ch. 22: transition metals and coordination compounds
TRANSCRIPT
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.