1.1 Why Study Organometallic Chemistry?
Introduction
 coordination chemistry
 study of compounds formed b/w metal ions and other neutral or
1. Introduction
negatively charged molecules (ligand)
 e.g. [Co(NH2CH2CH2NH2)2ClNH3]2+[Cl-]2.
 metal complex
 a charged species consisting of metal ion bonded to one or more
groups of molecules (ligand)
1.1 Why Study Organometallic Chemistry?
1.1 Why Study Organometallic Chemistry?
Introduction
 organometallic compounds
 metal-carbon bonds (e.g., WMe6)
 interface b/w classical organic and inorganic chemistry in dealing with
the interaction b/w inorganic metal species and organic molecules
 metal-organic compound
 organic fragment is bound only by metal-heteroatom bonds [e.g.,
Ti(OMe)4].
 nanoscience and nanotechnology
 benefiting with the use of such compounds as the most common
precursors for nanoparticles
Introduction
 green chemistry
 object of minimizing both energy use and chemical waste in industry
and commerce
 atom economy
 reactions are chosen that minimize the formation of by-products or
unreacted starting materials
1.2 Coordination Chemistry
1.3 Werner Complexes
Stereochemistry
Introduction
 coordination compound of transition metal ions
 most common type of complex
 complex MLn
 ML6
 aqua ions [M(OH2)6]2+ (M = V, Cr, Mn, Fe, Co, or Ni).
 octahedral coordination geometry
 organometallic chemistry
 based on one of the Pythagorean regular solids
 subfield of coordination chemistry
 complex contains an M-C or M-H bond [e.g., Mo(CO)6]
 organometallic species tend to be more covalent
 metal is often more reduced
1.3 Werner Complexes
Stereochemistry
 complex : assembly of metal and ligands
 may have a net ionic charge: [PtCl4]2−)
 together with the counterions: K2[PtCl4]
1.3 Werner Complexes
Stereochemistry
 ligands can donate one lone pair to each of two or more metal ions
 gives polynuclear complexes: compound 1.2
 bridging group: represented by μ as in [Ru2(μ-Cl)3(PR3)6]+
 in some cases, both the cation and the anion may be complex:
 Magnus’ green salt [Pt(NH3)4][PtCl4]
 square brackets, [ ]:
 enclose the individual complex molecules or ions
(L = PR3)
1.3 Werner Complexes
Chelate Effect
1.3 Werner Complexes
Chelate Effect
 Chelate: ligands can have more than one donor atom, ligands most
 chelate ligands: bidentate or polydentate
commonly donate both lone pairs to the same metal to give a ring
compound (Greek word for “claw”): 1.3
 ethylenediamine (NH2CH2CH2NH2, en)
cryptates
sepulchrates
1.3 Werner Complexes
Chelate Effect: Thermodynamic Data for Cd2+ Complexes
[Cd(CH3NH2)4]2+and [Cd(en)2]2+
 most prominent when the total ring size is five or six atoms;
 smaller rings: strained,
 for larger rings
 second complexing atom is farther away
 formation of the second bond may require the ligand to contort
 more complete understanding of chelate effect
 requires the determination of the enthalpies and entropies of these
reactions
1.3 Werner Complexes
Chelate Effect
 formation of the tris chelate releases six NH3 molecules:
 total number of particles increases from four to seven: creates entropy
and so favors the chelate form
 each chelate ring usually: about 105 in the equilibrium constant for
reactions such as Eq. 1.1.
 equilibrium constants for complex formation: formation constants
 the higher the value, the more stable the complex
1.3 Werner Complexes
Chelate Effect: Thermodynamic Data for Cd2+ Complexes
[Cd(CH3NH2)4]2+and [Cd(en)2]2+
1.3 Werner Complexes
Werner Complexes
 pioneered by Nobel Prize winner Alfred Werner
(1866-1919).
 received the Nobel Prize in 1913 for his
유사
큰 차이
큰 차이
coordination theory of transition metal-amine
큰 차이
complexes.
 At the start of the 20th century, inorganic
chemistry was not a prominant field until Werner
studied the metal-amine complexes such as
[Co(NH3)6Cl3].
1.3 Werner Complexes
Werner Complexes
Werner Complexes
Werner Complexes: Observation of Co(NH3)1.3
6Cl3
Compounds
 Werner recognized
 the concept of two valences
 primary valence: oxidation number
 secondary valence: coordination number
 existence of several forms of cobalt-ammonia chloride
 have different color and other characteristics
 chemical formula has three chloride ions per mole
 number of chloride ions that precipitate with Ag+ ions per formula is
not always three : only ionized chloride ions will form precipitate with
e.g.: [Co(NH3)6]Cl3
 3 Cl– : primary valence
 6 NH3: secondary valence
silver ion
 to distinguish ionized chloride from the coordinated chloride
 formulated the complex formula
 explained structure of the cobalt complexes
1.3 Werner Complexes
1.3 Werner Complexes
Werner Complexes: Complex Cation [Co(NH3)6]3+
Werner Complexes
 Alfred Werner proposed
 cobalt ammines (ammonia complexes) the metal ion is surrounded by
six ligands in an octahedral array as in 1.6 and 1.7.
1.3 Werner Complexes
1.3 Werner Complexes
Werner Complexes
Werner Complexes: Number of Isomers of [Co(NH3)6Cl3]
Jørgensen’s model
yellow
purple
green or violet
1.3 Werner Complexes
Werner Complexes
1.3 Werner Complexes
Werner Complexes: Possible Hexacoordinate
Isomers
for [Co(NH3)4Cl2]+
 in 1907, Werner succeeded in making the elusive purple isomer of
[Co(NH3)4Cl2]+
 via the carbonate [Co(NH3)4(O2CO)] in which two oxygens of the
chelating dianion are necessarily cis
1.3 Werner Complexes
1.3 Werner Complexes
Werner Complexes
 Werner resolved optical isomers of [Co(en)2X2]2+
 only an octahedral array can account for the optical isomerism
Werner Complexes: cis- and trans-Isomers
1.3 Werner Complexes
Werner Complexes: Possible Hexacoordinate
Isomers
for [Co(NH3)4Cl2]+
1.3 Werner Complexes
Werner Complexes: Possible Hexacoordinate
Isomers
for [Co(NH3)4Cl2]+
1.3 Werner Complexes
1.3 Werner Complexes
Werner Complexes: Possible Structures for [Pt(NH3)2Cl2]
Werner Complexes
 responded by resolving a complex (1.12)
containing only inorganic elements.
 extraordinarily high specific rotation of
36,000◦
 required
1000
recrystallizations
to
resolve
 won the chemistry Nobel Prize for this
work in 1913
1.4 The Trans Effect
The Trans Effect
1.4 The Trans Effect
The Trans Effect: Initial Attack by Entering Group
at a
Square-Planar Pt(II) Center
Evidence for Associative Reactions
 In the 1920s, Chernaev discovered trans effect
 certain ligands, Lt
 facilitate the departure of a second ligand, L, trans to the first
 facilitate their replacement or substitution, by an external ligand
 are more effective at this labilization
1.4 The Trans Effect
1.4 The Trans Effect
The Trans Effect
The Trans Effect
In 1926, Chernyaev: introduced the concept of trans effect in Pt
chemistry
NH3
- In reactions of sq-py Pt(II) compounds
- ligands trans to chloride are more easily replaced
- than those trans to ligands such as ammonia
NH3
Pt
Cl
NH3
Order of trans effect
CN-
~ CO ~ C2H4 > PH3 ~ SH2 > NO2- > I- > Br- > Cl- > NH3 ~ py > OH- > H2O
1.4 The Trans Effect
The Trans Effect: Stereochemistry and the trans
Effect in
Pt(II) Reactions
1.4 The Trans Effect
The Trans Effect: Stereochemistry and the trans
Effect in
Pt(II) Reactions
1.4 The Trans Effect
1.4 The Trans Effect
The Trans Effect: p-Bonding Interaction between
Metal d
2
Orbital and Suitable Orbitals of Ligand L
The Trans Effect: Sigma-Bonding Effect: trans influence
Order of trans effect based on the relative s-donor properties of the ligands
H- > PR3 > SCN- > I- ~ CH3- ~ CO ~ CN- > Br- > Cl- > NH3
Stronger the Pt-T bond
Weaker the Pt-X bond
stronger
weaker
Ground state is higher in energy
Smaller activation energy
1.4 The Trans Effect
1.4 The Trans Effect
The Trans Effect: Activation Energy and the trans Effect
The Trans Effect: Pi-Bonding Effects
Order of trans effect based on the p-acceptor ability of the ligands
Order of trans effect
CN-
~ CO ~ C2H4 > PR3 ~ H- > CH3- ~ SC(NH2)2 > C6H5- > NO2- ~ SCN- ~
I- > Br- > Cl- > NH3 ~ py ~ OH- ~ H2O
p-bond formation between metal and ligand
transition state is lower in energy
Smaller activation energy
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands
 formation constants for different metal ion (acid)-halide ligand (base)
combinations, where large positive numbers mean strong binding
Poor trans effect
s-bonding effect: weaken bond
p-bonding effect: stabilization of TS
Low ground state
higher ground state
Lower transition state
High transition state
(trans influence)
(trans effect)
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands: Hard-Soft Acid-Base (HSAB)
1. Relative solubilities of halides. solubilities of silver halides in water
decrease, going down the column of halogens in the periodic table:
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands: Hard-Soft Acid-Base (HSAB)
2. Coordination of thiocyanate to metals.
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands: Hard-Soft Acid-Base (HSAB)
3. Equilibrium constants of exchange reactions.
 numerous ions and other groups can act as ligands, forming bonds to
metal ions.
 thiocyanate, SCN-, has the capacity to bond through either its S or N
 when it bonds to a large, highly polarizable metal ion such as Hg2+, it
attaches through sulfur ([Hg(SCN)4]2-);
 when it bonds to smaller, less polarizable metals such as Zn2+, it
Is it possible to predict the relative magnitudes of such equilibrium constants?
attaches through nitrogen ([Zn(NCS)4]2-
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands: Hard and Soft Bases
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands: Hard and Soft Acids
1.6 The Crystal Field
1.5 Soft Versus Hard Ligands
Soft Versus Hard Ligands: Hard and Soft Acids
The Crystal Field
 was originally developed in the 1930s to describe the electronic structure of
metal ions in crystals
 where they are surrounded by anions that create an electrostatic field
with symmetry dependent on the crystal structure
 energies of the d orbitals of the metal ions are split by the electrostatic
field
 approximate values for these energies can be calculated
1.6 The Crystal Field
1.6 The Crystal Field
The Crystal Field
 Crystal field theory (CFT):
 important advance in understanding the spectra, structure, and
magnetism of transition metal complexes
 how the d orbitals of the transition metal are affected by the ligands?
 is an electrostatic approach: to describe the split in metal d-orbital
energies: energy levels that determine the ultraviolet and visible spectra
 crystal field stabilization energy (CFSE), D, (or 10 Dq):
 energy difference b/w the actual distribution of electrons and that for
all electrons in the uniform field levels
The Crystal Field
1.6 The Crystal Field
1.6 The Crystal Field
High Spin Versus Low Spin: Magnetisms
High Spin Versus Low Spin
 in the d6 configuration: to give the low-spin form t2g6eg0
 if D is small enough: to give the high-spin form t2g4eg2.
paramagnetism
ferromagnetism
antiferromagnetism
ferrimagnetism
1.6 The Crystal Field
1.6 The Crystal Field
Inert Versus Labile Coordination
 labile:
 compounds that react rapidly
 lability: constantly undergoing change or something that is likely to
undergo change
 undergo reactions with t1/2 < 1 min: kinetically labile.
 inert or robust:
 compounds that react slowly
 inert is used to describe something that is not chemically active
 kinetically inert
 stable and unstable
 stable compound has large equilibrium constant for formation
Inert Versus Labile Coordination
 substitutionally or kinetically labile or inert:
 inert & labile Oh complexes
 labile exchange t1/2 less than 1 min. at RT
 inert exchange t1/2 more than 1 min. at RT
1.6 The Crystal Field
1.6 The Crystal Field
The Jahn-Teller Effect
Inert Versus Labile Coordination
 Jahn-Teller theorem states that
 degenerate orbitals cannot be unequally occupied
 to avoid these unfavorable electronic configurations, molecules distort
(lowering their symmetry) to render these orbitals no longer degenerate
1.6 The Crystal Field
The Jahn-Teller Effect
 an octahedral Cu(II) complex, containing a d9 ion
1.6 The Crystal Field
The Jahn-Teller Effect
 the expected degrees of Jahn-Teller distortion for different electronic
configurations and spin states are summarized in the following table:
 significant Jahn-Teller effects are observed in complexes of high-spin Cr(II)
(d4), high-spin Mn(III) (d4), Cu(II) (d9), Ni(III) (d7), and low-spin Co(II) (d7)
 Cu(II) complexes generally exhibit significant Jahn-Teller effects; the
distortion is most often elongation of two bonds
z axis에 위치한 오비탈이 영향을 크게 받는다.
1.6 The Crystal Field
Jahn–Teller Effect on a d9 Complex
 elongation, which results in weakening of some metal-ligand bonds, also
affects equilibrium constants for complex formation
 [trans-Cu(NH3)4(H2O)2]2+ is readily formed in aqueous solution as a
distorted octahedron with two water molecules at greater distances than
the ammonia ligands;
1.6 The Crystal Field
Jahn–Teller Effect on a d9 Complex
 formation constants for these reactions show the difficulty of putting the fifth
and sixth ammonias on the metal:
 4개까지 M-N 결합이 M-O 결합보다 더 강하여 치환이 이루어지나, 5개부터
신장이 일어나서 M-N 결합이 M-O 결합에 비해 강하지 않으므로 반응이
약하게 일어난다.
 liquid ammonia is the required solvent for [Cu(NH3)6]2+ formation
Stepwise Stability Constant for Displacement1.6ofThe
H2Crystal
O byField
2+
8
2+
9
NH3 from [Ni(H2O)6] (d ) and [Cu(H2O)6] (d )
1.6 The Crystal Field
Low-Versus High-Field Ligands
 colors of transition metal ions
large Jahn–Teller Effect
 arise from the absorption of light: dπ-dσ energy gap, D
 this gap: of the crystal field strength of the ligands
 high-field ligands
 CO and C2H4 give rise to a large value of D
No Jahn–Teller Effect
 Low-field ligands,
 H2O or NH3, can give such a low D
1.6 The Crystal Field
Low-Versus High-Field Ligands
1.6 The Crystal Field
Odd Versus Even dn Configurations
 spectrochemical series of ligands
 if a molecule has an odd number of electrons
 π-donor ligands such as halide or H2O tend to be weak-field
 if a mononuclear complex-one containing only a single metal atom
 π-acceptor ligands such as CO tend to be strong-field ligands
 not all of them can be paired up
 d7 (e.g., [Re(CO)3(PCy3)2]) guarantees paramagnetism
 in dinuclear complexes,
 the odd electrons on each metal may pair up
 however, as in the diamagnetic d7–d7 dimer, [(OC)5Re−Re(CO)5]
1.6 The Crystal Field
Odd Versus Even
 complexes with an even dn configuration
 can be diamagnetic or paramagnetic depending on whether they are
high or low spin,
 low-spin
diamagnetic
complexes
are
much
more
common
in
organometallic chemistry because the most commonly encountered
ligands are high field
1.6 The Crystal Field
Other Geometries
dn Configurations
1.6 The Crystal Field
Other Geometries: 7-Coordinate Structures
1.6 The Crystal Field
Other Geometries: 8-Coordinate Structures
1.6 The Crystal Field
Other Geometries
 Other Geometry
1.6 The Crystal Field
Other Geometries
 third-row metals
 tend to form stronger M−L bonds
 more thermally stable complexes
 are also more likely to give diamagnetic complexes
1.6 The Crystal Field
1.6 The Crystal Field
Other Geometries
Other Geometries
 the same metal and ligand set in different oxidation states
 low oxidation states are usually accessible only with strong-field ligands
that tend to give a high D
1.7 The Ligand Field
1.6 The Crystal Field
Isoconfigurational Ions
 transition metals tend to be treated as a group rather than as individual
elements
 variable valency of the transition metals leads to many cases of
isoconfigurational ions
 dn ions of the same configuration (e.g., n = 6) show important similarities
independent of the identity of the element.
 d6 Co(III) is closer in properties to d6 Fe(II) than to d7 Co(II)
The Ligand Field
 crystal field picture gives a useful qualitative understanding, but, once
having established what to expect,
 we turn to the more sophisticated ligand field model,
 really a conventional molecular orbital, or MO
 picture for accurate electronic structure calculations
1.7 The Ligand Field
s Donor Ligands
1.7 The Ligand Field
The Ligand Field: Metal AO Matched by Symmetry with LGO
for Oh Complex
 these ligands have an electron pair capable of being donated directly
metal AO
LGO
toward an empty (or partly empty) metal orbital.
N1(s1 + s2 + s3 + s4 + s5 + s6)
1.7 The Ligand Field
1.7 The Ligand Field
The Ligand Field: Metal AO Matched by Symmetry with LGO
The Ligand Field: Metal AO Matched by Symmetry with LGO
for Oh Complex
for Oh Complex
metal AO
metal AO
LGO
LGO
N2(s1 - s3)
N3(s2 - s4)
N4(s5 - s6)
N6(s1 - s2 + s3 - s4)
N5(2s5 + 2s6 - s1 - s2 - s3 - s4)
1.7 The Ligand Field
The Ligand Field: Sigma-Donor Interactions with
Metal d
Orbitals
 Eg (dx2-y2 and dz2) orbitals match the Eg ligand orbitals.
1.7 The Ligand Field
The Ligand Field: Sigma-Donor Interactions with
Metal s,
p Orbitals
 valence s and p orbitals of the metal have symmetry that matches the two
 symmetries match
remaining irreducible representations:
 an interaction between the two sets of Eg orbitals to form a pair of bonding
 s matches A1g
orbitals (eg) and the counterpart pair of antibonding orbitals (eg*)
 there are no ligand orbitals matching the T2g symmetry-whose lobes point between
the ligands-so these metal orbitals are nonbonding
 the set of p orbitals matches T1u
 symmetry match,
 A1g interactions lead to the formation of bonding and antibonding orbitals
(a1g and a1g*),
 T1u interactions lead to formation of a set of three bonding orbitals (t1u) and
the matching three antibonding orbitals (t1u*)
1.7 The Ligand Field
The Ligand Field: Sigma-Donor Interactions with
Metal s,
p, and d Orbitals
1.7 The Ligand Field
The Ligand Field
1.9 Back Bonding
1.9 Back Bonding
Back Bonding: p-Bond Formation in a Linear L-M-L Unit
Back Bonding
 ligands such as NH3
 good σ donors but are not significant π acceptors
 ligands such as CO
 good π acceptor
 π-acid ligands are of very great importance in organometallic chemistry
1.9 Back Bonding
1.9 Back Bonding
p Acceptor Ligands
Back Bonding
 the electron pair on the ligand is
stabilized by the formation of a
bonding molecular orbital, and the
empty metal d orbital is destabilized
in the formation of an antibonding
orbital.
Back Bonding
Back Bonding: s and p Interaction between CO 1.9
and
a
Metal Atom
1.9 Back Bonding
Back Bonding: Effects of p Bonding on Do using a d3 Ion
stabilize t2g and unstabilize
t2g* by bonding, decrease DO
stabilize t2g by bonding, increase DO
1.9 Back Bonding
Back Bonding
1.9 Back Bonding
Back Bonding
 IR spectroscopy: π∗ orbital of CO that is involved in the back bonding
 free CO, 2149 cm−1;
 H3B−CO(s-only), 2178 cm−1.
 Metal complexes: weakening of the C−O bond: A net positive charge
raises ν(CO), and a net negative charge lowers it.
 Cr(CO)6, ν(CO) = 2000 cm−1
 V(CO)6−, 1860 cm−1
 Mn(CO)6+, 2090 cm−1.
 Cr(tren)(CO)3, 1880 cm−1: electron donating effect of trien
1.9 Back Bonding
Back Bonding
Frontier Orbitals
 A comparison of isoelectronic complexes or ligands can be useful in
making analogies and pointing out contrasts.
 picture for CO holds with slight modifications for a whole series of π
acceptor (or soft) ligands,
 isoelectronic complexes
 Series of compounds such as
1.9 Back Bonding
 alkenes, alkynes, arenes, carbenes, carbynes, NO, N2,and PF3
V(CO)6−,
Cr(CO)6,and
Mn(CO)6+
 they have the same number of electrons distributed in very similar
 a filled orbital that acts as a σ donor
 an empty orbital that acts as a π acceptor.
 the highest occupied (HOMO) and lowest unoccupied molecular orbitals
structures.
 isoelectronic ligands:
(LUMO) of L
 CO and NO+ or CO and CN−
 The HOMO of L: a donor to the LUMO of the metal (dσ)
 The LUMO of the ligand accepts back donation from a filled dπ orbital of
the metal
1.9 Back Bonding
Frontier Orbitals
 frontier orbitals,
1.9 Back Bonding
π-Donor Ligands
 ligands such as OR−, F−,and Cl−
 HOMO and LUMO of each fragment
 are π donors
 nearly always dominate the bonding
 as a result of the lone pairs that are left after one lone pair has formed
 strong interactions between orbitals require
 not only that the overlap between the orbitals be large
 but also that the energy separation be small
the M−L σ bond
1.9 Back Bonding
p Donor Ligands
1.9 Back Bonding
π-Donor Ligands
 halide ions may participate in this type of interaction, with an electron pair
donated in an p-fashion to an empty metal d orbital
 the electron pair on the ligand is stabilized by the formation of a
bonding molecular orbital,
 and the empty metal d orbital is destabilized in the formation of an
antibonding orbital
1.9 Back Bonding
π-Donor Ligands
1.10 Electroneutrality
Elcetroneutrality
 In 1948 Pauling proposed the powerful electroneutrality principle.
 the atoms in molecules arrange themselves so that their net charges fall
within rather narrow limits, from about +1 to −1 overall
 tends toward a preferred charge, which differs according to the
electronegativity of the element concerned.
 nonmetals, such as C, N, or O, tend to be closer to −1
 metals, such as Li, Mg, and Fe, tend to be closer to +1
1.10 Electroneutrality
1.10 Electroneutrality
Elcetroneutrality
Elcetroneutrality
 implies that as far as electroneutrality arguments go, an element will bond
best to other elements that have complementary preferred charges
 electropositive element prefers an electronegative one: NaCl and TiO2
 elements with an intermediate electronegativity tend to prefer each
other: HgS and Au metal
 isolated Co3+ ion
 is not a electroneutral species, as it has an excessively high positive
charge
 seek good electron donors as ligands, such as O2− in Co2O3, or NH3, in
the ammine (NH3) complexes
 isolated W(0) atom
 too electron rich for its electronegativity
 it will prefer net electron-attracting ligands such as CO that can remove
electron density → forms p back bonding
1.10 Electroneutrality
Trends with Oxidation State
 the d orbitals will be much more strongly stabilized than the others on going
from the atom to the ion
 this is why the atomic electron configuration for the transition metals
involves s-orbital occupation (e.g., Co, d7s2)
 the configuration of the ion is d6, not d4s2
1.10 Electroneutrality
Schematic Energy Levels for Transition Elements
1.10 Electroneutrality
1.10 Electroneutrality
Periodic Trends
Trends with Oxidation State
 the more electron rich (the more reduced, or low oxidation state) the metal
 alter the orbital energies as we go from left to right in the transition series
complex, the less positive will be the charge on the metal
 for each step to the right, a proton (+ neutron) is added to the nucleus
 destabilize the d orbitals
 this extra positive charge stabilizes all the orbitals
 make them more available for back donation
 the earlier metals are more electropositive: easier to remove electrons
from their less stable energy levels
1.10 Electroneutrality
Periodic Trends
 this time the order is 3d ∼ 4s > 3p
 because the s orbital, having a maximum electron density at the
1.10 Electroneutrality
Periodic Trends
 a large difference between a d0 state and a d2 state, both common in the
early transition metals [e.g., d0 Ti(IV) and a d2 Ti(II)]
nucleus, is more stabilized by the extra protons that we add for each
 d0 oxidation state cannot back bond because it lacks d electrons,
step to the right in the periodic table, than are the p orbitals, which have
 d2 state often has an exceptionally high back-bonding power because
a planar node at the nucleus.
 the d orbitals are stabilized because of their lower principal quantum
number (e.g., 3d vs. 4s and 4p for Fe)
early in the transition series the d orbitals are relatively unstable for the
reasons mentioned above.
 d0 Ti(IV) species (C5H5)2TiCl2: does not react with CO at all
 d2 Ti(II) fragment, (C5H5)2Ti, forms a very stable monocarbonyl,
(C5H5)2Ti(CO), with a very low ν(CO), indicating very strong back
bonding
1.10 Electroneutrality
1.10 Electroneutrality
Periodic Trends
Periodic Trends
 as we go down a group from the first-row transition element to the second
 this trend also extends to the third row,
 as the f electrons that were added to build up the lanthanide elements
row,
 the outer valence electrons become more and more shielded from the
nucleus by the extra shell of electrons that has been added
 more easily lost
are not as effective as s, p, or even d electrons in shielding the valence
electrons from the nucleus
 a smaller change on going from the second-to the third-row elements
 the heavier element will be the more basic
than was the case for moving from the first row to the second
 more electronegative
 high oxidation states will be more stable
1.10 Electroneutrality
Periodic Trends
 Cr(VI) in Na2CrO4 and Mn(VII) in KMnO4; both are powerful oxidizing
agents
 their stable analogs in the second and third rows, Na2MoO4,Na2WO4,and
KReO4, which are only very weakly oxidizing
1.10 Electroneutrality
Periodic Trends
 lanthanide contraction
 increase in covalent radii is larger on going from the first to the second
row than it is on going from the second to the third
1.11 Types of Ligand
1.11 Types of Ligand
π Complexes
Types of Ligand
 most ligands form the M−L σ bond by using a lone pair
 π Complexes
 :PR3 or pyridine: lone pairs are often the HOMO and the most basic
 ethylene has no lone pairs, yet it binds strongly to low-valent metals
 the HOMO is the C=C π bond,
electrons in the molecule
 classical Werner coordination complexes always involve lone-pair donor
 these electrons form the M−L σ bond
 This type of binding is represented as (η2-C2H4)
ligands
 two other types of ligand
 where η represents the hapticity of the ligand,
 π ligand: C2H4
 defined as the number of atoms in the ligand bonded to the metal
 σ ligand: H2
1.11 Types of Ligand
M−L Bonding
M−L Bonding
1.11 Types of Ligand
Linear p Systems: p-Ethylene Complexes: Bonding
in
Metal-Alkene
1.11 Types of Ligand
Linear p Systems: p-Ethylene Complexes: Bonding
in
Metal-Alkene
1.11 Types of Ligand
Linear p Systems: p-Ethylene Complexes: Bonding
in
Metal-Alkene
 ethylene donates electron density to the metal in a s fashion, using its p
bonding electron pair.
 electron density can be donated back to the ligand in a p fashion from a
metal d orbital to the empty p orbital of the ligand
1. C-C distance in complex 137.5 pm in comparison with 133.7 pm in free
ethene.
2. stretch of C=C in complex 1516 cm-1 in comparison with 1623 cm-1 in free
ethene.
 The lengthening of this bond can be explained by a combination of the two
factors involved in the synergistic s donor-p acceptor nature of the ligand:
(a) donation to the metal in a s fashion reduces the electron density in a
filled p bonding orbital within the ligand, weakening the C-C bond;
(b) the back-donation from the metal to the p* orbital of the ligand also
reduces the C-C bond strength by populating this antibonding orbital
1.11 Types of Ligand
Linear p Systems: p-Ethylene Complexes: Rotation
of an
2
h -C2H4
1.11 Types of Ligand
Components of Metal-Dihydrogen Bonding
 stable H2 complexes centered on metals likely to be relatively poor donors
 high oxidation states
 Surrounded by ligands that function as strong electron acceptors like
CO or NO
1.11 Types of Ligand
1.11 Types of Ligand
σ Complexes
Components of Metal-Dihydrogen Bonding
 σ Complexes
 Molecular hydrogen has neither a lone pair nor a π bond, yet it also binds
as an intact molecule to metals: [W(η2-H2)(CO)3L2].
 related σ complexes are formed with C−H, Si−H, B−H, and M−H bonds.
 the basicity of electron pairs decreases in the following order because
being part of a bond stabilizes electrons.
 lone pairs > π-bonding pairs > σ-bonding pairs,
 usual order of binding ability is therefore as follows:
 lone-pair donor > π donor > σ donor.
1.11 Types of Ligand
M−L Bonding
M−L Bonding
1.11 Types of Ligand
Types of Ligand
Phosphines(PH3)
Pi Acceptor Orbitals of Phosphines (revised view)
3s
3pz
3py
3px
3dyz
3py
3dxz
3px
1.11 Types of Ligand
Ambidentate Ligands
Ambidentate Ligands
1.11 Types of Ligand
Ambidentate Ligands
Ambidentate Ligands
1.11 Types of Ligand
Ambidentate Ligands
1.11 Types of Ligand
Spectator Versus Actor Ligands
Ambidentate Ligands
1.11 Types of Ligand
Spectator Versus Actor Ligands
1.11 Types of Ligand
Facial-, Meridional-Isomers: [Ma3b3] Type
Facial isomers have
three identical ligands
on one triangular face.
Meridional isomers
have three identical
ligands in a plane
bisecting the molecule.
1.11 Types of Ligand
Facial-, Meridional-Isomers: [Ma3b3] Type
1.11 Types of Ligand
Spectator Versus Actor Ligands
1.11 Types of Ligand
Spectator Versus Actor Ligands
1.11 Types of Ligand
Facial-, Meridional-Isomers: [Ma3b3] Type
1.11 Types of Ligand
Spectator Versus Actor Ligands
1.8 Types of Ligand
Spectator Versus Actor Ligands