CH302  –  INORGANIC  CHEMISTRY  III  Part  B  –  Organometallic  Chemistry  

Dr  T.  Gadzikwa  

EAN  RULE  EffecJve  Atomic  Number  

§  TransiJon  metals  have  9  valence  orbitals  (1  s  +  3  p  +  5  d)  

§  Upon  bonding  to  a  ligand  set,  there  will  be  a  total  of  9  low  lying  orbitals  

§  The  low  lying  MOs  can  accommodate  up  to  18  valence  electrons  -­‐the  18-­‐electron  rule  

§  With  18  valence  electrons,  1st  row  transiJon  metals  have  the  effec5ve  atomic  number  of  Kr  

§  Complexes  stable  when  they  have  18  valence  electrons  

§  ExcepJons  exist.  However  rule  provides  useful  guidelines  to  the  chemistry  of  organometallic  complexes  

2  

EAN  RULE  Determining  Number  of  Valence  Electrons  

§  Two  ways  of  coun5ng  electrons  in  organometallic  compounds  

§  In  the  donor  pair  method  the  ligand  is  considered  to  donate  electron  pairs  to  the  metal.  Charge  of  ligand  and  formal  oxidaJon  state  of  metal  need  to  be  determined  

 

 

§  In  the  neutral  ligand  method  the  ligand  is  considered  to  donate  electron  pairs  to  the  metal.  Considers  the  number  of  electrons  that  would  be  donated  if  the  ligand  was  neutral  

3  

MClH3C CH3

ClMn+ +$$$2Cl($$$+$$$2CH3(

MClH3C CH3

ClM0 +$$$2Cl$$$$+$$$2CH3

MClH3C CH3

Cl

EAN  RULE  

Donor  Pair:  Determining  OxidaJon  State  §  In  the  donor  pair  method,  determining  the  formal  oxida5on  state  

of  the  metal  is  key  1.  Overall  charge  of  the  compound  

     

2.  Total  charges  of  the  ligands  if  they  were  free:    

 3.  Overall  charge  –  ligand  charges:  

   Oxida&on  state  =  -­‐1  –  (-­‐8)  =  +7      

4  

O

MnO O

O

MnO4-

O

Mn+7O O

O

Overall  charge  =  -­‐1  

Total  ligand  charges  =  4  x  -­‐2  =  8    Mn O2-

O2-O2-

O2-

EAN  RULE  

Determining  OxidaJon  State  §  Examples  

5  

Mo

OC

Mo

CO

OC

CO

Cr

OC CO

COOC

Cl

TaHH 0  –  5  =  +5  

0  –  4  =  +4  

0  –  2  =  +2    

+2/2  =  +1  

Ta H-

H-

Mo

CO

MoCO

CO

CO

Cr

CO CO

COCO

CCl

EAN  RULE  

Charge:

Electrons:

Cl CO Fe

LIGANDS METAL

C O

Donor  Pair  Method  

§  Example:  (η5-­‐C5H5)Fe(CO)2Cl  

6  

MClH3C CH3

ClMn+ +$$$2Cl($$$+$$$2CH3(

FeCl CO

CO

-1 -1 0 0

6 2 2 2

+2  

II  

EAN  RULE  Periodic  Table  

7  

2, or 1 units in length. Therefore, the resultant, R, for these combinations can be written as | l 1 ! l 2 |, | l 1 ! l 2 " 1| or | l 1 " l 2 |.

In heavier atoms, a different type of coupling scheme is sometimes followed. In that scheme, the orbital angular momentum, l , couples with the spin angular momentum, s , to give a resultant, j , for a single electron. Then, these j values are coupled to give the overall angular momentum for the atom, J . Coupling of angular momenta by this means, known as j - j coupling, occurs for heavy atoms, but we will not consider this type of coupling further.

In L-S coupling, we need to determine the following sums in order to deduce the spectroscopic state of an atom:

L l S s m L L L L L Si i i# # $ # # " " " # ! "Σ Σ M , 1, 2, , 0, , , ,… … | | … | |J L S

Note that if all of the electrons are paired, the sum of spins is 0, so a singlet state results. Also, if all of the orbitals in a set are fi lled, for each electron with a positive value of m there is also one having

6C

12.011

1H

1.0079

2He

4.0026

3Li

6.941

4Be

9.0122

5B

10.81

7N

14.0067

8O

15.9994

10Ne

20.179

11Na

22.9898

12Mg

24.305

13Al

26.9815

14Si

28.0855

15P

30.9738

16S

32.06

17Cl

35.453

18Ar

39.948

36Kr

83.80

35Br

79.904

34Se

78.96

33As

74.9216

32Ge

72.59

31Ga

69.72

30Zn

65.38

29Cu

63.546

28Ni

58.69

27Co

58.9332

26Fe

55.847

25Mn

54.9380

24Cr

51.996

23V

50.9415

22Ti

47.88

21Sc

44.9559

20Ca

40.08

19K

39.0983

37Rb

85.4678

38Sr

87.62

39Y

88.9059

40Zr

91.22

41Nb

92.9064

42Mo

95.94

43Tc

(98)

44R

101.07

45Rh

102.906

46Pd

106.42

47Ag

107.868

48Cd

112.41

49In

114.82

50Sn

118.69

51Sb

121.75

52Te

127.60

53I

126.905

54Xe

131.29

86Rn

(222)

85At

(210)

84Po

(209)

83Bi

208.980

82Pb

207.2

81Tl

204.383

80Hg

200.59

79Au

196.967

78Pt

195.09

77Ir

192.22

76Os

190.2

75Re

186.207

74W

183.85

73Ta

180.948

72Hf

178.48

57La*

138.906

56Ba

137.33

55Cs

132.905

87Fr

(223)

88Ra

226.025

89Ac*

227.028

104Rf

(257)

105Ha

(260)

106Sg

(263)

107Ns

(262)

108Hs

(265)

109Mt

(266)

110Ds

(271)

111Rg

(272)

58Ce

140.12

59Pr

140.908

60Nd

144.24

61Pm

(145)

62Sm

150.36

63Eu

151.96

64Gd

157.25

65Tb

158.925

66Dy

162.50

67Ho

164.930

68Er

167.26

69Tm

168.934

70Yb

173.04

71Lu

174.967

103Lr

(260)

102No

(259)

101Md

(258)

100Fm

(257)

99Es

(252)

98Cf

(251)

97Bk

(247)

96Cm

(247)

95Am

(243)

94Pu

(244)

93Np

237.048

92U

238.029

91Pa

231.036

90Th

232.038

9F

18.9984

*LanthanideSeries

*ActinideSeries

8 10

IA1

IIA2

IIIA13

IVA14

VA15

VIA16

VIIA17

VIIIA18

IIIB3

IVB4

VIB6

VIIB7

IB11

IIB12

VB5

VIIIB9

! FIGURE 2.8 The periodic table of the elements.

2.6 Spectroscopic States 57

Fe  =  s2d6,  8  electrons  FeII  =  s0d6,  6  electrons    

EAN  RULE  

Charge:

Electrons:

Cl CO Fe

LIGANDS METAL

C O

Donor  Pair  Method  

§  Example:  (η5-­‐C5H5)Fe(CO)2Cl  

8  

MClH3C CH3

ClMn+ +$$$2Cl($$$+$$$2CH3(

FeCl CO

CO

-1 -1 0 0

6 2 2 2

+2  

II  

EAN  RULE  Neutral  Ligand  Method  

§  Example:  (η5-­‐C5H5)Fe(CO)2Cl  

9  

FeCl CO

CO

Total    

18  

MClH3C CH3

ClM0 +$$$2Cl$$$$+$$$2CH3

Fe0

Charge:

Electrons:

Cl CO

LIGANDS METAL

0 0 0 0

C O

0

5 1 2 2 8

EAN  RULE  

Common  Ligands  and  their  Electron  Counts  

10  

32 GENERAL PROPERTIES OF ORGANOMETALLIC COMPLEXES

model played a dominant role because the oxidation state of the types of com-pound studied could almost always be unambiguously defined. The rise of thecovalent model has paralleled the growth in importance of organometallic com-pounds, which tend to involve more covalent M−L bonds and for which oxidationstates cannot always be unambiguously defined (see Section 2.4). We have there-fore preferred the covalent model as being most appropriate for the majority ofthe compounds with which we will be concerned. It is important to be conversantwith both models, however, because each can be found in the literature withoutany indication as to which is being used, so you should practice counting underthe other convention after you are happy with the first. We will also refer to anyspecial implications of using one or other model as necessary.

Electron Counts for Common Ligands and Hapticity

In Table 2.2 we see some of the common ligands and their electron counts on thetwo models. The symbol L is commonly used to signify a neutral ligand, whichcan be a lone-pair donor, such as CO or NH3, a π-bond donor, such as C2H4, ora σ -bond donor such as H2, which are all 2e ligands on both models. The symbolX refers to ligands such as H, Cl, or Me, which are 1e X ligands on the covalentmodel and 2e X− ligands on the ionic model. In the covalent model we regardthem as 1e X· radicals bonding to the neutral metal atom; in the ionic model, weregard them as 2e X− anions bonding to the M+ cation. Green2 has developeda useful extension of this nomenclature by which more complicated ligands canbe classified. For example, benzene (2.1) can be considered as a combination ofthree C=C ligands, and therefore as L3.∗ The allyl group can be considered as a

TABLE 2.2 Common Ligands and Their Electron Counts

Ligand Type Covalent Model Ionic Model

Me, Cl, Ph, Cl, η1-allyl, NO (bent)a X 1e 2eLone-pair donors: CO, NH3 L 2e 2eπ-Bond donors: C2H4 L 2e 2eσ -Bond donors: (H2) L 2e 2eM−Cl (bridging) L 2e 2eη3-Allyl, κ2-acetate LX 3e 4eNO (linear)a 3e 2ea

η4-Butadiene L2b 4e 4e

=O (oxo) X2 4e 2eη5-Cp L2X 5e 6eη6-Benzene L3 6e 6e

a Linear NO is considered as NO+ on the ionic model; see Section 4.1.bThe alternative LX2 structure sometimes adopted gives the same electron count.

∗Undergraduates will need to become familiar with organic “line notation,” in which only C−Cbonds are shown and enough H groups must be added to each C to make it 4-valent. For example,2.6 represents MCH2CH=CH2.

Neutral Ligand Donor Pair

EAN  RULE  Metal-­‐Metal  Bonds  

§  An  M-­‐M  bond  is  a  covalent  bond  in  which  each  metal  donates  1e-­‐  §  The  metal  gets  +1  electron  for  each  M-­‐M  bond  in  both  counJng  

methods,  eg:  

11  

Total    

18  

Charge:

Electrons:

LIGANDS METAL

5 x CO Mn Mn

5 x 0

5 x 2

0 0

1 7

EAN  RULE  PracJce  

12  

EAN  RULE  Applicability  of  the  Rule  

§  Useful  for  predicJng  the  number  of  ligands  that  will  bind  to  a  metal  

 

§  However,  the  majority  of  metal  complexes  do  not  saJsfy  the  18-­‐electron  rule  

§  Octahedral  complexes  fall  into  three  categories  regarding  18-­‐electron  rule  

13  

Chapter 1: Structure and Bonding: 6

Prepared by Dr. Kevin Shaughnessy, Spring 2011 The University of Alabama, Tuscaloosa, AL

interactions with the carbene carbon)

µ-CYR or CR2

2 -2 4

Imide (M=NR) 2 -2 4 Oxide (M=O) 2 -2 4

Peroxide (terminal or bridging) 2 -2 4

Alkylidine or carbyne, terminal 3 -3 6

µ-Alkylidine

3 -3 6

Nitride 3 -3 6

Application of the 18-electron rule:

The 18-electron rule can be used as predictor for the number of ligands a particular metal will coordinate.

V(CO)6 Cr(CO)6 (CO)5Mn-Mn(CO5 Fe(CO)5 (CO)3Co(µ-CO)2Co(CO)3 Ni(CO)4

d5

17 e d6

18 e d7

18 e d8

18 e d9

18 e d10

18 e The 18 electron rule works best for low-valent metals with small ligands that are strong σ donors and/or π acceptors (i.e., H- and CO). These ligands give large Δ, thus there is a strong preference for filling the dπ orbitals (requiring 18 electrons), and are small enough to allow the metal to be coordinatively saturated.

For complexes that follow the 18 electron rule, it can be used to predict reactivity as we will see throughout the semester.

Note, that just like the octet rule, the 18-electron rule is not an absolute requirement. There are many exceptions.

Common exceptions to the 18 electron rule:

d8 metals: The d8 metals (groups 8 - 11) have a tendency to form square-planar 16 electron complexes. This tendency is weakest for group 8 (Fe(0), Ru(0), and Os(0)) and is very strong for groups 10 and 11 (Pd(II), Au(III)). Square planar, 16 electron complexes of of d8 metals results in completely filled orbitals except the high energy dx 2 − y 2 (see MO discussion below).

MR

RR

CR M

M

O OM

M C RR

CM M

M

M N

Me Pd Me

PMe3

PMe3

Me3P Rh Cl

CO

CO

Me3P Au CH3

CH3

CH3

EAN  RULE  

§  Low  valent  metals  with  small,  high-­‐field  ligands,  eg  carbonyl  or  hydride  compounds  

 

A:  18-­‐Electron  Rule  Obeyed  

14  

Cr(CO)6  

The 18-electron rule and its exceptions (1) Octahedral complexes ‡ can be split into 3 categories: category number of electrons description

A 18 18-electron rule obeyed B 12-18 18-electrons not exceeded C 12-22 18-electron rule disobeyed

A. Those that obey the 18 electron rule Complexes with strong π-acceptor ligands (e.g. [V(CO)6]-, [Cr(CO)6], [W(CO)6], [Mn(CO)6]+) - t2g strongly bonding ∴filled - eg strongly antibonding (due to synergic bonding) ∴empty

3d (eg + t2g)

4s (a1g)

4p (t1u)

∆O

a1g

eg*

t2g

t1u

eg

a1g*

t1u*

6 CO

large

(a1g + t1u + eg)

π* orbitals of CO (t2g)

Cr

t2g*

strongly bonding(filled)

strongly anti-bonding(empty)

- Complexes of strong π-acceptor ligands ‡ tend to obey the 18-electron rule

irrespective of their coordination number (e.g. [Fe(CO)5], [Fe(CO)4]2-, [Ni(PPh3)4]).

- Note: d8 and d10 configurations are exceptions to this rule (see later).

Low  valent:  Already  have  some  d  electrons  ∴  18  electrons  easy  to  reach    Small  ligands:  As  many  as  needed  can  fit  around  metal  ∴  18  electrons  easy  to  reach    Strong  π-­‐acceptor  ligands:  Bonding  orbitals  low,  must  be  filled  for  stability  ∴  <18  electrons  difficult  to  maintain    Large  Δo:  A  lot  of  energy  required  to  fill  eg*  ∴  18  electrons  difficult    to  exceed  

EAN  RULE  B:  18-­‐Electron  Rule  Not  Exceeded  

§  2nd  and  3rd  row/high  valent  TM  complexes  (high  in  the  spectrochemical  series  of  metal  ions)  

§  σ-­‐  or  π  donor  ligands  (low  to  medium  in  the  spectrochemical  series)  

 

15  

W(Me)6  (W6+,  d0,  12  e)    

B. Octahedral complexes with 12-18 electrons 2nd and 3rd row TM complexes (high in the spectrochemical series of metal ions) with σ-donor or π-donor ligands (low to medium in the spectrochemical series) - t2g non-bonding or weakly anti-bonding (because the ligands are either σ-donors

or π-donors) ∴ t2g can contain from 0 to 6 electrons - eg fairly strongly antibonding (because 2nd/3rd row TMs bond more effectively to

the ligands)∴eg empty

e.g. [ZrF6]2- (Zr4+, d0, 12 e), [PtF6]2- (Pt4+, d6, 18 e), [OsCl6]2- (Os4+, d4, 16 e), [WMe6] (W6+, d0, 12 e), [Zr(OH2)6]3+ (Zr3+, d1, 13 e)

3d (eg + t2g)

4s (a1g)

4p (t1u)

∆O

a1g

eg*

t2g

t1u

eg

a1g*

t1u*

6 Me

(a1g + t1u + eg)

W

Octahedral, σ-donor ligands (e.g. [WMe6])

moderate ∆o for 2nd or 3rdrow transition metal

t2g non-bonding

eg* is reasonablyantibonding for a 2nd or 3rd row transition metal

High  valent:  Have  few  d  electrons  ∴  18  electrons  difficult  to  reach    Ligands  not  π-­‐accepJng:  t2g  non-­‐bonding,  not  a  good  acceptor  ∴  no  drive  to  reach  18  electrons    Moderate  Δo:  Reasonable  energy  required  to  fill  eg*  ∴  18  electrons  difficult  to  exceed  

EAN  RULE  C:  18-­‐Electron  Rule  Disobeyed    

§  Metal  alracts  as  many  L  to  try  to  counterbalance  its  +ve  charge  (think  neutral  ligands)  

§  1st  row  TM  complexes  (low  in  the  spectrochemical  series  of  metal  ions)  

§  σ-­‐  or  π  donor  ligands  (low  to  medium  in  the  spectrochemical  series)  

 

16  

[Cu(OH2)6]2+  (Cu2+,  d9,  21  e)    

C. Octahedral complexes with 12-22 electrons 1st row TM complexes (low in the spectrochemical series of metal ions) with σ-donor or π-donor ligands (low to medium in the spectrochemical series) = 1st row transition metal coordination complexes {e.g. [TiF6]2- (Ti4+, d0, 12 e), [Co(NH3)6]3+ (Co3+, d6, 18 e), [Cu(OH2)6]2+ (Cu2+, d9, 21 e)} - t2g non-bonding or weakly anti-bonding (because the ligands are either σ-donors

or π-donors) ∴t2g can contain from 0 to 6 electrons - eg only weakly antibonding (because 1st row TMs don’t bond as effectively to the

ligands) ∴eg can contain from 0 to 4 electrons

3d (eg + t2g)

4s (a1g)

4p (t1u)

a1g

eg*

t2g

t1u

eg

a1g*

t1u*

(a1g + t1u + eg)

π orbitals (t2g)

Cu

Octahedral, π-donor ligands (e.g. [Cu(OH2)6]2+)

t2g*

6 OH2

small ∆o

eg* is only moderatelyantibonding for a

1st row transition metal

t2g only weaklyanti-bonding

Many  ligands:  Large  electron  contribuJon  from  ligands  ∴  18  electrons  easy  to  exceed    Small  Δo:  Lille  energy  required  to  fill  eg*  ∴  18  electrons  easy  to  exceed  

EAN  RULE  Tetrahedral  Complexes  

§  Cannot  exceed  18  e-­‐s:  no  low  lying  MOs  that  can  be  filled  to  obtain  tetrahedral  complexes  with  >18  electrons  

§  AddiJonally,  TM  complex  with  maximum  d10  will  receive  8  electrons  from  4  ligands,  for  a  total  of  18  electrons  

§  ∆t  small  (≈4/9  ∆o),  no  preference  for  e  or  t2  orbitals  to  be  filled  (can  have  8-­‐18  electrons)  

 

17  

Ni(PPh3)4  (Ni0,  d10,  18-­‐electron  complex)    

Tetrahedral Complexes

• Tetrahedral complexes cannot exceed 18 electrons because there are no low

lying MOs that can be filled to obtain tetrahedral complexes with >18 electrons. In addition, a transition metal complex with the maximum of 10 d-electrons, will receive 8 electrons from the ligands ‡ a total of 18 electrons.

• ∆t is small (~4/9 ∆o), so there is no particular preference for the e or t2 orbitals

to be filled (can have 8-18 electrons) – similar to class C octahedral complexes. Tetrahedral - e.g. [Ni(PPh3)4] (Ni0, d10, 18-electron complex)

3d (eg + t2g)

4s (a1g)

4p (t1u)

a1

e

t2*

t2

a1*

t2*

Ni

(a1g + t1u + eg)

∆t ∆t ~ 4/9 ∆o

4 PPh3

EAN  RULE   d8,  Square-­‐Planar  Complexes    §  d8  metals  (groups  8  -­‐  11)  tend  to  form  square-­‐planar,  16  electron  

complexes    

§  Square-­‐planar,  16  electron  complexes  results  in  completely  filled  orbitals  except  the  high  energy  dx2  −  y2    

§  Thus,  16  electrons  produces  a  more  stable  complex  

 

18  

The  electron  pairs  from  the  4L  occupy  the  bonding  orbitals.  

The  dxy,  dxz,  dyz  and  dz2  orbitals  are  either  weakly  bonding,  non-­‐bonding,  or  weakly  anJbonding.  

The  dx2-­‐y2  orbital  is  anJ-­‐bonding,  and  if  filled,  will  weaken  the  σ  bonds  with  the  ligands.  

M ML4 4L

d

s

p

σdxy (b2g)

dxz, dyz (eg)

dz2 (a1g)

dx2-y2 (b1g)

EAN  RULE  

§  Tendency  is  weakest  for  group  8  (Fe0,  Ru0,  and  Os0)  and  is  very  strong  for  groups  10  and  11  (PdII,  AuIII)  

 

§  StabilisaJon  in  square-­‐planar  geometry,  increases  with  Δο  

§  For  Group  10,  LFSE  overcomes  desire  for  addiJonal  coordinaJon  (increased  bond  energy)  

§  Increasing  charge,  increases  Δο  (as  does  going  down  the  rows)  

d8,  Square-­‐Planar  Complexes    

19  

36 GENERAL PROPERTIES OF ORGANOMETALLIC COMPLEXES

20e. For the 18e rule to be useful, we need to be able to predict when it will beobeyed and when it will not.

The rule works best for hydrides and carbonyls because these are stericallysmall, high-field ligands. Because they are small, as many generally bind as arerequired to achieve 18e. With high-field ligands, ! for the complex will be large.This means that the d∗

σ orbitals that would be filled if the metal had more than18e are high in energy and therefore poor acceptors. On the other hand, the dπ

orbitals that would have to give up electrons if the molecule had less than 18eand are low in energy because of π bonding by CO (or, in the case of H, becauseof the very strong σ bond and the absence of repulsive π interactions with lonepairs). The dπ level is therefore a good acceptor, and to be stable, a complexmust have this level filled (otherwise the electrophilic metal will gain electronsby binding more CO, or the solvent or some functional group in the ligands untilthe 18e configuration is attained).

Conversely, the rule works least well for high-valent metals with weak-fieldligands. In the hexaaqua ions [M(H2O)6]2+ (M = V, Cr, Mn, Fe, Co, Ni), thestructure is the same whatever the electron count of the metal and so must bedictated by the fact that six H2O’s fit well around a metal ion. H2O has two lonepairs, one of which it uses to form a σ bond. This leaves one remaining on theligand, which acts as a π donor to the metal and so lowers !; H2O is therefore aweak-field ligand. If ! is small, then the tendency to adopt the 18e configurationis also small because it is easy to add electrons to the low-lying d∗

σ or to removethem from the high-lying dπ .

An important class of complexes follow a 16e, rather than an 18e, rule becauseone of the nine orbitals is very high lying and is usually empty. This can happenfor the d8 metals of groups 8–11 (Table 2.3). Group 8 shows the least andgroup 11 the highest tendency to become 16e. When these metals are 16e, theynormally adopt the square planar geometry, but large distortions can occur.3 Someexamples of 16e complexes of this sort are RhClL3, IrCl(CO)L2, PdCl2L2, and[PtCl4]2−, [AuMe4]− (L = 3◦ phosphine).

TABLE 2.3 The d 8 Metals that can Adopt a 16eSquare Planar Configuration

Group

8 9 10 11

Fe(0)a Co(I)b Ni(II) Cu(III)c

Ru(0)a Rh(I)b Pd(II) —Os(0)a Ir(I)b Pt(II) Au(III)

a These metals prefer 18e to 16e.bThe 16e configuration is more often seen, but 18e complexesare common.cA rare oxidation state.

EAN  RULE  

§  M(0)  d8  metal  ions  tend  to  obey  the  18-­‐electron  rule,  M(II)  d8  metal  ions  almost  never  do,  and  M(I)  d8  metal  ions  someJmes  do.  

 

d8,  Square-­‐Planar  Complexes    

20  

   M(0)                          M(I)                                    M(II)    Fe(0),  Ru(0),  Os(0)                  Co(I),  Rh(I),  Ir(I)                            Ni(II),  Pd(II),  Pt(II)    

     Obey              SomeJmes  obey                          Almost  never  obey    M  =  Fe,  Ru,  Os   M  =  Co,  Rh,  Ir   M  =  Ni,  Pd,  Pt  

C MCC

C

C

O

O

O

O

O

Trigonal bipyramidal18 electrons

Ph3P MPPh3

PPh3

H

C

OTrigonal bipyramidal

18 electrons

MHPh3P

PPh3CO

Square Planar16 electrons

MClPh3P

PPh3CO

Square Planar16 electrons

EAN  RULE  Early  TransiJon  Metals  Complexes  

§  E.g.,  with  d0  metals  it  is  oten  not  possible  to  fit  the  number  of  ligands  necessary  to  reach  18  electrons  around  the  metal  

 

21  

(4) Steric Effects and Early Transition Metal Compounds

• Steric effects can produce low-coordinate (not many ligands) complexes which

often have <18 electrons.

ScIIIN

NN

SiMe3

SiMe3

SiMe3

SiMe3

Me3Si

Me3Si

Cl MII

Si

SiMe3Me3Si

SiMe3

Si

SiMe3Me3Si

SiMe3

M = Cr2+ d4 10 electronsM = Mn2+ d5 11 electronsM = Fe2+ d6 12 electrons

M = Sc3+ d0 6 electrons

• For early transition metals (e.g. with d0 metals) it is often not possible to fit the

number of ligands necessary to reach 18 electrons around the metal.

ScIII Cl

PMe3

Me

WVI MeMeMe Me

Me

CO

V0 COOCOC CO

CO

W6+ d0 12 electrons V0 d5 17 electrons Sc3+ d0 16 electrons

EAN  RULE  Periodic  Trends  

 

22  

Transition-metal Organometallics2

Transition metals are never on time

Early

Middle

Latenearly  always  <  18e  

mostly  18e  

18e  or  16e  (square  planar)  

EAN  RULE  Steric  Effects  

§  Bulky  ligands  can  prevent  the  approach  of  the  full  complement  of  ligands  that  would  allow  the  metal  to  achieve  the  18  electron  configuraJon  

 

23  

(4) Steric Effects and Early Transition Metal Compounds

• Steric effects can produce low-coordinate (not many ligands) complexes which

often have <18 electrons.

ScIIIN

NN

SiMe3

SiMe3

SiMe3

SiMe3

Me3Si

Me3Si

Cl MII

Si

SiMe3Me3Si

SiMe3

Si

SiMe3Me3Si

SiMe3

M = Cr2+ d4 10 electronsM = Mn2+ d5 11 electronsM = Fe2+ d6 12 electrons

M = Sc3+ d0 6 electrons

• For early transition metals (e.g. with d0 metals) it is often not possible to fit the

number of ligands necessary to reach 18 electrons around the metal.

ScIII Cl

PMe3

Me

WVI MeMeMe Me

Me

CO

V0 COOCOC CO

CO

W6+ d0 12 electrons V0 d5 17 electrons Sc3+ d0 16 electrons

EAN  RULE  Strong  Oxidants  or  Reductants    

§  Many  18  electron  complexes  can  be  reduced  or  oxidised  to  give  17  or  19  electron  complexes  

§  Such  compounds  are  oten  good  oxidising  or  reducing  agents  (i.e.  they  want  to  get  back  to  being  18-­‐electron  compounds).    

 

24  

(5) Strong oxidants or reductants

• Many 18 electron complexes can be reduced or oxidised to give 17 or 19

electron complexes. Such compounds are often good oxidising or reducing

agents (i.e. they want to get back to being 18-electron compounds).

FeII FeIII

FerroceneFeII, d6, 18 electronsOrange, diamagnetic

Ferrocenium cationFeIII, d5, 17 electrons

Dark blue, paramagneticOxidantPerfectly Happy !

+ e-

- e-

CoIII CoII

CobaltoceneCoII, d7, 19 electrons

Dark purple, paramagneticStrong Reductant

Cobaltocenium cationCoIII, d6, 18 electronsYellow, diamagnetic

Perfectly Happy !

+ e-

- e-