General Chemistry I 1 BONDING IN TRANSITION METAL COMPOUNDS AND COORDINATION COMPLEXES 8.1 Chemistry...

General Chemistry I 1

BONDING IN TRANSITIONMETAL COMPOUNDS ANDCOORDINATION COMPLEXES

8.1 Chemistry of the Transition Metals

8.2 Introduction to Coordination Chemistry

8.3 Structures of Coordination Complexes

8.4 Crystal Field Theory: Optical and Magnetic

Properties

8.5 Optical Properties and the Spectrochemical

Series

8.6 Bonding in Coordination Complexes

8CHAPTER

General Chemistry I

General Chemistry I 2

347

Emerald

3BeO∙Al2O3∙6SiO2

with someAl3+ replaced

by Cr3+

General Chemistry I 3

8.1 CHEMISTRY OF THE TRANSITION METALS

348

General Chemistry I 4

- Decreasing radii for small Z transition atoms

→ Increase in Zeff

- Increasing radii for large Z transition atoms

→ Increase in electron-electron repulsion

349

General Chemistry I 5

Lanthanide contraction: bad shielding by 4f orbitals

→ the radii of the 6th period ~ the 5th period

→ decrease in atomic and ionic radii by increasing Z

along the 6th period

349

General Chemistry I 6

350

melting point: function of the bond strength in solids

- roughly correlated with the number of unpaired e-

General Chemistry I 7

351

Enthalpy of hydration of M2+ ions

M2+(g) → M2+(aq): Hhyd

= Hof(M2+(aq)) – Ho

f(M2+(g))

Lowering of Hhyd from a line

→ due to crystal field stabilization

Anomalies of Mn

→ due to the stable half-filled d shell

General Chemistry I 8

Oxidation states

more common oxidation state

Increasing tendency toward higher oxidation states among

heavier transition elements in the same group:

Fe (2,3) → Ru (2,3,4,6,8), Ni(2,3) → Pd(2,4)

351

General Chemistry I 9

Hard and Soft Acids and Bases

Pearson (1963)

~ Extension of Lewis’ definition –

electron pair acceptor (acid) and donor

(base) – by adding categories ‘hard’

and ‘soft.’

~ 'Hard' species: small, high charge

states, low electronegativities, weakly

polarizable

~ 'Soft' species: large, low charge

states, high electronegativities, strongly

polarizable

~ ‘Borderline’ species

Ralph Pearson (US, 1919 - )

353

General Chemistry I 10

354

General Chemistry I 11

HgF2(g) + BeI2(g) → BeF2(g) + HgI2(g) s/h h/s h/h s/s

354

Prediction of chemical reactivities of inorganic reactions

~ Preferred direction: hard acid/hard base or soft acid/soft

base

AgBr(s) + I–(aq) → AgI(s) + Br–(aq) s/b s s/s b

EXAMPLE 8.2 Predict whether the following reactions will occur.

(a)CaF2(s) + CdI2(s) → CaI2(s) + CdF2(s)

(b)Cr(CN)2(s) + Cd(OH)2(s) → Cd(CN)2(s) + Cr(OH)2(s)

NO

YES

General Chemistry I 12

8.2 INTRODUCTION TO COORDINATION CHEMISTRY

355

Formation of Coordination Complexes

Werner’s investigation:

Compound 1: CoCl36NH3 (orange-yellow)

Compound 2: CoCl35NH3 (purple)

Compound 3: CoCl34NH3 (green)

Compound 4: CoCl33NH3 (green)

Alfred Werner (Swiss,1866-1919) Nobel prize in chemistry(’13)

Treatment with HCl → did not remove NH3

AgNO3 + Cl- → AgCl(s) in the ratio of 3 : 2 : 1 : 0

General Chemistry I 13

Conductivity measurements:

Compound 1: [Co(NH3)6]3+(Cl–)3 → Conductivity of Al(NO3)3

Compound 2: [Co(NH3)5Cl]2+(Cl–)2 → Conductivity of Mg(NO3)2

Compound 3: [Co(NH3)4Cl2]+(Cl–) → Conductivity of NaNO3

Compound 4: [Co(NH3)3Cl3] → Nonelectrolyte

→ Concept of “coordination sphere”

around the central metal ion

inner and outer coordination sphere

→ Formation of an octahedral complex

356

In the above complexes, NH3 and Cl- that are attached to Co are called LIGANDS

General Chemistry I 14

357

anhydrous CuSO4CuSO4∙5H2O

→ [Cu(OH2)4]SO4∙H2O

General Chemistry I 15

Monodentate ligands

mono “one” and dens “tooth”

357

General Chemistry I 16

Bidentate ligands

Chelating ligands: chelate (G. chele, “claw”)

[Co(EDTA)]– ~ 1 hexadentate[Pt(en)3]4+ ~ 3 bidentates

358

(‘en’)(‘ox’)

General Chemistry I 17

Naming coordination compounds

1) Single word for a coordination complex ~ [prefix-ligand-metal]

2) Cation first followed by anion ~ K[…] or […]Cl

3) Ending with the suffix “-o” for anionic ligand, chlorido (Cl),

no change for neutral ligands except aqua (H2O), ammine (NH3),

carbonyl (CO). Note: “chloro” for Cl in a compound ligand

4) Prefixes for the number of ligands ~ di-, tri-, tetra-, penta-, hexa-, …

(bis-, tris-, tetrakis-, … for ligands with di- (etc) in their names)

5) Alphabetical ordering for many ligands

6) Roman numeral (oxidation state) in (..) after the name of metal

~ […cobalt(III)]Cl or K[…ferrate(III)]

anionic complex ions: the ending “-ate”

359

General Chemistry I 18

359

General Chemistry I 19

Ligand substitution reactions

[Ni(OH2)6]2+(aq) + 6 NH3(aq) → [Ni(NH3)6]2+(aq) + 6 H2O

360

Another example

Cu(H2O)62+(aq)

Pale blue

HCl(aq) NH3(aq)CuCl4(aq)

_

GreenCu(NH3)6

2+(aq)Deep blue

General Chemistry I 20

‘Inert’ coordination complex:

thermodynamically unstable, kinetically

stable (inert)

‘Labile’ coordination complex:

thermodynamically unstable, kinetically

unstable (labile)

3 33 6 3 2 6 4[Co(NH ) ] ( ) 6H O ( ) [Co(H O) ] 6NH ( )aq aq aq

2 23 6 3 2 6 4[Co(NH ) ] ( ) 6H O ( ) [Co(H O) ] 6NH ( )aq aq aq

takes a week

takes a matter of seconds

361Difference between ‘inert’ and ‘labile’

Ene

rgy

Reaction

Ene

rgy

Reaction

General Chemistry I 21

8.3 STRUCTURES OF COORDINATION COMPLEXES361

Octahedral complexes with geometrical isomers(complexes of type MA2B4 (or MA2B2; B is bidentate)

cis-[Co(NH3)4Cl2]+

trans-[Co(NH3)4Cl2]+

cis-[CoCl2(en)2]+

trans-[CoCl2(en)2]+

General Chemistry I 22

Octahedral complexes with mer / fac isomers

(Complexes of type MA3B3)

mer-isomer: Similar ligands define a meridian

of the octahedron

fac-isomer: Similar ligands define a face of an octohedron

-- all three groups are 90° apart.

mer-Co(NH3)3(Cl)3 fac-Co(NH3)3(Cl)3

362

General Chemistry I 23

Tetrahedral complexes

~ Dominant for four-coordinate complexes

~ No geometrical isomers for tetrahedral

complexes of MA2B2

363

Square planar complexes

~ Au3+, Ir+, Rh+, Ni2+, Pd2+, Pt2+

~ cis-[Pt(NH3)2Cl2]

(anticancer drug, ‘cisplatin’)

~ trans-[Pt(NH3)2Cl2]

Linear geometry

~ Ions with d10 configuration: Cu+, Ag+, Au+,

Hg2+

General Chemistry I 24

Optical isomers are molecules that rotate plane polarized light

Enantiomers (Gk. άτιος, “opposite”, and μέρος, “part or porti

on”) are optical isomers whose structures are non-

superimposable mirror images (they lack reflection-rotation

symmetry)

Chiral center (chirality [G. χειρ (kheir), "hand"] ~ handedness) is

a central atom around which enantiomers are formed

A racemic mixture has equal amount of enantiomers (net

rotation of plane polarized light = 0)

Chiral Structures366

General Chemistry I 25

E.g. enantiomers of the [Pt(en)3]4+ ion

E.g. enantiomers of all-cis [Co(NH3)2(H2O)2Cl2]+

366Octahedralcomplexes oftype MA3

(A is bidentate)

Octahedralcomplexes oftype MA2B2C2

General Chemistry I 26

EDTA (ethylenediaminetetraacetate) ion

Hexadentate ligand, sequestering metal ions

Antidote for lead poisoning, preserves freshness of oil

367

General Chemistry I 27

8.4 CRYSTAL FIELD THEORY: OPTICAL AND MAGNETIC PROPERTIES

367

Crystal Field Theory 

~ Ionic description of metal-ligand bonds

~ Ligands are treated as point charges approaching

the central metal ion

Octahedral coordination complexes

Degeneracy of d-orbitals lifted into two groups :

2 2 2, and , , xy yz zz x yd d d d d

General Chemistry I 28

Crystal Field Theory

• Ligands such as a halide or oxide are regarded as an electrostatic,

point charge, or point dipole type, which set up an electrostatic field.

A

B

o = crystalfield splittingenergy

metal d orbitals sphericalcharges

octahedralenvironment

Cr3+

367

General Chemistry I 29

368

General Chemistry I 30

369

Fig. 8.17 An octahedral crystalfield increases the energies ofall five d orbitals, but the increaseis greater for the dz and dx - y

orbitals.2 2 2

General Chemistry I 31

• Electron configuration of octahedral complexes d1-d3

by Hund’s rule

370

General Chemistry I 32

-From d4 to d7 octahedral complexes there are two possibilities,illustrated for d4 (E.g. Mn(III) complexes)

e-e repulsion low-spin configuration

ligand-ligand repulsion

If o is large (strong-field ligands), t2g4 has a lower energy.

: low-spin complex, minimum number of unpaired e-

If o is small (weak-field ligands), t2g3eg

1 has a lower energy.

: high-spin complex, maximum number of unpaired e-

370

high-spin configuration

eg

t2gt2g

eg

Low spin (t2g4) configuration High spin (t2g

3eg1) configuration

E

General Chemistry I 33

Fig. 8.18. Electron configuration for (a) high spin (large o) and (b) low spin (small o) octahedral crystal field splitting energies for Mn(III) complexes

Weak field configuration Strong field configuration H2O weak field ligand CN– strong field ligand

369- Example: d4 octahedral complexes of Mn(III)

5 x degenerated orbitals (3d4)

35

25

oo

o

eg

t2g

dxy dyz dxz

dz dx - y2 2 2Mn(H2O)63+

HIGH SPIN

5 x degenerated orbitals (3d4)

35

25

o

o

o

eg

t2g

dxy dyz dxz

dz dx - y2 2 2

Mn(CN)63-

LOW SPIN

General Chemistry I 34

370

Crystal Field Stabilization Energy (CFSE)

The amount by which the (otherwise equal) energy levels for the d electrons of a metal ion are split by the electrostatic field of the surrounding ligands in a coordination complex.

Energy difference between electrons in an octahedral crystal field and those in the hypothetical spherical crystal field.

General Chemistry I 35

Square planar crystal field

370

sp > 1.6 0

General Chemistry I 36

Tetrahedral crystal field

371

t = 4/9 o

General Chemistry I 37

Fig. 8.20. Correlation diagram showing the relationships among d-orbital energy levels in crystal fields of different symmetries.

372

General Chemistry I 38

Magnetic properties

Paramagnetic compounds

~ One or more unpaired electrons

~ Large, positive magnetic susceptibility

~ Attracted by the magnetic field

→ “weigh” more

~ Prevalent among transition-metal complexes

Diamagnetic compounds ~ All of the electrons are paired

~ Small, negative susceptibility

~ Repelled by the magnetic field

373

Magnetic susceptibility ~ Strength of a sample’s interaction with a magnetic field

General Chemistry I 39

8.5 OPTICAL PROPERTIES AND THE SPECTROCHEMICAL SERIES

374

Transition-metal complexes

~ absorb visible light → colorful

E.g. [Co(NH3)5Cl]2+ ion absorbs greenish yellow light (~530 nm)

Only red and blue light transmitted

→ purple (complementary color)

Wavelength of the strongest absorption, max

d10 complex ~ colorless (no absorption, all d-levels are filled)

High-spin d5 complex ~ weak absorption (spin flip required)

o max, so /E h h hc

General Chemistry I 40

Cr(CO)6 [Co(NH3)5(OH2)]Cl3 K3[Fe(C2O4)3] K3[Fe(CN)6] [Co(en)3]I3

Colors of the hexaaqua complexes of metal ions prepared from their nitrate salts.

E.g. [Co(H2O)6]2+

375

General Chemistry I 41

Spectrochemical series

~ An ordering of ligands according to their ability to cause

crystal field splittings.

Spectrochemical series for ligands

2 3

Weak-field ligands (high spin) Intermediate-field ligands Strong-field ligands (low spin)

I Br Cl F ,OH : NCS en COH O N ,CNH

Spectrochemical series for metal ions

Mn2+ < Ni2+ < Co2+ < Fe2+ < Fe3+ < Co3+ < Mn4+ < Pd4+ < Ir3+ < Pt4+

Crystal field theory cannot explain the spectrochemical series!

376

General Chemistry I 42

8.6 BONDING IN COORDINATION COMPLEXES

377

Valence bond theory

dsp3 hybrid orbitals

~ linear combination of one s, three p atomic orbitals

and the dz2 atomic orbital

~ five equivalent new hybrid orbitals

~ trigonal bipyramid, PF5, CuCl5–

General Chemistry I 43

d2sp3 hybrid orbitals

~ linear combination of one s,

three p atomic orbitals

and dz2, dx2-y2 orbitals

~ six new hybrid orbitals

~ octahedron, SF6

378

General Chemistry I 44

Molecular orbital theory

Ligand field theory ~ Failure of CFT and VB theories to explain the spectrochemical

series

~ MO description for ligands

Construction of MOs for octahedral complexes (of 1st row

D-block metals)

~ Interaction between the metal 4s orbital with six ligand orbitals

→ s and s* orbitals

~ Interaction between three metal p orbitals with three ligand orbitals

→ triply degenerate p and p* orbitals

~ Interaction of the dz2 and dx2-y2 orbitals with ligand orbitals

→ a pair of d and d* orbitals

378

General Chemistry I 45

Fig. 8.27. Formation of bonding MOs from overlap of metal and ligand orbitals.

379

General Chemistry I 46

Bonding MOs

Nonbonding MOs

380

MO correlation diagramfor octahedral Cr(III)complex ([CrCl6]3-): bonding only

Antibonding MOs

General Chemistry I 47

Formation of and * bonds

(1) Interaction between an empty metal d orbital with a filled atomic

ligand p orbital. E.g. 3p orbitals of Cl–

(2) Interaction between a filled metal d orbital with an empty ligand *

antibonding molecular orbital. E.g. CO, CN–

→ metal-to-ligand (M-L) donation or backbonding

- and * MOs: M d orbital - L p orbitalor M d orbital - L * orbital

381

General Chemistry I 48

(3) Overlap of each of the metal nonbonding dxy, dyz,

and dxz orbitals with four ligand p orbitals

→ Formation of three pairs of

bonding and antibonding

MOs, t2g and t2g*.

Fig. 8.30. Bonding MO by constructive overlap of a metal dxy orbital with four ligand p orbitals.

382

General Chemistry I 49

Order of bonding strengths for different ligands

Weak-field ligands (small o)

→ Overlap between occupied p() bonding orbitals of

ligands (Br–, Cl–, CO) with t2g orbitals of metal

→ Increase in energy of t2g and decrease in o

Strong-field ligands (large o)

→ Overlap between unoccupied * antibonding orbitals of

ligands (CO, CN–) with t2g orbitals of metal

→ Lowering of energy of t2g orbitals by back-bonding (M→L) Intermediate-field ligands ~ H2O, NH3

383

General Chemistry I 50

Fig. 8.31. (a) (ML) [or (b) (ML)] donation showing a reduction (or increase) in Δo compared with that from bonding alone.

(a) Slight increase in energy of t2g electrons (in t2g* MOs)

(b) Significant lowering in energy of t2g electrons

due to back-bonding → Electrons of t2g MOs are delocalised

into unoccupied *(L)

383

eg eg*

E t2g*

t2g

Partially filledmetal d orbitals Filled ligand

p () orbitals

Empty ligandp (*) orbitals

donor (M L)ligands

eg eg*

t2g*

t2g

Partially filledmetal d orbitals

Filled ligandp () orbitals

Empty ligandp (*) orbitals

acceptor (M L)ligands

(a) (b)

General Chemistry I 51

Fig. 8.32. Effect of bonding on the energy-level structure for octahedral coordination complexes.

Summary of the MO picture (Ligand Field Theory) of bonding in octahedral coordination complexes

Cl-, Br- ligandse.g. [CrCl6]3–

IIlustrated for V2+,Cr3+,Mn4+

(d3)

CO, CN–, NO+

Ligands e.g.Mn(CN)4

384

H2O, NH3 ligands

e.g. [V(H2O)6]2+

General Chemistry I 52

10 Problem Sets

For Chapter 8,

2, 8, 18, 26, 32, 44, 46, 58, 64, 66