Metal Complexes and Organometallics - Winter Semester 2024-2025 PDF

Module-2 Metal Complexes and Organometallics Winter Semester 2024-2025 1  Contents…..(6 hours) Inorganic complexes - structure, bonding and applications Organometallics – introduction, stability, structure and applications of m...

Module-2 Metal Complexes and Organometallics Winter Semester 2024-2025 1  Contents…..(6 hours) Inorganic complexes - structure, bonding and applications Organometallics – introduction, stability, structure and applications of metal carbonyls and ferrocene Metals in biology: haemoglobin and chlorophyll- structure and property 2  Inorganic Complexes Inorganic/coordination complex is a molecule containing one or multiple metal centers that is bound to ligands (atoms, ions, or molecules that donate electrons to the metal). These complexes can be neutral or charged. The examples are: Neutral Complexes: [CoCl3(NH3)3], K4[Fe(CN)6], etc. Cationic Complex : [CO(NH3)6]3+ and Anionic Complex : [CoCl4(NH3)2]− Selected examples of metal complexes with names: [Co(NH3)5Cl]Cl2 --- Chloropentaamminecobalt(III) chloride [Cr(H2O)4Cl2]Cl --- Dichlorotetraaquochromium(III) chloride K[PtCl3(NH3)] --- Potassiumtrichloroammineplatinate(II) [PtCl2(NH3)2] --- Dichlorodiammineplatinum(II) [Co(en)3]Cl3 --- tris(ethylenediamine)cobalt(III)chloride [Ni(PF3)4] --- tetrakis(phosphorus(III)fluoride)nickel(0) 3  Structure and Bonding  Double Salt: Double Salts Vs Coordination Compounds Ferric alum (NH4)2SO4.Fe2(SO4)3.24H2O In water: NH4+, SO42-, Fe3+ Co-ordination Compounds NH3 3+ 3Cl– Fe(CN)2 + 4KCN Fe(CN)2.4KCN H3N NH3 (counter ion) Co H3N NH3 NH3 H 4K+ + [Fe(CN)6]4- ligand (coordination sphere) N M H H N forms a coordinate covalent bond to the metal 5 Ligands  Molecule or ion having a lone electron pair with an atom (donor) that can be donated to a metal atom forming a dative bond is called a Lewis base.  coordinate covalent bond: metal-ligand bond  monodentate : one bond to metal ion  bidentate : two bonds to metal ion  polydentate : more than two bonds to a metal ion possible 6 Chelating Agents EDTA-Na2 Bind to metal ions removing them from solution. Phosphates are used to tie up Ca2+ and Mg2+ in hard water to prevent them from interfering with detergents. Sodium tripolyphosphate Na5P3O10 EDTA-Metal complex Important biomolecules like heme and chlorophyll are porphyrins four pyrrole rings (five-membered closed structures containing one nitrogen and four carbon atoms) linked to each other by methine groups (―CH=). Lewis Acid Base Theory - Gilbert N. Lewis, 1920s  Lewis Acid/Base reactions:  Lewis Base: electron pair donor;  Lewis Acid: electron pair acceptor  Ligands: Lewis bases ; Metals: Lewis acids ; Coordinate covalent bonds  Metal Complexes - Formation of a complex was described as an acid - base reaction according to Lewis  Sidgwick’s Effective atomic number (EAN) rule is based on the octet theory of Lewis and this is the first attempt to account for the bonding in complexes.  In general, EAN of metal = Atomic Number − Oxidation state + CN × 2 [CN = Coordination number Total number of electrons (EAN) in the metal ion and those donated by ligands is equal to that of the next higher noble gas. Example of hexammine Rhodium (III) ion, [Rh(NH3)6]3+. Atomic number of Rhodium = 45 ; Rhodium is present in the oxidation state of +3. ∴Total number of electrons in Rh+3 = 45 − 3 = 42 Since each NH3 ligand contributes 2 electrons to the Rhodium ion, electrons contributed by 6 NH3 ligands = 6×2 = 12 ∴ The EAN of Rh+3 in the complex = 42 + 12 = 54 Atomic number (54) corresponds to the atomic number of Xenon, according to Sidgwick the complex will be stable. Werner Coordination Theory Werner's Theory: Alfred Werner, Swiss chemist put forward a theory to explain the formation of complex compounds. trans- cis- Limitations: 1. Bonding within coordination sphere. 2. Square planar (or) Tetrahedral  Primary valencies of Co in CoCl3 is 3, and oxidation state +3 Only negative ions and are ionizable. These are written outside the coordination sphere. These are non-directional and do not give any geometry to complex compound Example: [Co(NH3)6]Cl3, number of primary valencies 3, oxidation state +3 The secondary valency of metals is either by negative ions or neutral molecules or both. It represents the coordination number of the metal and are non-ionizable Secondary valencies are written inside the coordination sphere. These are directional in nature and give definite geometry to the complex. Example: [Co (NH3)6]Cl3 coordination number is 6. 10 11 Valence Bond Theory (Linus Pauling, 1931) Valence bond theory predicts that the bonding in a metal complex arises from overlap of filled ligand orbitals and vacant metal orbitals. It does not explain the color indicated by coordination compounds, the thermodynamic/kinetic stabilities of coordination complexes. Also, it does not differentiate 12 Limitations between weak/strong ligands.  Exceptions to the 18e “Rule” Early Transition Metals Middle Transition Metals Late Transition Metals 16e- and sub-16e- configurations are 18e- configurations are common 16e- and sub-16e- configurations are common common Coordination geometries higher than Coordination geometries of 6 are Coordination geometries of 5 and 6 relatively common common lower are common: d8 = square planar Tetrahedral Geometry Tetrahedral copper complex [CuCl4]2- 3d 4s 4p Cu ground state 3d94s2 Cu2+ 4 e– pairs by Cl– ions One unpaired electron - paramagnetic and attracted by magnets— High spin complexes Square Planar Geometry Square planar nickel complex [Ni(CN)4]2- 3d 4s 4p Ni (3d84s2) Ni2+ [Ni(CN)4]2- dsp2 All paired electrons – diamagnetic - weakly repelled by magnets – Low spin compelxes Octahedral sp3d2 Geometry Gives [CoF6]3– four unpaired electrons, which makes it paramagnetic and is called a high-spin complex. Ground state Co= (3d74s2) Octahedral d2sp3 Geometry [Fe(CN)6]3- Fe: (3d64s2) 3d 4s 4p Fe+3 [Fe(CN)6]3- CN– Strong ligand d2sp3 Examples 18 Bonding in Coordination Compounds  Many of the properties of metal complexes are dictated by their electronic structures. Crystal field theory (CFT)  Electronic structure can be explained by an ionic model that attributes formal charges on to the metals and ligands. This forms basis of crystal field theory (CFT), which is considered as the core concept in inorganic chemistry.  Consider bonding in a complex to be an electrostatic attraction between a positively charged nucleus and the electrons of the ligands.  Electrons on metal atom repel electrons on ligands.  Focus particularly on the d-electrons on the metal ion.  Ligand field theory (LFT) and the molecular orbital theory (MO) are considered sophisticated models as compared to CFT. LFT explains complexes, wherein, the interactions are covalent.  CFT Assumptions ⮚ Interaction between the metal ion and the ligands are purely electrostatic (ionic) ⮚ Ligands are considered as point charges ⮚ Ion-ion interaction, if the ligand is negatively charged and ion-dipole ⮚ Interaction between electrons of interaction, if the ligand is neutral the cation and those of ligands are entirely repulsive. This is ⮚ Electrons on the metal are under responsible for splitting of d repulsive from those on the orbitals. ligands ⮚ CFT does not consider the overlapping between metal and ⮚ Electrons on metal occupy those ligand orbitals. d-orbitals farthest away from the direction of approach of ligands. ⮚ The d-orbitals lose their degeneracy due to the approach of ligands during the formation of complex. 20 Octahedral Field  The d-orbitals in the isolated gaseous  discrete point charges (ligands) are metal are degenerate. allowed to interact with the metal, the  If a spherically symmetric field of negative degeneracy of the d orbitals is removed. charges is placed around the metal, these  Not all d orbitals interact with the six orbitals remain degenerate, but all of them point charges to the same extent. are raised in energy as a result of the repulsion between the negative charges  Those orbitals along the axes will be on the ligands and in the d orbitals. destabilized more than that of the orbitals that lie in between the axes.  Octahedral Complex and d-Orbital Energies ⮚ For the Oh point group, the dx2-y2, dz2 orbitals belong to the eg irreducible representation and xy, xz, yz belong to the T2g representation. ⮚ The splitting extent of these two sets of orbitals is denoted by ∆0 or 10 Dq. As the barycenter must be conserved on going from a spherical field to an octahedral field, the t2g set must be stabilized as much as the eg set is destabilized. 22  For d1-d3 systems: Hund‘s rule predicts that the electrons will not pair and occupy the t2g set.  For d4-d7 systems (2 possibilities) : Either pairing the electrons in t2g set (low spin or strong field) or electrons in eg set, higher in energy, but do not pair (high spin or weak field).  Pairing energy (P) and eg-t2g splitting (∆0 or 10 Dq) Orbital occupancy for high- and low-spin complexes of d4 through d7 metal ions high low spin: high low spin: spin: strong- spin: strong- weak- field ligand weak- field ligand field field ligand ligand Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. [Co(NH3)6]3+ > [Co(NH3)6]2+ [Co(NH3)6]3+ = 23,000 cm-1 (3d) [Rh(NH3)6]3+ = 34,000 cm-1 (4d) [Ir(NH3)6]3+ = 41,000 cm-1 (5d) ⮚ Summary ⮚ Pairing energy and CFSE (If pairing energy is greater – High spin) – Smaller the pairing energy low spin. ⮚ Weak field ligands – Cl-, OH-, F- causes smaller splitting (High spin) - ∆0 – small  Strong field ligands – CN-, CO causes larger splitting (High spin) - ∆0 – large  Higher charge – smaller size  short M-L distance  stronger interaction energy  High ∆0  Down the group, larger size, greater interaction with ligands – Higher CFSE 26 Spectrochemical series (strength of ligand interaction) Effect of ligand on splitting energy Increasing Δ Cl- < F- < H2O < NH3 < en < NO2- < CN- Increasing Δ Low spin – color variations shown with increasing CFSE (Cr3+ = 24-3-18 = d3) Spectrochemical Series – Nature of ligands  For a given ligand, the color depends on the oxidation state of the metal ion. I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO WEAKER FIELD STRONGER FIELD  Complexes of cobalt (III) show the shift in color due SMALLER D LARGER D to the ligand.  (a) CN–, (b) NO2–, (c) phen, LONGER  SHORTER  (d) en, (e) NH3, (f) gly, (g) H2O, (h) ox2–, (i) CO3 2– For a given metal ion, the color depends on the ligand. Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.  Color of Complexes Depending upon the CFSE, the light of difference wavelengths will be absorbed to effect d-d transitions and hence the complexes have different colors depending upon the wavelength of light absorbed. d-d transitions take place by the absorption of visible radiation by transition metal ions and transmitted light is colored with the complementary color of the color of the light absorbed. e.g., when red-light is removed from the white light, the eye sees its complementary color – blue-green. 29  [Co(NH3)6]3+ [Co(en)3]3+ [Co(NO2)6]3- are orange yellow while [CoF6]3- [Co(H2O)6]3+ are blue Co3+ = 27-3-18 = d6 NH3, en and NO2 are strong field ligands (more splitting – (t2g)6 (eg)0 - Large ∆o) To favour d-d transition, the energy required is high. Hence absorbs violet color or blue color and will appear yellow or orange. F- and H2O are weak field ligands (less splitting - (t2g)4 (eg)2 - small ∆o) Because of small ∆o, lower energy radiations are required for d-d transition. The complex absorbs yellow or orange (lower energy) and appear as blue. Some limitations of CFT are as follows: This theory only considers the d-orbitals of a central atom. The s and p orbits are not taken into account in this study. The theory fails to explain the behaviour of certain metals, which exhibit large splitting while others exhibit minor splitting. For example, the theory provides no explanation for why H2O is a stronger ligand than OH–. The theory excludes the possibility of pi bonding. This is a significant disadvantage because it is found in many complexes. The orbits of the ligands have no significance in the theory. As a result, it cannot explain any properties of ligand orbitals or their interactions with metal orbitals. 31  Charge transfer spectra If the splitting of d orbitals results from the effect of point charges, we would expect the anionic ligands to exert greatest effect. But most of these anionic ligands lie on the weak field end of spectrochemical series. Moreover OH- lies below the neutral water molecule. NH3 produces a strong field than H2O. But their dipole moments are in the reverse order contradictory to the original assumption of purely electrostatic interaction between metal ion and ligands. Some covalent interactions also exist (LFT) OH- having more negative charge than H2O gives a lower d orbital splitting. CN-, bipy and NO2- produce unusually high splitting but they are weak bases. This can be understood on the basis of two types of pi bonding. LMCT, MLCT. OH- having 3 unshared pair of electrons, greater tendency to back donate electrons to the metal than water having only 2 unpaired electrons. When this occurs, the net effective positive charge on the metal ion would decrease and hence the extent of splitting also decrease. Electronically unsaturated ligands CN-, bipy, CO (having fairly low lying MOs capable of accepting electrons) – metal to ligand. Hence the removal of electrons from filled d orbitals of the metal leaves the metal with high positive charge, causing greater degree of splitting than the ligands which do not form a pi bond such as NH3. Thus the splitting energy is highly dependent on the metal ion charge. 32 Applications of Coordination Compounds  Coordination compounds are of great importance.  Play many important functions in the area of analytical chemistry, metallurgy, biological systems, industry and medicine. Catalysis Extraction of metal ions Analytical chemistry (development of numerous analytical methods) Hardness estimation - Biological importance Medicinal application Industrial application Extraction / Purification of Metals  Extraction  processes of metals, like those of silver and gold, make use of complex formation.  These noble metals are extracted from their ore by the formation of cyanide complexes - dicyanoargentite(I) - [Ag(CN)2]– and dicyanoaurate (I) - [Au(CN)2]– in the presence of oxygen and water, from which the metallic forms can be separated by the addition of zinc.  Ag2S + 4NaCN  2 Na[Ag(CN)2] + Na2S  2 Na[Ag(CN)2] + Zn  Na2[Zn(CN)4] + 2Ag↓  Purification of metals can be achieved through formation and subsequent decomposition of their coordination compounds. For example, impure nickel is converted to [Ni(CO)4], which is decomposed to yield pure nickel (the Mond process). Detection of Complex Formation In the qualitative methods of analysis, complex formation is of immense importance in the identification by color change and separation of most inorganic ions.  Formation of Precipitate Ni2+ + 2 HDMG [Ni(DMG)2] + 2H+  Ni2+ and Pd2+ form insoluble colored precipitates with dimethyglyoxime Industrial Applications Coordination compounds are used as catalysts for many industrial processes. Examples rhodium complex, [(Ph3P)3RhCl], a Wilkinson catalyst – Catalytic hydrogenation of alkenes. Wilkinson catalyst Alkene Alkane Articles can be electroplated with silver and gold much more smoothly and evenly from solutions of the complexes, [Ag(CN)2]– and [Au(CN)2]– than from a solution of simple metal ions. In black and white photography, the developed film is fixed by washing with hypo solution which dissolves the non decomposed AgBr to form a complex ion, [Ag(S2O3)2]3–. Prussian blue – Mixture of hexacyanoFe(II) and Fe(III) - Fe4[Fe(CN)6]3 inks, blueprinting, cosmetics, paints (commercial coloring agents) Hardness of water  Hardness of water is estimated by titration with the sodium salt of EDTA.  During titration, the calcium and magnesium ions in hard water form the stable complexes, Calcium EDTA and Magnesium EDTA.  Hardness of water is estimated by simple titration with Na2EDTA.  The selective estimation of these ions can be done due to difference in the stability constants of calcium and magnesium complexes. 38 Organometallics – Introduction, stability, structure and applications of metal carbonyls and ferrocene 39  “Organometallic chemistry ?” Organic chemistry: Inorganic chemistry:  more or less covalent C-X bonds  primarily ionic M-X bonds, dative M-L bonds  rigid element environments  variable and often fluxional environments  fixed oxidation states  variable oxidation states Organometallics are more covalent Knowledge of inorganic and coordination chemistry is useful to understand geometries, electron counts and oxidation states of organometallic compounds Organometallics are more covalent and often less symmetric than coordination compounds, so orbital symmetry arguments are not as important Organic Chemistry: C-C / C-H : nearly covalent Cδ+-Xδ-: polar (partly ionic) reactivity dominated by nucleophilic attack at C SN2 and SN1 like reactivity Organometallics: C is the negative end of the M-C bond ("umpolung") reactivity dominated by electrophilic attack at C or nucleophilic attack at M associative and dissociative substitution at M 41 What are Organometallics?  An area which bridges organic and inorganic chemistry.  A branch of coordination chemistry where the complex has one or more metal-carbon bonds. C always is more electronegative compared to M.  The leading journals of the field define an "organometallic" compound as one in which there is a bonding interaction (ionic or covalent, localized or delocalized) between one or more carbon atoms of an organic group or molecule and a main group, transition, lanthanide, or actinide metal atom (or atoms)  Following longstanding tradition, organic derivatives of metalloids such as boron (B), silicon (Si), germanium (Ge), arsenic (As), tellurium (Te) are also included in this definition. Zeise’s Salt- The first transition metal  Discovery 1827 organometallic compound:  Structure ~ 150 years later First -bonded Organometallic Compound- Diethyl Zinc: 3 C2H5I + 3 Zn (C2H5)2Zn + C2H5ZnI + ZnI2 42  Nomenclature Hapticity: A group of contiguous atoms of a ligand are coordinated to a central atom. η - Indicated by the Greek character ' ‗ M M M M M Nomenclature of Ligands ―eta-x‖ was originally developed to indicate how many carbons of a π-system were η x coordinated to a metal center. Hapticity is another word used to describe the bonding mode of a ligand to a metal center. An η5-cyclopentadienyl ligand, for example, has all five carbons of the ring bonding to the transition metal center. ηx values for carbon ligands where the x value is odd usually indicate anionic carbon ligands (e.g., η5-Cp, η1-CH3, η1-allyl or η3-allyl, η1-CH=CH2) The # of electrons donated (ionic method of electron counting) by the ligand is usually equal to x + 1 Even ηx values usually indicate neutral carbon π-system ligands (e.g., η6-C6H6, η2- CH2=CH2, η4-butadiene, η4-cyclooctadiene) Number of electrons donated by the ligand in the even (neutral) case is usually just equal to x. η5-Cp η3-Cp η3-allyl η1-allyl 44  Nomenclature Denticity: If the coordinating atoms are not contiguous (not connected to each other), the ―Kappa‖ Κ-notation is used dppe - Ethylenebis(diphenylphosphine) k1-dppe M PPh2 PPh2 - bridging ligand 3 Organometallic compounds are classified into three types. (i) Sigma (σ) bonded organometallic compounds: In these complexes, the metal atom and carbon atom of the ligand are joined together with a sigma bond, For Example: (a) Grignard reagents, R–Mg–X where R is an alkyl or aryl group, and X is a halogen. (b) Zinc compounds of the formula R2Zn such as (C2H5)2Zn (ii) Pi (π) bonded organometallic compounds: These are the compounds of metals with alkenes, alkynes, benzene and other ring compounds. In these complexes, the metal and ligand form a bond that involves the π-electrons of the ligand. (iii) Sigma and π-bonded organometallic compounds Metal-carbonyl compounds formed between metal and carbon monoxide possess both σ-and π- bonding. Generally, oxidation state of metal atoms in these compounds is zero. Stability of Organometallic Compounds  In general terms, the stability of an organometallic compound may refer to either its thermal stability, or resistance to chemical attack (by air and moisture). Obviously, these different types of stabilities would depend both on thermodynamic as well as kinetic factors. The organometallic compounds are generally hydrolysed via nucleophilic attack by water, which is facilitated by: (1) the presence of empty low-lying orbitals on the metal (2) the polarity of metal-carbon bonds. Rate of hydrolysis is dependent on M-C bond polarity – greater the polarity, faster will be the rate 47  18 Electron Rule Main Group Elements : Octet rule ; Organometallic compounds : 18 Electron Rule – Electronic structures of many compounds are based on a total valence electron count of 18 on the central metal atom (EAN rule of Sidgwick) Provides useful guidelines to the chemistry of many organometallic complexes Majority of the organometallic compounds are with the metals from the middle of the d-block. Rule breaks down for early and late d-block metals. Changes in the number of valence electrons has a profound influence on the bonding structure and reactions of a compound. Exceptions -16electron complexes are common for Rh(I), Ir(I), Pd(0) and Pt(0)  The 18-electron Rule or Effective atomic number (EAN)  The 18e rule is a way to help us decide whether a given d-block transition metal organometallic complex is likely to be stable. Not all the organic formulas we can write down correspond to stable species. Recall: Second row elements (B, C, N, O, F) have 4 valence orbitals (1s + 3p) so they can accommodate up to 8 valence electrons--the octet rule.  For example, CH5 requires a 5-valent carbon and is therefore not stable. Stable compounds, such as CH4, have the noble gas octet, and so carbon can be thought of as following an 8e rule.  The 18e rule, which applies to many low-valent transition metal complexes, follows a similar line of reasoning. The metal now has one s, and three p orbitals, as before, but now also five d orbitals. We need 18e to fill all nine orbitals; some come from the metal, the rest from the ligands. Therefore, we can expect that the low-lying MOs can accommodate up to 18 valence electrons-- The 18-Electron Rule.  The rule states that ―thermodynamically stable transition metal organometallic compounds are formed when the sum of the metal d electrons and the electrons conventionally considered as being supplied by the surrounding ligands equals 18” 49 50 51 Counting electrons of a metal complex To count the electrons of a metal complex, one must: a) note any overall charge on the metal complex b) know the charges of the ligands bound to the metal center (ionic ligand method) c) know the number of electrons being donated to the metal center from each ligand (ionic ligand method) Similarly for a transition metal complex, the electron count is the sum of the metal valence electrons + the ligand centered electrons.  Covalent Model: # e = # metal electrons (zero valent) + # ligand electrons - complex charge Metal: The number of metal electrons equals its column number (i.e., Ti = 4e, Cr = 6e, Ni = 10e) Ligands: In general L donates 2 electrons, X donates 1 electron.  Ionic Model: # e = # metal electrons (dn) + # ligand electrons Metal: Determined based on the number of valence electrons for a metal at the oxidation state present in the complex Ligands: In general and L and X are both 2 e donors.  Complexes with 18 e- count are referred to as saturated. 52  Complexes with count lower than 18e- are called unsaturated.  Electron Counting Methodology Covalent Model Ionic Model All 18 electron organometallic compounds are generally Stable Counting electrons: is the process of determining the number of valence electrons about a metal center in a given transition metal complex. To determine the electron count for a metal complex: 1) Determine the oxidation state of the transition metal center(s) and the metal centers resulting d-electron count. To do this one must: a) note any overall charge on the metal complex b) know the charges of the ligands bound to the metal center (ionic ligand method) c) know the number of electrons being donated to the metal center from each ligand (ionic ligand method) 2) Add up the electron counts for the metal center and ligands  Classes of Ligands: Four generic classes of ligands: L: a neutral electron pair donor (CO, PR3) L-: an anionic electron pair donor (X-, H-) L+: Cationic electron pair donor (Nitrosyl, NO+) M-M: neutral 1 electron donor 55  Exceptions to the 18e “Rule” Early Transition Metals Middle Transition Metals Late Transition Metals 16e- and sub-16e- configurations are 18e- configurations are common 16e- and sub-16e- configurations are common common Coordination geometries higher than Coordination geometries of 6 are Coordination geometries of 5 and 6 relatively common common lower are common: d8 = square planar  Below is a list of common organometallic ligands and their respective electron contributions. (Ionic method) Neutral 2e donors Anionic 2e donors Anionic 4e donors Anionic 6e donors PR3 (phosphines) X- (halide) C3H5- (allyl) Cp- (cyclopentadienyl) CO (carbonyl) CH3- (methyl) O2- (oxide) O2- (oxide) alkenes CR3- (alkyl) S2- (sulfide) alkynes Ph- (phenyl) NR2- (imide) nitriles H- (hydride) CR22- (alkylidene) RnE- (silyl, germyl, alkoxo, OR- (alkoxide, bridging amido etc.) ligand) SR- (thiolate, bridging ligand) NR2- (inorganic amide, bridging ligand) PR2- (phosphide, bridging ligand) Example 1 Please note that we are using the Ionic Method of electron-counting. 95% of inorganic/organometallic chemists use the ionic method. The ionic method assigns formal charges to the metal and ligands in order to keep the ligands with an even # of electrons and (usually) a filled valence shell. Synthetically, the ionic method generally makes more sense and the one that we will use in this course. 1) There is no overall charge on the complex 2) There is one anionic ligand (CH3−, methyl group) 3) Since there is no overall charge on the complex (it Now we can do our electron counting: is neutral), and since we have one anionic ligand Re(+1) d6 present, the Re metal atom must have a +1 charge to 2 PR3 4e- compensate for the one negatively charged ligand. 2 CO 4e- The +1 charge on the metal is also its oxidation state. CH3− 2e- So the Re is the in the +1 oxidation state. We denote CH2=CH2 2e- I this in two different ways: Re(+1), Re(I), or Re. Total: 18e- 59 Other examples 16 e- 16 e- 18 e- 16 e- 16 e- 60  Metal-Carbonyls As one goes from a terminal CO- bonding mode to μ2-bridging and finally μ3-bridging, there is a relatively dramatic drop in the CO stretching frequency seen in the IR.  Standard Bonding Modes 2e- neutral donor 2e- neutral donor 3e- neutral donor 61  Types of CO-Metal bonding interactions Formation of σ-bond:  The overlapping of empty hybrid orbital on metal atom with the filled hybrid orbital on carbon atom of carbon monoxide molecule through lone pair electrons results into the formation of a M←CO σ-bond.  Formation of π-bond by back donation:  This bond is formed because of overlapping of filled dπ orbitals or hybrid dpπ orbitals of metal 62 atom with antibonding pi orbitals on CO molecule. Structure of Ni(CO)4 63 64 65 Applications 1. Metal carbonyls are used in several industrial processes. Perhaps the earliest application was the extraction and purification of nickel via nickel tetracarbonyl by the Mond process. 2. Fe(CO)5 is used for the preparation of inductors, pigments, in the production of radar- absorbing materials, and in thermal spraying. 3. Metal carbonyls are used in a number of industrially important carbonylation reactions. In the oxo process, an alkene, hydrogen gas, and carbon monoxide react together with a catalyst (such as HCo(CO)4- Cobalt tetracarbonyl hydride) to give aldehydes (hydroformylation). H2 + CO + CH3CH=CH2 → CH3CH2CH2CHO 4. Several other Metal-Carbonyl complexes have been employed in the hydrocarboxylation and hydrogenation reactions. Dicobalt octacarbonyl [Co2(CO)8] can be used for hydrosilylation of olefins also. 5. Many organometallic complexes are the sources for the pure metal particles/ metal coatings using Chemical Vapour Deposition (CVD) process. 66 Structure and Bonding in Ferrocene Mössbauer spectroscopy indicates that the iron center in ferrocene should be assigned the +2 oxidation state. Each cyclopentadienyl (Cp) ring should then be allocated a single negative charge. Thus ferrocene could be described as iron(II) bis(cyclopentadienide) Fe2+[C5H5- ]2. The number of π-electrons on each ring is then six, which makes it aromatic according to Hückel's rule. These twelve π-electrons are then shared with the metal via covalent bonding. Since Fe2+ has six d-electrons, the complex attains an 18-electron configuration, which accounts for its stability. In modern notation, this sandwich structural model of the ferrocene molecule is denoted as Fe(η5-C5H5)2. Crystallography reveals that the cyclopentadienide rings are in staggered conformation. Hybridization: d2sp3 Magnetic Nature: Diamagnetic 67 Applications of Ferrocene 1. Fuel additives: Ferrocene and its derivatives could be used as antiknock agents in the fuel for petrol engines. They are safer than previously TEL. 2. Pharmaceutical: Ferrocene derivatives have been investigated as drugs e.g. one drug has entered clinic trials, Ferroquine (7-chloro-N-(2-((dimethylamino)methyl)ferrocenyl)quinolin-4- amine), an antimalarial. Ferrocene-containing polymer-based drug delivery systems have been investigated. 3. Solid rocket propellant: Ferrocene and related derivatives are used as powerful burn rate catalysts in ammonium perchlorate composite propellant. 4. As a ligand scaffold: Chiral ferrocenyl phosphines are employed as ligands for transition- metal catalyzed reactions. Some of them have found industrial applications in the synthesis of pharmaceuticals and agrochemicals. 68 Applications of Ferrocene as a Fuel additive, a smoke suppressant and a chiral catalyst precursor Ferrocene powder Ferrocene crystals Ferox Gas & Diesel Fuel Additive is a catalyst that is an eco-friendly fuel additive and horsepower booster. It allegedly increases mileage from between 10 and 20% while also significantly reducing harmful emissions. 69  Metals in biology Contents……Metals in biology (haemoglobin, chlorophyll- structure and property) 70 Chlorophyll- Structure and Property  Structure of Chlorophyll Chlorophylls are green pigments with polycyclic, planar structures resembling the protoporphyrin system present in haemoglobin In chlorophyll, Mg2+ is the metal centre The four inward-oriented nitrogen atoms of the porphyrin ring in chlorophyll are coordinated with the Mg2+ All chlorophylls have a long phytol side chain, esterified to a carboxyl-group substituent in ring IV Chlorophylls also have a fifth five membered ring not present in heme The heterocyclic five-membered ring system that surrounds the Mg2+ has an extended polyene structure, with alternating single and double bonds Such polyenes characteristically show strong absorption in the visible region of the electromagnetic spectrum Chlorophylls have unusually high molar extinction coefficients (higher light absorbance) and are therefore particularly well-suited for absorbing visible light during 71 photosynthesis  Chloroplasts always contain both chlorophyll a and chlorophyll b  Both are green, their absorption spectra are sufficiently different that they complement each other‘s range of light absorption in the visible region  Both chlorophyll a & b absorb in the blue and red region so that the remaining green region is transmitted – hence chlorophylls are green in colour  Most plants contain about twice as much chlorophyll a as chlorophyll b  During photosynthesis, light energy is captured by pigments in the light-harvesting complex (LHC) proteins and transferred to the reaction centers of the thylakoid membrane in green plants. . 72 Role of Mg in chlorophyll  Without Mg2+ the corrin ring is fluorescent – i.e. the absorbed light energy is emitted back immediately  With Mg2+ chlorophyll becomes phosphorescent  In the case of fluorescence, the absorbed light energy is lost immediately – will not be used for chemical reaction  In the case of phosphorescence, there will be excited state of finite life time and the energy can be used for chemical reactions  The Mg2+ coordination increase the rigidity of the planar ring: The energy loss as heat due to vibration of the ring during light absorption is prevented 73 Photosynthesis Reaction Two types of photosystems cooperate in the light reactions 74 A Photosynthesis Road Map Nicotinamide Adenine Dinucleotide Phosphate Hydrogen (NADPH) Adenosine 5'-triphosphate (ATP) 75 Hemoglobin Hb  141 Amino acid  146 Amino acid Mb 153 Amino acid Hb is not an exact Four units of Hb tetramer of Mb 3 major types of Hb Hb A (Adult) Hb F (Fetal) Hb S (Sickle cell) 76  Each of these subunit polypeptides contains a heme group—an iron atom at the center of a poryphyrin ring—which reversibly binds a single O2 molecule in the ferrous state (Fe2+).  Whereas free heme binds O2 irreversibly and is converted to the ferric state (Fe3+) in the process, Hb can reversibly bind O2 because the valence state of the iron atom is protected by encapsulating the heme in the globin protein fold  Each tetrameric (α2β2) Hb can therefore reversibly bind four O2 molecules.  Oxygenation changes the electronic state of the Fe2+ heme iron, which is why the color of blood changes from the dark, purplish hue characteristic of venous blood to the brilliant scarlet of arterial blood. 77  The organic component of the heme group— the protoporphyrin—is made up of four pyrrole rings (A, B, C & D) linked by methine bridges to form a tetrapyrrole ring. Four methyl groups, two vinyl groups, and two proprionate side chains are attached.  The iron atom at the center of the protoporphyrin is bonded to the four pyrrole atoms.  Under normal conditions the iron is in the ferrous (Fe2+) oxidation state. The iron atom can form wo additional bonds, one on each side of the heme plane, called the fifth and sixth coordination sites.  The fifth coordination site is covalently bound by the imidazole side chain of the globin chain (the ―proximal histidine,‖ α87 and β92).  The sixth coordination site of the iron ion can bind O2 or other gaseous ligands (CO, NO, CN−, and H2S) 78  Role of distal histidine: Makes O2 to bind in a bent fashion and makes it difficult for CO to bind in a linear fashion.  An isolated heme binds CO 25000 times as strongly as O2 in solution. In the living system binding affinity for oxygen is reduced considerably. For CO to bind strongly, it has Tense (T) state Relaxed (R) state to bind linearly which is made difficult by distal histidine 79  Max Perutz - molecular structure of haemoglobin in 1959. Haemoglobin is a tetrameric protein. The main type of haemoglobin in adults is made up of two subunits each of ‘𝜶’ and ‘𝝱’ polypeptide chains. Each polypeptide chain is linked to a heme prosthetic group. 𝜶 subunit – It is made up of alpha polypeptide chain having 141 amino acid residues. 𝝱 subunit – It is made up of beta polypeptide chain having 146 amino acid residues. Heme group – It is an iron-containing prosthetic group, which is attached to each polypeptide chain. It contains iron in the centre of the porphyrin ring. In the quaternary structure, there is a strong interaction between 𝜶 and 𝝱 subunits. The subunits are bound together by mostly hydrophobic interactions, hydrogen bonding and a few ion pairs or salt bridges. Haemoglobin is present in two conformations, i.e. R state and T state. Oxygen has more affinity to R state and deoxyhaemoglobin is primarily present in T state Binding of oxygen is also regulated by the partial pressure of oxygen. In the lungs where pO2 is high, oxygen binds with Hb and in tissues, where pO2 is low, oxygen is released. Binding of the first oxygen molecule to the heme unit of one subunit of the deoxyhaemoglobin (T-state) causes conformational changes leading to an increase in the affinity, thereby the second molecule binds more rapidly. This type of binding is known as allosteric binding, where binding at one site affects the affinities of the remaining binding sites. 80