Module 2 Coordination Compounds and Organometallics Lecture Notes PDF

Module-2: Metal Complexes and Organometallics Fall Semester 2022-2023 September - 2022 1  Contents….. Inorganic complexes - structure, bonding and applications...

Module-2: Metal Complexes and Organometallics Fall Semester 2022-2023 September - 2022 1  Contents….. Inorganic complexes - structure, bonding and applications 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)3Cl3] --- tris(ethylenediamine)cobalt(III)chloride [Ni(PF3)4] --- tetrakis(phosphorus(III)fluoride)nickel(0) 3 Double Salt Complexes Fe(CN)2.4KCN Ferric alum (NH4)2SO4.Fe2(SO4)3.24H2O Potassium ferrocyanide Lose their identity in solution Retain their identity in solution When dissolved in water the solution The ferrocyanide ion exist as shows the properties of NH4+, SO42-, distinct entities both in the solid Fe3+ and solution  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 ligand 4K+ + [Fe(CN)6]4- (coordination sphere) N M H H N forms a coordinate covalent bond to the metal 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. EDTA-Metal complex Important biomolecules like heme and chlorophyll are porphyrins  Werner’s Work First attempt to explain the bonding in coordination complex-1893 This imaginative theory was put forward before the electron had been discovered by J.J. Thompson in 1896 Werner won the Noble prize for chemistry in 1913 All his studies were made using simple reaction chemistry without modern instrumental techniques. He explained the nature of the bonding in complexes and concluded that in complexes metals exhibit two kinds of valency The primary valency is the number of negative ions which are equivalent to the charge on the metal ion. The secondary valency is the number of ligands that are attached or coordinated to metal ion. 8 Werner Coordination Theory Werner's Theory: Alfred Werner, Swiss chemist put forward a theory to explain the formation of complex compounds. Limitations: 1. Bonding within coordination sphere. 2. Square planar (or) Tetrahedral It is unable to account for coordination compounds’ magnetic, colour, and optical properties. It didn’t explain why coordination compounds aren’t formed by all components. The directional features of bonds in coordination compounds were not explained. Lewis Acid Base Theory - Gilbert N. Lewis, 1920s ❖ Lewis Acid/Base reactions: Base: electron pair donor; 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 Rule ❖ 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.  Effective Atomic Number Each ligand donates an electron pair to the metal ion, thus forming a coordination bond. EAN predicts the number of ligands in many complexes successfully. Exceptions If the original metal has odd number of electron pair does not result noble gas configuration. Attaining the noble gas configuration is a significant factor but not necessary condition for complex formation It is also necessary to form a symmetrical structure like tetrahedral, square planar, octahedral irrespective of the number of electrons involved. 11 Effective Atomic Number- Exceptions 12  Valence Bond Theory (Linus Pauling, 1931) This theory was developed by Pauling. Coordination compounds contain complex ions, in which ligands form coordinate bonds to the metal. Thus the ligand must have a lone pair of electrons, and the metal must have an empty orbital of suitable energy available for bonding. The theory considers which atomic orbitals on the metal are used for bonding. The shape and stability of the complex are predicted. The theory has two main limitations. Most transition metal complexes are coloured, but the theory provides no explanation for their electronic spectra. The theory does not explain why the magnetic properties vary with temperature. For these reasons it has largely been superseded by the crystal field theory. 13  Shapes of d orbitals. The order of the electron orbital energy levels, starting from least to greatest, is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p 14 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 15 Limitations between weak/strong ligands. 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  Crystal field theory The crystal field theory is now much more widely accepted than the valence bond theory. It assumes that the attraction between the central metal and the ligands in a complex is purely electrostatic. If the ligand is a neutral molecule such as NH3, the negative end of the dipole in the molecule is directed towards the metal ion. The electrons on the central metal are under repulsive forces from those on the ligands. Thus the electrons occupy the d orbitals furthest away from the direction of approach of ligands. In the crystal field theory the following assumptions are made. 1. Ligands are treated as point charges. 2. There is no interaction between metal orbitals and ligand orbitals. 3.The d orbitals on the metal all have the same energy (that is degenerate) in the free atom. However, when a complex is formed the ligands destroy the degeneracy of these orbitals, i.e. the orbitals now have different energies. 20 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 ⮚ Electrons on metal occupy those overlapping between metal and d-orbitals farthest away from the ligand orbitals. direction of approach of ligands. ⮚ The d-orbitals lose their degeneracy due to the approach of ligands during the formation of complex. 22  The directions in octahedral complex. In an isolated gaseous metal ion, the five d orbitals do all have the same energy, and are termed degenerate. If a spherically symmetrical field of negative charges surrounds the metal ion, the d orbitals remain degenerate. However, the energy of the orbitals is raised because of repulsion between the Crystal field splitting of energy levels in field and the electrons on the metal. an octahedral field. In most transition metal complexes, either six or four ligands surround the metal, giving octahedral or tetrahedral structures. In both of these cases the field produced by the ligands is not spherically symmetrical. Thus the d orbitals are not all affected equally by the ligand field. 23 The approach of six ligands along the x, y, z, -x, -y and –z directions will increase the energy of the dx2-y2 and dz2 orbitals (which point along the axes) much more than it increases the energy of the dxy, dxz and dyz orbitals (which point between the axes). The size of the energy gap between the t2g and eg levels is called Crystal Field Stabilization Energy (CFSE) 24  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. 25 Ligands which cause only a small degree of crystal field splitting are termed as weak field ligand. Ligands which cause large splitting are called strong field ligands Weak field ligands I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO Strong field ligands 26 High- and low-spin complexes: (a) d4 high-spin arrangement (weak ligand field); (b) d4 low-spin arrangement (strong ligand field). 27 CFSE and pairing energy for some complexes 28 [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) 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 ▪ 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 Δ LARGER Δ to the ligand. ❖ (a) CN–, (b) NO2–, (c) phen, LONGER λ SHORTER λ (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. Relationship between the Colour observed in some Coordination Entities Wavelength of Light absorbed 32 Examples 33 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 p 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. 34 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. Organometallics – Introduction, stability, structure and applications of metal carbonyls and ferrocene 40 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 41 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 42  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 44  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” 45 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. 46  Complexes with count lower than 18e- are called unsaturated.  Ligands and the number of electrons they contribute for the metal-carbon bond CH3 in CH3-Mg-Br; C6H5 in (C6H5)3As; CH2=CH-CH2 in CH2=CH-CH2-MgBr Contributes one electron in M-C bond. Olefins in π olefinic organometallics contributes two electrons Butadiene in π butadiene organometallic complexes contributes four electrons Cyclopentadienyl group in π cyclopentadieny organometallics contributes 5 electrons. Each terminal CO contributes two electrons and each bridging CO contributes only one electron to the valency shell of the metal atom The 18 electron rule implies that the number of electrons acquired by the metal through covalent or coordinate bonding with the ligands plus the number of original electrons in (n-1)d, ns and np orbitals of the metal prior to its complexation should be equal to 18 in any of the stable complexes of the metal. This rule is helpful in rationalizing coordination numbers and studying structures of simple metal carbonyls and organometallic compounds 47 48 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- this in two different ways: Re(+1), Re(I), or Re. I Total: 18e- 49 Other examples 16 e- 16 e- 18 e- 16 e- 16 e- 50  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 51  Metal-Carbonyls Three types (two of which are important) 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 d orbitals of metal atom with antibonding π* orbitals on CO molecule. 52 Structure of Ni(CO)4 53 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, as dietary supplements in the production of radar-absorbing materials in the stealth technology, 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) 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. 54 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+[C 5H5- ] 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 55 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. 56 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. 57  Metals in biology Contents……Metals in biology (haemoglobin, chlorophyll- structure and property) 58  Metals in biology They can form complexes with proteins and other biologically active substances They can exhibit various oxidation states which are readily formed and interconverted in the environment of the cell. This interconversion of oxidation states is uniquely suited to biochemical functions. They are either very good catalysts themselves or they act as cofactors in reactions involving enzymes. Biological role of Iron As an electron carrier in plants, animals and bacteria (cytochromes) For electron transfer in plants and bacteria (Ferredoxins) For oxygen storage in muscle tissue (myoglobin) As an oxygen carrier in the blood of fish, birds, mammals (Hameoglobin) 59 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 60 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  Chlorophyll is always associated with specific binding proteins, forming light-harvesting complexes (LHCs) in which chlorophyll molecules are fixed in relation to each other, to other protein complexes, and to the membrane. 61 Role of Mg in chlorophyll  Without Mg2+ the chlorin 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 chlorin ring: The energy loss as heat due to vibration of the ring during light absorption is prevented 62 Photosynthesis Reaction Two types of photosystems cooperate in the light reactions Nicotinamide Adenine Dinucleotide Phosphate Hydrogen (NADPH) Adenosine 5'-triphosphate (ATP) 63 A Photosynthesis Road Map Nicotinamide Adenine Dinucleotide Phosphate Hydrogen (NADPH) Adenosine 5'-triphosphate (ATP) 64 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) 65  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. 66  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 67  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 68