CHEM 152 Coordination Chemistry Lectures 2022 PDF

CHEM 152 : COORDINATION CHEMISTRY LECTURES BY : PROF. N. K. ASARE- DONKOR  OBJECTIVES At the end of the course students will be able to:  Predict properties of diatomic molecules using molecular orbital theory  Name and write formulas for complex ions and coordination compounds  Identif...

CHEM 152 : COORDINATION CHEMISTRY LECTURES BY : PROF. N. K. ASARE- DONKOR  OBJECTIVES At the end of the course students will be able to:  Predict properties of diatomic molecules using molecular orbital theory  Name and write formulas for complex ions and coordination compounds  Identify and draw the various types of isomers (including geometric and optical isomers)  Apply the various theories to bonding in coordination compounds. Distortion from regular geometry-symmetry, Magnetic and magnetic properties of Transition Metals.  Describe applications of coordination compounds (emphasis on biological systems)  COURSE OUTLINE 1.Chemical Bonding II: Molecular Orbitals of diatomic molecules 2. Coordination Chemistry:- Development of Coordination Chemistry, Nomenclature, Stereochemistry, Isomerism in complexes, Stability of complexes, Chelate effect 3. Theories of bonding in Coordination Compounds: VBT, MOT CFT/LFT, Crystal Field stabilization energies, Distortion from regular geometry-symmetry, Magnetic and magnetic properties of Transition Metals. 4. Introductory Bioinorganic Chemistry  RECOMMENDED TEXT BOOKS Shriver, D. and Atkins, P., Inorganic Chemistry, 4th Edition Housecroft, C. and Sharpe, A. G., Inorganic Chemistry, 5th Edition. Pearson Education Limited. Miessler, G. L. and Tarr, D. A.,Inorganic Chemistry, 5th Edition Swaddle, T. W., Inorganic Chemistry: An Industrial and Environmental Perspective Mark Weller, Jonathan Rourke, Tina Overton and fraser Armstrong 7th edition. Inorganic Chemistry, Oxford University Press, great Claredon Street, Oxford, United Kingdom.  Silberberg,M. S. sand Amates P. Chemistry – The Molecular Nature of Matter and Change with Advanced Topics, Eighth Edition. The McGraw-Hill Education, 2 Penn Plaza, New York  NivaldoJ. Tro, Chemistry: A Molecular Approach 5th Edition Hoboken, NJ. Pearson Education, Inc.  Advanced Inorganic Chemistry by F. A. Cotton, Geoffrey Wilkinson et al. 5 Mode of Delivery 1. Direct Delivery in lecture rooms 2. On-line Lectures 3. Online lecture notes / power point materials Mode of Assessment  Class Assignments  Impromptu’ Quizzes and Tests  Online Assignments/Tests  Class tests  Mid-Semester Examinations  End of Semester Examinations  Previous Knowledge  Basic Inorganic Chemistry (i) Atomic Structure, Qualitative wave mechanics (ii) Periodic Table and periodicity, Reactivity Parameters (iii) Chemical Bonding 1: ionic bond, covalent bond, Dative (coordinate) bond, valence bond theory, Resonance, Multiple bond, shapes of Molecules, Hybridization of non-transition elements (iv) Forces within molecules, Bond strength, Bond energy, Polarity, Continuity of bonds (v) Crystal structure, X-ray diffraction  Chemical bonding II: Molecular Orbitals of Diatomic molecules Introduction The three main types of bonding are: Covalent bonding between atomic pairs (two-centre bonds) – The localized bond. Delocalized (multicentre) covalent bonding Ionic bonding. In the localized bond approach which is the simplest bonding in any molecule or complex: (i) The electrons involved in bonding remain localized between pairs of atoms. (ii) The bonding in the whole structure is the sum of the individual bonds between the pairs of atoms. Theories under the localized bond approach are the Lewis electron-pair bond theory, the hybridization theory, the valence shell electron-pair repulsion (VSEPR) theory and the valence bond theory. Although there is satisfying correspondence among these theories, each constitute a separate approach and address a distinctly different aspect of the localized bond problem. Since each of these theories has advantages and limitations, we must learn to move readily from one theory to the other, depending on which bonding or structural features we seek to explain.  Introduction to the Molecular orbital Theory The Molecular Orbital Theory does a good job of predicting electronic spectra and Para magnetism, when VSEPR and the VB Theories do not. The MO theory does not need resonance structures to describe molecules, as well as being able to predict bond length and energy. The major draw back is that we are limited to talking about diatomic molecules (molecules that have only two atoms bonded together), or the theory gets very complex. The MO theory treats molecular bonds as a sharing of electrons between nuclei. Unlike the VB theory, which treats the electrons as localized balloons of electron density, the MO theory says that the electrons are delocalized. That means that they are spread out over the entire molecule.  Molecular Orbital Theory The goal of molecular orbital theory is to describe molecules in a similar way to how we describe atoms, i.e. in terms of orbitals, orbital diagrams and electronic configurations. Each line in the diagram represents an orbital. The molecular orbital volume encompasses the volume of the molecule. The electrons fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms. 12  In the case of diatomic molecules, the interactions are easy to see and may be thought of as arising from constructive interference (same sign) of electron waves (orbitals) of two different atoms, producing bonding molecular orbital and destructive interference (opposite sign) of the electron waves, producing an antibonding molecular orbital. This approach is referred to as LCAO-MO (Linear Combination of Atomic Orbitals to Produce Molecular Orbitals) 13  Atomic orbitals Heisenberg Uncertainty Principle states that it is impossible to define what time and where an electron is and where it is going next. This makes it impossible to know exactly where an electron is travelling in an atom. Since it is impossible to know where an electron is at a certain time, a series of calculations are used to approximate the volume and time in which the electron can be located. These regions are called Atomic orbitals. These are also known as the quantum states of the electrons.  Only two electrons can occupy one orbital and they must have different spins states, +½ spin and -½ spin (easily visualized as opposite spin states) Orbitals are grouped into subshells This field of study is called quantum mechanics  Atomic Subshells These are some examples of atomic orbitals: S subshell: (Spherical in shape) There is one S orbital in an s subshell. The electrons can be located anywhere within the sphere centered at the atom’s nucleus P subshell: (Shaped like two balloons tied together) There are 3 orbitals in a p subshell that are denoted as px, py, pz orbitals. These are higher in energy than the corresponding s orbitals.  D subshell: The d subshell is divided into 5 orbitals (d xy, dxz, dyz, dx2- y2, dz2). These orbitals have very complex shapes and are higher in energy than the S and p orbitals.  Atomic and Molecular Orbitals In atoms electrons occupy atomic orbitals, but in molecules they occupy similar molecular orbitals which surround the molecules. Two 1s atomic orbitals combine to form two molecular orbitals, one bonding () and one antibonding (*). This is an illustration of the molecular orbital of H2 Note that each electron from each atom is being “shared” to form a covalent bond. This is an example of orbital mixing. 18 The He “dimer” 19 Examples of Sigma Bond Formation 20 Molecular Orbital Diagram (H2) 21 Molecular Orbitals from p AOs 22 Molecular Orbitals from p AOs 23 Diatomic Molecules: MO diagram for O2 24 MOs from O2 25  MO diagram for F2 Note: 1. there is no mixing of AO’s of same symmetry from a single F atom because there is a sufficient difference in energy between thethe 2. 2s and 2p orbitals π-type in F. by the combination of the P and P MO’s formed x y orbitals make degenerate sets (i.e. they are identical in energy) 26 27  MO diagram for B2 In the MO diagram for B2, there are several differences from that of F2. Most importantly, the ordering of the orbitals is changed because of mixing 28 Molecular orbital diagrams for Li2 to F2  The separation between the ns and np orbitals increases with increasing atomic number This means that as we go across the second row of the periodic table, the amount of mixing decreases until there is no longer enough mixing to affect the ordering This happens at O2. At O2, the ordering of the 2σg and the 1πu MO’s changes As we go across the second row with increasing atomic number, the effective nuclear charge (and electronegativity) of the atoms increases This is why the energies of the analogous orbitals decreases from Li2 to F2 31 Molecule Li2 Be2 B2 C2 N2 O2 F2 Ne2 Electrons 6 8 10 12 14 16 18 20 Net bonds 1σ 0 (1σ) 1σ, 2π 1σ, 1π 1σ 0 Bond order 1 0 1 2 3 2 1 0 Unpaired 0 0 2 0 0 2 0 0 electrons Bond length(Å) 2.67 n/a 1.59 1.24 1.01 1.21 1.42 n/a Bond Energy 105 n/a 289 609 941 494 155 n/a (Kjmol-1) Diamagnetic(d)/ d n/a p d d p d n/a Paramagnetic(p) The trends in bond length and energy, can be understood The trends in bond length and energy, can be understood from the size of each atom, the bond order and by examining the orbitals that are filled Bond Order is given as by examining the orbitals that are filled om the size of each atom, the bond order and by examining the orbitals that are filled The trends in bond length and energy, can be understood from the size of each atom, the bond order and by examining the orbitals that are filled 33  Electronic Configurations Li2 = KK σ2s2 Be2* = KK σ2s2σ*2s2 C2 = KK σ2s2σ*2s2σ2p2π2p1π2p1 N2 = KK σ2s2σ*2s2σ2p2π2p4 O2 = KK σ2s2σ*2s2σ2p2π2p4π*2p2 (O2 = KK σ2s2σ*2s2σ2p2π2p4π*2px1 π*2py1) F2 = KK σ2s2σ*2s2σ2p2π2p4π*2p4 Molecular Orbitals for Heteronuclear Diatomic Molecules Molecular orbital diagram Hydrogen Fluoride (HF)  Note: 1. The F 1s orbital is at much lower energy than the H 1s orbital (because of the higher nuclear charge) 2. The F 1s2 electrons are core electrons – Their energy does not change when HF is formed. 3. The H 1s and the F 2p valence electrons go into molecular orbitals with new energies. 36 Molecular orbital diagram Molecular orbital diagram for Carbon monoxide (CO) Iodine monochloride (ICl) Molecular orbital diagram Nitric oxide (NO) 38  Coordination compounds/complexes  The d block metal form coordination complexes with molecules and ions.  Coordination complexes What is the electronic basis of the colour of metal complexes?  Coordination complex: A structure containing a metal (usually a metal ion) bonded (coordinated) to a group of surrounding molecules or ions (ligands). Ligand (ligare is Latin, to bind): A ligand is a molecule or ion that is directly bonded to a metal ion in a coordination complex A ligand uses a lone pair of electrons (Lewis base) to bond to the metal ion (Lewis acid) Coordination sphere: A metal and its surrounding ligands Note: religion is derived from Latin: religare, to bind tightly Coordination Complexes: Three common structural types Octahedral: Square planar Tetrahedral Most important What determines why a metal takes one of these shapes?  Lewis acids and bases A Lewis base is a molecule or ion that donates a lone pair of electrons to make a bond Examples: NH 3 O H2 Cl - F - Electrons in the highest occupied orbital (HO) of a molecule or anion are the best Lewis bases A Lewis acid is a molecule of ion that accepts a lone pair of electrons to make a bond + 3+ 2+ n+ Examples: H Co Co M Molecules or ions with a low lying unoccupied orbital (LU) of a molecule or cation are the best Lewis acids  The formation of a coordinate complex is a Lewis acid-base reaction Lewis base: NH 3 Coordination complex: Lewis base (electron pair donor) coordinated to a Lewis acid (electron pair acceptor) Coordination complex: Ligand (electron donor) coordinated to a metal (electron acceptor) Lewis acid: Co3+ The number of ligand bonds to the central metal atom is termed the coordination number  The basic idea is that the ligand (Lewis base) is providing electron density to the metal (Lewis acid) The bond from ligand to metal is covalent (shared pair), but both electrons come from the ligand (coordinate covalent bond) In terms of MO theory we visualize the coordination as the transfer of electrons from the highest occupied valence orbital (HO) of the Lewis base to the lowest unoccupied orbital (LU) of the Lewis acid Lewis base Le wis b as e Le wis ac id HO LU Lewis acid NH 3 Co3+  Types of Ligands: Bidentate (two tooth) Ligands Some common bidentate (chelates): (en) Ethylenediaminetetraacetate ion (EDTA): a polydentate chelating ligand Chelate from Greek chela, “claw” EDTA wraps around the metal ion at all 6 coordination sites producing an exceedingly tight binding to the metal  Alfred Werner: the father of the structure of coordination complexes The Nobel Prize in Chemistry 1913 "in recognition of his work on the linkage of atoms in molecules by which he has thrown new light on earlier investigations and opened up Alfred Werner new fields of research especially in Switzerland University of Zurich inorganic chemistry" Zurich, Switzerland b. 1866 (in Mulhouse, then Germany) d. 1919  Werner’s Theory Alfred Werner suggested in 1893 that metal ions exhibit what he called primary and secondary valences.  Primary valences were the oxidation number for the metal (+3 on the cobalt at the right).  Secondary valences were the coordination number, the number of atoms directly bonded to the metal (6 in the complex at the right). 49  Properties of Some Ammonia Complexes of Cobalt(III) Many coordination compounds are brightly coloured. Different coordination compounds from the same metal and ligands can give quite different numbers of ions when they dissolve. The central metal and the ligands directly bonded to it make up the coordination sphere of the complex. 50  In CoCl3 ∙ 6 NH3, all the six NH3 groups are bonded to the cobalt and the 3 chloride ions are outside the coordination sphere. In CoCl3 ∙ 5 NH3, the five NH3 groups and one chlorine are bonded to the cobalt, and the other two chloride ions are outside the coordination sphere. In CoCl3 ∙ 4 NH3, the four NH3 groups and two chlorines are bonded to the cobalt, and the other one chloride ion is outside the coordination sphere. 51 Werner proposed putting all molecules and ions within the coordination sphere in brackets and those “free” anions (that dissociate from the complex ion when dissolved in water) outside the brackets. 52  This approach correctly predicts there would be two forms of CoCl3 ∙ 4 NH3. The formula would be written [Co(NH3)4Cl2]Cl. One of the two forms has the two chlorines next to each other. The other has the chlorines opposite each other. 53  The Kepert Model In the light of the success of VSEPR theory in predicting the shapes of molecular species of the p-block elements we might reasonably expect the structures of the complex ions: [V(H2O)6]3+ (d2), [Mn(H2O)6]3+ (d4), [Co(H2O)6]3+ (d6), [Ni(H2O)6]2+ (d8) and [Zn(H2O)6]2+ (d10) to vary as the electronic configuration of the metal ion changes. However, each of these species has an octahedral arrangement of ligands. Thus, it is clear that VSEPR theory is not applicable to d- block metal complexes. 54 23-54  We turn the Kepert model which rationalizes the shapes of d-block metal complexes ([MLn], [MLn]m+ or [MLn]m-) by considering the repulsions between the groups L. Lone pairs of electrons are ignored. For coordination numbers between 2 and 6, the following arrangements of donor atoms are predicted: 2 linear 3 trigonal planar 4 tetrahedral 5 trigonal bipyramidal or square-based pyramidal 6 octahedral 55 23-55  The Kepert model presumes that: 1. Ligands geometry is not determined by the presence of lone pairs on the central atom 2. Ligand-ligand repulsion determines geometry 3. Ligand geometry may be constrained by other factors such as – Ligand back bone, Chelating ligands, tripodal ligands – Specific orbital overlap (d8 square planar systems) 56  Geometry and shape of Coordination Compounds  The coordination number of coordination compounds bonds formed by the metal ion and the ligands vary from 2-9, with 6 being most common. Coordination Number = 2,  Give rise to then linear, Rare for most metals  Common for d10 systems (Cu+, Ag+, Au+, Hg2+) Coordination Number = 3  Trigonal planar  Rare for most metals  Is known for d10 systems (example HgI3‑) 57 Coordination Number = 4  Tetrahedral, or square planar  Tetrahedral structure is observed for non-transition metals, BeF42- and d10 in ions such as ZnCl42-, FeCl4-, FeCl42-  Square planar is found with second and third row transition metals with d8 Rh+, Pd2+ Coordination Number = 5  Trigonal bipyramid  Square pyramidal 58 Coordination Number = 6,  Octahedral and prismatic (rare) Coordination Number = 7  Relatively uncommon, pentagonal bipyramidal,  Second and third row transition metals, lanthanides, and actinides Coordination Number = 8  Relatively common for larger metal ions, common geometry antiprism and dodecahedron Coordination Number = 9  For larger metal ions, geometry tricapped trigonal prism [Nd(H2O)9]3+ 59  Summary of Coordination geometries 60 23-60  Nomenclature Cation - Anion: Salts: name cation before anions i.e., [Co(NH3)5Cl]Br2, we name [Co(NH3)5Cl]2+ complex ion before bromides counter ions. Complex: Within complex ion, the ligands are named in alphabetical order before the metal i.e., pentaaminechlorocobalt(II) (note that penta is an indication of the number of NH3 group, and not considered in the alphabetizing of the ligand). Ligand: Anionic ligands end in -o and neutral ligands are named based on their molecular name (exceptions are aqua H2O, amine: NH3) Greek prefixes are used to indicate number of ligands, di-, tri-, tetra-, penta-, hexa-. Exception occurs when ligand already has Greek prefix in its name, The prefixes bis-, tris-, tetrakis-, pentakis, & hexakis. are used instead.  i.e., Ir(bpy)3 -trisbipyridineiridium (III) (bipyridine already has bi in its name). If the complex is an anion, then its name ends with suffix -ate. Furthermore, oxidation state of the metal is given in Roman numerals in parenthesis at the end of the name. Writing Formula of complexes Rules: 1. The cation is written before the anion. 2. The charge of the cation(s) is balanced by the charge of the anion(s). 3. For the complex ion, neutral ligands are written before anionic ligands (negative charge), and the whole ion is placed in brackets. 62 Writing names from Formula of complexes Procedure : 1. The cation is named before the anion. 2. Within the complex ion, the ligands are named, in alphabetical order. 3. Neutral ligands generally have the molecule name. Anionic ligands drop the -ide and add -o after the root name. 4. Numerical prefixes denote the number of a particular ligand. 5. Oxidation state of metal ion is in Roman numeral in parenthesis – Stock notation/convention. 6. Alternatively, the proportion of constituents may be given by means of stoichiometric prefixes, or the charge on the ion can be designated by the Ewens-Bassett number - Ewens-Bassett convention. 7. For anionic complex, the end of the metal name is replaced by –ate i.e. name ends -ate. 8. Complex cations and neutral molecules are given no distinguishing ending. 9. Formulas and names may be supplemented by italized prefixes cis, trans, fac, mer, etc. Name from formula a) K3[Au(CN)4] Potassium tetracyanoaurate(I) or (3-) b) K[Co(NH3)2 (C2O4)2] Potassium diaminedioxaloCobaltate(III) c) [Cr(en)2F2]NO3 Bis(ethylenediamine)difluorochromium(III) nitrate 64 d) K2[OsCl5N] Potassium pentachloronitridoosmate(VI) or potassium pentachloronitridoosmate (2-) e) Na3[Ag(S2O3)2] Sodium bis(thiosulphato)argentate(I) or (3-) f) [Pt(py)4][PtCl4] tetrakis(pyridine)platium(II) tetrachloroplatinate(II) or (2+) and (2-) Formula from Name a) Tetraamminechromium(III) nitrate [Cr(NH3)4] (NO3)3 b) dichlorobis(ethylenediamine)platinum(IV) bromide [PtCl2(en)2]Br2 c) bis(ethylenediamine)zinc(II) tetraiodomercurate(II) [Zn(en)2][HgI4]  Naming metals in anionic complexes Iron: Ferrate Copper: Cuprate Lead: Plumbate Silver: Argentate Gold: Aurate Tin: Stannate Osmium: Osmate Cobalt: Cobaltate Amtimony: Antimonate Rhenium: Rhenate Platinum: Platinate Rhodium: Rhodate  Ligands Consider [Ag(NH3)2]+ Ligand (contains the donor atom, directly bonded to metal) :NH3 - ligand occupy one site in coordinate sphere (monodentate) examples (Monodentate) N3- , X-, CN- , OH-, NH3 , pyridine, H2O Polydentate ligand - known as chelating agents - ligand which has several donor sites that can multi-bond (coordinate) metal simultaneously (chelates) i.e. en, oxalate, 1.10 phenanthroline, carbonate, bipyridine [EDTA]4- or (ethylenediaminetetraacetate), phenylpyridine Examples of Typical mono and poly dentate Ligands 68 69 Names of Some Common ligands 70  ISOMERISM One of the interesting aspects of the chemistry of coordination compounds is the possibility of the existence of isomers. Isomers of a compound contain the same numbers and types of atoms, but they have different structures. Several types of isomerism have been demonstrated, but only a few of the most important types will be described here 71  Isomers Structural isomers are molecules that have the same number and type of atoms, but they are attached in a different order. Stereoisomers are molecules that have the same number and type of atoms, and that are attached in the same order, but the atoms or groups of atoms point in a different spatial direction. © 2015 Pearson Education, Inc.  Structural Isomerism Ligand / Linkage Isomerism: Same complex ion structure but point of attachment of at least one of the ligands differs. [Co(NH3)4(NO2)Cl]Cl and [Co(NH3)4(ONO)Cl]Cl Linkage Isomers [Co(NH3)5(NO2)]Cl2 [Co(NH3)5(ONO)]Cl2 Pentaamminenitrocobalt(III) Pentaamminenitritocobalt(III) chloride chloride 73 Ionization Isomerism Complex salts which show ionization isomerism are composed in such a way that a ligand and a counter ion change their places Although the compounds [Pt(en)2Cl2]Br2 and [Pt(en)2Br2]Cl2 have the same empirical formulas, they are quite different compounds. Coordination Isomerism Coordination isomerism refers to cases where there are different ways to arrange several ligands around two metal centers. If in a complex salt both anion and cation are complexes there can be an exchange of ligands between cation and anion 74 Hydrate Isomerism This is a special case of the ionization isomerism. Here water molecules are present as ligand in one case and as water of crystallzation in the second case For example, [Cr(H2O)4Cl2]Cl.2H2O and [Cr(H2O)5Cl]Cl2.H2O have the same formula, but they are obviously different compounds. [Cr(H2O)6]Cl3 [CrCl(H2O)5]Cl2.H2O [CrCl2(H2O)4]Cl.2H2O 75  Polymerization Isomerism In coordination chemistry, it is possible to have two or more compounds that have the same empirical formula but different molecular weights. For example, [Co(NH3)3Cl3] consists of one cobalt ion, three ammonia molecules, and three chloride ions. This is the same ratio of these species as is found in [Co(NH 3)6][CoCl6], which has a molecular weight that is twice that of [Co(NH 3)3Cl3]. Other compounds that have the same empirical formula are [Co(NH3)5Cl] [Co(NH3)Cl5] and [Co(NH3)4Cl2][Co(NH3)2Cl4]. In coordination chemistry, it is possible to have two or more compounds that have the same empirical formula but different molecular weights. For example, [Co(NH3)3Cl3] consists of one cobalt ion, three ammonia molecules, and three chloride ions. This is the same ratio of these species as is found in [Co(NH3)6][CoCl6], which has a molecular weight that is twice that of [Co(NH3)3Cl3]. Other compounds that have the same empirical formula are [Co(NH3)5Cl][Co(NH3)Cl5] and [Co(NH3 )4Cl2 ][Co(NH3 )2 Cl4]. 76 Stereoisomerism 1 Atoms or groups arranged differently spatially relative to metal ion 77 20_444 Cl Cl H3N NH 3 H3N NH3 Co Co H3N NH 3 H3N Cl Cl NH 3 Cl Cl Co Co Cl Cl (a) (b) Cl 20_444 Cl H3N NH 3 H3N NH3 Co Co H3N NH 3 H3N Cl Cl NH 3 Cl Cl Co Co Cl Cl (a) (b) 78  Facial and Meridional isomerism An octahedral complex with formula, MX3Y3 e.g. Co(NH3)3Cl3] has two isomers 79  Stereoisomerism 2 Optical isomerism: For structures that do not possess a plane of symmetry, the mirror images are not superimposable. Known as chiral structures, such molecules rotate a beam of polarized light. If the beam is rotated to the right (when looking along the beam in the direction of propagation), the substance is said to be dextrorotatory (or simply dextro) and indicated by (+). Those substances that rotate the plane of polarized light to the left are said to be levorotatory or levo and indicated as(-). 80  Optical isomerism: Have opposite effects on plane-polarized light (no superimposable mirror images) 20_446 Polarizing filter Tube containing Unpolarized sample light  Polarized light Rotated polarized light 20_448 Mirror image of right hand Left hand Right hand 20_450 Cl The trans isomer and Cl Isomer II cannot be its mirror image are superimposed exactly N N identical. They are not N N on isomer I. They are Co isomers of each other. Co not identical structures. N N N Cl Cl Cl Cl Cl N N N N Cl N N trans Co cis Co Co N N N N N Cl Cl Isomer I N Isomer II N Isomer II has the same structure as the mirror (a) (b) image of isomer I. 20_449 N N N Mirror image Co N N of Isomer I N N N N N N N Co Co N N N N Isomer I Isomer II N N  Theories of Bonding in Coordination Compounds 86  Basis for Bonding Theories Models for the bonding in transition metal complexes must be consistent with observed behavior. Specific data used include stability (or formation) constants, magnetic susceptibility, and the electronic (UV/Vis) spectra of the complexes. 87  Valence Bond theory  Indicates hybridization in octahedral complexes.  For the first row transition metals, the hybridization can be: d2sp3 (using the 3d, 4s and 4p orbitals), or sp3d2 (using the 4s, 4p and 4d orbitals).  The valence bond approach is inadequate because it fails to explain the electronic spectra and magnetic moments of most complexes. 88 89  Crystal Field Theory In crystal field theory, the electron pairs on the ligands are viewed as point negative charges that interact with the d orbitals on the central metal. The nature of the ligand and the tendency toward covalent bonding is ignored. 91 d Orbitals 92  Ligands, viewed as point charges, at the corners of an octahedron affect the various d orbitals differently. 93 94  The repulsion between ligand lone pairs and the d orbitals on the metal results in a splitting of the energy of the d orbitals. 95  d Orbital Splitting __ __ e g dz2 dx2-y2 0.6∆o ∆o __ __ __ __ __ Spherical field 0.4∆o __ __ __t 2 dxy dxz dyz g Octahedral field 96  In some texts and articles, the gap in the d orbitals is assigned a value of 10Dq. The upper (eg) set goes up by 6Dq, and the lower set (t2g) goes down by 4Dq. The actual size of the gap varies with the metal and the ligands. The colours exhibited by most transition metal complexes arises from the splitting of the d orbitals. As electrons transition from the lower t2g set to the e set, light in the visible range is absorbed. 97  The splitting due to the nature of the ligand can be observed and measured using a spectrophotometer. Smaller values of ∆o result in colours in the green range. Larger gaps shift the colour to yellow. 98 THE ELECTROMAGNETIC SPECTRUM 99  The magnitude of the splitting (ligand effect) Strong Weak field field The energy gap between t2g and eg levels is designated Do or 10Dq The magnitudes of Do depends upon the metal ion and the attaching ligand. Magnitudes of Do are typically :– 100 - 400 kJ/mol (8,375 - 33,500 cm-1) 11 kJ/mol = 83.7 cm–1  Placing electrons in d orbitals Strong field Weak field Strong field Weak field d1 d2 d3 d4  Placing electrons in d orbitals d5 d6 d7 1 u.e. 5 u.e. 0 u.e. 4 u.e. 1 u.e. 3 u.e. d8 d9 d10 2 u.e. 2 u.e. 1 u.e. 1 u.e. 0 u.e. 0 u.e. When the 4th electron is assigned it will either go into the higher energy eg orbital at an energy cost of D0 or be paired at an energy cost of P, the pairing energy. d4 Strong field = Weak field = Low spin High spin (2 unpaired) (4 unpaired) P < Do P > Do Notes: the pairing energy, P, is made up of two parts. 1) Coulombic repulsion energy caused by having two electrons in same orbital  Crystal–Field Stabilisation Energy (CFSE) CFSE(Oh) = (–0.4x + 0.6y )Do + nP Added to this is a pairing energy term of 2P which accounts for the spin pairing associated with the two pairs of electrons in excess of the one in the high-spin configuration.  Pairing Energy, P (P) The pairing energy, P, is made up of two parts. 1) Coulombic repulsion energy caused by having two electrons in same orbital. Destabilizing energy contribution of Pc for each doubly occupied orbital. 2) Exchange stabilizing energy for each pair of electrons having the same spin and same energy. Stabilizing contribution of Pe for each pair having same spin and same energy P = sum of all Pc and Pe interactions  The Spectrochemical Series Based on measurements for a given metal ion, the following series has been developed: Spectrochemical Series for ligands I-