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

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

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

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

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.

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

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

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.

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.

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

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

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

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

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

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

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-)

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.

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

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-1 [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 = 112.8 pm M—CO carbonyl bond length = 115 pm

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 + 4 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)6 + PPh3 Cr(CO)5(PPh3) + CO Re(CO)5Br + en Re(CO)3(en)Br + 2 CO

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]

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

Kinetics of Ligand Exchange

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

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

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

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

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

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-

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)8 + H2 2 HCo(CO)4 Reduction of carbonyl complex followed by addition of H+ Co2(CO)8 + 2 Na 2 Na[Co(CO)4] 2 Na[Co(CO)4] + H+ HCo(CO)4 Dihydrogen Complexes First Complex characterized in 1984 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