Outline of Topics on Coordination

You must be familiar with common , their charge, and dn for each T-metal Ti V Cr Mn Fe Co Ni Cu Zn +2 ion 2 3 4 5 6 7 8 9 10 1. Structures of the basic Coordination Geometries 2 thru 8 2. Stereoisomers in Octahedral and Square Planar Complexes 3. VBT, CFT, and MO Theory in brief /applications 4. Applications of LFSE , magnetism, and spectra 5. Mechanisms of Octahedral Substitution in Co(III) and Cr(III) 6. Mechanism of Square Planar Substitution in Pt(II) the trans effect 7. Electron Transfer Reactions - Marcus theory. Nobel Prizes related to 1901 van’t Hoff LeBel Tetrahedral carbon/ stereochemistry (1867) 1913 Alfred Werner Coordination chemistry (1893) 1912 RMgX reagent 1954 L. Pauling nature of chemical bond 1962 Hemoglobin structures 1963 Ziegler & Natta Titanium catalysts for stereoreg polymerization 1964 Dorothy Crowfoot Hodgkin Vitamin B12 structure (Co) 1973 G. Wilkinson and E. O. Fischer ferrocene and sandwiches 1979 H. C. Brown, G. Wittig hydroboration, P ylids 1983 electron transfer reactions of metal complexes 1985 H. Hauptman and J. Karle Direct methods to solve phase problem 1987 Lehn, Pedersen, Cram supramolecular chem/crown ether/host guest 1992 Rudy Marcus adiabatic theory of electron transfer 1996 Kroto, Curl, Smalley C60 2001 Nyori, Knowles, Sharpless chiral metal catalysts Rh, Ti 2005 Grubbs, Chauvin, Schrock metathesis / carbenes Ru, Mo 2010 Heck, Suzuki, Negishi organometallic catalysis PdP4 Kolbe on van’t Hoff http://ursula.chem.yale.edu/~chem125/125/history/Kolbe.html

I have recently published an article in Journal für praktische Chemie (14 , 288 ff.) giving as one of the reasons for the contemporary decline of chemical research in Germany the lack of well-rounded as well as thorough chemical education. Many of our chemistry professors labor with this problem to the great disadvantage of our science. As a consequence of this, there is an overgrowth of the weed of the seemingly learned and ingenious but in reality trivial and stupefying natural philosophy. This natural philosophy, which had been put aside by exact science, is at present being dragged out by pseudoscientists from the junk-room which harbors such failings of the human mind, and is dressed up in modern fashion and rouged freshly like a whore whom one tries to smuggle into good society where she does not belong.

A J. H. van't Hoff who is employed at the Veterinary School in Utrecht appears to find exact chemical research not to his taste. He deems it more convenient to mount Pegasus (evidently loaned from the Veterinary School) and to proclaim in his "La chimie dans l'espace" how, to him on the chemical Parnassus which he ascended in his daring flight, the atoms appeared to be arranged in the Universe. Werner Theory of Coord. Complexes 1893 at age 26 # AgCl ppt # ions from Conductivity Werner

CoCl 3 - 6 NH 3 3 4 luteo [Co(NH 3)6]Cl 3

- 5 NH 3 2 3 purpeo [Co(NH 3)5 Cl]Cl 2 1 - 4 NH 3 1 2 violo/praeseo [Co(NH 3)4Cl 2]Cl *

IrCl 3 - 3 NH 3 0 0 Ir(NH 3)3Cl 3

PtCl 4 -6 NH 3 4 5 [Pt(NH3)6]Cl 4

- 5 NH 3 3 4 [Pt(NH 3)5Cl]Cl 3

- 4 NH 3 2 3 [Pt(NH 3)4Cl 2]Cl 2 *

-3 NH 3 1 2 [Pt(NH 3)3Cl 3]Cl *

-2 NH 3 0 0 Pt(NH 3)2Cl 4 *

-KCl - NH 3 0 2 K[Pt(NH 3)Cl 5] 2 - 2 KCl 0 3 K 2[PtCl 6]

1 cis or trans-dichlorotetraamminecobalt (III) chloride 2. potassium hexachloroplatinate(IV) CHEM 3030 LAB EXPTS

1. Synthesis of [Co(NH 3)4(CO 3)]NO 3 , [Co(NH 3)5Cl]Cl 2 IR, Conductivity

2. [Cr(NH 3)6](NO 3)3 liq NH 3, UV-Vis

3. Magnetic Susceptibility Gouy and Evans Methods

2+ 4. Linkage of Co(NH 3)5(NO 2) UV-Vis, IR

5. X-ray Structure of an Iron Macrocyle in P2 12121 SHELX software

6. Coordination Chemistry of Nickel UV-Vis , IR, or magnetism Common Ligands HS 184 , 204 3rd - monodentates: aqua, halides, NH 3, CN , PR 3, thf, py, dmso, NCS - bidentates: en, acac -, oxalate 2-, bipy, phen, diphos, glycinate - - - 2- , dmgh mono or bi : RCOO CO 3 polydentates : dien, trien, porphyrin 2-, Pc 2-, edta 4-, 18crown, cyclam

Paramagnetism HS 579 673 3rd µ in Bohr magnetons = sqr(n(n+2) where n = # unpaired e’s 1. 1.73 2. 2.84 3. 3.87 4. 4.89 5. 5.9 BM spin only values Bonding Approaches to Coordination Complexes

1. Lewis. Electron pair bond concept. Ligands are e-pair donors, metals e- pair acceptors. dative bond -both electrons in bond come from .

2. VBT (Pauling) Vacant hybrid orbitals on metal are generated from LCAO suited to each geometry ( tetrahedral sp 3 etc) Bonds are formed by e-pair donation from ligand orbital into vacant hybrid orbitals on metal. 3. CFT. Bethe 1930. The degeneracy of d orbitals is lifted by the electrostatic field of ligands (taken as point charges) in various

symmetries. In octahedral symmetry t 2g (-4Dq) and eg (+6Dq) separated by ∆o = 10 Dq. Bonding is presumed to be purely ionic.

4. MOT LCAO-MO’s are generated by combining symmetry adapted LGO’s with metal orbitals. Bonding and antibonding combos. Electrons then fill levels from bottom up according to Hund’s rules. Bond Theory Pauling 1930’s hs-555/ 639

AO’s lack the directional character necessary for bond formation in many geometries. Hybrid orbitals are linear combinations of AO’s with suitable directionality . - 10 4 linear sp Ag(CN) 2 d (sp) 0 BM 2 5 2 6 trigonal sp Fe(N(SiMe 3)2)3 d (sp ) 5.9 BM 3 2- 8 3 8 tetrahedral sp NiCl 4 d (sp ) 2.8 BM 2 2- 8 2 8 square planar dsp PtCl 4 d (dsp ) 0 BM 3 0 3 10 5 coordinate dsp PF 5 d (dsp ) 0 BM 2 3 3+ 6 12 octahedral d sp (inner) Co(NH 3)6 d (hyb) 0 BM 3 2 2+ 8 12 sp d (outer) Ni(NH 3)6 d (hyb) 2.84 BM 3 3 4- 2 14 7 coordinate d sp V(CN) 7 d (hyb) 2.84 BM 4 3 4- 2 16 8 coordinate d sp W(CN) 8 d (hyb) 0 BM 5 3 2- 0 18 9 coordinate d sp ReH 9 d (hyb) 0 BM In CFT the energies of the d orbitals are obtained from perturbation theory using the electrostatic potential V and the d orbital wavefunctions.

Edz2 = ∫ ψdz2 Voct ψdz2 = +6Dq and Exy = -4Dq etc. Dq = 1/6 ze 2 r4/a 5 where r 4 is the mean radius 4 of the electron and a is the metal ligand bond length.

CFSE ( in units of Dq) see p-563 of HS z2 x2-y2 xy xz yz One on z 5.14 -3.14 -3.14 0.57 0.57 Linear 10.28 -6.28 -6.28 1.14 1.14 trigonal -3.21 5.46 5.46 -3.85 -3.85 trigonal bipy 7.07 -0.82 -0.82 -2.71 -2.71 square -4.28 12.28 2.28 -5.14 -5.14 tetrahedral -2.67 -2.67 1.78 1.78 1.78 square pyram 0.86 9.14 -0.86 -4.57 -4.57 octahedral +6.00 +6.00 -4.00 -4.00 -4.00 Spectrochemical series - orders ligand in terms of increasing Dq

I- < F - < OH - < ox 2- ~ H2O < NH3 < en < bipy < phen < CN - < R - < CO Dq values in cm -1 and pairing energies (avg PE/electron) hs 559/642

+2 ions Ti V Cr Mn Fe Co Ni 6F - 730

6 H 2O 1240 1400 780 940 930 850

6 NH 3 1020 1080 3 en 910 1100 1150 PE 23,500 25,500 17,600 22,500

+3 ions Ti V Cr Mn Fe Co Rh 6F - 1700 1500 1400 1310 2264

6 H 2O 2030 1785 1740 2100 1370 1820 2720 3 ox 2- 1800 1700 1400 1800 2600

6 NH 3 2160 2290 3410 3 en 2190 2400 3460 6 CN - 2660 3500 3220 4490 PE 28000 30,000 21,000 ∆ Predicting Hi or Low spin from PE and 0. Pairing energy for Co +3 PE = 21,000 cm -1 per electron

3- ∆ CoF 6 0 = 10Dq = 13,000 ∆ cfse for low spin = 24Dq - 4 Dq = 20 Dq =26,000 cm -1 cost of pairing 2 electrons = 42,000 cm -1 cost exceeds benefit .

3+ ∆ -1 Co(NH 3)6 0 = 10Dq = 23,000 cm ∆cfse = 20Dq = 46,000 cm -1 benefit exceeds cost

*This calculation assumes that PE is independent of ligand. A better criterion is to compare ∆/B for the complex with the crossover point in the Tanabe Sugano diagram at ∆/B = 20 3- ∆ 5 CoF 6 0/B = 13,000/763 = 17 high spin T2g 3+ ∆ 1 Co(NH 3)6 0/B = 23,000/530 = 43 low spin A1g Examples of CFSE effects. ∆ 1. Stability constants, lattice energies, Hhyd, etc (Fig 20.26-28 of HS) show a big M or W pattern across the transition series. (V for strong field) weak field cfse = 4,8,12,6,0,4,8,12,6,0 Dq for d 1 thru d 10 max at d 3 and d 8 strong field cfse = 4,8,12,16,20,24,18,12,6,0 Dq for d 1 thru d 10 max at d 6

2. Structural preferences (using cfse’s in units of Dq) cfse d3 hs d6 ls d6 hs d7 ls d8 d9

OCT 12 4 24 8 12 6 TET 3.56 2.67 8.90 5.34 3.56 1.78 SQUARE 15.56 5.14 21.56 10.28 24.56 12.28 low spin d 6 and d 3 are always octahedral and the most inert of the first t-series Co 3+ and Cr 3+ . Pt 2+ d8 square Pt +4 d6 and oct. Dq is large for CN - but small for halides, oxygen donors. ∆tet = 4/9 ∆oct and 3 rd series> 2nd > 1 st All 3 rd and 2 nd are low spin All amine or O-donor complexes of 1 st t-series are hi-spin except Co(III)

CFT Energetics 1000 cm -1 = 1 kK = 11.96 kJ/mol

3+ -1 Ti(H 2O) 6 absorbs at 500 nm or 20,000 cm or 240 kJ/mol

10 Dq = 20,000 cm -1 or Dq = 2000 cm -1 cfse = 4 Dq = 8000 cm -1 or 96 kJ/mol

3+ 3+ Ti (gas) → Ti(H 2O) 6 (aq)

∆Hhyd = -5400 kJ/mol

∆Hhyd includes the 6 Ti -water bonds plus the solvation. cfse thus makes up only a small fraction of the total energy Jahn-Teller Theorem - A non-linear molecule in an orbitally degenerate ground state (T or E) will distort to remove the degeneracy.

Significant distortion in metal complexes are observed for E states. Axial elongation or compression is observed for octahedral cases.

4 2+ d Cr CrBr 2 (s) Cr-Br bond 4 at 2.54 2 at 3.00 Å

9 2+ 2 2 d Cu CuCl 2 4 at 2.3 , 2 at 2.95 Å ( x -y hole) 2 CuF 2 4 at 2.08 , 2 at 1.95 (z hole)

While T states are predicted to distort, the effect seems to be too small to detect. What MO Theory tells us about T-metal chemistry in O h 1. Bonding is delocalized over metal and all ligands.

A1g bonding MO is a 7 centre 2-electron bond

φbonding = c 1 φ4s + c 2{φa + φb + φc + φd + φe + φf } where φ4s is the metal 4s orbital and φa-f are the ligand donor orbitals. for pure covalent bond c 1 = c 2 = √2)/2 for more ionic bond c 1 << c 2 2. 10Dq depends on sigma and pi bonding

3. The 6 sigma bonds are A 1g , T 1u , and Eg * 4. The eg metal d orbitals in O h are antibonding eg

5. Pi bonding LGO’s are T 1g , T 1u , T 2g , T 2u

for pi acceptor ligands the metal t 2g from CFT is π * for pi donor ligands the metal t 2g from CFT is π

Note spectrochemcial series - - - 2- - - I < F < OH < ox ~ H 2O < NH 3 < en < bipy < phen < CN < R < CO π-donors weak σ strong σ / π-acceptors MO for Octahedral Complex Sigma only Pi Bonding a) pi donor I - b) pi acceptor CO

Diagram shows interaction along one axis, T 2g in O h

Bonding and Symmetry Basics

Geometry VB hybrids pt grp LGO’s sigma pi

Linear sp z D∞h σg σu πu πg

3 Tetrahedral sp Td a1 t2 t1 t2 e

2 Square Planar d x2-y2 sp xy D4h a1g b1g eu in eu b2g a2g

out eg b2u a2u 2 3 Octahedral d sp Oh a1g t1u eg t1u t2g t2u t1g

AO symmetry : O h s (a 1g ) p (t 1u ) d (t 2g eg)

D4h s (a 1g ) pxy (eu) pz (a 2u ) d z2 (a 1g ) dxz,yx (eg) dxy (b 2g ) d x2-y2 (b 1g )

Td s (a 1) p (t 2) d (e, t 2)

Note that the symmetry adapted sigma LGO’s belong to the same irred. reps as the AO’s used to form the hybrid orbital sets in VBT Examples of things that ain’t .

1. Co(NH 3)6Cl 3 is 6-coordinate not 9

2. BeCl 2 is tetrahedral in the solid state 3. NaCl is 6-coordinate in the solid state.

2+ 4. Ca (aq) is 6 coordinate in water

5. Mn(CO) 5 is a dimer octahedral Mn 2(CO) 10

6. P 2O5 is actually P 4O10 + - 7. PCl 5 (s) is actually [PCl 4] [PCl 6] + - 8. N 2O5 (s) is [NO 2] [NO 3] nitronium nitrate 5+ 2- 2- 9. “BiCl” contains 2 Bi 9 4 BiCl 5 and 1 Bi 2Cl 8 more precisely Bi 24 Cl 28 with a tricapped trigonal prism and sq pyramids “You can’t readily assume the coordination number from the formula” Factors influencing Coordination and Geometry

1.Maximize number of bonds and Bond energy 2. Minimize L-L repulsion 3. LFSE if not d 0 or d 10 4. chelate ring strain, conditions, other factors

Example: 6 vs. 4 coordinate 1 vs. 2 & 3 Tetrahedral vs. Square 2 vs. 3 Do not delude yourself into thinking there are simple principles which allow you to decide matters where 1 kcal/mol can tip the balance either way . Pentacoordination Four Coordinate 1. VSEPR or ligand repulsion dominant

Trigonal Bipyramid D3h Tetrahedral 3- PF 5 Ni(CN) 5 , Fe(CO) 5 CH 4 Ni(CO) 4 Pt(PPh 3)4

2. LFSE dominant

Square Pyramid C 4v Square Planar 8 3- 8 9 2+ 2+ 3+ d Ni(CN) 5 d , d Cu Pt Au pentacoordination was once thought rare but is now rather common. It is not usually possible to predict which 5 or 7 coordinate geometry might be preferred. Neither is stereochemically rigid enough to study isomers.

Thus studies are limited mostly to sq Pt(II) and Oct d 3 and d 6 [Cr(en) 3][Ni(CN) 5] Raymond, Inorg Chem 7, 1362 (1968)

Ax 2.17 Å Eq 1.87 Å Ax 1.83 Å Eq 1.91 Å cfse= 2(9.04 Dq) cfse = 2(7.07 Dq) 19 F NMR of PF 3Cl 2 Holmes, Inorg Chem 3, 1748 (1964)

JPF =1051 Hz

JFF = 124 Hz Some Novel Nickel (II) Complexes 2+ Ex 1 Lifschitz Salts [Ni(stien) 2] square  Ni(stien) 2X2 oct

Ni(stien) 2X2 is sometimes yellow and diamagnetic and sometimes blue and paramagnetic. - - blue : X = Cl or CH 3COO or in humid air or in dmso soln. - - yellow X = CF 3COO or ClO 4 or dry or in CCl 4 soln stien = meso stilbenediamine NH 2CH(Ph)CH(Ph)NH 2

Nyburg and Wood, Inorg Chem 3, 468 (1964)

Ni(stien) 2(CH 3COO) 2 µ = 3.13 blue pwd

Ni(stien) 2(CCl 3COO) 2 diamagnetic yellow pwd

Ni(stien) 2(ClCH 2COO) 2 - 2/3 EtOH 4/3 H 2O µ = 1.76 yellow-grn xtal 2+ triclinic P-1 Z=3 [Ni(stien) 2(X) 2] at (0,0,0)* (1/2,0,1/2)* and [Ni(stien) 2] at (0,0,1/2)

Ni(stien) 2(ClCH 2COO) 2 -4H 2O blue xtal 2+ P2 1/c Z =2 [Ni(stien) 2(H 2O) 2] at ( 0,0,0) and (0,1/2,1/2) Ex 2 Kilborne and Powell, J. Chem Soc A 1168 (1970)

Ni(PPh 3)2Br 2 tetrahedral µ = 3.2 BM Ph = phenyl

Ni(PR 3)2Br 2 square µ = 0 BM R = alkyl

Xtals of Ni(PPh 2(CH 2Ph)) 2Br 2 have 3 Ni per unit cell P-1. Which Ni complex lies on a special position in P-1?

2 are tetrahedral and 1 square (trans) µ = 2.7 BM. Tet Ni-P 2.31 Å Square Ni-P 2.26 Å Ni-Br 2.35 Å Ni-Br 2.30 Å

2 2 In tet form the antibonding t2* is occupied. In square b 1g * (x -y ) vacant. Thus the bonds are stronger in the d 8 square form as long as steric repulsion is not a factor. By changes in the nature of the phosphine and halide ligands it is possible to tip the balance either way. Ex 3 Holm, JACS 92, 1855 (1970) Ni 2+ complexes of mixed arylalkylphosphines show an equilibrium between square and tetrahedral forms. Holm

has studied some 30 of the general form NiP 2X2 where P = P(CH 2Ph)(Ph-Z) 2 (Z para substituent) X = halide

NiP 2X2 (SQUARE)  NiP 2X2 (TET) 1 3 µ = 0 BM B1g µ = 3.2 BM T1

Z= CF 3, X = Cl 11% tet µ = 1.1 BM Z= NMe 2, X = I 92% tet µ = 3.02 Vis (xy  x2-y2) 20kK vis 3 d-d (5.5, 11.8, 20 kK) ∆Go = +1.1 kcal/mol ∆Go = -1.3 kcal/mol for various subst. ∆Ho = 1 to 2 kcal/mol ∆So = 2 to 4 eu cfse, bonds favor square entropy favors tet {S = Rln(9) } π π Z= CF 3 favors -acceptor higher cfse Z= NMe 2 -donor, lower cfse X = Cl small steric effect X = I large steric effect favors tet +,o,- Ex 4 Byrne, JACS 109 1282 (1987) Co(NOR) 4 Chem Comm 1491 (1986) Alkyl complexes of the norbornyl ligand

(C 7H11 ) provide a rare example of a low spin tetrahedral complex of Co (III)

+ cation Co(V) µ = 0 o neutral Co(IV) µ = 1.73 BM - anion Co(III) µ = 3.18 BM alkyl ligands lie high in the spectrochemical series ( strong σ-donors ) Norbornyl is a bulky ligand thus favoring tetrahedral structure.

Co(IV) structure is Pmn2 1 (#31) z =4 with Co on a mirror. EX 5. Tetrahedral Pd(II) HS 671 Yeo Chem Comm 1477 (1999)

µ = 2.5 BM paramagnetic bite angle O-Pd-O = 105 0 -how does this change d orb energies? chelate is diphenylphospinoxido ferrocene Lower coordination numbers 2 , 3 + - + Ag(NH 3)2 , HgI 3 , Cu(tu) 2Cl , MnL*2 , AuL 3 FeL*3 - • bulky ligands L= PPh 3 , L* = N(Si(Me) 3)2 • filled d shells Ag + Cu + Au + Higher coordination numbers 3- 4- 7 pentagonal bipy ZrF 7 Zr(acac) 3Cl V(CN) 7 2- + 2- capped trig prism TaF 7 Mo(CNR) 6I Mn(EDTA)(H 2O) 3- capped oct NbOF 6 -3 4- 8 sq anti prism Ti(NO 3)4 , TaF 8 dodecahedron Mo(CN) 8 cube CsCl 2- 9 tricapped trig prism ReH 9 , • vacant d shells • small ligands - -1 -2 -1 • chelates with tight bites NO 3 , acac , ox , trop T-shaped 3 Coordinate [RhP 3]ClO 4 Reed, JACS 99, 7076 (1977)

RhP 3Cl + TlClO 4 → [RhP 3]ClO 4 + TlCl in dry CH 2Cl 2 + in even weakly coordinating solvents one obtains square [RhP 3S]

diamagnetic d 8 cannot

be trigonal planar D 3h (why).

Angles 98, 103, 160 o Rh-P 2.21, 2.24 Å

triclinic space group P -1 ; a = 11.925 , b = 14.466 c = 15.553 8; α = 92.637 , β = 88.931 , γ = 112.744 (15)'; Linear bis- (trimethylsilylamido)iron and Trigonal tris Lappert, Inorg Chem. 27, 1782 (1988),

Bis Fe(II) in gas phase by electron diffraction Fe-N 1.84 Å N-Bridged Fe(II) Dimer in solid state µ = 3.52 BM. Tris Fe(III) as solid by X-ray Fe-N 1.91Å (high spin) SOFTWARE FOR STRUCTURE VIEWING Free Mercury graphic package from Cambridge http://www.ccdc.cam.ac.uk/products/mercury/

Data input as CIF or SHELX RES file unit cell dimensions, space group SYMM codes, and atom locations in fractional coordinates. view asymmetric unit in variety of modes, stick, ball, spacefill etc. computes and displays X-ray powder pattern packing display hkl planes measure distances, angles, mean planes identify H-bond and close contacts

WINGX free software for solving structures http://crick.chem.gla.ac.uk/~louis/wingx/download.html requires X-ray diffraction data in an HKL file - see Expt 5. Some recent structures viewed via the MERCURY software.

[Et 4N] 2 [Fe(CN) 2(CO) 3 P2 1/n z=4 IC 42,5046 (2003) [Cu(CH 3CN) 2 ] [BPh 4] C2/c Z = 4 Cu on -1 site III [Yb (PY) 5I2] I (0.5 PY) Z =8 in P2 1/c 7 coord Yb 2 independent Yb’s per unit cell 2+ Au 2(dpim)] [ClO 4]2 in P2 1/n Z =4 IC 42, 8430 (2003) Au(dpim)Cl in P2 12121 Z=4

chiral C1 sq

SQ

III Tb (W) 4(µ-SQ) 2] [NO 3] W P2 1/c Z =4 square antiprism

SQ is the enolate anion of 3-methyl-4-hydroxycyclobut-3-ene-1,2-dione , W= H 2O III La (SQ) 3W6 -W in P2 1/c Z =4 9 coordinate La

B. Alleyne et al, Inorg Chem., 40, 1045 (2001)