Ch. 22: Transition Metals and Coordination Compounds

Dr. Namphol Sinkaset Chem 201: General Chemistry II

I. Chapter Outline

I. Introduction II. Properties of Transition Metals III. Coordination Compounds IV. Bonding in Coordination Compounds V. Biomolecules

I. Introduction • Transition metal chemistry is colorful! • Cr3+, Fe2+, Cu2+.

I. Introduction • Coordination compounds involving transition

metals often have brilliant colors. Examples include Co(H2O)6

2+ and CoCl64-. But why is one pink and the other blue?

• In this chapter, we examine what is responsible for the colors seen in coordination compounds.

• We will need a different bonding theory to adequately explain properties of metal coordination compounds.

II. General Properties of Transition Metals

• In 1st semester G-chem, the transition metals are generally ignored.

• We review electron configurations of the transition metals.

• Additionally, we will summarize the atomic size, ionization energy, electronegativity, and oxidation states of these metals.

II. Electron Configurations

• Recall that the ns orbital fills before the (n-1)d orbitals.

• However, the ns orbital empties before the (n-1)d orbitals.

• Also, there are some strange fillings.

II. Sample Problem 22.1

• Write ground state electron configurations for Os and Nb2+.

II. Atomic Size • As expected, size

generally decreases across the period.

• As expected, size increases from 1st transition row to 2nd.

• Atoms in 3rd transition row are not larger due to lanthanide contraction.

II. 1st Ionization Energy

• As w/ main group, 1st IE increases as go across a period.

• 3rd row transition metals have higher Z and are same size as 2nd row. Thus, 1st IE’s require

more energy.

II. Electronegativity

• In general, EN increases across period, like main group.

• However, EN increases between 1st and 2nd period.

• Au has an exceptionally high EN.

II. Oxidation States

III. Coordination Compounds

• Recall that a complex ion has a central metal ion bound to ligands.

• When the complex ion combines with counterions, it forms a coordination compound.

• Coordination compounds have a primary valence and a secondary valence.

III. Coordination Compounds

• Primary valence is the oxidation state of the metal.

• Secondary valence is the # of ligands, a.k.a. the coordination number.

• Formula for this example is [Co(NH3)6]Cl3.

III. The Complex Ion

• The metal-ligand complex can be thought of as a Lewis acid-base adduct.

• The bond is called a coordinate covalent bond.

III. Ligands • There are monodentate, bidentate, and

polydentate ligands. • These terms refer to the # of lone pair e-’s

that can be donated to the central metal.

III. Bidentate and Polydentate

III. Common Ligands

• Note that some monodentate ligands can coordinate in different ways.

• e.g. carbon monoxide, cyanide, thiocyanate.

III. Common Geometries

• Coordination #’s can range from 2 to 12.

• The most common are 4 and 6, however.

III. Nomenclature

1) Name cation first then anion. 2) Name ligands in alphabetical order then

metal. Neutral ligands keep their name (some

exceptions like carbon monoxide carbonyl).

-ide ligands become -o. -ate ligands become -ato. -ite ligands become -ito.

III. Nomenclature

3) Denote # of ligands w/ Greek prefixes. If ligand has a Greek prefix in its name,

use bis-, tris-, or tetrakis- to indicate #. 4) If metal is in cation, use normal metal

name. If metal is in anion, add -ate suffix to the root of metal’s name.

5) Indicate oxidation state of metal w/ Roman numeral in parentheses after metal name.

III. Some Common Names

III. Sample Problem 22.2

• Determine the correct name or formula for the compounds below. a) K2[Ni(CN)4] b) [Co(H2O)4Cl2]Cl c) K[AuCl4] d) [Co(en)3]Br3 e) potassium hexacarbonylvanadate(-1) f) copper(II) hexacyanoferrate(II) g) hexaamminecobalt(III) chloride

III. Structure & Isomerization

• Because of the many bonding sites on the central metal, there is much more variation in structures of coordination compounds.

• Isomers are compounds that have the same formula but different structures.

• We can create a hierarchy of the different types of isomers.

III. Types of Isomers

III. Coordination Isomers

• Perhaps the easiest to understand, it’s when a coordinated ligand exchanges places with an uncoordinated counterion. e.g. [Co(NH3)5Br]Cl vs. [Co(NH3)5Cl]Br

III. Linkage Isomers

• As mentioned previously, some ligands have more than one site through which they can coordinate.

III. Example Linkage Isomers

• [Co(NH3)5NO2]2+ vs. [Co(NH3)5ONO]2+.

III. Geometric Isomers

• These isomers occur when ligands are bonded to different coordination sites. Have cis-trans isomers (same side,

opposite side) in square planar (MA2B2) octahedral complexes (MA4B2). Have fac-mer isomers (facial, meridian) in

octahedral complexes (MA3B3).

III. Cis-trans Isomerization

III. Fac-mer Isomerization

III. Optical Isomers

• Optical isomers are nonsuperimposable mirror images of one another. Most common analogy are right and left

hands. • Molecules or ions that have this quality

are called chiral, and the isomers are called enantiomers.

III. Example Optical Isomers

IV. Bonding in Coordination Compounds

• The common geometries found in coordination compounds can described w/ valence bond theory (VBT).

• VBT is inadequate to describe color and magnetic properties, however. We need a new theory, crystal field theory,

to explain properties of coordination compounds.

IV. Valence Bond Theory

• Recall that valence bond theory involves hybridization of atomic orbitals.

• The coordinate covalent bond forms from the overlap between a completely filled atomic orbital and an empty atomic orbital.

• The metal has the empty orbitals and the ligands have the filled orbitals.

IV. Metal Hybridizations

IV. Crystal Field Theory

• VBT is good for geometries, but nothing else.

• Crystal field theory (CFT) focuses on what happens when e-’s on the ligands approach the central metal.

• When these e-’s come in, they repel e-’s in the unhybridized orbitals of the metal. This results in destabilization of the metal’s

unhybridized d orbitals.

IV. CFT for Oh Complexes

IV. d Orbital Splitting • Orbitals in direct line w/ ligands will get

destabilized most, resulting in two sets of d orbitals in the complex.

IV. Colors of Complex Ions • Colors of solutions of complex ions

arise from electrons transitioning between the split d orbitals.

IV. Measuring the Splitting

• Thus, if we take the absorption spectrum, we can see what wavelength is absorbed by the e- as it moves up.

• The splitting is equal to the energy of the photon via the equation Ephoton = hν = hc/λ.

• We look at the spectrum of Ti(H2O)63+

as an example.

IV. Splitting in Ti(H2O)63+

IV. Crystal Field Splitting Energy

• Δ is called the crystal field splitting energy (CFSE).

• Different ligands will result in greater CFSE.

• If the energy difference is large, it’s a strong-field complex.

• If the energy difference is small, it’s a weak-field complex.

IV. Spectrochemical Series • Observations of the splitting in different

metal complexes allow for prediction of which ligands will split the d orbitals most.

• This list is the spectrochemical series, going from strong-field to weak-field ligands.

• CN- > NO2- > en > NH3 > H2O > OH- > F-

> Cl- > Br- > I-. • Additionally, high-charge metal cations

promote strong fields.

IV. Magnetic Properties

• As you know, magnetic properties depend on existence of unpaired e-’s.

• According to Hund’s rule, e-’s will maximize spin before pairing. This is because pairing spins costs energy.

• If the CFSE is small enough, the energy cost of pairing is higher than moving up to higher energy orbitals.

IV. Fe2+ in Strong & Weak Fields • How many d e’s are in Fe2+? Fe is d8, but we subtract 2 e-’s for charge. Thus, Fe2+ has 6 d e-’s.

Low-spin complex High-spin complex

IV. Sample Problem 22.3

• How many paired electrons would you expect in the following complexes? Are they diamagnetic or paramagnetic?

a) [FeCl6]3- b) [Co(CN)6]4-

IV. Tetrahedral Complexes • The splitting in tetrahedral complexes is

opposite that of octahedral. • Almost all tetrahedral complexes are high-

spin due to less ligand-metal interactions.

IV. Square Planar Complexes • Square planar complexes have a complex

splitting pattern. • Square planar complexes occur in d8 metals

like Pt2+, Pd2+, Ir+, and Au3+. • They are normally low-spin.

V. Useful Metal Complexes • Living system contain many molecules that

have metal complexes. • Transition metals are important components

of life.

V. Porphyrins

• A porphyrin is a ligand that has a planar ring structure w/ four nitrogen atoms that can coordinate to a central metal.

V. Hemoglobin

V. Chlorophyll

• As you can see, it has the porphyrin ligand coordinated to magnesium.

• Different from hemoglobin, the complex is not surrounded by a protein.

V. Anticancer Drugs

• Cisplatin is an effective anticancer drug.

• The trans isomer is ineffective.

• It is thought that the trans isomer cannot bind correctly to cancer cell’s DNA.