An Introduction to Transition Metal Chemistry

Properties Common to the Transition Elements 1. The free elements conform to the metallic bonding model. Lattices: close-packed or body-centered cubic High Thermal & Electrical Conductivities ; Malleable & Ductile. 2. Multi-Oxidation States (Exception: zinc & cadmium) 3. Paramagnetic: Gr St: Unpaired Electrons NMR, ESR (electron spin resonance), and magnetic susceptibility 4. Electronic Transitions(of free elements and their compounds & complexes)

Infrared, Visible or ultraviolet region for the visible cases, of course, colored species are the result. 5. Strong Tendency to Form Complexes: Cations & neutral atoms: Lewis acids Complexes (common): 2-6 bases (Ligands)

Oxidation State Tendencies and Their Causes Oxidation State: Preferences on ionization energy & bond energy Two perspectives: 1. the maximum (the most oxidized) state & 2. the preferred (most stable) state The first half of the 3d series (Sc thru Mn): Maximum state loss (or, more accurately, the participation in chemical bonding) of all valence electrons. Eg: Sc2O3, TiO2, VO2+, CrO3 and MnO4 Beyond Mn: Highest state decreased Eg: FeO42, Co2O3nH2O, NiO2, CuO+ and ZnO. Decrease from FeVI to ZnII correlates with the number of vacancies than with the number of occupancies in the valence orbitals of these elements.

Maximum Oxidation States: Binary oxides and fluorides (Table 1) Cr & Mn:Oxides higher state than fluorides. Coordination effect: MnVII requires 7 F & only 4 O

Most common state is MII (difluorides) of all elements except Sc & Ti Electron Configurations of the M2+ cations: [Ar]4s03dx A variation of the inert pair effect: s orbital is empty / s orbital w.

nonbonding electron pair

The preferred state: The most stable (toward disproportionation or thermal dissociation) and/or the least reactive Highly dependent on ligands Metals in high OxSt are hard stabilized by hard bases (e.g. F and O donors) Lower OxSt stabilized by soft bases (S2 and I)

Table 1: When bonded only to O & F higher states (MIII and/or MIV) / preferred by early elements; MII is favored later in the series Correlates with ionization energies Energy to remove 2 & 3 electrons from the 3d elements Fig 1 Steady increase going across the period in both cases Removal of electrons: Sc/44.10 eV ; Mn / 56.75 eV (Mn the first element for which the MII state is preferred). Hence, Mn3+ is destabilized relative to its neutral atom by 12.65 eV (about 1220 kJ/mol) more than is Sc3+ versus Sc The relative destabilization is even greater for Co, Ni, Cu, Zn; (Fe exception due to pairing energy effect Volt-equivalents oxidation state preferences in aqueous solution Table 2: Data for the 3d metals Overall trends are similar to those observed for binary fluorides and oxides: MIII is favored by the early elements, while MII dominates later in the series

Comparisons of the 3d, 4d, and 5d Elements Ionization energies 4d elements < 3d congeners Eg: (IE1:IE3) Tc = 52.08 eV, ~8% < Mn. Lower IEs and reduced steric interactions (longer bonds): Favor high OxSts. Most stable fluorides of Nb, Mo, Tc, Ru: NbF5, MoF6, TcF6, RuF5 (Compare these to the corresponding 3d elements in Table 1) Ag is exception: AgI state 5d elements similar to their 4d congeners (TaF5, WF5, etc.) Some cases: still higher OxSts stabilized (ReF7, OsF6, lrF4, and PtF4).o Aqueous Eo known with less certainty for the 4d and 5d elements o Expected trends hold o Tendency: decrease OxSts across each period & increase OxSts down each

group.

Occurrences, Isolation, and Uses of the Free Elements 1. Abundances tend to decrease going down each family 2. Lighter elements normally bonded to O in their natural deposits The heavier, softer metals generally prefer sulfur 3. OxSt preferences reflected in ores and minerals of these metals Going across the first transition series, the relative natural abundances follow the order Sc > V $ Cr < Mn > Co < Ni > Cu $ Zn Free elements obtained by chemical reduction Eg: Titanium (the KroIl Method) TiO2 + 2 Cl2 + 2 C TiCl4 + 2 CO Ti + 2MgCl2 (1) (2)

TiCl4 + 2 Mg (1000o C)

TiC14 (volatile, bp = 136C) easily separated from impurities such as FeC13 by distillation. High-temperature chemical reduction then yields the free metal.

Iron (the Bessemer Blast Furnace): CaCO3 ( CaO + CO2 CaSiO3 2 CO 3 Fe(l) + 4 CO2 (3) (4) (5) (6)

SiO2 + CaO ( 2 C + O2 ( Fe3O4 + 4 CO (

Nickel (the Mond Method): 2 NiS + 3 O2 ( NiO + H2 ( Ni + 4 CO 50oC Ni(CO)4(g) > 200oC 2 NiO + 2 SO2 Ni + H2O Ni(C0)4(g) Ni + 4 CO (7) (8) (9) (10)

Coordination Compounds and Complex Ions Chemistry of the transition elements: Metal (Lewis acidbase interactions) with 2 or more donor ligands coordination compounds or complexes; if charged, they are complex ions Organometallic Chemistry: metal--metal or metalnonmetal bonds (not dative) yet are referred to as coordination compounds or complex ions General equation complex formation is Mm+ + n Lx [ML]mnx (11)

Enormous number of possibilities: Variable metal oxidation states Great variety of ligands available Different possible coordination numbers Metal Comlexes: Number of Ligands & Geometric Structure

Low-Coordinate Geometries Two ligands / central metal: linear or an angular manner ; linear more common Idealized symmetry (ie, the point group of highest possible order): Dgh or Cgv The monovalent cations of Grp 11 form numerous dicoordinate complexes: CuCl2, [Ag(NH3)2]+, and [Au(CN)2] linear Complexes of HgII (which is isoelectronic with AuI), such Hg(CH3)2 and Hg(Br)SCN, are known as well Metal has 10 valence eConfigurations . . . ns0(n l)d10 & . . . ns1(n l)d9 energetically close Two plausible hybridization schemes: For s0d10 : electron pairs donated by ligands reside in linear combination of the ns and npz orbitals sp hybridization (linear geometry) sdz2 hybridization also feasible: such hybridization is actually favored (it removes charge from the region between the metal and ligands (Fig 2)

Stoichiometry is not predictive of coordination (especially in the solid state)

Eg: FeF2 CN = 2!? In crystalline FeF2 & FeF3: CN 6-coordinate octahedral interstices of fluoride ion sub-lattices Complexes metal CN = 3 uncommon Cu+ as the metal ion K+[Cu(CN)2] 1:2 metal:ligand stoichiometry there are two bridging and one terminal cyanide ligands, with the CuCN framework creating a helical chain Fig 3 Sterically large ligands trigonal coordination Eg: (Me3Si)2N & (Me3Si)3C form trigonal complexes with 3d elements Four bulky groups around metal create steric repulsions sufficient destabilize the system M in the plane of the three donor atoms, but in Sc[N(SiMe3)2]3 the scandium is above the plane (C3v local symmetry.

Tetrahedral. Square Planar, and Intermediate Geometries8 CN = 4 : Tetrahedral (for four equivalent ligands, giving local Td symmetry and LML bond angles of 109.5) Square planar (ideal symmetry D4h, 90o angles) Infinite number of possible geometries between these limiting cases; all have idealized D2d symmetry and bond angles between 90 and 109.5 (Fig 4) Intermediate geometries (flattened, squashed, or distorted tetrahedra) less

common than either the tetrahedral or square planar structures Eg: Complex anions in Cs2[CuX4] (X = Cl and Br), which have XCuX bond angles between 100 and 103.

The geometry (tetracoordinate complex): Steric and Electronic Factors Tetrahedron superior / steric d electrons determines electronic factors

Square planar geometry favored by d8 metal ions electrons Eg: PdII, PtII, and AuIII (Both geometries are common for NiII) Steric effect more important for nickel than for its 4d and 5d

congeners because of its smaller size Ni complexes:difference in stability between the tetrahedral & square planar geometries is very small Eg: NiX2(PJR)2 (X = Cl , Br, and I; R = alkyl) Gradual transition from square planar to tetrahedral geometry

increasing steric bulk (Table 4)o Several of the intermediate species such as NiBr2(PJ2Et)2, can be

produced in either of the two forms.

Certain ligands dictate (or at least strongly favor) square planar

geometry

One such donor is terpyridine, a tridentate ligand containing three aromatic rings (Fig 5a)

Better examples: tetradentate porphine molecule and its derivatives,

called porphyrins. The conjugate T system must remain planar to maintain effective pT overlap The result is a natural hole for the metal ion, created by the planar

arrangement of donor nitrogens The metal may be either in the ligand plane (D4h site or out of the

plane (C4v). Porphyrins are found in many biologically important compounds (Bioinorganic Chemistry).

Geometric and Optical Isomerism

Stereoisomers spatial arrangement of atoms: geometric and optical isomerism CN=4 square planar complexes: (cistrans) isomers Ma2b2 No isomerism is possible for tetrahedral Ma2b2 Historical Importance: existence of two different Pt(NH3)2C12 used by Alfred Werner (father of coordination chemistry) to prove that (unlike carbon) transition metals often have square planar geometries The cis isomer of Pt(NH3)2Cl2 (cis-platin) is an antitumor agent, while the trans isomer is inactive.I

somerism affected by chelating ligands: bidentate ethylenediamine (en, H2NCH2CH2 NH2) substitute for two NH3 ligands to give Pt(en)Cl2

Only cis isomer isolated because en is not long enough for its nitrogens to occupy trans

positions Regardless of geometry complexes: bidentate ligands spanning non-cis coordination sites rare Tetrahedral geometry: four different ligands creates an asymmetric metal optical isomerism Few tetracoordinate metal complexes optically pure enantiomers: rapid racemization at or

below room temperature.o Stereoisomerism thru ligands: via trivalent donor atoms have three different substituents o Become asymmetric upon coordination to a metal o Ligands: PH(Me)J & N-monomethyl derivative of ethylenediamine, MeNHCH2CH2NH2.

Trigonal Bipyramidal, Square Pyramidal, and Intermediate Geometries CN =5 Uncommon. Produce less stable complexes CN = four or six Decomposition: ML5 ML4 + L (12)

Complexes truly 5-coordinate: 2 limiting structures plus infinite number of intermediate cases Trigonal bipyramid (TBP): D3h Square pyramid (SPY): C4v Main group atoms 10 valence electrons : TBP electron geometries Less clear-cut for metal complexes: energy difference between the TBP and SPY structures is often small Eg: 2[Cr(en)3][Ni(CN)5]3H,O slightly distorted TBP geometry, and the other, SPY geometry.

o Solid-state structure contains two different [Ni(CN)5]3 anions, one having a

The ideal SPY geometry is essentially never encountered. In square pyramidal complexes, the metal ion is normally raised out of the basal plane (typically, by 3050 pm) toward the axial group:L

M L L L L

This changes the SPY bond angles to (calculated) optimum values of about

104o (axialbasal) and 87o (basalbasal) from the 90o ideal. Bond angles for the ideal TBP and SPY and for the optimized SPY geometries are summarized in Fig 7

TBP and SPY geometries become indistinct thru their interconversion Interconversion (facile) fluxional behavior in solution Eg: Fe(CO)x(PF3)5x prepared and exhibit TBP geometry Isomerism is possible when x = 1- 4, since there are two non-

equivalent ligand positions (axial and equatorial) In solution: all three geometric isomers of Fe(CO)3(PF3)2PF 3 CO OC Fe CO PF 3 PF 3 OC Fe CO CO CO PF 3 OC Fe PF 3 CO PF 3

Octahedral and Distorted Octahedral Complexes CN 6: Most common Complimentary Geometry: trigonal prism Stereoisomerism Octahedrals

Ma4b2 isomers cis and trans Ma3b3 isomers fac (for facial) and mer (for meridional) Fig 9 , Fig 10 & Table 6 Distortions:

rigonal Tetragonal (or Jahn-Teller) Fig 11

Optical isomerism: Common for octahedral complexes having chelating ligands Eg: hexadentate EDTA ligand and tris(ethylenediamine) complexes. [Co(en)3]3+ belongs to the D3 point group Has neither a mirror plane nor an inversion center: optically active. (Actually, optical activity results from the lack of an improper axis of rotation. Since S1 = W and S2 = i, the minor planeinversion center (guide identifies nearly all such cases, and so is often used in place of the more rigorous test.) The two enantiomers are shown in Fig 10; the Greek letters and for right- and left-handed spirals looking down the threefold axis, respectively.

High-Coordinate Geometries CN 7 less common General rule structural diversity increases with increasing coordination Three well-established geometries for CN = 7:

Pentagonal bipyramid (idealized D5h symmetry) Capped octahedron (C3v) Capped trigonal prism (C2v) Fig 12

o

Examples for polydentate ligands stabilizing CN 7 Eg: [Co(crypt)]2+ [Co(SCN)4]2

CN = 8 cube, square antiprism, and the dodecahedron (Fig 15) (like the trigonal prism and antiprism, common in extended lattices but rare in discrete complexes) Square antiprism Relate to cube like octahedron to trigonal prismtwisting one face (for the 8-coordinate case, by 45o) generates the alternative structure. Energy differences among these geometries are small, interconversions are usually

facile, fluxionality is common, and the separation of isomers is often impossibleo Most of the known 8-coordinate complexes have large metal ions (4d, 5d,

lanthanide, or actinide elements) and/or either small or macrocyclic ligands. Oxyanions such as NO2, NO3, IO3, and SO42 are often found in

octacoordinate complexesO X O M

Structural Isomerism Coordination Isomerism [Co(NH3)6]3+[Cr(CN)6]3 and [Cr(NH3)6]3+[Co(CN)6] 3 both have the molecular formula CoCrC6H18N12 Other salts having the same formula: [Co(NH3)5(CN)]2+[Cr(NH3)(CN)5] 2 In two different oxidation states:

Pt2H12C16N4 are the salts [Pt(NH3)4]2+ [PtCI6] 2 [Pt(NH3)4Cl2]2+[PtCl4] 2

Ionization and Hydrate Isomerism Ionization isomerism: Pt(NH3)3(Br)(NO2) [(H3N)3PtBr]+NO2 and [(H3N)3PtNO2]+Br Two PtII salts: [Pt(NH3)4C12]Br2 and [Pt(NH3)4Br2]Cl2.

Linkage Isomerism Ambidentate can act as Lewis base thru two or more kinds of donor atoms [(H3N)5CoNO2]2+ and [(H3N)5CoONO]2+ nitro & nitrito Ambidentate: SCN- & SeCN2Se N d Se N Se N N Se d N Se N Se