1. Lecture 14 — Organometallic Ligands and Bonding
Organometallic Basics
An Organometallic Complex contains at least one M—C bond
Includes ligands: CO, NO, N2, PR3, H2
Both s and p bonding between M and C occur
History
Zeise’s Salt synthesized in 1827 = K[Pt(C2H4)Cl3] • H2O
Confirmed to have H2C=CH2 as a ligand in 1868
Structure not fully known until 1975
Ni(CO)4 synthesized in 1890 (Ludwig Mond of Brunner Mond & Co.)
Grignard Reagents (XMgR) synthesized about 1900
Accidentally produced while trying to make other compounds
Utility to organic synthesis recognized early on

2. 4) Ferrocene synthesized in 1951 (several researchers)
a) Modern organometallic chemistry begins with this discovery
Many new ligands and reactions produced ever since
5) Organometallic chemistry has really been around for millions of years
a) Naturally occurring cobalimins contain Co—C bonds
b) Vitamin B12

3. Geometric isomers of octahedral complexes MX3Y3
fac-[CoCl3F3]3–
mer-[CoCl3F3]3–

4. Geometric isomers of octahedral complexes MX2Y4
cis-[CoCl3F3]3–
trans-[CoCl3F3]3–
Also holds true for square planar complexes MX2Y2

5. Naming Coordination Compounds
A complex is a substance in which a metal atom or ion is associated with a group of neutral molecules or anions called ligands. Coordination compounds are neutral substances (i.e., uncharged) in which at least one ion is present as a complex.
A. To name a coordination compound, no matter whether the complex ion is the cation or the anion, always name the cation before the anion. (This is just like naming an ionic compound.)
B. In naming the complex ion:
Name the ligands first, in alphabetical order, then the metal atom or ion. Note: The metal atom or ion is written before the ligands in the chemical formula.

2. The names of some common ligands are listed in Table 1.

6. Table 1. Names of Some Common Ligands
Anionic Ligands Names   Neutral Ligands Names
Br- bromo   NH3 ammine
F- fluoro   H2O aqua
O2- oxo   NO nitrosyl
OH- hydroxo   CO carbonyl
CN- cyano   O2 dioxygen
C2O42- oxalato   N2 dinitrogen
CO32- carbonato   C5H5N pyridine
CH3COO- acetato   H2NCH2CH2NH2 ethylenediamine

7. More Nomenclature
3. Greek prefixes are used to designate the number of each type of ligand in the complex ion, e.g. di-, tri- and tetra-. If the ligand already contains a Greek prefix (e.g. ethylenediamine) or if it is polydentate ligands (ie. can attach at more than one binding site) the prefixes bis-, tris-, tetrakis-, pentakis-, are used instead. The numerical prefixes are listed in Table 2.
Table 2. Numerical Prefixes
Number Prefix Number Prefix Number Prefix
1 mono 5 penta (pentakis) 9 nona (ennea)
2 di (bis) 6 hexa (hexakis) 10 deca
3 tri (tris) 7 hepta 11 undeca
4 tetra (tetrakis) 8 octa 12 dodeca

8. More nomenclature
4. After naming the ligands, name the central metal. If the complex ion is a cation, the metal is named same as the element. For example, Co in a complex cation is call cobalt and Pt is called platinum. If the complex ion is an anion, the name of the metal ends with the suffix –ate. For example, Co in a complex anion is called cobaltate and Pt is called platinate. For some metals, the Latin names are used in the complex anions e.g. Fe is called ferrate (not ironate).
Table 3: Name of Metals in Anionic Complexes
Name of Metal Name in an Anionic Complex
Iron Ferrate
Copper Cuprate
Lead Plumbate
Silver Argenate
Gold Aurate
Tin Stannate
5. Following the name of the metal, the oxidation state of the metal in the complex is given as a Roman numeral in parentheses.

9. Example of nomenclature
[Cr(NH3)3(H2O)3]Cl3 Answer: triammine triaqua chromium(III) chloride Solution: The complex ion is inside the parentheses, which is a cation. The ammine ligands are named before the aqua ligands according to alphabetical order. Since there are three chlorides binding with the complex ion, the charge on the complex ion must be +3 (since the compound is electrically neutral). From the charge on the complex ion and the charge on the ligands, we can calculate the oxidation number of the metal. In this example, all the ligands are neutral molecules. Therefore, the oxidation number of chromium must be same as the charge of the complex ion, +3.

10. Ligands and Nomenclature
Common Organic Ligands
Binding Modes
Bridging is possible with organometallic ligands

11. Different numbers of atoms of the organometallic ligand may be involved in bond
The letter h (eta) represents the number of atoms the ligand contacts

12. Stable electron configurations for organometallic compounds: The 18-electron (or 16-electron) rule
Counting Electrons
The octet rule governs organic and simple ionic compounds: s + 3p orbital
The 18-electron rule governs organometallics (with many exceptions)
s + 3p + 3d orbitals
Donor ligands provide the electrons other than the d-electrons

13. 3) The “Donor Pair” method of electron counting
a) Common organometallic ligands are assigned an electron count and charge; those that are commonly ions are treated as such
b) The overall charge on the complex must equal the total charge on ligands plus the charge on the metal; this helps determine d-electron count of metal
c) Add up all electrons from metal d orbitals and ligands to find total e- count

14. Examples of Electron Counting
Cr(CO)6
Total charge on ligands = 0, so charge on Cr = 0, so Cr = d6
6 CO ligands x 2 electrons each = 12 electrons
Total of 18 electrons
(h5-C5H5)Fe(CO)2Cl
Total charge on ligands = 2-, so Fe2+ = d6
(h5-C5H5- = 6) + (2CO x 2 = 4) + (Cl- = 2) = 12 electrons
Charged complex: [Mn(CO)6]+
Total ligand charge = 0, so Mn+ = d6
12 electrons from 6 CO ligands gives a total of 18 electrons
M—M Bond: (CO)5Mn—Mn(CO)5
Each bond between metals counts 1 electron per metal: Mn—Mn = 1 e-
Total ligand charge = 0, so Mn0 = d7
5 CO ligands per metal = 10 electrons for a total of 18 electrons per Mn

15. Justification for and exceptions to the 18-electron Rule
MO Theory predicts that 18 electrons fill bonding orbitals
This number is more stable than more (filling antibonding orbitals) or less Do

16. When is the 18-electron rule most valid?
a) With octahedral complexes of large Do.
Ligands are good s-donors and good p-acceptors (CO)
Exceptions to the 18-electron rule are common
Weak field ligands with small Do make filling eg* possible ( > 18e-)
p-donor ligands can make t2g antibonding ( < 18 e-)

17. 5). Square Planar Complexes (d8) follow a 16-electron rule
5) Square Planar Complexes (d8) follow a 16-electron rule. 18 electrons would destabilized the complexes by filling the high energy dx2-y2 orbital.

18. Carbonyl Complexes (CO)
Bonding
Review of CO Molecular Orbitals
HOMO resides mostly on C = s-donation
LUMO resides mostly on C = p-acceptance
Reinforce each other and provide strong bonding
Bonding of CO to a Metal

19. Characteristics of CO complexes
Infrared Spectroscopy
Free CO stretch n = 2143 cm-1
Cr(CO)6 CO stretch n = 2000 cm-1 because p-back donation from metal weakens the CO bond by adding e- to antibonding p* orbital
Negative charge on complex further weakens CO bond:
[V(CO)6]- n = 1858 cm [Mn(CO)6]+ n = 2095 cm-1
d) Bridging CO further weakened by extra p-back donation (e- count = 1/M)
X-Ray Crystallography
Free CO bond length = pm
M—CO carbonyl bond length = 115 pm

20. Synthesis and Reactions of CO Complexes
Carbonyl complexes of most metals exist.
Most obey the 18-electron rule
Bridging decreases down the periodic table as d-orbitals become larger.
Synthesis
Direct reaction: Ni CO Ni(CO)4 Toxic, used to purify Ni
Reductive Carbonylation: CrCl3 + CO + Al Cr(CO)6 + AlCl3
Thermal/Photochemical: 2 Fe(CO)5 + hn Fe2(CO)9 + CO
Reactions: useful for the synthesis of other compounds by substitution of CO
Cr(CO) PPh Cr(CO)5(PPh3) CO
Re(CO)5Br en Re(CO)3(en)Br CO

22. Short note on thermodynamics
Equilibrium constants
Given the equilibrium reaction MLn–1 + L MLn, we can define the complex formation constant for adding that extra ligand is
Kfn = [MLn]/[MLn-1][L]
The overall formation constant will be related to M + n L MLn and
Kf = bn = [MLn]/[M][L]n
Dissociation constants may be defined similarly; the reverse of the first equation gives Kdn = [MLn-1][L]/[MLn]

23. Thermodynamics and kinetics
II. The relationship between equilibrium constants K and rate constants k is, for kf the reaction A + B AB, is Keq = kf/kr kr To study the reactivity of ligands on a metal, it is critical to know how easily the ligand molecule is removed from the metal and replaced by another molecule of the same ligand; this is called an exchange reaction, and isotopically-labelled atoms are used to monitor the rate of exchange kex M(H216O)5(H218O) + H216O M(H216O)6 + H218O

24. Kinetics of Ligand Exchange

25. The dn configurations which benefit the most from ligand field stabilization energy

26. high spin d5 d3 kex= 2 x 10-6 sec-1 kex= 2 x 107 sec-1

27. LFSE
Cr2+ kex= 2 x 107
hs d4 ion
–6Dq
Odd # eg* electrons
Cu2+ kex= 2 x 107
d9 ion
–6Dq
Odd # eg* electrons
weakened bond
• Jahn Teller distorted transition-metal complexes
tend to undergo ligand substitution
RAPIDLY= substitution labile

28. electron is stabilized by –1/2δ1
distortion is favored
for high spin d4

29. electron is stabilized by –1/2δ1
distortion is favored
for d9

32. Ligands Similar to CO
CN- (cyanide) is isolectric to CO
Stronger s-donor and slightly weaker p-acceptor than CO
More stable with M+ due to –1 charge than M0 (which favors CO)
Considered a classical ligand rather than organometallic for this reason
NN (dinitrogen) is isoelectric to CO
Weaker s-donor and weaker p-acceptor than CO, so doesn’t bind well
Very important in Nitrogen Fixation, so much research centers on complexes
NO+ (nitrosyl) is isoelectric to CO
Similar to CO in s-donor and p-acceptor properties
Electron counting scheme considers linear NO complexes as 2 e- NO+
Electron counting scheme considers bent NO complexes as 2 e- NO-

33. Hydride and Dihydrogen Complexes
Hydride Complexes
M—H bonding is s-donation only from H- (2 electron, -1 charge)
Synthesis
Reaction with H2: Co2(CO) H HCo(CO)4
Reduction of carbonyl complex followed by addition of H+
Co2(CO) Na Na[Co(CO)4]
2 Na[Co(CO)4] + H HCo(CO)4
Dihydrogen Complexes
First Complex characterized in Mo(CO)3(PR3)(H2)
Bonding
s-donation from H2 s molecular orbital
p-acceptance from H2 s* molecular orbital
H—H bond is weaker and longer than free H2 (82-90 pm vs. 74 pm)
Electron-rich metals can completely rupture the H2 bond by p-back donation
Other good p-acceptor ligands on the metal helps stabilize the H2—M complex