Ch. 23 Transition Metals and Coordination Chemistry

The Transition Metals

*d-block elements

*Debate about inclusion

*Lanthanides La (lanthanum)-Lu (lutetium) stacked here; debate about inclusion

Minerals

*Most metals, including transition metals, are found in solid inorganic compounds known as minerals. Usually, these are oxides, hydroxides, sulfides, or carbonates

*Minerals are named by common, not chemical, names

Transition Metal Properties

*Partially-filled d orbitals --> multiple oxidation states, -3 to +8

*S-orbitals are filled first (except Cr, Cu) but emptied first for cation: (Review BLBMWS p. 306) Fe^0, [Ar]4s^23d^6, Fe^2+, [Ar]3d^6

*Magnetic Properties

-paramagnetic --> unpaired electrons

*Colors from electron transitions between d-orbitals

*Coordination compounds with 4-9 ligands

Colors and Magnetism General Properties

*Most transition metal ions are colorful and paramagnetic (except for most with d^0 and d^10 configuration)

Common Oxidation States

*Transition metals often have multiple oxidation states

-Negative values are found in some organometallics

*M(II) oxidation state is common (loss of two s electrons)

Atomic Radii

*As one goes from left to right across a row, we see a decrease, then an increase in the radius of transition metals

*Two factors: Zeff & metallic bonding strength

*While Zeff increases across a row, so does number of nonbonding electrons

Horizontal Trends in Properties of Period 4 Elements

*Atomic radius decreases from left to right

*Electronegativity increases from left to right

*First ionization energy increases from left to right

Why there trends in the periodic table?

*Atomic size remains fairly constant as atomic number increases, because electrons are added to inner shells, shielding outer s-eectrons from nuclear charge efficiently

*With electronegativity and ionization energy, as a positive charge is added to the nucleus, an inner electron is added that shields the outer s-electrons, so the ease to add or remove an electron remains fairly constant as we fill the d-shells

*Lanthanide contraction: Filling 4f orbitals causes increase in Zeff resulting in contraction in size

Transition-Metal Complexes

*Transition metals can have molecules or ions bound to them

*These give rise to complex ions or coordination compounds

Components of Coordination Compounds

*Central transition metal or transition metal ion

*Ligands

-Molecules (neutral) and/or anions (charged) that bind to the central metal

Have one or more donor atoms that each donate a lone pair of electrons to the metal atom/ion

*Coordinate covalent bond

-Special case of covalent bonding (not ionic)

*Complexes

-Do not dissociate into ions in solution, rather they act as a single molecule or ion

Metal-Ligand Bond

*Transition metals or their ions are Lewis acids

*Ligands are Lewis bases (e.g. NH3, Cl)

*Coordination complexes are formed by Lewis acid-base interactions resulting in a coordinate covalent bond

Complex ion charge

Sum of the charges on the metal and ligands

Coordination number

*Number of ligand donor atoms bound to metal

*A given metal in a particular oxidation state has specific coordination numbers governed by:

-Metal size

-Ligand size

-Ligand type

-Ligand charge

*Example: Iron(III) can bind to 6 fluorides, but it can only accommodate 4 of the larger chlorides

Geometry

(Shape) of a complex depends on coordination number and metal ion

Common Ligands

*Ligand types:

-Monodentate: coordinate to one site on the metal

-Bidentate: one ligand coordinates to two sites on the metal at once

-Polydentate: one ligand coordinates to more than two sites on the metal at once

Common Monodentate Ligands

*Monodentate ligands coordinate to one site on the metal donating one electron pair at a time

-Notice how nitrite and thiocyanate ions can donate electrons on 2 different atoms BUT only 1 pair of electrons is involved at a time

For the exam, know the ligands in table 23.4 and 23.5

Bi- and Polydentate Ligands

*Bidentate example:

-h2N-CH2-CH2-NH2

-2 atoms (nitrogens) donate e- pairs to the metal ion at the same time

*Polydentate example:

-ethyleneediaminetetraacetate (EDTA^4-) supplies up to 6 donors

For the exam, know the ligands in table 23.4 and 23.5

Chelates

*Bidentate and polydentate ligands are also called chelating agents. Metal ion is held tightly by ligand

- chele (Greek) means a pincer-like claw

*Here, EDTA uses its two N and four O donor atoms to wrap around a metal ion and form a chelate complex

*Chelates are particularly stable and (generally) have larger formation constants, Kf, than complexes with mono dentate ligands

Chelate effect

*Increased stability and rate of formation for a metal complex because of bi- or polydentate ligands

Chelates in Biological Systems: Hemoglobin

*The iron in hemoglobin carries O2 through the blood

*Carbon monoxide and cyanide are poisonous because they will bind more tightly to the iron than will oxygen

Nomenclature Rules

1) Cation is named before the anion in salts

-NaCl sodium chloride

-K3[Fe(CN)6] potassium hexacyanoferrate (III)

2) In a complex, ligands are named first, metals last

-[PtCl6]^2- hexachloroplatinate (IV)

3) Ligands are named in alphabetical order by ligand, not by number prefix

-tetraamminedichloro

-tribromotrinitro

4)Lighted names: anions end in -o; most neutral ligands are named as the molecule

-pyridine, ethylenediamine, cyano, iodo

-Exceptions: water (aquo, aqua), carbon monoxide (carbonyl), ammonia (ammine)

5) Number of each ligand is indicated by Greek prefixes

-di, tri, tetra, penta, hexa, hepta, octa, nona, deca, undeca, dodeca

6) Metal oxidation state is indicated by Roman numerals

-K4[Fe(CN)6] potassium hexacyanoferrate (II)

-[Co(en)3]Br3 tris(ethylenediamine)cobalt(III) bromide

7) Anionic metal complex names end in -ate:

-Na3[Co(NO2)6] sodium hexanitrocobaltate

What is the name of [Co(NH3)5Cl]Cl2?

A. chloropentaaminecobalt(III) chloride

B. pentaamminechlorocobalt(III) chloride

C. chloropentaamminecobalt(II) chloride

D. pentaamminechlorocobalt(II) chloride

B. pentaamminechlorocobalt(III) chloride

Isomerism

*It is possible to assemble a coordination complex in several different ways

*Isomers

Isomers

*Compounds with the same formulas but different arrangements of atoms

*Two types

-Structural isomers (coordination-sphere isomers, linkage isomers)

-Steroisomers (geometric isomers, optical isomers)

Coordination-sphere Isomers

*Structural isomer

*Different ligands occupy the coordination sphere but overall formula is the same

*Example: the complex with empirical formula CrCl3O6H12 exists as three different complexes:

-[Cr(H2O)6]Cl3

-[Cr(H2O)5Cl]Cl2 x H2O

-[Cr(H2O)4Cl2]Cl x 2H2O

*Cl^- is a ligand or counter ion

*H2O is a ligand vs. hydration shell

Linkage Isomers

*Structural isomer

*The ligand is bound to the metal by a different donor atom

*The formula for the ligand is written starting with the donor atom

*Example: the ligand NO2^- vs. ONO^-

Geometric Isomers

*Stereoisomer

*Square planar and octahedral complexes can have cis or trans isomers

Optical Isomers

*Stereoisomers

*Optical stereoisomers = enantiomers

*Molecules or ions that are not superimposable on their mirror image are termed chiral

*Why are they "optical" isomers?

-Optical isomers rotate polarized light to either the right or left

- d - dextrorotatory (right handed)

- l - levorotatory (left handed)

-Racemic mixture: amount of d = amount of l

Determining Optical Isomers

*Structures which are mirror images, but can rotate on top of one another, are the same compound just drawn different ways

*Only if you CANNOT get them to line up by drawing from another angle are they optical isomers

*Practice drawing 3D molecules in 2D. (Just because you can draw them differently doesn't mean they are isomers!!)

*For example, this cis platinum complex can be drawn in 2D eleven other ways, all of which are the same molecule

*Convince yourself that this complex is NOT chiral

[Ag(SCN)2]^-1 and [Ag(NCS)2]^-1 are ____.

A. geometric isomers

B. coordination-sphere isomers

C. linkage isomers

D. optical isomers

E. identical ions

C. linkage isomers

Color and Magnetism

*Color of a complex depends on:

-Identity of the metal ion

-Oxidation state of the metal ion

-Identity of the ligands bound to the metal

Absorption Spectroscopy

*Sample is irradiated by light of various wavelengths (i.e. various energies)

Color

*The absorption of light by solutions of transition metal compounds results in color because of the remaining, unabsorbed (transmitted) light

*If all yellow light is absorbed, the solution is purple (complementary color)

*Similarly, if all green is absorbed, the solution will appear red

Which peak in this spectrum corresponds to the lowest-energy transition by an electron in a chlorophyll molecule?

A. 420 nm

B. 500 nm

C. 610 nm

D. 650 nm

D. 650 nm

the higher the wavelength, the lower the energy

_______ is a property of substances in which all of the electrons are paired.

A. Diamagnetism

B. Paramagnetism

C. Ferromagnetism

D. Antiferromagnetism

A. Diamagnetism

Magnetism in Transition Metals

*Magnetism can be used to analyze d-electron populations

*Three major types of magnetic behavior

-Diamagnetic

-Paramagnetic

-Ferromagnetic

Diamagnetic

No atoms or ions with magnetic moments

Paramagnetic

Magnetic moments unaligned without a magnetic field. When a magnetic field is applied, spins align and it is attracted to a magnet

Ferromagnetic

Coupled magnetic centers aligned in common direction without a magnetic field (permanent magnet)

Magnetism of Coordination Compounds

*Many coordination compounds are paramagnetic

*Example:

-What is the electron configuration of Co^3+? [Ar]3d^6

-How many unpaired electrons in Co^3+?

4s shell is empty, 3d shell has one paired electron pair, and 4 unpaired electrons

-How do we explain that [Co(CN)6]^3- is diamagnetic and [CoF6]^3- has 4 unpaired electrons

Crystal-Field Theory

*A model for bonding in transition-metal complexes that helps explain experimental observations of color and magnetism

*Involves the interaction of ligands with the d-orbitals of the metal

Five d-Orbitals

*The five d-orbitals in an uncompelled metal have the same energy

*d-orbitals in a "ligand field" experience repulsion by the ligand electrons differently, based on their respective symmetries relative to the ligand axes

Crystal Fields and d-Orbitals

*Strong orbital overlap (yellow)

*Weaker overlap (red)

*e- - e- repulsion between ligands and d-orbitals is energetically unfavorable

Ligands and d-Orbitals:

Octahedral Crystal Field

*dx^2y^2 and dx^2 are anti bonding with higher energy levels and low spin

*dx, dxz, dye are bonding with lower energy levels and high spin

-e --> doubly-degenerate

-t --> triply-degenerate

*The notation comes from "Group Theory"

Splitting of d-Orbital Energies by Octahedral Ligand Field

*Free metal ion --> metal ion plus ligands (negative point charges with spherical symmetry) --> In octahedral crystal field, delta is the splitting energy often called delta(o)

*The energy gap between the two sets of orbitals depends on the metal and, to a larger extent, on the ligand

Spectrochemical Series

*Ligands can be arranged in an order based on ligand field strength. This series was empirically derived using visible light absorption spectroscopy

*I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO

*weaker field --> stronger field

*less splitting --> more splitting

Light Absorption

*The energy gap between d orbitals often corresponds to the energy in a photon of visible light absorbed by the lower energy d electron (a d-d transition)

*How do I calculate the corresponding splitting energy, delta, in J?

-Use E = hc/lambda

Ligand Effect on Electron Configuration of Octahedral Complexes

*When energy gap is small, it costs less E to promote electrons into the higher-energy orbital than it does to pair them up: high-spin

*When energy gap is large, the E cost is lowest when electrons are paired up in lower-energy: low spin

Crystal Field Splitting and Magnetism

*Number of unpaired electrons determines whether the complex is paramagnetic (some unpaired electrons) or diamagnetic (all paired electrons)

*Since ligands affect whether a complex is high or low spin, ligands affect # of unpaired electrons in a complex

*More unpaired electrons = a greater magnetic moment. Ligands have a dramatic effect on magnetic properties of coordination compounds

*This effect can be used in reverse!

-If one measures the magnetic properties of a complex with a known metal (e.g. Co) and unknown ligands, the results would allow qualitative predictions on the d-orbital energy splitting by the ligands (i.e. are they strong or weak field ligands)

Ligand Field Stabilization

*Ligand field stabilization is the energy obtained by putting electrons in lower energy orbitals

*Thus Cr^3+ (d30, Co^3+ (d6, low spin), and Fe^2+ (d6, low spin) are especially inert as octahedral complexes

*Complexes which undergo rapid ligand exchange are labile. Those which undergo slow ligand exchange are inert

d-orbitals in Tetrahedral Complexes

*Ligands affect dxy, dxz, dyz orbitals on diagonals; dx2y2 and dz2 are along the x, y, and z axes, so they are less perturbed

*delta(t) < delta(o)

*Less E to promote electrons than Epairing; therefore all tetrahedral complexes are high spin

d-d Electronic Transitions Give Rise to Color

*Sapphires and rubies are both forms of Al2O3, corundum

*Their different colors come from impurities of Fe and Ti (sapphire) or Cr (ruby)

d-d Electronic Transitions

*Metal complex can absorb a particular wavelength of light (energy), leading to an "electronic transition", an excitation of an electron from one orbital to another

Ligands Change delta, and Thus Color

*The greater the delta, the shorter the wavelength of the absorbed light

Charge-Transfer Transitions

*Some transition metal complexes with no d electrons (d0) are highly colored. Why?

*Ligand-to-metal charge transfer (LMCT); has fair higher molar absorptivity than seen for d-d transitions in dn complexes

ethylenediamine(en)

bipyrimidine (bipy or bpy)

diethlenetriamine