TRANSITION METAL COMPLEXES Chapter 25, VB/CF Handout

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1 TRANSITION METAL COMPLEXES Chapter 25, VB/CF Handout The energy of a covalent bond is largely the energy of resonance of two electrons between two atoms the resonance energy increases in magnitude with increase in the overlapping of the two atomic orbitals involved in the formation of the bond, the word overlapping signifying the extent to which regions in space in which the two orbital wave functions have large values coincide...consequently it is expected that of two orbitals in an atom the one which can overlap more with an orbital of another atom will form the stronger bond with that atom, and, moreover, the bond formed by a given orbital will tend to lie in that direction in which the orbital is concentrated. Linus Pauling, Nobel Prize in Chemistry in 1954 for "his research into the nature of the chemical bond and its application to the elucidation of the structure of complex substances" and Nobel Prize in Peace in 1962 oxyhemoglobin has an oxygen molecule bound to iron heme has oxygen, water, or carbon monoxide as 6th ligand EX 1. What is the electronic geometry about the Fe 2+ center above on the left? the Zn 2+ center on the right? ruby (Cr 3+ ) carbonic anhydrase enzyme in red blood cells Properties of Transition Elements (25.1, 25.2) The color of many gemstones is attributable to a transition element or a small concentration of a transition element impurity. emerald (Cr 3+ ) alexandrite (Cr 3+ ) FIG I. Atomic Radii for the Transition Elements garnet (Fe 2+ ) peridot (Fe 2+ ) turquoise (Cu 2+ )

2 - 2 - EX 2. Why are the atomic radii of the transition elements in the 6th period nearly identical to those in the 5th period? FIG II. Oxidation States for the Transition Elements EX 3. a) Why is 2+ a common oxidation state of most transition elements? b) Why is 1+ a common oxidation state for Ag? Coordination Compounds (25.3) metal Lewis acid (accepts an electron pair) ligand Lewis base (donates one or more electron pairs) and forms a coordinate covalent bond; a ligand is a neutral molecule or anion, rarely a cation coordination number number of ligands coordinated type of ligand classified according to number of coordinating III lone A pairs FIG Coordination monodentate bidentate polydentate monodentate bidentate

3 - 3 - FIG III. A Coordination Complex charge and oxidation number (must keep track of the electrons!) central metal and ligands directly bonded constitute first coordination sphere of complex; enclose in brackets anion: complex ion - [FeCl 4 ] - Fe 3+ with 4 Cl - ligands neutral: [Os(CO) 5 ] Naming Coordination Compounds (25.3) cation: coordination compound - [Co(NH 3 ) 5 Cl]Cl 2 [Co(NH 3 ) 5 Cl] 2+ complex ion with Co 3+, 5 NH3 and 1 Cl - ligand cation named first, anion second ligands listed first, alphabetically without regard to prefix; use di-, tri-, tetra-, penta-, hexa- for more than one kind of a ligand; bis(ligand), tris(ligand), tetrakis(ligand), pentakis(ligand), hexakis(ligand) if ligand name already has a numerical prefix metal listed last with the oxidation number in parentheses in roman numerals anions append the suffix -ate to the name of the metal; six metals use the Latin word for the metal in an anionic complex Ligand Names to Recognize ligand formula name monodentate fluoride ion F - fluoro chloride ion Cl - chloro bromide ion Br - bromo iodide ion I - iodo cyanide ion CN - cyano thiocyanate ion SCN - thiocyanato hydroxide ion OH - hydroxo water H 2 O aqua ammonia NH 3 ammine carbon monoxide CO carbonyl bidentate oxalate ion C 2 O 2-4 (ox) oxalato ethylenediamine H 2 NCH 2 CH 2 NH 2 (en) ethylenediamine Alfred Werner Swiss chemist who received Nobel Prize in 1913 for elucidating bonding in coordination compounds.

4 - 4 - FIG IV. Bidentate Ligands Fe Cu Pb Ag Au Sn EX 4. Give the formulas for a) pentaaaquachlorochromium(iii) ion FIG V. Hexadentate Ligand b) tetraamminecopper(ii) sulfate c) hexaamminechromium(iii) hexachloroiridate(iii) d) dichlorobis(ethylenediamine)cobalt(iii) chloride e) potassium hexacyanoferrate(ii) FIG VI. [CoEDTA] - Ion (Ribbon Model) (chelates Ca 2+ ) in bathroom cleaners, shower sprays; anticoagulant for blood; used in chelation therapy for treating mercury and lead poisoning; food preservative

5 - 5 - Coordination Compounds: Valence Bond (VB) Model (25.5, VB/CF handout) hybridization from valence bond theory (note two structures have a steric number of 4); concentrate on properly counting the electrons in the metal coordination sphere linear (SN=2) EX 5. [Cu(CN) 2] - Name: trigonal planar (SN=3) EX 6. [HgI 3] - Name: tetrahedral (SN=4) EX 7. [Co(Cl) 4] 2- Name: a blue copper protein

6 - 6 - square planar (SN=4) EX 8. [Pt(en) 2] 2+ Name: Pt: [Xe]4f 14 6s5d 9 Pt 2+ : [Xe]4f 14 5d 8 trigonal bipyramidal (SN=5) EX 9. PCl 5 Name: octahedral (SN=6) EX 10. [Co(NH 3) 6] 3+ Name:

7 - 7 - Structure and Isomerization (25.4) isomers: same formula but differ in structure and properties structural bonds differ coordination ion exchange between ligand and counter ion [Cr(NH 3 ) 5 SO 4 ]Br and [Cr(NH 3 ) 5 Br]SO 4 linkage ligand can bind in more than one way either C (cyano) or N (isocyano) end of :C N: stereoisomers (our main concern) spatial arrangement of atoms differs, bonding the same geometric (cis/trans) FIG VII. Isomers of [Pt(NH 3 ) 2 Cl 2 ] FIG VIII. Isomers of [Pt(NH 3 ) 4 Cl 2 ] + EX 11. Which below are identical to (1) and which are geometric isomers of (1)? The anti-cancer chemotherapy drug cisplatin introducing a 26 o kink into a piece of double-stranded DNA, beginning the process of apoptosis (cell death). EX 12. One isomer in FIG VIII absorbs light at 560 nm and the other absorbs at 680 nm. Which isomer absorbs at 560 nm? Why?

8 - 8 - optical isomers enantiomers (nonsuperimposable mirror images) molecules that have enantiomers (such as enzymes) are chiral occurs if complex 1) has a chiral ligand, 2) is tetrahedral with 4 different groups, or 3) is octahedral with no mirror plane FIG IX. Enantiomers of [Co(en)3] 3+ red bonds in front of Co and other three are behind EX 13. Draw the enantiomer of the following FIG X. Enantiomers of all cis [Co(NH3)2(H2O)2Cl2] +

9 - 9 - Coordination Compounds Bonding and Properties from Three Models (25.5) A summary of the three approaches to model and explain transition metal complexes Valence Bond Theory (24.5, VB/CF handout) 1. shape and bonding hybridization explains 2. magnetism sometimes incorrect 3. color and spectra inadequate explanation Crystal Field (CF) Theory (24.5, VB/CF handout) 1. magnetism - crystal field splitting (ΔE) explains a) octahedral: d orbitals split (Δo) into t2g/eg set Δo increases with increasing ON (M 2+ < M 3+ < M 4+ ) Δo increases down a group (3d < 4d < 5d) Δo depends upon the type of ligand b) tetrahedral: d orbitals split (Δt) into t2/e set Δt = 4 / 9 Δo => almost always high spin splitting reversed from octahedral c) square planar: d orbitals split into 4 (1 degenerate) and 1 high in energy => often d 8 2. color need empirical spectrochemical series but cannot explain as CF completely ignores nature of ligand 3. bonding no insight since ligand/metal interactions are assumed to be only electrostatic (i.e., bonding is purely ionic) Molecular Orbital Theory 1. what is analogous to the "stable octet" (# of valence electrons) => bonding, nonbonding, antibonding MOs 2. explains spectrochemical series a) σ only bonding ligands b) σ and π bonding ligands π donor donates electrons to metal via a filled ligand orbital (commonly a p) interacting with an empty metal orbital π acceptor accepts electrons from metal via an empty ligand orbital (commonly π*) interacting with a filled metal orbital four-leaf clover between axes shape of p z orbital but both lobes are positive, has a negative collar four-leaf clover on x,y axes The Five d Orbitals EX 14. How many nodes does a 3d orbital have? How many are radial? How many are angular?

10 The Crystal Field Model (25.5, VB/CF handout) ionic bonding ligands, (-) point charges, e - /e - repulsion with metal d orbitals crystal field (CF) splitting energy, Δ Assume each of the d orbitals below holds a pair of electrons. Bring in another pair of electrons along each of the x, y, and z axes in the above figure. Determine where the greatest electron repulsions will occur. These are the highest energy orbitals in the crystal field. FIG XI. Metal in an Octahedral Crystal Field The above d orbitals are degenerate (have the same energy) for the free transition metal ion without any ligands. However, they split in the crystal field yielding the energy levels given below for the three most common transition metal geometries. dx2-y2 high in energy, usually not occupied => often d 8 high spin (low spin very rare) low and high spin negative charges uniformly distributed over surface of sphere big = 4/9 Δ o FIG XII. CF Splitting in Square Planar, Tetrahedral, and Octahedral Crystal Fields

11 FIG XIII. Patterns of Filling d Orbitals in Octahedral Complexes Only four d electron configurations have the option of high or low spin. high spin weak field ligand π donor ligand small Δo synonyms low spin strong field ligand π acceptor ligand large Δo synonyms FIG XIV. High and Low Spin Octahedral Mn(III) Complexes Mn 2+ : 3d 5 FIG XV. High and Low Spin Octahedral Fe(II) Complexes Fe 2+ : 3d 6

12 FIG XVI. Tetrahedral and Square Planar Ni(II) Complexes Tetrahedral complexes have an inverted octahedral electron configuration with Δt = 4 / 9 Δo => almost always high spin and commonly d 10. BIG Ni 2+ : 3d 8 Cu 2+ : 3d 8 Square planar complexes have an orbital high in energy and generally unoccupied => usually low spin; favored by d 8 ions in 4th and 5th periods. The Color of Complex Ions and Spectroscopy (25.5) EX 15. Explain the effect or cause of the following with respect to the color of a transition metal complex a) complex has cis/trans isomers b) when it appears colorless c) when it has only a faint color d) nature of coordinating ligands e g Δ o = Eeg Et2g = E photon = hν = hc/λ FIG XVII. Crystal Field Splitting in Cr(III) Complexes FIG XVII gives a spectrochemical series: Δo splitting due to ligand, F - < H 2 O < NH 3 < CN -

13 The Molecular Orbital Model Sigma Bonding bonding cannot be fully ionic construct molecular orbitals metal d, s, p ligand p or sp 3 relationship of d orbitals to x, y, z axes orbitals closer in energy interact more strongly FIG XVIII. Metal/Ligand Sigma Bond Orbital Interactions in Octahedral Complexes Sigma bonding interaction between two ligand orbitals and metal d z 2 orbital Sigma bonding interaction between four ligand orbitals and metal d x 2 - y 2 orbital FIG XIX. MO Diagram (for [CrCl6] 3- ) 18 ELECTRON RULE (metal ~ low ON) EX 16. Consider the general MO diagram on the left without any electrons. a) What is the minimum number of electrons required to bond 6 ligands? Which MOs would be occupied? b) How many electrons are required to fill all of the bonding and nonbonding MOs? How many electrons are required to fill the valence shell of a transition element? Give the electron configuration. c) What is the maximum number of electrons required to fill all of the bonding and nonbonding MOs and the lowest energy antibonding MOs? (somewhat oversimplified rationale) some examples: 12-electron species: [TiF6] 2-18-electron species: Cr(CO)6 22-electron species: [Zn(en)3] 2+ metal atomic orbitals complex molecular orbitals ligand lone pair orbitals

14 Pi Bonding understand reason for spectrochemical series simple MO diagram is for σ bonding Can ligands also form π bonds to the metal? Look at t 2g ("nonbonding" d xy, d yz, d xz ) consider p and π*orbitals of ligand (those not involved in σ bonding) FIG XX. Pi Bonding with Ligand p Orbitals Pi bonding interaction between four ligand orbitals and metal d xy orbital The ligand sigma bonds to the metal. But when the ligand has additional electrons, especially in a p orbital, it can also form a pi bond to an unoccupied nonbonding metal d orbital: pi donor ligand The size of Δo decreases as the nonbonding metal electrons are pushed to higher energy by the ligand pi electrons. In addition to the normal sigma and pi bonding, one now has a pi bond with a nonbonding metal d orbital (lower right below)

15 There is another form of pi bonding. Remember the MO of carbon monoxide. The lowest unoccupied molecular orbital is the π* MO which has just the correct shape to doubly connect with the dxy, dyz, or dxz (the three four-leaf clover nonbonding d orbitals which are off axis ). FIG XX. Pi Bonding with Ligand p Orbitals Here, too, the ligand sigma bonds to the metal. But when the ligand has an empty π* anti-bonding MO it can also form a pi bond with a filled metal non-bonding d orbital: pi acceptor ligand The size of Δo increases as the nonbonding metal electrons are delocalized into the π* MO, lowering their energy. This interaction is frequently referred to as pi backbonding as it transfers charge back to the ligand.

16 Effect of Ligand on Octahedral Crystal Field Splitting (Δo) Pi bonding primarily effects the energy of the t2g transition metal complex MOs (the off axis nonbonding metal dxy, dyz, and dxz atomic orbitals). Consequently, the energy difference between the three t2g orbitals and the eg set of bonding metal dx2-y2 and dz2 atomic orbitals, known as the crystal field splitting (Δo), changes. Pi donor ligands are weak field ligands that have lone pair electrons not involved in σ bonding (such as those in a p orbital) whose orbital can overlap with an empty metal t2g orbital. This increases the electron density around the metal and raises the energy of the t 2g set. The size of the Δo energy gap decreases and a high spin complex results. On the other hand, pi acceptor ligands are strong field ligands that have an empty orbital (such as a π*) which can overlap with a filled metal t2g orbital. Delocalization decreases the electron density around the metal, lowering the energy of the t2g orbital. This increases the size of Δo and results in a low spin complex. Spectrochemical Series I - < Br - < Cl - < F < OH - < ox (high spin, weak ligand field, small Δo, π donor) (filled p) H2O < NH3 < en (intermediate) CN - CO (low spin, strong ligand field, large Δo, π acceptor) (empty π*) EX 17. Draw examples of the two types of π bonding to a metal. Use the proper metal orbital(s) and label them and the proper ligand orbital(s) and label them also. Draw the π donor and π acceptor examples with a coord-inate system showing the proper orientation.