii)Pairing E. (d1) (d2) (d3)

Therefore, there are two important parameters to consider:

• The Pairing energy (P) [is a repulsive energy]
• eg to t2g Splitting (referred to as Δ0, 10Dq)
• For both the high spin (H.S.) and low spin (L.S.)

situations, it is possible to compute the CFSE.

Δo vs. P (pairing energy repulsive energy)

Δo is dependent on L & M

3d < 4d < 5d M2+ < M3+ < M4+

• Nature of the ligands
• The charge on the metal ion
• Whether the metal is a 3d, 4d, or 5d element

Dependence: Δo

π-bases < weak π-bases < no π-effect < π-acids

Ligands: Weak field ligands ; small splitting (Δο ~7000 – 30000 cm-1)

Strong field ligands; large splitting (Δο > 30000 cm-1)

CFSE: Octahedral complex

Tetrahedral Field: Considerations

Tetrahedral molecule inside a cube; metal at the center. Two ‘e’ orbitals: face of the cube The ligands occupy the four alternate corners. Three ‘t2’ orbitals: center of the edges of the cube

Tetrahedral complexes

Tetrahedral Field: Considerations So, t2 is nearer to the ligand hence higher energy compared to e-orbitals.

<e-orbitals-metal-ligand = 109o28’ / 2 = 54o44’ <t2-orbitals-metal-ligand = 109o28’ / 3 = 35o16’

Tetrahedral field

Δt = 4/9 Δo

Trend for ΔT

Octahedral vs Tetrahedral Field

Spinels- Use of CFSE

Spinels- Use of CFSE

Spinels- Use of CFSE

Spinels- Use of CFSE

Special case of d8 Octahedral

Jahn-Teller Distortion

Jahn-Teller Distortion

Non-linear unsymmetrical molecule: Higher energy

• J. T. distortion

Lower degeneracy/ lower energy.

Tetrahedral, Octahedral and Square Planer

71

A Fe(II) HS-LS compound: Colour change

From Basic Science to Real time applications: Story on HS-LS complexes (Not for exam) Room Temperature

TC TC

T / K MT / cm3 mol-1 250 350 300

Red White

3

• O. Kahn, C. Jay

and ICMC Bordeaux

72

Display Device

(2) (3)

Display

Compound in Low spin state (Thin Layer) LS-HS transition tuneable with light: Applications